Adsorption on Silica Surfaces

MARCEL DEKKER, INC.

NEWYORK BASEL

This book is printed on acid-free paper. Marcel Dekker, Inc. 270 Madison Avenue, New York, NY 10016 tel: 212-696-9000; fax: 212-685-4540 aster

Marcel Dekker AG Hutgasse 4, Postfach 8 12, CH-4001 Basel, Switzerland tel: 41-61-261-8482; fax: 41-61-261-8896 http://ww~.dek~er.com The publisher offers discounts on this book when ordered in bulk quantities. For more information, write to Special Sales/Professional Marketing at the headquarters address above.

Neither this booknor any part may be reproducedor transmitted in any form or by any means, electronic or mechanical, including photocopying, microfilming, and recording, or by any information storage and retrieval system,without permission in writing from the publisher. Current printing (last digit) 10987654321

Silicon is the earth’s most abundant element. Silicon dioxide, or silica, appears in the form of crystalline minerals (principally quartz), stones, etc. It is prepared industrially on a large scale in the form of powder of small particlesize and high surface area. Silica has numerous applications. For instance, quartz enters into the preparationof ceramics, whereas silica powder has many applications in our lives. Silica belongs to the limited number of finely divided mineral oxides whose surface properties have been studied extensively for years. ono graphs and specialized conferences are currently devoted to the study of silica. It is quite difficult to describe in a single volume the vast amount of information available in the literature. For this volume, we chose to concentrate our efforts ona given subject-the study of adsorptionon silicasurfaces. Adsorptionon silica is fundamentaltoimportantapplications, dernonstrated by the invaluable behavior of silica a support in chromatographic phases, catalysts, etc. More recently, silica has become an efficient additive or reinforcement agent for rubber used in the preparation of tires. Its presence decreases rolling resistance, thus saving gasoline and consequently reducing air pollution. The unique adsorption properties and reactivity of silicain part explain this peculiar behavior. This is just one timely example that demonstrates the necessity for further investigation for a better understanding of the surface and adsorption characteristics of silica. The book gives, an introduction, historical overview and discussion of the state-of-the-art of silica research. The first chapters allow nonspecialists to update their knowledge of silica surface characteristics.

he topics discussed include: (1) the surface chemistry of silica essentially ed to the existence of hydroxyl or silanol groups, which differ in nature gy, revealed by modern high performing techniques such IR spectroscopies; (2) morphology at the nanoscale and mesoscale, y quantitative image analysis and molecular computer modeling; (3) porosity determined by thermogravimetry methods; (4) surface energy sical interaction potential, shown by inverse gas chromatography, ilica is commonly exposed to water in the formof vapor or liquid and thus it is important to learn more about the state and special properties of adsorbed water molecules. Indeed, subsequent solute adsorption processes have to take into account the electrical properties of the silica-water interface and adsorbed thisgeneral introd~ction to silica, adsorption on nonmodified and silica surfaces becomes the major concern of the book. Adsorption from the gas phase, currently widely described in the literature, is in this book discussed in relation to ion-bombarde~silica samples. Thenature of theadsorptionmedium,eitherwaterororganicliquids, stronglyinfluencestheadsorptionprocess.oreover, bents, such heavymetalcationsandsurfactants, need to be considered separately from ~ ~ c r u ~ u l e cspecies, ~ l a r such polyelectrolytes and polymers used for the stabilization of silica suspension-supported enzymes and natural macrom~lecules finding interesting applications in biology. All these situations are examined in the book. ne chapter focuses on adsorption from organic solutions, whichis of particularsignificance,especiallyforliquidchromatographyseparationtechniques.Adsorptionprocessesdeterminetheperformance modified silica used chromatographicsupport.Forimportantpracticalapplications such the ~ a n u f a c t u r i n gof polymeric (silicone) materials containing silica, hot melts for adhesives, etc., the fundamental study of the adsorption from the melt deserves special attention, and two chapters are devoted to this aspect. reas adsorption on silica is beneficial for many applications, the manipon and or inhaling of finely divided, high-surface-area silica particles may generate health problems of great concern, discussed here. Finally, examples are provided detailing timely studies. First, the preparation of s so porous silica with controlled surface chemistry and subsequent adsorption characteristics is presented, and then the role of silica particles in phase transformation water and heterogeneous nucleation processes in the atmosphere is addressed. n summary, this book contains a wealthof information on adsorption that will be useful not only for newcomers in the field of silica science but also for those with general interest in adsorption phenomena. I am convinced that much progress will be made in the near future, following the d e v e l o p ~ e n of t moreexacttheoriesandtheappearance of moresophisticatedanalytical

methods (e.g., nano~icroscopies) allowing scientists to examine adsorption at the molecular level along with powerful computer s i ~ u l a t i o n , I twill not be very long before another comple~entaryvolume on this subject is necessary. I would like to thank all the contributors, who prepared excellent chapters. I acknowledge the dynamism and constant support of the publish in^ team at Marcel Deltker, Inc.

This Page Intentionally Left Blank

Preface iii Contri~utors xi 1. Brief Historical Review andCurrentState-of-the-Art Unger Dipika

of Silica

1

Infrared and N M R Characterization of the Silica Surface 9 Barry A . D. 3.

Silica SurfaceChemicalProperties G a ~ aM l . S. El Shajii

35

4.

Adsorptionon Silica Surfaces 63 Vla~i~ir

5.

Nanoscale and Mesoscale Morphology of Henri Van

6.

Ther~ogravi~etric Approach for Determining Porosity of Silica Gels 16’7

Silica Surfaces

119

Jncek

7. Surface Energetics of Silica Investigated by Inverse Gas Chromatography 205 Eugine Henri Vergelati vii

olecular Dynamics Simulation of Polyatomic Molecules in ode1 Silica Pores 243

Aleksander

ehaviorat Silica Surfaces 277 Jonathan P. and M . Dove 10. SurfaceandInterfaceStructure Jacques 11.

of Silicas 297

e ChargeandZetaPotential of Silicain ic SolventsandWater 343

Mixtures of

~arek

12. AdsorptiononIon-~ombarded Gianfranco ~erofolini,L.

Silica

rption of Heavy Metal Cations on

369

Silica

399

ification of Silica-Water Interfacial ehavior by Adsorption rfactants,Polymers,andTheirMixtures 441

Zhang olynler and Polyelectrolyte ~dsorption-Stability of ilica Suspensions 463 16. Adsorption and Chemisorption of Enzymes and Other Natural Macromolecules on Silicas 523

~ertykh

17. Adsorption on Silica Surfaces from Solution and Its Impact onChromatographicSeparationTechniques 565 D~ika V. ~ h w a l d ,

18. Adsorption of Polydimethylsiloxane Chains on Plane Silica Surfaces Liliane Liger, Hubert Hervet, ~ a r t i a Deruelle l 19. Adsorption of PDMS on BareFumed Silica Surfaces 621 Jean-Pierre Cohen

645

20.

Modulation of Silica Pathogenicity by SurfaceProcesses Bice and W i l ~ i aE.~ Wallace

21.

DesigningSurfaceChemistryinMesoporous Glen E. Fryxell Jun

22.

Fumed Silica as a Host for Study of Large Surface-to-~olume Ratio Problems in Finely Divided Aqueous Systems: ~ m ~ l i c a t i o nfor s theAtmosphere 689

Bogdan Index

747

Silica

665

This Page Intentionally Left Blank

Institut de Chimie des Surfaces et Interfaces, Mulhouse, France Department of Physics, Universityof Helsinki, Helsinki, Finland Institute of Physics,University

of Silesia,

Poland STMicroelectronics,Agrate,Italy Department

of Physics, University Joseph

Fourier, Grenoble, France

V. Lomonosov

Department of Chemistry, M. Moscow State University, Moscow, Russia

e Rhodia-Silicone,Rhodia,SaintFons,France School of EarthandAtmospheric Institute of Technology, Atlanta, Georgia

Sciences, Georgia

Institutfur ~norganische ChemieundAnalytischeChemie, Johannes Gutenberg-Universit~t, Mainz, Germany Department of Chemistry,Faculty

of Science, Ain

Department of Material Sciences, Battelle Pacific Northwest National Laboratory, Richland, Washington Department of Inorganic, Physico-Chemical, and Material Chemistry, University of Torino, Torino, Italy Department of Chemistry,SimonFraserUniversity,Burnaby, British Columbia, Canada Department of Adsorption and Planar Chromatograp~y, Maria Curie Sklodowska University, Lublin, Poland Institut fur Anorganische Chemie und Analytische Chemie, Johannes Gutenberg-Universitat, Mainz, Germany Laboratoire de Physique des Fluides France, Paris, France

Organiscis, Coll6ge de

r Applied Geology and Geochemistry, Northwest National Laboratory, Richland, Washington

InstitutfurTechnischeChemie,TechnischeUniversitat Munchen, Garching bei Munchen, Germany Department of Electrochemistry, Technical University of Lublin, Lublin, Poland, and Department of Physical Chemistry, Abo Akademi, Abo, Finland ar InstitutfurAnorganischeChemieundAnalytischeChemie, Johannes Gutenberg-Universitat, Mainz, Germany r

LaboratoiredePhysiquedesFluidesOrganisks,

Department of Material Sciences, Battelle Pacific National Laboratory, Richland, Washington

Coll6ge de

Northwest

EniChem, Centro Ricerche di Novara, Novara, Italy Department of Chemistry, University of Ottawa, Ottawa, Ontario, Canada

r InstitutdeChimiedesSurfacesetInterfaces,Mulhouse, France Universitk Franche Comtk-UFR de Techniques, BesanCon, France

Sciences et

Department of Chemistry, University of Perugia, Perugia, Italy HenryKrumbSchool New York, New York

of Mines,ColumbiaUniversity,

Institute of Surface Chemistry, National Academyof Sciences, Kiev, Ukraine Institut AnorganischeChemieundAnalytischeChemie, Johannes Gutenberg-Universitat, Mainz, Germany e* Centre de Recherche sur la Matiire Diviske, CNRS, and Universitk d'Orlkans, Orlkans, France Rhcine Poulenc Industrialisation, CR Carriires, Saint Fons, France U.S.NationalInstitute Health, Morgantown, West Virginia

OccupationalSafetyand

lskii Institute of SurfaceChemistry,NationalAcademy Sciences, Kiev, Ukraine

of

Henry Krumb School of Mines, Columbia University, New York, New York

"Current Laboratoire de Physico-Chirnie, Structurale et Macrornolkculaire, Ecole SupCrieure de Physique et Chirnie Industrielles, Paris, France.

This Page Intentionally Left Blank

Anorganische fur Institut Chemie und Analytische Chemie, Johannes ~utenberg-Universitat,Mainz, Germany

I.

1

Classical Period

11. Microparticulate Porous 111. SilicaAdvances Sol-Gel in

Silicas for Chromatographic Separations 2 2 Chemistry

IV. Mesostructured 2 Silica Materials V. Perspectives References

4 4

Following the invention of silica sols and gels in the 1920s and the manufacture of pyrogenic silica in the 1940s, finely dispersed and porous silicas became a subject of intensive research and development during the period 1950 to 1970. A number of internationally reputed scientists from both academic and industrial backgrounds havemadesubstantialcontributionsto this field, including R. K. Iler ( ~ i l m i ~ g t o 'USA), n, H. Bergna [2] ( ~ i l ~ i n g t o USA), n, A. V.Kiselev and his school (Moscow, Russia), C. Okkerse (Vlaardingen, The Netherlands), D. Barby [3] (Warrington, US), J.T.Fripiat(Orleans,France), A. Morrow(Ottawa, Canada), K. S. W. Sing (Fairfield, UK), and H. K. Ferch (Hanau,Germany), among others. During this period the formation of colloidal silicas, pyrogenic silica and silica xerogels, and precipitated silicas was studied in great depth with the aim of manufacturing products of technical importance with reproducible properties. These

materials have found widespread application as fillers, lubricants, and adsorbents. The annual production amounts to about onemillion metric tons. Comprehensive studies were also undertaken to characterizethe silica surface in order to assess its adsorption behavior and its chemical reactivity. This led to a family of silica-based derivatives with controlled surface functionalities ranging from hydrophobic to polar surface characteristics.

The aforementioned achievements prepared the way for refined silica adsorbents which were chemically tailored and hence employed in chronlatographic separations, mainly in the liquid phase [4].Between theyears 1970 and 1980, n-alkyl silanized mesoporous spherical silicas were developed, which created the basis for the advancement of so-called reverse-phase high-performance liquid chromatograPLC) [5-71. This separation method is currently widely employed for the high~resolutioll separation of complex mixturesin the chemicaland p ~ a r ~ a c e u t i c a l industries, as well as environmental, food, and toxicological analysis. The main factor contributing to its success was the development of synthetic pathways in which both the pore structural properties andparticle size distribution of the silica spheres were adjusted and controlled during the manufacturing process. Apart from hydrophobic silanized silicas, other chemically bonded silicas have been synthesized with aspecifically desired type, density, and topographyof ligands [S]. These types of tailored adsorbents are not only employed for high-resolution analytical separation but also in process chromatography where they are used for the purification and isolation of value-added synthetic and natural products.

In the 1980s a renaissancein silica chemistry appeared whichwas rooted in the field of sol-gel chemistry [S]. During the course of this period, nanostructured silicas were synthesized as nanoparticles, monoliths, and coatings. These achievements were primarily accomplished due to the use of high-resolution spectroscopic and other related techniques, which enabled researchers to elucidate the mechanism of formation at nanoscale dimensions and, thus, to analyze the surface properties in greater detail than was previously done. This field is still a major focus of research and development.

Initiated by theadvances of the te~plate-drivensynthesis of silica-rich, highly shape-selective zeolites, such as in the 1970s, researchers of the Mobil Oil Corporation recently developed anew family of ordered mesoporous silica [lo-121. This materials breakthrough was accomplished by using surfactant liquid crystals (with longn-alkyl chains) as structure-directing agents to produce a new generation of silicate and alu~inosilicate mesoporous molecularsieves (MMS) in the 1990s [3-

151 A thorpugh controlof pore size was attained with pore diameters tunable from 15 to 300 A and narrow poresize distributions. Furthermore, the materialspossess very high specific surface areas 700 m2/g) and pore volumes. reck and his team at Mobil Oil Corporation designated this family of mesoporous silicas as M41S, which consists of hexagonal phase (pbm) referred to as MCM-41, a cubic phase (Ia3d) known as M ~ ~ - 4and 8 , alamellar phase knownas MCM-50 [12]. These materials were synthesized in basic solutions in which silicate oligoanions (I-) are structured by the cationic surfactant (S"). Extensive research led to the conclusion that, during the synthesis of these materials, mesostructured silica/organic assemblies were formed owing to electrostatic interactions, which in turn formed periodic structures. ~ostsyntheticremoval of the surfactant species yielded an ordered mesoporous silica, namely ~ C ~or ~- C ~~ - l 4This 8 . discovery has, indeed, inspired scientists on a global scale to research creatively into the design, synthesis, modification, and application of these materials 16-18] Stemming from the pioneering research of Mobil scientists, a whole new range 1s have been synthesized ranging from imperfectly ordered [19], KIT [20], and MSU [21-241, to orderedsilicas, such as and SBA [28,29]. Allthese novel materials require a structuredirecting agent for their formation the latter organic molecules can be cationic [30-351, anionic, gemini 1361, orneutralsurfactants [37-391 and larger triblock copolymers [40-44]. When the template is changed, the surfactant (S) and inorganic (I) interaction is modified (S"I", S'X-I", SoIo,S-I'); the properties of the resulting material are affected, such as the porewall chemistry, surfacesilanol concentration, thermal and hydrothermal stability, and pore size. The painstaking investigation into the synthesis mechanism of the assembly of these materials enabled researchersto realize that factors such as the temperature, precursors, additives, surfactant, and concentration of reagents were certainly cting the chemistry at the inorganic/organic interface in the reaction mixture and hence the physicochemical nature of the product.Several synthesis mechanisms have been developed throughout the world by employing a number of techniques and various synthesis media [45-511. ~nderstandingthe mechanism of formation of these mesostructured materials laid the foundations for thesynthesis of silicas with highly defined pore size, pore shape, and pore dimensionality; hence, materials for which the adsorption behavio was thoroughly studied [52,53]. Moreover, these mechanistic concepts were successfully applied to other oxidic mesoporous adsorbents.As the wallsof these materials did not consist of crystalline material, research over the last five to six years has shown that one can dictate the framework composition from silicates to metallosilicates [54,55] (e.g., Al, V, Ti-silicates) to nonsilicates such as Nb, Al, Zr, Zr-P, and Ti-P oxides, and sulphates and aluminophosphates [56-651. One of the most recent advances has been the use of block coEolymers 140-441 for the preparation of mesoporous materials with pores 300 A. Furthermore, scientists have been able to fabricate mesoporous silica thin films [66-691, spheres [38,70-773, fibers [78], and monoliths [79] using sol-gel chemistry and emulsion chemistry [80-821.

er

So it appears that materialsscientists, in the truesense of the word, are researching in new realms of creativity to design mesoporous materials of controllable size, morphology, and chemical composition with pore sizes spanning the entire mesopore range for numerous current and potential applications including catalysis, sensors, separations, and optoelectronics.

With the knowledge accumulated so far it seems realistic to suggest that the scale of tailoring the pore sizes of high regularity will extend from the microporous to that of the whole mesoporous range. In addition to this, the macroscaledimensions, i.e., particle formation, can be controlled in such a way as to produce monodispersed nanoparticles, submicron-sized particles, and particles in the millimeter size range. Monolithic silicas with bimodal and multimodal pore size distributions are in their initial phase of processing using organosilicate composites as synthetic intermediates. Silica surfaces can also be functionalized and imprinted to yield a high selectivity and specificity with respect to molecular recognition. The advances in modern silica chemistry will have a substantial impact on the development of miniaturized high-resolution analyticalsystems, sensors, and other devices.

1. R. K. Iler, The Chemistry Silica, Wiley Interscience. New York, 1979. 2. H. E. Bergna. The Colloid C~emistrysilica? Advances in Chemistry Series 234, American Chemical Society, Washington D.C., 1994. D. Barby, in C~aracterization Powder S ~ r f a c e s D. Parfitt and K. S. W. Sing, eds.), Academic Press, London, 1976, pp. 353-421. 4. K. K. Unger, Porous Silica Its Prop~rtiesand Use as a Support in C o ~ ~ m Liquid n Elsevier, Amsterdam, 1979. Chro~atograp~zy, K.K. Unger, Packings and S t a t i o ~ a rPhases ~ in Clzromatograp~icT e c h ~ i ~ u e ~ s , Marcel Dekker, New York, 1990. 6. R. P. Scott, Silica and Bonded Phases, John Wiley and Sons, New York, 1993. 7. M. J.J. Hetem, Chemically Silica Swfaces in Chromatography, ~undamental Study,Hiithig Verlag, Heidelberg, 1993. 8. E. F. Vansant,P.vanderVoort,and K. Vrancken, Characteri~atio~ ~ o d ~ c a t i oofnthe Silica S ~ r j a c eElsevier, , Amsterdam, 1995. 9. C. Brinker and G. W. Schereer, Academic Press, London, 1990. 10. T.Kresge, E. Leonowicz, W. J. Roth, J. C. Vartuli, and J. Beck. Nature, 359710 (1992). 11. J. Beck, C. T,Chu, I. D. Johnson, C. Kresge, E. Leonowicz, W. J. Roth, and J. C. Vartuli, US.Patent 5,108,725 to Mobil Oil Corporation (1992). 12. J. Beck, J. C. Vartuli, W. J.Roth, M. E. Leonowicz,C. Kresge, K. D. Schmitt,C. W. Chu, D. H. Olson, E. W.Sheppard, S. B. McCullen, J. B. Higgins, and J. L. Schlenker. J. Am. Chem. Soc. 114:10834 (1992).

13. J. C. Vartuli, C. T. Kresge, M. E. Leonowicz, A. S. Chu, S. B. McCullen, I. D. Johnson, and E. W. Sheppard. Chem. Mater. 6: 1816 (1994). 14. J. C. Vartuli, C. T. Kresge, E. Leonowicz, S. Chu, S. B. McCullen, I. D. Johnson, and E. Slzeppard. Chem. Mater. 6:2070 (1994). 15. C. Vartuli, K. D. Schmitt, C. T. Kresge, J. Roth, M. E. Leonowicz, S. B. McCullen, S. D. Hellring, J. S. Beck, J. L.Schlenker, D. H. Olson,and E. Sheppard. Stud. Surf. Sci. Catal. 8453 (1994). 16. A. Corma and D. Kumar, in Trends ~ f f t e ~ i fCf l~s e ~ i(R. ' ~Catlow t ~ ~and A. Cheetham,eds.), NATO AS1Series,KluwerAcademicPublishers,498:403 997). 17. A. Sayari and P. Liu. Microporous Mater. 12:149 (1997). 18. D. M. Anotelli and J. Y. Ying. Curr. Opin. Colloid Interface Sci. 1:523 (1996). 19. P. T. Tanev, and T. J. Pinnavaia. Science 267:865 (1995). 20. Y. H. Yu, Y. Sun, and Gao. Cat. Lett. 47: 167 (1997). 21. S. A.Bagshaw, T. Kemmitt, andN. B. Milestone. Microporous Mesoporous Mater., 22:419 (1998). 22. S. Bagshaw, F. DiRenzo, and F. Fajula. Chern. Comm. 18:2209 (1996). 23. X. M. Zhang, R. Zhang, J. S. Suo, and S. B.Li. Chem. Lett. 8:755 (1988). 24. P. T. Tanev, Ye Liang, and T. J. Pinnavaia. Am. Chem. Soc. 119:8616 (1997). 25. T. Ishikawa, M. Matsuda, Yasukawa, K. Kandori, S. Inagaki, T. Fukushirna, and S. Kondo. J. Chem. Far. Trans. 92:1985 (1996). 26. K. Hamaguchi and H. Hattori. React. Kinet. Catal. Lett. 61:13 (1997). 27. S. Inagaki and Y. Fukushima. Microporous Mesoporous Mater. 21567 (1998). 28. W. Z. Zhou, H. M, A. Hunter, P. Wright, Q. F. Ge, and J. M.Thomas. J. Phys. Chern. B 1025933 (1998). Y. Zhao, Q. S. Huo, J. L.Feng,B. F. Chmelka, and C. D. Stucky. J. Am. 29. Chem. Soc. 120:6024 (1998). 30. G. D. Stucky, Q. Huo, Firouzi, B. F. Chmelka, S. Schacht, I. G. Voigt-Martin, and F. Schiith. Stud. Surf. Sci. Catal. 105:3 (1996). 31. Q. Huo, D. I. Margolese, U. Ciesla, D. G. Dernuth, P. Feng, T. Gier, P. Sieger, A. Firouzi, B. F. Chrnelka, F. Schiith, and C. D. Stucky. Chem. Mater. 6:1176 994). 32. Q. Huo, D. I.Margolese, U. Ciesla, P. Feng, T. E. Gier, P. Sieger, R. Leon, P. Petroff, F. Schiith, and G. Stucky. Nature 368:317 (1994). 33. C. Glinka, J. Nicol, G. D. Stucky, E. Ramli, D. Margolese,and Q. Huo, in in P Q ~ Q U ~S f f t e r i f f lMater. s; Res. Proc., (S. Kornarneni, D. M. Smith, and J. S. Beck, eds.), Materials Research Society, Pittsburgh, P. A. 1995, 371, p. 47. 34. L. M. Bull, D. Kumar, S. P. Millar, T. Besier, M. Janicke, C. D. Stucky, and B. F. Chmelka. Stud. Surf. Sci. Catal. 84:429 (1994). Monnier, F. Schiith, Q. Huo, D. Margolese, D. Kumar, M. 35. G. D. Stucky, Krishnamurty, P.Petroff, A. Firouzi, M. Janicke, and B. F. Chmelka. Mol. Cryst. Liq. Cryst. 240:240 (1994). 36. Q. Huo, R. Leon, P. Petroff, and G. D. Stucky. Science 268: 1324 (1995). 37. R. Mokaya and W. Jones. J. Chem. Soc. Chern. Comm. 8:981 (1996). 38. M. Griin, Lauer, and K. K. Unger. Adv. Mater. 9:254 (1997).

39. E. Prouzet and T. J. Pinnavaia. Angew. Chem. Int. Ed, Engl. 36:516 (1997). 40. A. Bagshaw, S. A. Prouzet, and J. Pinnavaia. Science269:1242(1995). 41. A. Tuel and S. Gontier. Chem. Mat. 8:114(1996). 42. P. T. Tanev and T. Pinnavaia. Science 271:1267 (1996). 43. M. Templin, A. Franck, A. D. Chesne, H. Leist, Zgang, R. Ulrich, U. Schadler, and U.Wiesner, Science 278:1795 (1997). 44. D. Zhao, J. Feng, Q. Huo, N. Melosh, C. H. Fredrickson, B. F. Chmelka, and G. D. Stucky. Science 279:548 (1998). 45. A. Monnier, F. Schiith, Q. Huo, D. Kumar, D. Margolese, R. S. Maxwell, G. D. Stucky, M. Krishnamurty, P. Petroff, A. Firouzi, M. Janicke, and B. F. Chmelka. Science 261: 1299(l 993). 46. A. Firouzi, D. Kumar, L. Bull, T. Besier, P. Sieger, Q. Huo, S. A. Walker, J. A. Zasadzinski, C. Glinka, J. Nicol, D. Margolese, C. D. Stucky, and B. F. Chmelka. Science 267:1139 (1995). 47. A. Firouzi, F. Atef, A. G. Oertli, G. D. Stucky, and B. F. Chmelka. J. Am. Chem. Soc. 119:3596 (1997). 48. C. U. Chen, S. Q. Xiao, and M. E. Davis. Microporous Mater. 4: 1 (1995). 49. C. U. Chen, H. X. Li, and E. Davis. Microporous Mater. 2: 17 (1993). 50. C. Y. Chen, S. L. Burkett, H. X. Li, and M. E. Davis, Microporous Mater. 2:27 (1993). 51. A. C. Voegtlin, F. Ruch, J. L. Guth, J. Patarin, and L. Huve. Microporous Mater. 9:95 (1997). 52. P. J. Branton, P. G. Hall, and K. S. W. Sing. J. Chem. Soc. Chem. Comm. 16:1257 (1993). 53. P. L. Llewellyn, Y. Grillet, F. Schiith, and H. Reichert. Microporous Mater. 3:345 (1994). 54. Kolodziejski,A.Corma, M.T.Navarro,and J. Pkrez-Pariente.Solid State Nucl. Magn. Reson., 2:253 (1993). 55. M. Janicke, D. Kumar, G. D. Stucky, and B. F. Chmelka. Stud. Surf. Sci. Catal. 84:243 1994). 56.A.Imhof and D. Pine. Nature, 389:948(1997). 57. F. Vaudry, Khodabandeh, and M. E. Davis. Chem. Mater. 1451 (1996). 58. A. Bagshaw and T. J. Pinnavaia. Angew. Chem. Int. Ed. Engl. 35:1102 (1996). 59. D. M. Antonelli and J. Y. Ying. Angew. Chem. Int. Ed. Engl. 34:2014 (1995). 60. D. M. Antonelli and J. Y. Ying. Angew. Chem. Int. Ed. Engl. 35:426 (1996). 61. M. Antonelli and J. Y. Ying. Chem. Mater. 85374(1996). 62. Z, R. Tian, W. Tong, J. Y. Wang, N. G. Duan, V.V. Krishan, and S. L. Suib. Science 267:926 (1997). 63. T. Holland, P, K. Isbester, C. F. Blanford, E. J. Munson, and A. Stein. J. Am. Chem. Soc. 119:6796 (1997). 64. Zhao, Luan, and L. Kevan. Chem. Comm. 11:1009 (1997). 65. U.Ciesla, Schacht, G. D. Stucky, K. K. Unger, and F. Schiith. Angew. Chem. Int. Ed. Engl. 35:541 (1996). 66. H. Sang, N. Coombs, 1. Sokolov, and G. A. Ozin. Nature 381:589 (1996). 67. I. A. Aksay, M. Trau, I. Honma, N. Yao, L. Zhou, P. Fenter, P. M.Eisenberger, and S. M. Gruner. Science 2732392 (1996).

68. S. H. Tolbert, T. E. Schaffer, J. Feng, K. Hansrna, and G. D. Stucky. Chern. Mater. 9:1962 (1997). E. Martin, M. Anderson, J. G. Odinek, and P.P. Newcomer.Langmuir 69. 13:4133 (1997). 70. Q. Huo, Feng, F. Schuth, and G. D. Stucky. Chem. Mater. 9: 14 (1997). 71. P. J. Bruinsrna, A. Y. Kim, J. Liu, and S. Baskaran. Chern. Mater. 9:2S07 (1997). 72. S. Schacht, Q. Huo, I. G. Voigt-Martin, G. D. Stucky, and F. Schuth.Science 273:768 (1996). 73. M. Grun, Lauer, K. K. Unger, A. Matsomoto, and T. Tutsurni. Adv. Mater. 3:254 (1997). 74. G. Buchel, M.Grun, K.K. Unger, A. Matsumoto, and K. Tsutsurni. Supramolec. Sci. 5:2S3 (1998). 7s. K. Schumacher, M.Grun and K. K. Unger. Microporous Mesoporous Mater., in press. 76. K. Schurnacher, C. du Fresne von Hohenesche, K. K. Unger, R. Ulrich, A. du Chesne, U. Wiesner, and H. W. Spiess. Adv. Mater., submitted. 77. M. Grun, K. K. Unger, A. Matsurnoto, and K. Tutsurni. Microporous Mesoporous Mater., in press. 78. Q. Huo, D. Zhao, J. Feng, K. Weston, S. K. Buratto, and G. D. Stucky. Adv. Mater., 9974 (1997). 79. M. T. Anderson, J. E. Martin, J. G.Odinek, P. P. Newcomer, and J. Wilcoxon. Microporous Mater 13 (1997). 80. D. Velev, T. A. Jede, and R. F. Lobo. Nature 389:447 (1997). 81. S. Schacht, Q. Huo, I. G. Voigt-Martin, G. D. Stucky, and F. Schuth. Science 273:768 (1996). 82. M. Ogawa. Chern. Soc. Chem. Comrn. 1O:1149 (1996).

This Page Intentionally Left Blank

Department of Chemistry, Universityof Ottawa, Ottawa, ent of Chemistry, Simon Fraser University, Burnaby, British C o l ~ l ~ b iCanada a,

Introduction I. TI. ~ibrationalSpectroscopy General spectral features and surface silanol groups B. Low-wavenumber spectral data C. Siloxane sites 111. NMR Spectroscopy 'H spectroscopy B. 2 9 ~spectroscopy i C. l70spectroscopy References

9 10 10 15 18 20 23 26 29

30

In a bookwhich is concerned with adsorption on silica surfaces it is first necessary to describe the nature of the silica surface, and several sections of this chapter will deal with aspectsof this problem. In this chapter, we will describe two of the most widely used spectroscopic techniqueswhich have been used to characterize thesilica surface andadsorbed species thereon;vibrationalspectroscopy and solid-state P" spectroscopy.Theformerhas beenused foroverfortyyears and is so pervasive that it is now widely used to study adsorbed species on surfaces under several guises, such as trans~issioninfrared (IR), Rarnan, diffuse reflectance IR (DRIFTS), photoacoustic IR, and total reflectance IR. On the other hand, solidstate NMR is relatively new as applied to studies of surfaces, the methodology is

rrow

still under active development, and some appropriate background material will be presented. The two techniquesare far more powerful when combined in asingle study than the sum of their contributions would suggest. This synergy results from the complementary information which each technique can produce, which in turn can point to additional experiments for the other technique. Thus, whereas vibrational spectroscopy can provide detailed information regarding hydrogen bonding, physical adsorption, and chemisorption,it cannot readily determine if a species on the surface is mobile, nor can it give a quantitative measure of the number of adsorbed species of various types without cumbersome calibrations to determine extinction coefficients which themselvescan vary accordingto surface coverage.NMR, on the otherhand,canprovidequantitativedata relative to the absolutenumbers of adsorbed species and it can provide information relative to the motional freedom of these species. Both methods yield qualitative data relative to the structure of adsorbed species, via the well-known group frequency conceptin vibrational spectroscopy, and chemical shift data in NMR. This can beused to determine the nature of the functional groups present either on the surface of silica itself, or in adsorbed species. Apartfromits historical position ofbeing developedahead of solid-state NMR,IR spectroscopyremains so populartodayas an analytical technique because current Fourier transform IR (FTIR) spectrometers are very inexpensive and an IR spectrum under favorable circumstances can be recorded in about one second. Although some special sampling techniques are required, these again are inexpensive and it is very easy to obtain spectra of samples under vacuum or in the presence of a controlled atmosphere, however reactive that atmosphere is. On theotherhand, solid-state NMR requires expensive equipment(abouttenfold more so thanFTIRinstrumentation)andthe sensitivity is much less, often requiring data acquisition times lo2 to lo4 longer. Further, if samples are to be sampledinvacuum or undercontrolledatmosphericconditions,then NMR probeshave to bespecially constructed to withstandthe rigors of spinning sealed-in-glass samples at about 3 kHzusingthemagicanglespinning(MAS) method. This chapter will give a brief overview of the utilization of both methods for studying the surface of silica. Some examples of the power of the combined use of these methods havebeen published

Vibrational spectroscopy has had a profound effect on our understanding of the surface chemistry of silica. Indeed, it has been the model system for the use of IR spectroscopy for probing the surface chemistry of oxides, the first studies having been carried out in late 1950. The books written by Hair [3] and Little [4] in the 1960s have become “classics” with respect to the early use of IR spectroscopy for studying the surface properties of silica and adsorbed species thereon.

The most common method which has been used to record the IR spectrum of silica has been to compact a high-surface-area silica (200 to 300 m2/g) into a selfsupporting disk which contains from to 10 mg of SiOz per cm2. By using this preparation technique the disk can then be conveniently mounted in a suitable vacuumchamberwheresubsequentvacuumactivationcan be carried out in order to remove adsorbed water or other contaminants, and adsorption experiments at the gas-solid interface can easily be carried out. This chapter will only be concerned with infrared studies of silica in vacuum and/or in the presence of an adsorbate. The IR spectrum of a typical self-supporting disk of a pyrogenic or fumed silica after heating under vacuum for1 h at 150°C is shown in Fig. 1A. Pyrogenic or fumed silicas (some trade names are Aerosiland Cab-0-Sil) are made by the flame hydrolysis of SiCI4 at 1000°C. These nonporous silicas have a low bulk density, and adsorbed water canbe removed by evacuation at 20°C. However, the spectral properties are identical when evacuation is carried out at 150°C; evacuation at the latter temperature is preferred in order to remove any trace impurities which maybe present. The spectrum is characterized by a sharp absorption band at 3747 cm-' with a broad tail to low wavenumber having a maximum near 3550 cm-' (see Fig. 2A for an expanded view). The sharp peak is due to the OH stretching vibration of isolated non-hydrogen-bondedSiOH or Si(OH)2 groups and the broad tail having a peak near 3550 cm" is due to these groups when they are H-bonded. The weak broad features between 2000 and 1300

0

A,IR spectrum of a self-supporting diskof silica, and B, a thin film of In both cases the silica had been heated under vacuum at 150°C for l h.

0

c m -1

IR spectra of Si02 disks: A, after activation under vacuumat 150°C, and B, 450°C; C, the difference spectrum, i.e., subtraction of curve B from A.

cm"' are due to overtones and combination modes of bulk Si02 vibrations and the region of total absorption between 1300-1000cm" and from 850-750 cm-l are due to absorption of IR radiation by bulk modes of Si02. The latter vibrational modes also are common properties of silica gel and precipitated silica, although there are differences in the region associated with theOH stretching modes of the various silanol groups. These differences will be discussed later after we have considered the silanol spectral region of a fumed silica in more detail. The IR spectrum in the OH stretching region shown on an expanded wavenumber scale in Fig. 2A. Additional weak features near 3720 and 3650 cm" can be discerned and we will comment on these later. The hydrogen-bonded hydroxyls start to condense to liberate water when silica is heated above 15OoC,the reaction being [5]:

I

Si

The majority of the H-bonded silanols condense in this way when silica isheated at 450°C under vacuum, and a typical IR spectrum is shown in Fig. There is an intensified peak at 3747 cm" due to isolated silanols and an asymmetric tail to low wavenumber.

The spectral changes are more clearly seen from “difference” spectra, that is, subtracting the spectrum for the 450°C activated sample from the 150°C one, as shown in Fig. 2C. Peaks going down represent new features which are created upon heatingfrom l50 to 450°C and peaksgoing up representspectral features lost during the heat treatment. Consider a chain of H-bonded silanols which contain an odd number of OH groups. In the example of three OH groups below, the terminal OH which is not H-bonded absorbs at 3720 cm-l, whereas those which are H-bonded absorb at about 3550 cm-’. Heating results in the process:

which creates one siloxane site, one water molecule, and an isolated SiOH group (3747 cm”). The net effect is an increase in the number of isolated SiOH groups and an increase in the intensity of the 3747 cm-’ IR band; this would not occur if the chain contained an even number of OH groups. Although IR spectroscopy is admirably able to distinguish between hydrogenbonded and free silanols, it is not able to differentiate between single or geminal hydroxyls; NMR, however, can make this distinction and this aspect will be discussed later in Section 111. A fully hydroxylated silica of whatever origin contains about 4.6 OH groups of all types per nm2. Heating to about 450°C reduces this number to about l .4/nm2. has It been known for many years13-51 that as long as the silica is not heated above about 450”C, the processes illustrated in reaction schemes (l) and (2) are reversible(wewill discuss later what happens if higher temperatures are used). There is a weak spectral feature near 3650 cm-l which can be seen in Fig. and 2C. This is known to be due to SiOH groups which are perturbed by interparticle contact; such OH groups are inaccessible to many reactants, depending on their steric size[7,8]. This effect is illustrated in Fig. 3,which showsthe stretching region of a silica which had been heated under vacuum at 150°C (Fig. 3A), the spectrum after complete reaction with TiC14 (Fig. 313) and the difference spectrum (Fig. 3C). The last spectrum (3C) is that of the silanol groups which have reacted. The reaction withTiC14 is veryfast and is complete (i.e., there no further spectral change) after about 60 [7,8]. The basic initial reactions are: Si0H

TiCI,

cl+

SiOTiCl,

HCl

(3)

Si0 TiC1, I-

TiC1,

(4)

Si0H can be seen in Fig. 3, although all of the isolated silanols (3747 cm”) and most of the H-bonded silanols (3550 cm-’) react, there is a large residual peak near 3650 cm-’, due to the silanols which are inaccessible to this reactant.

00

spectrum of a disk of fumed silica after heating under vacuum at 150°C for 1 h. B, IR spectrum of A after complete reaction with excess Tiel4. C, difference spectrum curve minus curve B showing the spectral changes as a result of the reaction with TiC14.

~e have previously shown that the number of residual inaccessible silanols (the intensity of the 3650 cm-' peak) after complete reaction with different reactive hydrogen sequestering agents increases as the steric size of the reactant increases [7,8]. Silanol groups also undergo H/D exchange with deuterated alcohols (ROD) and in the same studywe examined theeffect of the steric size of R on the extent of exchange of these perturbed SiOH or Si(OH)2 groups. Although fumed silicas are favored for IR spectroscopic studies, gels and precipitated silicas are more commonlyused for N M R studies because of their greater bulk density, and therefore, greaterconcentration of surface species perunit

volume. Indeed, muchof the workto be described later in the contextof NMR,and in later chapters, is relative to gels and precipitated silicas. The groups of Burneau and Lavalley [9-1 11 in France have carried out extensive vibrational investigations of the silanol groups on fumed silica, silica gel, and precipitated silica. Thenumber of silanol groups which are inaccessible toa reactant is generallygreateron gels and precipitatedforms ofsilica thanon fumed silicas, but this is mainly due to the mesoporous nature of these materials. Moreover, we haveconfirmed that this is true evenin the special case of a nonporous precipitated silica which was studied comparatively relative to fumed silica [8]. Although gels and precipitated silicas show spectral features which are similar to fumed silica when activated under vacuum at 450°C or higher (namely a relatively sharp IR band at 3747cm" due to isolated silanols), for room temperature or 150°C activation there is generally a much more intense "band" due to and inaccessible silanols than is observed in the caseof a fumed silica [8illustrates this point for a precipitatedsilica, and should be compared with Fig. 3 insofar as the identical manipulation was carried out, Fig. 3being relative to a fumed silica, Fig. 4 for a precipitatedsilica. Note that the region associatedwith the inaccessible silanols (3650 cm") and the H-bonded silanols (3550 cm-') is much more intense both before and after reaction with Tic& for the (particularly compare the spectra after complete reaction, Figs reader is referred to the literature for a more detailed IR comparison of these materials [8-121.

Much of the earlier IR work which was related to the adsorption of molecules on silica was concerned with studying the disappearance of features associated with SiOH stretching vibrations, and, if applicable, looking at the appearance of new features due to adsorbed species [3-51. This has been relatively straightforward if theadsorbedmoietycontainsa "light" functionalgroup,generallyone which containsahydrogenatom.Therefore,adsorbed specieswhich contain CH, or NH, functionalities are easily detected because their CH or NH stretching and deformation vibrations lie above 1300 cm". Below this, self-supporting disks of silica are opaque to IR radiation except for two regions of partial transparency between 1000 and 850 cm-' and from about '750 to 550 cm" (see Fig. M ) . (The situation is more extreme for disks of other oxides, e.g., Ai2@, Ti02, or Zr02, whichare totally absorbing below 1000 to 800 cm-l withoutanywindows of transparency 131.) Accordingly, the traditional methods of IR transmission spectroscopy do not permit access to the low-wavenumber spectral region, a region where most of the important functionalities other than hydrogen are expected to absorb, for example, most metal-oxygen modes, or modes (Y is a transition metal or S, P, etc.). Thus, when TiC14 adsorbs on silica to yield a SiOTiC13 surface species, the SiOTi modesareexpected to lie nearl000cm-' and the TiC13 stretching modes are expected to lie near 500 to 400 cm"' and both spectral regions are inaccessible to IR radiation when using the self-supporting disk IR transmission method.

IR spectrum of a disk of precipitated silica after heating under vacuum at 150°C for h. B, IR spectrum of after complete reaction with excess Tiel4. C, difference spectrum curve minus curve B showing the spectral changes as a result of the reaction with Tiel4.

It can be argued thatothermethodspermit access to the low-wavenumber spectral region, such as Raman spectroscopy and diffuse reflectance spectroscopy ( ~ R I F T S )It . is beyond the scope of this review to comment on the deficiencies of these methods; thereare manywhen compared to the well-understood transmission method, both from the experimental and theoretical points view. In 1984 one of us developed a new method for obtaining IR trans~ission spectral data for adsorbed species on silica [14]. With the then adventof FTIR methods and the ability to carry outspectral subtraction, we reasoned that if the quantity SiOz could be reduced to the extent that even the strongest bulk SiOSi mode was not totally absorbing, then one shouldbe able to detect new vibrational modeseven in spectral regionswherethe silicais strongly absorbing. The reduction in the quantity of adsorbent was achieved by forming a thin film of silica on an optically transparent substrate. The latter was initially NaCl (low-wavenumber limit about 600 cm-'), but more recently this has included ZnSe (500 cm"), KBr (400 cm"),

AgCl (300 cm"), CsI (200 cm-') and silicon 200cm">. Wide-rangemid-IR spectrometers have at best a CsI beam splitter, and the practical low-wavenumber limit of trallsmission is 200 cm-'. (Other beam splitters will permit lower wavenumber access, but at the expense of not being able to observe spectral features in the 4000 to 1000 cm" region, and for surface studies, it is essential to also be able to observe spectra in the latter spectral region.) We will show a recent example of the utility of the method. TheIR spectrum of a thin filmofsilicais shown in Fig. 1B. In spite of therebeingstronginfrared absorptionsat 1100,800, and 480 cm", spectral subtraction of this spectrum from that which is observed following chemisorption of a reactant can yield spectral features which were formerly obscured when using a self-supporting disk. We will illustrate thepower of this methodfromastudy of theadsorption of P(CH3)2Cl on silica, a reaction which eventually yields a stable SiOP= O(CH3)2 surface species [l]. (In that study, shifts and 31P solid-state NMR wereused to confirm the identity of the surface species.) Figure 5 (top) shows the spectra (1500-750 cm") of silica before (solid curve) and after (dotted curve) the reaction, respectively. Figure 5 (bottom) shows the difference spectrum at about ten times the original absorbance scale. Allof the peaks are better resolved and the peaksat 1244 cm" (peak A) and 1045 cm-' (peak B), which previously could notbe clearly seen, are characteristic of P = O and SiOP

1600

860

Top: thin-film spectra of silica before (solid line) and after (dotted line) the chemisorption of P(CH3)2CI. Bottom:differencespectrum (dotted minussolid) scaled about tenfold relative to the top spectra.

vibrational modes respectively. These bands also underwent the expected "0 shift when the reactant was adsorbed on an "0-exchanged silica [l]. The remaining modes, which are due to CH3 deformation and rocking modes, did not exhibit any shift after '*0exchange. The P= 0 and Si0P modes could never be detected using a self-supporting diskof Si02, andthis example serves to illustrate the utility and methodology behind the thin film, or "TF" technique. The method has been extensively used by Tripp and Hair 15-17] to study silane-modified silicas.

The surface of silica also contains siloxane bridges which are generally unreactive although it seems clear that some very reactive metal alkyls such as AlMe3 are capable of reacting as follows [18]: Si0Si

AlMq

Si0AlMe2

SiMe

(5)

However most of the reactive so-called hydrogen-sequestering agentswhich readily react with silanols do not react with the normal unstrained siloxane groups which are present at the surface. An exception occurswhen silica is heated under vacuumat temperatures greater than about 450°C. This causes the remaining low-wavenumber asymmetry of' the residual isolated silanol peak at 3747cm-' to disappear and the peak becomes progressively narrower and more symmetrical, as shown inFig. 6. The peak intensity also decreases as the activation temperature increases, and heating at about 1200°C willresult in complete dehydroxylation.It is been known for sometime that elimination of the isolated hydroxyls yields a surface that is increasingly hydrophobic and the rehydrationby reactions or (2) are increasingly irreversible [3,4]. Accompanying these changesis the growth of a pair of strong IR bands at 908 and 888 cm" (Fig. whose intensity increases as the temperature used to activate the silica increases, reaching a maximum when the silica is totally dehydroxylated [19221. The 908/888 cm" bands, and the weak shoulder near 932 cm", have been attributed to a highly strained four-membered siloxane ring 1231:

The site is highly reactive, facilitating the dissociative chemisorption of H 2 0 , 1,221. The reaction with ammoniaor methanol yields a new ber 3741 cm-') and SiNH2 or Si0CH3 respectively, whereas withwatertwogroupsaregenerated.Sucha dissociative chemisorption of these molecules does not occur on the normal siloxane sites which are present on the surface of silicas which have been heated at temperatures below 450°C. Finally the site also hasa Lewis acidcharacter because itcan facilitate the reversible coordination of pyridine, ammonia, or trimethylamille [20]. The number of these sites is small, being about 0. 15/nm2 for a totally dehydroxylated silica irtyfold fewer thanthe number of silanols on a fully hydroxylated silica. the IR signature of these sites canjust beseenwhen

Inf

0 .B

0.6

0.2

cm”

Infrared spectrum in the region of tbe isolated silanol group (3747 cm”) after heatinga disk in vacuum for1 h atthe indicated temperature in degrees Celsius.

silica is heated at about 500°C under vacuum, the temperatureat which silica starts to become increasingly hydrophobic. Therefore, although limited in number, the importance of these sites probably lies in their extreme reactivity and in discussing the nature of chemisorbed species formed on silicas where reaction temperatures are maintained in excess of 450°C for any length of time, it is important to realize that side reactions with these types of site, in addition to expected reactions with silanols, may occur. Thesesites have also been implicated in the anchoringof some unusual organometallic species on silica, and their importance in the use of silica as a catalyst support for surface organometallic chemistry must not be overlooked. For example,ithas recentlybeen shown [24] that rhenium derivatives can be attachedto silicawhen these sites are present, thea reactionwith 0 3 R e 0 yielding a pair of adjacent SiORe03 species, or with Me3SiORe03 toyield adjacent SiOMe3 and SiORe03. Lastly, Raman spectroscopy hasalso been able to show evidence for other cyclic siloxane sites on silica, the so-called “defect D2 and D l ” sites which are characterized by vibrations near 608 and 490 cm-’ which have been respectively attributed to six- and eight-member ring systems[25,26]. These sites are much less reactive than

1

0.8

0.6

L Q00

Infrared spectrum in the "window" region of partial transparency between 1000-800 cm-'afterheating at thesametemperatureasforFig. 5 showingthe appearance of the 908/888 cmw1 band&asthetemperature of activationunder vacuumincreases.From bottom to top, 1000°C, l lOO"C, and l1 50°C.

the highly strained four-member site andthereader literature for a discussion [25,26].

is referred to the original

can be used to study any nucleus of nonzero spin. In the context of silica es, the nuclei of interest are 'H, "Si, and 1 7 0 . The natural abundances of theseare respectively 99.99%, 4.7%,and0.04%.It is thusapparentthat 170 spectroscopy will be difficult unless enriched samples are used.Apart from nuclear abundance, theintensity an NMR signal is determined by the magnetogyric ratio of the nucleus, y, being proportional to its cube. As a result of this, the per-nucleus signal strengths for 'H, 2p§i, and 1 7 0 are in the ratio 1.0:0.0079:0.0025, and it is is by far the easiest nucleus to observe; indeed it is trivially easy with modern instrumentation to obtain the spectrum of 'H in the surface hydroxyl

layer of a high-area silica sample. Unfortunately, the useful information content of such a spectrum may be rather low infm). NMR of solids presents challenges not foundin NMR of liquids. In particular, the direct magnetic dipolar interaction between nuclei is averaged to zero by isotropic rotation in liquids. This is not the case in solids; the dipolar interaction is often the dominant spectral feature, and becauseit is anisotropic it leads to broad powder patterns with polycrystalline or amorphous samples. Similarly, chemical shift is anisotropic; in liquids it is averaged to its trace, but in solids it is not, and broad powder patterns can again arise. These problems are typically overcome by the experimental techniques of high-power decoupling, and magic angle spinning For an introduction to these techniquesand their chemical applications, see Fyfe [27]. For a more detailed discussion, see Schmidt-Rohr and Spiess [28]. A further difficulty is presented by nuclei which have spin (S for 1 7 0 ) . Such nuclei possess an electric quadrupole moment, and thus interact withelectric field gradients in their environment. This interaction is typically much larger than any other, and can lead to extremely broad powder patterns. For nuclei ofhalf oddinteger spin, the transition is not affected by the quadrupole interaction to first order [29] but the interactionis generally so strong that the second-order term yields significant broadening and line shift, in presently available fields [30]. resolution spectra of quadrupolar nuclei may be obtained via double-axis rotation [31,32], by dynamic angle spinning [33] or by two-dimensional multiple-quantum correlation spectroscopy [34]. The first two of these techniques arevery demanding experimentally, since theyinvolvehigh-speedspinning at two different angles, either simultaneously or sequentially. This is especially true for surface-chemical work where atmospheric integrity is important. The recently developed multiplequantum technique shows much promise, and the first surface studies are beginning to appear [35]. It should be noted that the second-order quadrupole interaction introduces an shift, which remains after anisotropies have been removed by one of the above techniques. This adds to the chemical shift, and the two can only be separated by experiments at different fields, or by estimation of the quadrupole term from a powder pattern. Multiple-quantum spectroscopyis an increasingly important NMR technique. It involvesthegeneration and indirect detection of multiple-quantumcoherence. “Coherence” isused here to mean an off-diagonalelement in thespindensity matrix. spin S has 2 s 1 possible values of the m quantum number, ranging from --S to S. Thus a system of k coupled spins has a total of (2s l)k spin quantum states, and is characterized by a ( 2 s (2s densitymatrix. Each quantum state may be characterized by the quantum number

i

If the density matrix contains a nonzero element connecting states and quantum coherence is said to exist in the spin system [36], where p familiar example is thesingle-quantumcoherencecreated between the states of a single spin by a 90” pulse.

a pand

Only single-quantum coherence can be directly created by RF irradiation, and detected in aspectrometer. ~ultiple-quantumcoherencecan be created by the interplay between irradiation,couplings,anddelays andcan be detected indirectly by reconversion to single-quantum coherence, after a period of evolution. The results are often presented as two-dimensional spectra. There are two potentially important uses of this in surface chemistry. ~ultiple-quantumcoherence can exist in spin systems only if spins are coupled. Thus its observation is direct evidence of such coupling ar-coupled spins in solids can be used to estimate the size of the ates This could be used, for example, probe to the existence pairs on Si02. The other application is tothe high-resolution spectr of quadrupolar nuclei, notedabove. a For odd-integral half spin, iple-quantum coherence between states of the same lml, e.g., the three-quantum coherence between and is not affected to first order by thequadrupolecoupling,buthasa different second-order dependence from the single-quantum 14) transition. Thus t~o-dimensionalcorrelation of the one- and three-quantum spectra permits the a high-resolution spectrum 1341. the great potential advantage over IR spectroscopy that intensity l~easurementis st~aightforward, inprinciple. The intensity of a signal is directly proportional to the total number of spins. Thus relative intensities are obtained from relative areas of resonance lines, and absolute intensities from comparison with a standard. A complication arises from thefinite time required forspin-lattic~ relaxation, the process of thermal equilibration of spins with their surroundings. Normally spectra other than 'H will require signal averaging of many (often thousands scans to obtain a satisfactory sigrral-to-noise ratio. In this situation, the observed signal will be proportional to (1 exp(-t/Tl)), where t is the repetition time for spectral scans, and is the spin-lattice relaxation time. (This assumes exponential relaxation, which is common, but not universal If t is not long compared to the intensity will be reduced, leading to incorrect absolute intensities, to incorrect relative intensities if different resonances are associated with different TIvalues, a not uncommon occurrence. ~nfortunately,there are many papers in the surface-chemical literaturewhere no effort has been made to demonstrate that a sufficiently long t value was used. For spin nuclei of low ~agnetogyricratio, such as 2gSi,the spin-lattice relaxation time can be embarrassingly long, leading to excessive time required for recording of a spectrum; for example delaysof 5 min were required to obtain accurate relative intensities in silicon carbide samples In such a situation it is common to utilize the technique of cross-polarization [40]. In this technique, magnetization is transferred from an abundant nucleus, usually 'H, to the observed nucleus X, e.g., 2gSi. Thishastwoadvantages: in favorable cases, the X nucleus signal is increased in theratio yH/yx, afactor offive inthe case, andthe experimental repetition time is determined not by of X, but by the usually much shorter of *H. For problemsinvolving surfaces, anotheradv cross-polarizationfrom surface-localized H, e.g., in thesurface makes possible the surface-specific observation of the X nucleus, because crosspolarization is not effective over distances exceeding a few angstroms.

i)

With the advantages of cross-polarization goes a disadvantage loss of quantitative accuracy. The signal enhancement is rarely the ideal maximum value quoted above. Cross-polarization takes place under spin-locked conditions, and rotatingframe relaxation of either H or may take place, leading to a reduction in magnetization during the time necessary for cross-polarization. The intensity of the observed signal is thus determined by a competition between cross-polarization and relaxation rates, and is a function of the experimental time interval allowed for cross-polarization. The cross-polarization rate depends on the dipolar coupling between H and X, which varies as rE&. If H and X are not directly bonded, this distance will depend on structural parameters that may notbe known. If the vector is moving, it is the average over the motion that will be relevant. If MAS is used, the rate canbe highly sensitive to the spinning speed [41], although this effect can be reduced by more complicated experiments [42]. Relaxation rates are dependent mainly on the proximity of other spins, particularly 'H, and paramagnetic impurities, and also on the motional frequencies of the relaxingnuclei. None of this inforlnation is likely to be known and there can obviously be large differences betweendifferent species on asurface. The result of this is that neither relative nor absolute intensities obtained with cross-polarization canbe considered reliable unless cross-polarization and relaxation rates have been measured, and corrections made for the effects of the latter.

Not surprisingly, the first reported N M R study on Si02 was the observation of the surface hydroxyls.O'Reilly et al. [43] observed the resonanceon a broadline instrument, and inferred from its Lorentzian shape that the surface protons constitute a dilute, randomly populated dipolar-coupled system [44]. The line shape has been further investigated by Freude et al. [45]. These authors also find Lorentzian lines, but the widths are not quantitatively accounted for by random filling of either a two- or three-dil~ensiollallattice. O'Reilly and coworkers found the line width to be independent of temperatureupto 280°C. Laterworkers [46] observed line narrowing above 150"C, which was attributed to proton diffusion on the surface. Certainly narrowing must occur when the motional frequency becomes comparable to the static line width. Possibly this occurs at different temperatures for different silica preparations. It has also been found [46] that physisorption of perdeuterated molecules leads to narrowing of thesurface proton line, even forsuch weakly interacting adsorbates as hexane. This mustbe due to induced motion of the protons, but the mechanism is unclear. Further advances in proton spectroscopy were not possible until the advent of line-narrowing techniques. Schreiber and Vaughan [47] used multiple-pulse methods on a static powder sample, and were thus able to measure the chemicalshift of surface protons, and its anisotropy. They found an absolute proton concentration of 1.6 nm"2 (773 K pretreatment), an isotropic chemical shift of 4 ppm and an anisotropy of 7 ppm. The data were fitted with an axially symmetrical shielding pattern, but slight deviations would not havebeen detected. This a small chemical shift anisotropy, compared with water, alcohols andionic hydroxyls [48] and may be indicative of rapid rotation about the Si- 0 bond.

Bernstein et al. [49] performed multiple-pulse measurements withsimilar results, and estimated that fast 0 -Si rotation with a Si 0 -H angle of 136-142" is required to explain the small observed anisotropy. These authors also performed the first MAS study of Si02, obtaining a protonline width of 150 Hz, on a 60 MHz instrument, and a measured isotropic shift of 2.5 ppm for the SiOH protons. The first MAS results with sealed samples, permitting atmospheric control,were obtained by Cay [SO]. For Si02gel evacuated at 450"C, a line width of 100 Hz was obtained on a60 MHz instrument. The isotropicshift was measured as 1.2 ppm. A remeasurement of the same sample on a 150 MHz instrument [51] found a line width of 115 Hz, and an isotropic shift of 1.8 ppm. The fact that the line width is only slightly dependent on the spectrometer field indicates that it mostly does not result from a dispersion of isotropic shifts (otherwise an increase by a factor of 2.5 would be expected). The main source of line width mustbe incomplete averagingof the homogeneous dipolar interaction [52] at the 3 kHz spinning speed employed; since the proton TIvalue is several seconds, it is unlikely that paramagnetic impurities play an important role. In contrast, pyrogenic silica evacuated to 800 or 1000°C shows [51,53] a narrow line at a shift of 1.8 ppm with width approximately proportional to field, about 0.3 ppm on both instruments. This shows that a small dispersion of isotropic chemical shifts can be observed, when the proton density is reduced sufficiently to allow complete suppressionof the dipolar couplingby MAS. An obvious step toward still better resolution is to combine MAS with multiplepulse line narrowing (CRAMPS) in an attempt toremove residual dipolar interactions. It should be notedthatmultiple-pulsetechniques also havethe effectof scaling chemical shift by a factor less than one 154,557, and it isby no means obvious that one will improve resolution by achieving a further narrowing that is greater than the shift contraction. This point has been investigated experimentally by Dec et al. [56] who find that CRAMPSgives a dramatic improvementin spectral quality for proton-rich organic solids, but only a slight improvement for untreated silica gel. From this work one would expect little or no improvement for asilica in which the proton density has been reduced by dehydration. The CRAMPS technique has been reviewed in detail by Maciel et al.1551. CRAMPS was first applied to silica by Rosenberger et al. [57]. A line width of 0.5 pprn was obtained on silica gel activated at 400°C. Only a single line was found, and these authors point out that, while separate lines of SiOH and Si(OH)2 are clearly visible in 29Si spectra present techniques do not seem able to resolve them in the ' H spectrum. (Lippmaa and coworkers [SS] observed two proton lines on Aerosil degassed at 730°C. The presentauthors have not been able to reproduce this result with Cab-0Sil, and other reports of this do not appear in the literature.) Subsequently, Macieland coworkers have appliedCRAMPS to silica gel [59,60] evacuated at room temperature and at500°C (see Fig. S), and to Cab-0-Si1 samples [61] treated at temperatures up to 650°C. Untreated silica gel shows a narrow peak at 3.5 ppm due to physisorbed water, together with a broad peak at about3.0 ppm, and a narrow peak at 1.7 ppm. The water peak is identified by the fact that it is removed on evacuation at room temperature. Evacuation at up to 200°C leaves the other two peaks intact. Evacuation at 500°C removes the broad peak, leaving the narrow one at 1.7 ppm. This behavior permits these peaks to be assigned as H-

'H CRAMPS spectra of silica gel subjected to increasingly stringent activation conditions. Top, untreated, bottom 9 h vacuum treatment at 500°C. (Reprinted with permission from Ref. 60. Copyright 1993, American Chemical Society.)

bonded (broad) and isolated (narrow) hydroxyl groups, by analogy with infrared results. Relaxation experiments were carried out [59] to determine the rate of dipolar spin flips among the protons. The results show that the hydrogen-bonded protons communicate among themselves on a time scale of l00 whereas, perhaps surprisingly, communication between these and the isolated silanols requires several milliseconds, suggesting that the average distances involved may be two to three times greater. The subsequent paper [60] deals mainly with the properties of physisorbed water. It does show, however, that the proton spectra of surface OH groups are essentially unchanged at temperatures below ambient. The CRAMPS results on Cab-0-Si1 E611 are similar to those for silica gel. Hbonded silanols are removed upon treatment at 450" C, leaving only an isolated silanol resonance at 2.0 ppm. For lower temperature treatments, the H-bonded resonance is somewhat narrower than observed on silica gel. Interestingly, an isolated OH signal is observed on surfaces covered withH20.This suggeststhat some OH groups are inaccessible to H20, a result confirmed by observation of H-Si cross-polarization on a sample exchanged with D20. Multiple-quantul~ proton NMR of Cab-0-Si1 hasbeen studied by Gerstein and coworkers [62,63]. In the first of these papers, which gives good exposition of the method,it found, surprisingly, thattreatment in liquid H20,followedby drying and Hz treatment at 400°C leads to a lower proton density, and smaller clusters of interacting protons, compared to sample in which the H20 treatment was omitted. In the subsequent paper, variety of treatment temperatures was used. For a sample treated at 120"C, the maximum observable order of multiple-

quantum coherence grew continuously with the excitation time, indicating a continuous distributionof surface protons.For samples treated at higher temperatures, a saturation of multiple-quantumorder was observed,corresponding to sevenproton clusters for 400°C treatment, and four-proton clusters for 500°C.

29SiNMR was introduced as an important surfxe-chemical technique by the classic paper of Maciel and Sindorf [64]. These authors studied the 29Sisignals arising fromcross-polarization by thesurfaceprotons ofsilicagel. Three lineswere observed (Fig. 9), at -91, 100, and -109 ppm. These were assigned, based on cross-polarization dynamics and by analogy with aqueous silicate systems, to silicons bearing two, one and no -OH groups, respectively. These species are often written in abbreviated form as Q2, Q3, and Signals were measured for a range of cross-polarization times, so that correction for rotating-frame relaxationand for differences in cross-polarization rates was possible. This leads to a ratio of 0.29 for the concentration of Si(OH)2 relative to SiOH. A figure is also reported for concentration of nonhydroxylated silicon, but it should be treated with care. Because of the r-3 dependence of dipolar coupling, cross-polarization is dominated by surface layer atoms. However, at long contact times, there will be increasing contributions from the second, and possibly deeperlayers. Such effects were recognized, but not included in the quantitative analysis of this paper.

29SiCP/MAS spectrum of silica gel. (Reprinted with permission from Ref. 64. Copyright 1980, American Chemical Society.)

A surface Si(OH), grouping might be expected to show a29Sichemical shift near -80 ppm. There is one example of this in the literature: Severin and Vankan 1651 prepared this species by treatment of partially dehydroxylated silica gel with followed by hydrolysis. The group was found to have low thermal stability, probably accounts for its absence from natural silicas. Brei [66] attempted to perform the same reaction on Aerosil, but failed to produce any silanetriol. In a subsequent paper by Sindorf and Maciel[67] silica gelsfrom several sources were compared, and their reaction with hexamethyldisilazane ( H M ~ Swas ) studied. The Me3Siline resulting from the reactionis well separated from native resonances. After determining the mass of Me3Si gravimetrically, its resonancewas used as an internal standard to determine absolute concentrations of SiOH and Si(OH)2 onthe surface. Thetotalhydroxylconcentration was nearlyconstant among the various gel samples, ranging from 4.3 to 5.2 The ratio of Si(OH):! to SiOHvaried by a factor of two among the different preparations, rangingfrom 0.09 to 0.20. Thediscrepancyincomparison to thehighervalue found in Ref. 64 was not discussed. It was found that the rate of reaction of Si(OH)2 with HMDS is three times that of SiOH. A paper by Linton et al. [68] on reactivity with trimethylchlorosilane also reports the geminal diols to be more reactive, by a factor of 17 at low fractional reaction. Sindorf and Maciel 1691 also investigated silicagel samplesevacuated at temperaturesashigh as 11OO"C, andsamplesprepared by rehydratingthese with liquid H20. ~nfortunately,cross-polarizationdynamics were not investigated above 500°C; the decreased proton density resulting from higher-te~peraturetreatmentmightchangetherelevantparameters,andonecannot be confident of quantitative accuracy the highest at temperatures used. Furthermore, no information wasgiven regarding possible loss of surface area at hightemperatures. It wasfoundthat initially geminalhydroxyls are removed more rapidly than isolated ones, leading to a minimum in their relative Concentration of about 12% for 400°C treatment temperature. From about 400700°C they are removed more slowly, leading to a relative maxil~umof 24% at 650"C, above which temperature their fractionalpopulation declines again. Treatment at 800°C (the highest temperature for which NMR data are reported) reduces both types of hydroxyl to approximately 15-20%of their initial concentrations. The rehydration appears to be kinetically controlled, and, as found by previousworkers,can beveryslow forsamplesdehydrated tothehighest temperatures.Rehydration toform geminal OH groupsappearsto be faster, since for samples rehydrated after treatment at 300"700"C, the geminal fraction is larger than in thestarting gel. Studiesinvolvingreactionwith HMDS [71] showedgoodagreement between absolute OH concentrationsdetermined by NMR and by mass loss. There is asmalldiscrepancy at low temperatures, which may be due to adsorbed H20. On a fully hydroxylated surface, all hydroxyls cannot be reacted with because of the size of the Me3Si group. If hydroxyls were removed randomly by high-temperature treatment, one would expect that eventually complete reaction would becomepossible. This is not observed, indicatingeither that dehydroxylation is not random, that the silylation reaction (which was carried out in the liquid

phase) has kinetic limitations or that all of the surface is not accessible to the reagent. Chuang et al. studied H-Si cross-polarization in silicagel samples that had been exchangedwith D 2 0 by variousprocedures.Suchspectra will show 29Si signals only from silicons that are close to an unexchanged proton, It was foundthat several percent(depending ontheexchangemethod) of protons could not be exchanged. These were attributed to "interior" silanols. striking feature of the CP spectra was the absence of a geminal diol peak, indicating that these species must befully deuterium-exchanged, and remote from the residual protons. It is clear from the results of Maciel and Sindorf that nonhydroxylated silicon cannot be reliably measured by cross-polarization. Fyfe and coworkers haveattemptedtoovercome this limitation by using 90" pulse excitation. This requires a measurement of the spin-lattice relaxation times of all species, and use of interpulse delays which are long in comparison with the longest Using the same silica gel as the Maciel group (Fisher type good agreement was obtained for the ratio of Si(OH), to SiOH. The paper can be criticized on the grounds that the authors did not assess the possible role of physisorbed H 2 0 and ambient O2 as relaxation agents for the surface layer, but not for the interior A longerrelaxationtimeforinterior(nonhydroxylated) silicon wouldmakethe apparent amount of nonhydroxylated species too small. In spite of this objection, elemental analysis (for C) of a sample silylated with Me3SiCl appears to indicate that correct results were obtained. In addition to silica gels, 29SiNMR has been applied to the study of pyrogenic silicas by several authors. In general, the lines are broader than in gels, and deconvolution procedures are often required to determine the different Si(OH)n groupings. It is known fromNMR studies of crystalline silicates that 29Sishifts are sensitive to structural influences, especially the Si-0Si bond angle. Thus the wider lines observed in pyrogenic silicas are generally attributed to greater structural randomness, associated with their high temperature of preparation. Lippmaa et al. [58] observed the presence of SiOH and Si(0H)2 on air-equilibrated Aerosil. Morrow and Gay studied Cab-0-Si1 which had been vacuum treated at temperatures up to 1OOO"C, and showed that Si(OH)2 groupings persist up to the highest temperatures. Unfortunately these papers do not present quantitative data. Lkonardelli et al. found that Aerosil has SiOH and Si(OH)2 concentrations similar to those found on a series of porous silicas (the preparation of this sample is not well described). Brei reported from cross-polarization studies that Si(OH)2 constitutes 15% of the hydroxylated silicon on Aerosil, but it is not clear whether adequate relaxation ~easurementswere made. Brei, like Sindorf and Maciel observed that a fraction of the OH groups cannot be silylated, after pretreatment of the surface up to 400°C. In addition, he showed that the unreacted silanols are associated withan infrared band at 3360-3380 cm"'. This band is known [8] to arise from hydroxyls perturbed by interparticle contact. Thus the question of whether thermal treatment randomly removes hydroxyls from noncontacting areas of the silica surface appears to remain open.

The study of Cab-0-Si1 by Liu and Maciel [61) also included 29Siobservations. Experiments included variable contact times, permitting correction of differences in cross-polarizationdynamics.An interesting result is thatuntreatedCab-0-Si1 shows a higher fraction of geminal diols than does silica gel. A ratio as high as 0.3 was reported, although this involved deconvolutionof a peak in which separate lines are not apparent, and no error analysis was given. As with silica gel, a fraction of the surface hydroxyls are not exchangeable with D20.However, this includes a substantial population of diols, not seen on gel samples. Several workers [25,78-811 have carried out 29Siexperiments on sol-gel-derived silicas. In general the spectral features are similar to those observed on other types of silica. These papers tend to focus more on the details of the hydrolysis and gelation processes, thanon thestructure of the finishedsilica. Thereappears [79,80] to be atrendtowardhigherproportions of Si(OH)2 with lower p hydrolysis, but the quantitative accuracy of these papers is uncertain.

As mentioned above, the only magnetic isotope of oxygen, 1 7 0 has spin and thus is quadrupolar. Besides introducing difficulties, the quadrupole interaction provides anothercharacterizationparameter,thequadrupolecouplingconstant, which can be observed in addition to the chemical shift. The obviously interesting surface group, which is not probed by other forms of NMR, is the siloxane bridge. The formation of such bridges, or other structures, in the dehydrosylation of silica would be a very interesting study. There is a problem with surface selectivity for quadrupolar nuclei, in that cross-polarization is less straightforward than with spin nuclei [82,83] and for unfavorable parameter values may fail, or lead to severe intensity distortions. Surface selectivity might,however, be obtained by other means, such as exchange [84] of surface OH with H2170, or by cross-relaxation with optically polarized inert gases (85,861. Liquid-phase '70spectroscopy has been reviewed by Kintzinger (871 and there are several published surveysof simple solid oxides [88,89] and of silicates. Much of the latter work, and applications to zeolites has been reviewed by Engelhardt and Michel [go]. Two crystalline phases ofsilica, cristobalite [32], and coesite [91] have been studied by hi~h-resolution1 7 0 spectroscopy. The latter possesses oxygen in five different environments with widely differing Si 0 Si bond angles. It is found that the quadrupole coupling constant, its anisotropy, and the chemical shift correlate with this bond angle. The quadrupole couplingvaries from l to 6.1 MHz as the bond angle changes from 137" to 180", whilethe chemicalshift changes from 58 to 29 ppm over the same range. There appearso far tobe only three publishedinvestigations [92-941 of high area amorphous silicas using 1 7 0 spectroscopy. Allof theseworkershaveprepared enriched silicas by reaction of SiCI4 with H2170, and the 1 7 0 is presumably uniformly distributed throughout the sample. In none of these cases is a surface area measurement reported. Geissberger and Bray [92] measured OH concentration by proton NMR, which might providean estimate of surface area. ~nfortunately their ratio of OH concentrations at degassingtemperatures of200°C and 500°Cis

markedly different fromwhat is normally observed,makingsuchestimates uncertain. Probably the surface area is 100 m2/g within a factor of two. Because of the uniform enrichment, the 90" pulse-excited spectra of these samples will be mainly characteristic of interior oxygens. These three papers report quadrupole couplings of 5.2, 5.8 and 5.0 Mhz respectively, none of which seems surprising, given the aboveresults for crystalline silicas. Walter et al. used H-Q cross-polarization to specifically observe the oxygen of surface OH groups; a distinctly different quadrupolecoupling of 4.0 MHz wasfound, similar tothat observed in a solid sample of triphenylsilanol. These results are of considerable interest, and indicate that a high-resolution study of surface-enriched silica would probably yield important results.

1. B. A. Morrow and S. J. Lang. J. Phys. Chem. 98: 133 19 (1994). 2. I. D. Gay, A. J. McFarlan, and B. A. Morrow. Phys. Chem. 95: 1360 (1991). 3. M. L. Hair, Infrare~Spectroscopy in Surjuce Cl~emistry,Marcel Dekker, N.Y., 1967. 4. L. H. Little, Infrare~Spectra A d s o r ~ Species, e~ Academic Press Inc., London, 1966. 5. B.A. Morrow. Stud. SurfaceSci. Cataly. 57A:A161(1990). 6. L. T. Zhuravlev.Langmuir3:316(1987). 7.B.A. Morrow and A. J. McFarlan. J. Non-Cryst. Solids 120:61 (1990). 8. B.A. Morrow and A. J. McFarlan. Langmuir 7:1695(1991). 9. A. Burneau, Barres, J. P. Gallas, and J. C. Lavalley. Langmuir 1364 (1990). 10. P. Gallas, J. C. Lavalley, A. Burneau, and Barres. Langmuir 7:1235(1991). 11. A. Burneau, B. Humbert, 0. Barres, J. P. Gallas, and J. C. Lavalley, in The Colloid Chemistry Silica (H. Bergna,ed.),AmericanChem.Soc.Adv.inChem. Series, 234, 1994, pp. 199-222. 12.B. A. Morrow and A. J. McFarlan. Phys. Chem. 96:1395 (1992). 13. D. T. Molapo, Ph. D. Thesis, University of Ottawa, 1998. 14. B. A. Morrow, C. P. Tripp, and R. A. McFarlane. J. Chern. Soc. Chem. Commun., 1282 (1984). 15. C. P. Tripp and M. L. Hair. Langmuir7:923(1991). 16. C. P. Tripp and M. L. Hair. Langmuir 8:1961 (1992). 17. C. P. Tripp and M. L. Hair. Phys. Chem. 975693 (1993). 18. M. Bartram, A.Michalske, and J. Rogers, Jr. J. Phys.Chem.95:4453 (1991). 19.B.A. Morrow and I. A. Cody. Phys.Chem.79:761(1975). 20. B. A. Morrow and I. A. Cody. J. Phys. Chem. 80:1995 (1976). 21.B. A. Morrow and I. A. Cody. J. Phys. Chem. 1998 (1976). 22.B. A. Morrow, I. A. Cody, and L. S. M. Lee. Phys. Chem. 80:2761 (1976). 23. T. A. Michalske and B. C. Bunker. J. Appl. Phys: 56:2686 (1984). 24. L. Scott and J.". Basset. J. Am. Chem. Soc. 116:12069(1994). 25. C. Brinker, R. J. Kirkpatrick, D. R. Tallant, B. C. Bunker, and B. Montez. J. Non-Cryst. Solids 99:418 (1988).

26. B. Humber, A. Burneau, J. P. Gallas, and J. C. Lavalley, J. Non-Cryst. Solids 143:75 (1992). 27. C. A. Fyfe, Solid State Chemists, C.F.C. Press, Guelph, Ont., 1983. ~ 28. K. Schmidt-Rohrand H. W.Spiess, Multidimensional S o l i d - S ~ ~ tNe M and Polymers, Academic Press, London, 1994. s ~ a g n e t i s Oxford ~, UniversityPress, 29. A.Abragam, The ~ r i ~ c ~ l eNuclear London, 1961, pp. 232-261. 30. C. A. Fyfe, G. C. Gobbi, J. S. Hartman, J. Klinowski, and J. M. Thomas. J. Phys. Chem. 86:1247 (1982). 31. Y. Wu, B. Q. Sun, A. Pines, A. Samoson,and E. Lippmaa. J. Magn. Reson. 89:297 (1990). 32. B: Chmelka, K. T. Mueller, A. Pines, J. Stebbins, Y. Wu, and J. W. Zwanziger. Nature 339:42 (1989). 33. K. T.Mueller, B. Q. Sun, G. C. Chingas, J. W. Zwanziger, T. Terao, and A. Pines. J. Magn. Reson. 86:470 (1990). 12779 34. A.Medek, J. S. Harwood, and L.Frydman. J. Am.Chem.Soc. (1995). K. Keskinen. J. Phys.Chem. 35. P. Sarv, C. Fernandez, J.-P. Amoureux,and 19223 (1 996). 36. R.R. Ernst, G. Bodenhausen, and A. Wokaun, in Principles Nuclear Mag~etic Resonance in One and ~imensions,Oxford University Press, Oxford, 1987,pp. 43 et seq. 37. J. Baum and A. Pines. J. Am. Chem. Soc. 108:7447 (1986). 38. R.A. Assink, M. Boslough, and R. T. Cygan. J. Magn. Reson. A 106:l 16 (1994). 39. G. R.Finlay, J. S. Hartman, M. F. Richardson, and B. L. Williams. J. Chem. Soc. Chem. Comm. 159 (1985). 40. A. Pines, M. G. Gibby, and J. S. Waugh. J. Chem. Phys. 59:569 (1973). Stejskal, J. Schaefer, and J. S. Waugh. J. Magn. Reson. 28:105 (1977). 41. E. Heidiger, B. H. Meier, and R. R. Ernst. Chem. Phys. Lett. 213:627 (1993). 42. 43. D. E. O’Reilly, H. P. Leftin, and W. K. Hall. J. Chem. Phys. 29:970 (1958). 44. See. Ref. 29, pp. 125-128. 45 D. Freude, D. Muller, and H. Schmiedel. Surf. Sci. 25:289 (1971). 46. B. Staudte. Z. Phys. Chern. (Leipzig) 255:158 (1974). 47. L. Schreiber and R. W. Vaughan. J. Catal. 58:413 (1975). 48. T.M. Duncan, A Compilation O~Chemical Shijt Anisotro~ies, The Farragut Press, Chicago,1990. 49. T.Bernstein, H. Ernst, D. Freude, I. Jiinger, B. Sauer, and B. Staudte. Z. Phys. Chem. (Leipzig) 262:1123 (1981). 50. D. Gay. Magn. Reson. 58:413 (1984). 51. Gay. Unpublished results. 52. M. Maricq and J. S. Waugh. J. Chem. Phys. 70:3300 (1979). 53. B. A. Morrow and I. D. Gay. J. Phys. Chem. 922569 (1988). 54. U. Haeberlen, High Resolution N M R in Academic Press, New York, 1976, pp. 61, 96, 106-113. G. E.Maciel, C. E.Bronnimann, and B.Hawkins.Adv.Magn.Reson.14:125 990).

rrow and

56. S. F. Dec, C. E. Bronnimann, R. A. Wind, and G. E. Maciel. J. Magn. Reson. 82:454 (1989). 57. H. Rosenberger, H. Ernst, G. Scheler, I. Junger, and R. Sonnenberger. Z . Phys. Chem. (Leipzig) 263346 (1982). 58. E.T. Lippmaa, A. V. Samoson, V. Brei, and Yu. I. Gorlov. Dokl. Akad. Nauk S.S.S.R. (Phys. Chem.) 259:403 (1981). E. Bronnimann, R. C. Zeigler, and G. E. Maciel. J. Am. Chem. Soc. 110:2023 59. (1988). 60. D. R. Kinney, 1.3. Chuang, and G. E. Maciel. J. Am. Cllern. Soc. 115:6786 (1993). 61. C. Liu and G. E. Maciel. J. Am. Chem. Soc. 103 996). 62. B. C. Gerstein, M. Pruski, and S.-J. Hwang. Analytica Chimica Acta 283:1059 (1993). 63. S.-J.Hwang, D. 0. Uner, T. S. King, M. Pruski, and B. C. Gerstein. J. Phys. Chem. 993697 (1995). 64. G. E. Maciel and D. W. Sindorf. J. Am. Chem. Soc. 102:7606 (1980). 65. J. W. Severin and J. J. Vankan. Philips J. Res. (1990). 66. V. V. Brei, J. Chem. Soc. Farad. Trans. 90:2961 (1994). 67. D. W. Sindorf and G. E. Maciel. J. Phys. Chern. 865208 (1982). 68. R. W. Linton, M. L, Miller, G. E. Maciel, and B. L. Hawkins. Surf. Interf. Anal. 7:196 (1985). 69. D. W. Sindorf and G. E. Maciel. J. Am. Chem. Soc. 105:1487 (1983). 70. R. K. Iler, in The C h e ~ i s t ~ y Silica, Wiley, New York, 1979, pp. 544-547. 71. D. W. Sindorf and G. E. Maciel. Phys. Chem. 875516 (1983). 72. Chuang, D. R.Kinney, and G. E. Maciel. J. Am. Chem. 11.523695 (1993). 73. C. A. Fyfe, G. C. Gobbi, and G. J. Kennedy. J. Phys. Chem. 89:277 (1985). 74. J. Klinowski, T.A. Carpenter, and J. M. Thomas. J. Chem. Soc. Chem. Comm. 956 (1986). 75. G. Engelhardt and R. Radeglia. Chem. Phys. Lett. 108:271 (1984). Non-Cryst.Solids 76. R. F. Pettifer, R. Dupree, I. Farnan,andU. Sternberg. 106408 988). 77. LConardelli, L. Facchini, C. Fretigny, P. Tougne, and A. P. Legrand. J. Am. Chem. Soc. 1146412 (1992). (1989). 78. A. Vega and G. W. Scherer. J. Non-Cryst. Sol. L. Walter, A. Wokaun, and A. Baiker. Mol. Phys. 71:769 (1990). 79 80. U. Darnrau, H. C. Marsmann, 0. Spormann, and P. Wang. J. Non-Cryst, 164 992). 81. R. H. Glaser, G. L.Wilkes, and C. E. Bronnimann, J. Non-Cryst.Sol.113:73 (1989). 82. A. J. Vega. J. Magn. Reson. 9650 992). 83. W.Sun, J. T. Stephen,L. D. Potter, and U. Wu. J. Magn.Reson. A 995). A. Morrow and A. Devi. Can. Chem. 48:2454 (1970). 84. 1 1604 85. T.Rddm, Appelt, R. Seydoux, E. L. Hahn, andA. Pines. Pllys. Rev. B 55: (1997). 86. E. Brunner, R. Seydoux, M. Haake, A. Pines, and J. A. Reimer. Magn. Reson. 130:145 (1998).

87. J.-P. Kintzinger, in N M R , Basic P r ~ n c ~ Z e s Progress, Vol. 17 (P. Diehl et al., eds.) Springer-Verlag, Berlin, 198 1. 88. G. L. Turner, S. E. Chung, and E. Oldfield. Magn. Reson. 64316 (1985). 89. T. J. Bastow, and S. N. Stuart. Chern. Phys. 143:459 (1990). 90. G. Engelhardt and D. Michel, inHig~-Resolurion Soli~-Srate o ~ s i l i ~ a tand es Zeolites, Wiley, Chichester, 1987, pp. 336-343. 91 P. J. Grandinetti, J. H. Baltisberger, I. Farnan, J. F. Stebbins, U. Werner, and Pines. Chern. Phys. 99:12341 (1995). E. Geissberger and P. J. Bray. J. Non-Cryst. Sol. 54121 (1983). 92. 93. T.H. Walter, G. L. Turner, and E. Oldfield. J. Magn. Reson. 76:106 (1988). 94. C. Jiiger, R. Dupree, S. C. Kohn, and M. G. Mortuza. J. Non-Cryst. Sol. 1.5595 (1 993).

This Page Intentionally Left Blank

~ e p a r t m e nof t Chemistry, Faculty of Science, Ain Shams University, Cairo, Egypt

Introduction I.

35

TI. Surface ~unctionality

36

III. HydroxylatedSurface ofSilica

38

IV.Types

40

of SurfaceHydroxyls

~ydroxylationin Relation to the Method of Preparation A.Pyrogenic silica B. Silicagel and precipitated silica

40 41 43

VI. Activity of Different Types of Surface Species A. Protolytic and ion exchange properties B. Hydrogenbonding C. Activity based on abstraction of proton from surface silanols

45 46 49 52

VII. ~ o d i ~ Surface ed of Silica

54

References

59

Silica, S i Q , is a giant atomic structurein which each silicon atom is bonded to four oxygen atoms, and eachoxygen atom to two silicon atoms in such a way that each silicon atom is at thecenter of a regular tetrahedronof oxygen atoms. Itis a solid of high melting point(1 density between 2 and 3 g/cm3 and refractive index in the range 1.5-1.6. The electronic con~gurationof the silicon atom suggests the idea of the formation of bondsbased on tetrahedrally orientedhybrid sp3 orbitals. Naturally occurring silica comprises quartz, tridyrnite, cristobalite, coesite, keatite,

i stishovite, and opal. These polymorphs exhibit a framework of tetrahedral motif; however, the pattern of linkage for each one is different, which appears in the different structural, physical, and chemical properties. Stishovite, a high-temperature form of silica, is unique among these polymorphs in being characterized by a skeleton in which silicon atoms are octahedrally 6-co-ordinated to oxygen atoms. Hybrid sp3 orbitals are presenteven if the silicon is bonded to three oxygen atoms only, but it is no longer stable when silicon remains bonded to only two oxygen atoms 111. It is known, however,that owing to free orbital silicon can co-ordinate octahedrally, provided that the ligand is small and slightly polarizable. In the case of silica the usual defects of Frenkel and Schottky type are less probable owing to the extremely high energy of chemical bonds. According to Mckinnis and Sutton [27 the breaking of Si-Q bonds becomes more frequent only at high temperatures in melts 120-250°C above the liquidus temperature of the composition. Oxygen, the bridge-forming element, has two unpaired electrons on the and 2p, orbitals, and the silicon atom possesses a vacant orbital. This usually results in more-or-less pronounced components which,if superpositioned over bonds in such a way to produce a uniform increase of the energy of all the bonds, the resultant structure is resistant to high temperatures. The presence of n levels in SiQ2 has been demonstrated by X-ray spectroscopy The four oxygen atoms in the tetrahedronwhich compete for theuse of the free orbiials of silicon reach equilibrium, resulting in equal Si-Q distances (1.62 f0.01A)of partial covalent (50%) character. TheSi- Q- Si angle changes onlyif the oxygen polarizability increased due to temperature rise, and the possible variation within wide limit results in the disappearance of the long-range order, as in the case in amorphous silica. TwoSiQ4groupshave neverbeen found to havemore than one oxygen atom in common. This is due to the stability of the electron configuration of the Si4' (complete octet) and its strong positive force field resulting from the highchargeandsmall size, whichleads to the stability of corner, shared SiQ4 tetrahedra.

It is generally acceptedthat surface ions are exposedto anasymmetrical forcefield. Their electronic orbitals being deformed, and the polarized ions havingoptical and chemical properties whichdiffer from those of the ions of the same type locatedin the symmetrical forcefield of the interior. What may notbe overlooked is the fact that the physics and chemistry of surfaces are dominated by polarization phenomena [4]. It is, therefore, not the chemical composition but primarily the surface configuration whichis the basis of the reactivity ofsilica. Surface functionality refers collectively to the chemical functional groups that terminate the bulk structure of the solid and their immediate environment. Surface groupsare constrained by the rigid three-dimensionalnetwork that backs upthe surface. Thus,some groups maybe forced into contact with each other,and in other cases some groups that would normally react with each other may be held where they cannot assume the proper orientationto react. At the surfaceof silica the structure terminates with siloxane links, Si- Q Si, as those communicated in the bulk, or silanol groups,

Si-OH. The latter is the result of incomplete condensation during the polymerization process through which the massive structureis formed and/or the tendency of the surface oxygemto formtwo bonds as other atoms in the interior. The proton is the most suitable candidate to beusedby surface oxygens to complete their bonding, which is supplied by water that is normally present in the media around the surface (during processing). Hydrous silica is therefore that formed in presence of water (liquid or vapor), and it is characterized by high population of surface hydroxyls. Anhydrous silicas are formed at high temperatures and recovered from the gas phase in dry atmosphere; its surface appears to accommodate a very small number of hydroxyls if water is omitted. Hydroussilica isthe most important when adsorption is considered, since these surface hydroxyls represent thekey of activity of silica in any process taking placeat the surface. The least important propertyof silica as a desiccant relies completely upon silica hydroxylation. The surface of hydrous silica can be visualized an extending domain of Si0- Si bonds and the relativelyhigh polarizability of the large oxygen ion compared with the small silicon ion makes it imperative that the surface is formed by oxygen ion alone. The implication is that at the uppermost layer of the extending joined tetrahedra, theoxygen atoms reside at a higher level than that accommodating the surfacesilicon atoms. When theseoxygens attach to protons to form hydroxyls on thesurfacethelatteraretherefore at higherpositions than thelinked silicons. Brunauer et al. [57 reported value of 129 f 8 ergs/cm2 at 23°C for the total energy of the silanol surface of silica which is only slightly higher than the total surface energyof liquid water, 118.5 ergs/cm2 This slight difference in surface energies made Brunauer et al. suggest that the chemical compositionof the surfaces, H 2 0 and Si02 should show considerable similarity if the hydroxyls on silica are on top and thesilicon atoms are below the OH groups. A similar suggestion was also made by Iler ["l. If the silicon atoms are 011 top, or even on the same level with the surface OHS it would be difficult to understand the close similarity between the surface energiesof water and hydrous silica. The recognition of silanol groups as surface functional groupswas easily achieved by using infrared spectroscopy;surfacegroupsare differentiated from internal groups by behavior of absorptionbandsas molecules are physically adsorbed.Anatomor nnolecule physically adsorbed (via H-bonding for example) near a surface group perturbs motions of atoms of the group, causinga shift in its stretching frequency as aresult of hydrogenbonding. Since forces responsiblefordisplacement of absorption bands act over distances of a few angstroms at most, displacementof an absorption band by a physically adsorbed vaporindicates that the groups causing the band are in exposed positions, i.e., surface species. Bands due to internal functional groups ordinarily do notchange in intensity or frequencyasvaporsare physically adsorbed. Isotopic band shift in the IR spectrum was also useful in this respect. In considering the vibration of atoms the chemical bondis often equated with the spring concept of the harmonicoscillator. The strength of the chemical bondis considered analogous to the strength of the spring in the simple harmonic oscillator, that is, the force constant. ~ u a n t u m - m e c ~ a ~ iconsiderations cal have shownthat the vibration of a diatomic molecule can be reduced to the motion of a single particle of reduced

mass, where l/pm l/ml l/m2, andm l , are the masses of the two atoms of the diatomic molecule. The relation between the force constant,k , and can be written as:

k In which c is the velocityof light (3 10" cm and i j is the wavenumber (frequency) at which the molecule has to vibrate. For two molecules of two different masses, one can write i j l / i j 2 considering k is the same in both S. If the OH group is treated as a simple diatomic harmoricoscillator, thus for and OD (D deuterium) we have (16/9 17/ 16)' (17/9)'/2 1.37, which means that the ratio between OH and OD frequencies is 1.37. Subjecting silica to D 2 0 and recording the infrared spectrum revealed the shift of the originally present band at 3750 cm-" (attributed to freely vibrating surface OH group) to appear at 2750 cm-'. The ratio 3750/2750 1.36 agrees well with the calculated one. This observation has led to the conclusionthat theisotopic band shift is due toexchange of the protonof the OH group withD of D20, and for this to occurthe OH group has to be located onthesurface as the nature of the exchange does not permit OH groups not on the surface to come into play. The use of infrared spectroscopy in combination with adsorption of different species was the subject of a very large body of research devotedeither to characterizing the surface of different silicas, or to studying the activity of the functional groups differently distributed onthesurface ofsilica. Computer-interfacedequipment made the technique invaluable in these respects. The work undertaken by Prof. Morrow and Prof. Davydov highlights this area.

Ideally, if the a r r a n g e ~ e n of t tetrahedra is to keep them stacked in an ordered design that is repeated in a regular manner all over the surface, then a long-range ordered structure has to result. Consequently, a regular distribution of the surface hydroxyls is to be expected in which all of them are equally separated making the surface of very low degree of heterogeneity. Naturally occurring polymorphs of silica exhibit variable degrees of crystallinity with which modeling of the surface structure may be a practical approach and the distribution of surface hydroxyls may beassessed depending on crystallographic data.In 1958, De Boer and Vleeskins [8] suggested a model basedon the (l 11) face of p-cristobalite or similar [aces of B-cristobalite. On such ideal surfaces, hydroxyls are separated by about 5 A andareina hexagonalarrFngement (Fig. that corresponds toan OH concentration of 4.55 per 100 A2, close to the value of4.6 they found experimentally in the case of various silica gels. Peri and Hensley [9] visualized the (100) face of p-cristobalite as silicon atoms each connected to a pairof hydroxyl groups located in rows (Fig. 2). Recent studies on silica gels using solid-state Si2' NM have shown that no single previously proposed model gives a completely adequate description of the surface structure of synthesized silica [lo]. The results indicated that silica surface is quite heterogeneous and may contain segments of surface resembling both the (1 11) and (100) faces of cristobalite. Synthesized silica, being

0 0 OH

Hexagonal arrangement of the ( I De Boer and Vleeskins. (From Ref. S.)

face of p-cristabolite as described by

aproduct of thepolymerization process, exhibits heterogeneity resulting from irregular packing of macromolecules as well as incomplete condensation between these macromolecules. At the surface the situation is further complicated by the unequal distribution of forces acting on surface atoms, in contrast to those in the bulk.Thesurfacestructure, therefore, dependsonthemethods of preparation being in wet media or in gas phase. In wet media, the packing of hydroxylated macromolecules and their mutual condensation are, due to high viscosity, subject to steric hindrances, which may result in incomplete condensation of neig~boring molecules. In gasphase,thelack of high viscosity introducesmorefreedom during formation of macromolecules. The absence of steric hindrances gives the molecules more flexibility during condensations, which results in a more ordered skeleton.

Rows of hydroxyls on the (100) face of p-cristabolite as described by Peri and Hensley. (From Ref. 9.)

4

Studies of different objectives (adsorption, hydration, dehydroxylation, and modification) have revealed that the surface hydroxyls on different silicas and even on samples of the same typeof silica are not of the same kind,and therefore are notof the same reactivity. The silicon atom on the surface may attach to one or two silanol groupsto completeitscovalency of four when its linking tothe bulk involves respectively three or two Si-0Si bridges. single Si-OHtype may be present as an isolated group where its nearest neighbors exert little or no effect on it! bonding. The distance between a pair of single isolated silanols is on average 5 as estimated from crystallographic data of crystalline silica. A single group may present near another one where both groups can mutually affect each other; the proqinenteffect is hydrogen bonding. Hydroxylswhich are separatedby more than 3.1 appear incapable of hydrogen bonding, while the strongest hydrogen Ponds must iFvolve some optimum distancewhich is considerably spaller than 3.1 A (2.4 to 2.8 A) [l l]. Silanols situated at distance shorter than2.8 A are called in the sense that they can interacteither mutually or simultaneously withan adsorbate molecule. Twosilanol groups linked to the same surfacesilicon atom are called which are not necessarily accessible to mutual hydrogen bonding, yet suitable to interact simultaneouslywith an incomingadsorbate.Theterms and are shorthand sometimes used for vicinal and geminal silanols. Two structures, Si 0 and Si (OH)3, can exist in theory, but no positive evidence has been presented to support their actual presence in any known cornpound.

The distribution of silanols on the surface a single vicinal or geminal type is a direct consequence of the way in which the building monomers are condensed during the process of polymerization, leading to the formed rigid structure. The geminal typeis the most evident pictureof incomplete condensation where twoOH groups escaped from being involved in the polymerization via the condensation process. Thecondensation ofsimple monomers, R,Si(OH),, x y 4and x 4, has been visualized to proceed through the hydroxyl groups belonging to individual molecules. This results in the evolution of two types of OH groups, those on the end Si atoms which are involved in further condensationgiving linear polymers, while the others on theSi atoms not at theend give rise to cross-linkage. The latter type of linkage causes the polymerized structure to extend in a threedimensional array which, upon solidification, forms an interlocking structure. This network constitutes the primary particles of which the massive material of silica is formed. The way in which these primary particles are formed determines to a large extent the physical and chemical properties of the resulting silica. Furthermore, the size and shape of these primary particles, which may control the density of their packing and the strength of the bonds between them, appears the basis on which the classification of different silicas depends.

Conveniently, silica gel, precipitated silica, and powders are the most important categories where adsorption is concerned. The patent literature tends to use the terms gel and precipitate in a way indicating that both types are similar. However, Barby [12] suggests that gel can bedefined as distinct froma precipitate and reported “a true precipitated silica is a dry silica with no long or short distance characteristic structure.” This definition classes precipitates as structureless silicas rather than a class of different manufacture. On the other hand, Weiser [l31 considers theonly difference between a precipitate and a gel is that a precipitate encloses onlypart of the liquid in which it is formed. Thedistinction between silica gel and silica powder also does not appear more evident. Silica gel may be considered as a coherent, rigid three-dimensional network of contiguous particles of colloidal silica, and silica powder may consist of small granules of silica gels or of coherent aggregates of submicron particles that are linked together in extremely weak networks. The outstanding book of Iler [l41 cites many other subdivisions of silica, and the interested reader should consultthis book for details. Both silica gels and powders are classified as microamorphous. Frondel [15] reported that amorphous silica is not truly amorphous but consists of regions of local atomic order, or crystals of extremely small sizes, which by careful X-ray diffraction studies appear to have the cristobalite structure. A broad, but perhaps less confusing, way of classification is one that classes silica in a high-temperature vapor-phase product as pyrogenicor fumed silica, and silica nucleated and grown from aqueous solutions supersaturated with monomeric Si(OH)4, and this includes gels and precipitates.

Anhydrous amorphous silica can be made in one of the following ways: 1. Vaporizing Si02 in an arc or plasma jet and condensing it in a stream of dry inert gas. 2. Oxidizing the morevolatile S i 0 in the vapor phase with air and condensing the Si02. 3. Oxidizing or hydrolyzing silicon compounds such as SiH4, SiC14 or HSiC1, in the vapor state with dry oxygen or in a hydrocarbon flame. The homogeneous medium causes the formation conditionsfor each silica particle to be the same, and a very narrow particle size distribution !i obtained with a diameter of 10-20 nm. The primary particles are only 10-20 A in diameter and closely packed, and the internal pores of the final particles are usually impervious to nitrogen but are detectable by water adsorption. It is concluded [l61 that the initially formed very small particles assume the translationalvelocities of large gas molecules and that final particle size is determined by collision and coalescence. The logarithm of final size is proportional to logarithm of growth time, and this isused to explain the observation that there is an increase in particle size with increase in silica concentration. The silica prepared in the vapor phase is generally of ahighdegree of hydrophobicity;however,thetypepreparedaccording to method 3 is the least hydrophobic of the three. This type of silica is commonly known as Cabosil@ in the USA (Cabosil@ registered trademark of Godfrey L.

Cabot, CO.), and Aerosil@ in Europe (Aerosil@ registered trademark of Degussa, erm many). Silanol groups at thesurface are predominantly isolated fromeach other. The high-temperature conditions make the rate of arrangement of silica tetrahedra enormously high, and since configuration with local order resembling that in the crystal must be the stable one, the tetrahedra tend to form a more ordered configuration. a consequence, a large fraction of surface silanols is situated on sites that are well separated from each other. Furthermore, pyrolysis of silicon compounds is likely to require a high temperature to achieve a sintered material, and this leads to a consequent decrease in adjacent OH groups through condensation during processing. Higher temperatures appear to favor the formation of singly isolated silanols. For instance, isolated OH groups are presentson thesurface of Aerosil OX 50 (low-surface-areaAerosilprepared ata higher temperature than other Aerosils) in a higher density than on higher-area counterparts Aerosil 200 and 300 (prepared at lower temperature than 50) The increase in particle size with increase in silica concentration favors the suggested arrangement of buildingtetrahedraasa result of preparationconditions. However,asthe resulting structure is not crystalline ona macroscale to be detected by ordinary x-ray diffraction, some incomplete condensation of silanol groups takes place, leading to the presence of a small number of geminal hydroxyls[10,18]. It is to be noted that chromatographic experiments indicated that silicasof different sourcesmayvary greatly inproperties andthat even batch-to-batch variations are quite common [19,20]. Thus, it seems that it is not valid to make statements on the silanol group types or density without taking into considerationtheaging of thesample.Storingasample of Aerosilforoneyear was found to result in a decrease in the number of isolated silanols and an increase in the total concentration of silanols as well as bridged groups [17]. This alteration is attributed to a cleavage of strained siloxane groups by absorbed water. It has been shown that Aerosil 200 contains 2.5 mg of water per gram solid in the form of physically adsorbed water just after preparation. Upon aging, this water causesrupturing of some siloxanes, transforming them to silanols, which increases the total density of silanol groups and decreases the density of isolated groups through the higher probability of hydrogen bonding. value of 1.8 OH/ nm2 for three-day-old Aerosil 200 (LiAlH4 method) corresponded to 2.65 OH/ nm2 for a one-year aged sample. With sufficiently aged Aerosil silicas(130-380 m2/g)the silanol group density is virtually independent of the specific surface area and amounts to an average value of about 2.5 OH/nm2. Again the value for a sample of Aerosil 50 is lower than this average. The concentration of surface hydroxyls measured by estrification with methanol at 200°C is found to be in therange 2-4OH/nm2 Carrottand Sing[22] reported ona slight aging of thepyrogenicare silica TIC800 overaten-yearperiodwitha loss of surfacearea andthedevelopment of interparticle mesoporosity. While it is concluded thatthe silanol density is independent of the specific surfacearea, the distribution of the groups as isolated or interacting is shown above to vary with surface area. These results indicate the strong dependence of the resulting surface ofsilica ontheconditions of preparationand storage. Therefore,the

decision about the extent of hydroxylation of a pyrogenic silica sample has to be made cautiously, taking into consideration the history of the sample.

Iler 141 drew attention to the fact that there is a close relation between formation of silica gel and of precipitate, and that factors like concentration and presence or absence of a coagulant, e.g., Na ions, are of major importance in this respect. A major difference from the pyrogenicsilica is that both gels and precipitates are wet media products. In wet media many factors intervene either separately or in an interlocking manner, making the situation more complicated than in the case of interactions in the gas phase. Different processes overlap and so they are not easily recognized. Factorslike pH, concentration,time, and temperature make the followup of the working reactions not particularly straightforward. In the silicic acid system, the formation ofgel has usually been ascribed to polymerization through the condensation of Si(OH)4 into siloxane chains, then branchingthe cross-linking toformathree-dimensionalmolecularnetwork. Flemming [23] observed that the temperature coefficients of polymerization were different in acidic and slightly basic solutions, suggesting that different mechanisms might be involved in these twodifferent pH regions. Brady et al. [24] found that the activation energy of polymerization (kcal/mol) is f0.5 at pH 5.5, 9.6 f0.3 at pH 8.5, and 14.6 f0.5 at pH 10.5 in the presence of KC1. At an early stage in the polymerization, the silicic acid is converted to spherical particles about 1-2 nm in diameter. These amorphous particles then grow in size and decrease in number by the phenomenon of Ostwald ripening. Under most conditions, further increase in molecular weight involves the chaining together of the particles into the growing microgel network until a gel is formed. The rate of condensation of monomeric silicic acid is a minimum at pH 2, the isoelectric point of silica in water; however, the process does not stop completely, also the gelling process does not suffer from extreme retardation. So it is not OH" or H' in solution which governs this condensation reaction, but it is the surface charge of the particles themselves; during this condensation the surface chargevaries. Silica particles do chain together into a uniform gel structure at this pH, and the structure is not any different from the formed one more rapidly at higher pH, e.g., 3-5. Since the formation of the SiSi bond between two colliding particles requires not only an SiOH group on one surface but also an ionized SiO- on the other surface, it appears that some ionization must occur even at the isoelectric point, pH 2. But because the overall net chargeis zero there must also be present an equal numberof positive charges of the kind present at lower pH values. The transformation of silica sols into firm homogeneous gels suggeststhat most, if not all, silica gel networks may be made up of discrete particles rather than chains of individual Si04 tetrahedra. It is not possible, however, to grow the silica particles to a large size before gelling. At low pH the silicic acid is stable long enough to cast the gel into a desired form and the particles ultimately remain less than 2-3 nm in diameter. At pH 4-10, except at low silica concentration, the presence of sodium ions causes very rapid precipitation or gelling, so that growth of discrete particles before gelling is impossible. Iler [25] showed that in gel formation the network probably extends through

the medium and forms a rigid structure before all the polysilicic acid units become attached. In a rapid precipitation or gelling, silica tetrahedra link together helterskelter with little opportunity to assume local crystalline order. This tends to produce a surface on which adjacent silanol groups areclose enough to each other and are therefore predominantly hydrogen bonded. Wetmediaprocessingmayproceed using colloidal silica, soluble silicate or hydrolysis of silicon alkoxides. The strong influence of pH was shown clearly in neutralized soluble silicate solutions [26] and colloidal silica and similar behavior is reported in nonaqueous silica gels derived from silicon alkoxides [28]. The similarity in aqueous and nonaqueous systems leads to thesuggestion that the surfacechemistryoperating in colloidal silica systems is also present in silicon alkoxide systems. In all cases porous silica is produced. When a gel is formed by mixing concentrated solutions, an immense number of nuclei are formed close to one another, so that theprimary particles are unable to reach as large a size as with more dilute reagents. Even if the latter gel is subsequently dried to the same proportion of water, the large size of its primary particles will give it a somewhat more weakly bonded structure and coarse texture. Furthermore, in the actof gelation the more concentrated, andhence more rigid, gel is unable to accommodate stresses set up between initiation and completion of gelation. Readjustments of internal linkages give rise to a spontaneous shrinkage or syneresis, which is most marked for concentrated gels. Once silica gelhas set it cannot be made to swell and repeptize by placing in pure water, i.e., it is irreversible. Similarly, when a gel shrinks on dehydration, losing avolumeapproximatelyequal to thevolume of liquid water removed, such shrinkage is irreversible. Irreversibility both of the gel formation and gel shrinkage is due toirreversibility of the condensation by which Si-0Si links are formed. The implication is that wet media processing results in porous silicas. The prevailing porosity, however, differs according to the way in which the prepared gel is subjected to thedrying step following its formation, Xerogels are derivedby natural evaporation in which the stress in the pores may be large enough to disrupt the structure. This type ofgelis 60% dense and has a reduction of 40 to in volume. When the drying process takes place in an autoclave so that the solvent is removed above its critical point (hypercritical technique), the resulting Aerogel lacks excessive cracking and stresses in the pores. The resulting gel is about 10% dense and shows little shrinkage. Eventually different textures result depending on the mode of drying, and the existence of different porosities has to have a marked effect on the final state of hydroxylation. Quantitative (total number) and qualitative (distribution of the OHS as isolated, vicinal, or geminal) differences in the extent of hydroxylation is to be expected.Hydroxyls which may be uniformly distributed on a flat surface as widely separated groups may exhibit different characteristics, and therefore reactivity, if they are contained in pores. According to Avnir and Farin (291, because of the existence of pores, clusters of OHS naturally appear in spite of a uniform distribution. As the density of isolated silanol groups decreases with increasing specific surface area, the acknowledged increase of the latter upon prevailingof microporosity supports the conclusionabout the effect of porosity on the type of resulting hydroxyls. The hypercritical technique used in the

preparation of aerogels causes thesegels to retain, to a greater extent, the porosity that is destroyed in case of xerogels. Indeed, aerogelswere found topossess a higher density of interacting hydroxyls (not isolated) than xerogels, as well as a higher total density of hydroxyls [30]. The shrinkage taking placein the case of xerogels is responsible, in part, for the lower total density of OH groups compared with the aerogels. different way of controlling the drying of a prepared gel is the use of drying control chemical agents CCAs) [31] including formamide, glycerol, and several organic acids, such as oxalic acid. This method is widely used with gels prepared from organosilicones. mentioned before, the gelling process takes place before the primary particles reach a large and frequent (common)size that results in pores of wide distribution range. The drying control agent acts on controlling the resulting pores, making them predominantly large (basic agent) or small (acidic agent) with a limited range of distribution in both cases. The effect is primarily on the formed primary particles that are large in the former case and small in the latter case. The increased population of interacting OHS in the case of small-pore gels is thus to be expected in contrast to the case of large-particle, large-pore gels. These latter gelsresemble thereforethepyrogenic silicas in which thegrowth of the primary particles results in non- or macroporosity, and the surface is characterized by a high density of isolated hydroxyl groups. It has to be noted, however, that interacting hydroxyls rnay predominate over theisolated ones on precipitated silica in spite of nonporosity [32], whichemphasizes therole of the preparation methodin determinin~the final picture of the resulting surface. The classification of solids as porous or non usually depends 011 the assessment of porosity using sorbate; nonetheless, a solid which is nonporous to NZrnay be porous to water. sconti~~lities in the solid structure which areeven smaller than those accessed by water may exist, but cannot be called pores, since there is no atom small enough to igure 3 schematically represents the act of gelation (a), leading to particles of ng intra- andinterparticle porosities (b). In (c process of condensationleads to a random distribution of 0 particle surface. The coalescence of two particles of the type results in the type shown in (d), and the surface that may result as a final product (not to scale) is shown in (e) The schem in represents the situation when a poreaccommodatesa large number of groupswhichareduetogeometric considerations accessible to hydrogen bonding in the major part, and (g) represents the different ways in which the condensation between Si04 tetrahedra may take place.

In bonding of thetype G- 0 the electrostatic repulsionamongbonded and nonbonded electron pairs on oxygen causes the the oxygen atom to value greater than 100" 1331. Thegroup whichitself has other atoms or fragrne ontheproperties of theasafunctionalgroup,itmay be that

Schematic representation of many aspects of silica gel and precipitate (for details see text).

other atoms attachedto M will also exert influences on the propertiesof OH. These effects may be through chemical bonding or cross-space, and in either case, the influence will be of electrostatic character [341. Zt might be expected that A XM-0 AX&H is electronegativity) either bond in M -0-H might break, suggesting that in such case MOH might act either as a base (break at M bond) or as an acid (breakat 0 -H bond). Thisis true in the caseof silica whereat the local environment of OH thegroup C is Si03 in which M isSi, thus 1.7, 1.4 taking 3.2,2.1, and 1.8 for 0, H,and Sirespectively. The Si-0 bond angle at oxygen is estimated to be113"

In wet media, pH plays an important role in determining the reactivity of solid surfaces. In acid solutions PZC, point of zero charge) protonation takes place, making the dangling hydroxyl positively charged, acting as an anion exchanger, and in basic solutions PZC) deprotonated sites predominate and cation adsorption can take place. For silica, the surface hydroxyls are characterized byweak acidity, where the acid dissociation constant of Si-OH is estimated to be [36-381. The high electron affinity of the Si ion causes thefree electron pairs on the

ilic

oxygen atomto beof relatively low activity, makingthe basicity also low. Protonation of Si-OH appears of very low probability; the proton association constant for the reaction 'H

Si

O H 5 Si

OH:

is characterized by log -1.9 [39], which indicates how low the probability of surface silanol protonation is. The estimated log K for the protonation of the siloxane link is extremely low, -16.9[39], which means that this group can be considered inert toward the process of protonation. Similarity of protolytic behaviors of silica and silicic acid has been emphasized [39,40], which contradicts the behavior of many metal (hydr)oxides. It has been argued 1411 that the less condensed packing of O(H) and the location of reactive groups above the surface maybe the origin of close similarity with the situation in solution. Also, it has been concluded that the titrable interface of silica/aqueous phase appearsdifferent from many other metal oxidesand hydroxides, and may be unique, and the treatment through the MUSIC approach showed that the charge attribution is different in the case of silica 1411. The charge density of the metal cation in oxide or hydroxide controls its protolytic characteristics, and this parameter, expressed as where is the metal charge and r its radius, was found to be related to the PZC through an empirical equation of the form:

in which (CN)(>is the coordination numberof oxygen in the metal oxideor hydroxide, and B are adjustable constants [42]. The successof such an empirical equation in predicting the PZC of17 metal oxides and hydroxides is illustrated in (Fig. 4) where the calculated values show agreement with experimental values through a straight line of slope 0.97 passingthroughthe origin. The equation predicts a value of PZC for silica which does not match the usual reportedexperimental values forsilica of different origins; the value forsilica lies near the straight line passing through the values for other samples; however, it belongs to the upper abscissa scale (negative values). This observation again emphasizesthat even when the protolytic behavior of manymetaloxides and hydroxidescan beassessed empirically, silica shows an odd behavior. The exchange of metal ions with the weakly acidic silanol groups of silica has been visualized as:

M"'

m(-SiO€€)

M(OSi-):-"

mH+

(3)

in which m is 1 or 2, where the exchangeis reversible in acid solutions [43-45], and the bonding in the formed surface complex is essentially ionic [46-49]. The Eigand ~ r o ~ e r t i e s the s ~ r f a c esi~anoEs arenot basically changed by the u t t ~ ~ h silica. e d On the other hand, factors like pH, charge density on the mental cation, the chelating effect of cations, and the stability constant of the metal-hydroxo complex, appear to be the controlling factors. The surface activity of silica toward adsorption of cations, together with its high specific surface have oriented the attention of the

-12

-8

-4

Relationbetweenexperimental and calculated PZCs using Eq. (2). The value for silica follows the upper abscissa. (Based on data from Ref. 42.) researchers in the field of catalysis toward its use as a catalyst support. The objective of a supportedcatalyst is to benefit from the high surface area of the support to accommodate a highly dispersed phase of the active substance, to achieve a longacting high activity, in contrast to the short halflife in the case of an unsupported catalyst. The supporting process involves the loading of the desired species on the surface of the support through an impregnationstep, followed by calcination, then reduction to reach to the metallic state. As the particle sizeof the phase to be reduced decreases, theafter-reductionproduct exhibits ahigh dispersity with expectedhigh activity. In most cases, however,silica-supported catalysts were not acknowledged as good catalysts, since silica primarily interacts weakly with the starting materials, for which, upon calcination, the formed weak bonds are broken, resulting in an aggregated phase that increases in size. At the reduction step this phase results in a poorly dispersedspecies with a consequent low catalytic activity. Nevertheless, the weaknessof interaction between silica hydroxyls and the starting material has been found to show some dependence on the distribution groups. In the system Mo-silica it was noted that the weakest interaction is that involved MOO:- moiety and a pair of vicinal OH groups which are widely separated, and as the separation decreases the interaction survives even after calcination at 500°C [SO]. The widely separated groups result in adsorbed M O species, the bonding of which does not withstand even temperatures lower than 400°C In this respect, pyrogenic silica is the least desirable type of silica where, firstly, its

surface area is usually lower than that of many porous silica gels or precipitates, and secondly, its surface is dominated by isolated (unpaired) silanols.

The hydrogen bond is more generally represented as A- H. where B is any CT or electron donor site (Lewis base), and A may be one of the following: C,N,P, S, Se, F, Cl, Br, I. One source of activity of the silica surface is the ability of the silanol groups to formhydrogen bonding either of homo-intermolecular or heterointermolecular type [52]. Homo-intermolecular hydrogen bonding takes place between adjacent silanol groups on the surface of silica in which one group acts as proton do$or, and the other as proton acceptor.As mentioned before, a distance less than 3 A is required to produce an interaction of this type. In some cases the distribution of silanols, according to preparation methods, enables the formation of an extending domain of H-bonded groups, or the presenceof a pair of OH groups which are perturbed veryweaklyby H-bonding.Infraredspectroscopycan easily now distinguish between all theseotypes of surface hydroxyls. Isolated silanol groups are those located far 5 A) from neighboring groups, and these groups are not able to form homo-intermolecular H-bonding except after heating to very high temperatures 1000°C) wherethe movementof bulk materials permitstheir presence near other, normally isolated, groups. However, the recognition of this case does not seem probable where the action of heat that caused the isolated groups to approach eachothereventually results in their involvement in condensation reaction. In porous silica, two isolated OH groups on opposite walls can interfere in hydrogen bonding due to geometric considerations. The H-bonded groups are the source of activity of silica surface toward thermal treatments. At temperatures above 2OO"C, surface dehydroxylation starts involving those groups entering in strong H-bonding, and above 400°C the more widely separated, but still able to interact in p groups come into play. The formed siloxane links in place of the condensed groups above 400°C are considered unsymmetrical and are of different reactivity fromthesymmetrical links that are formed below400°C. The latter are further shown to be similar to those in the bulk [54] as concluded from the consecutive increase in the intensity band at 470 cm"' due to bending mode of Si-0-Si in the DRIFT spectrum ofsilicagel heated up to 400°C (Fig. The intensity of the band at470 cm-' is inferred from its relative intensity if compared with the mostintense band at 1200-1000 cm-', and the differences in intensity have not been due to differences in sample mass as judged from the constancy of the overtone bulk band (1860 cm"). The spectral features of the Si 0 -Si bending band (470 cm") at different heating temperatures are represented in (Table 1). Adsorption of organicsubstancesfromnonaqueousmediaproceeds via the hereto type of H-bonding in which the silanol group on silica acts as hydrogen donor in most cases. The interaction is mainly controlled by the situation in media rather than by the activity of the silanols. Factors like functional group, geometry, size of the adsorbate molecule, and the nature of the used solvent are found to determine theresulting picture on thesilica surface. Benzene interacts through the electron cloud delocalized on the whole ring, while butanol reacts through its

too 1600

IiOO

800

DRIFT spectra ofsilicagel(Kieselgel60,Merck)heated (From Ref. 54.)

Spectral Features of the Si-0 at Different Heating Temperatures Temperature ("C) 100 200 300 400

up to 400°C.

Si Bending Band (470 cm

lo3 (a.u.)

Peak intensity 77 84 130 154

(cm") 70.4 73.0 75.3 75.3

group, which is a strongerinteraction than in the caseof benzene. Benzene isfound to be successively replaced by butanol as the butanol concentration in the liquid mixture is consecutivelyincreased upto 100% Someinorganic salts that dissolve to someextent in organicsolventswhena little water is present are found adsorbable on silica through interaction with surface silanols Hydrogen-bondedwater, when onlya little is present, formsathin layer that holds the salt, and when the surface hydroxyls are masked by different alcohols that are known to have high affinity to form H-bonding with surface Q adsorption of the salt occurs. On the other hand, an excess of water increases the solubility of the salt in the organic solventso that adsorption does not occur. The fullctionality of the adsorbate molecule may cause the interaction to proceed so that twosurface silanols areinvolvedthroughhydrogenbonding in which the silanols may act as proton donor and proton acceptor. With aromatic hydrocarbons the porosity of the used silica plays the most important role. The preferential adsorption on the OH groups that are of higher population in pores than in open regionsmakesthe systemof great potentiality when separation is considered. Obviously, the pore sizes have to be in the range that permits the accplnmodation of the adsorbate molecules. Purely microporous silicas (radius A) may beof little importance and/or unsuitable to the purpose of interest. Adsorption from the vapor phase lacks the intervention of factors like p precipitation, dissolution, or nature of solvent, yet the functionality, size, geometry, and pressure of adsorbate molecules together with porosity of silica remain working and controlling factors. When the functionality of the adsorbate enables both mono- and birnode bonding to occur, the interaction mode depends on the coverage. stronger mode of adsorption of 1,~-dimethoxyethaneinvolving both methoxy groupshydrogen-bonded to two silanols (vicinal) constitutes the dominant mode of interaction at low surface coverage In the presence of excess the mode of interaction changes mainly to the weaker one, involving one from an isolated O to onemethoxygroup, and completecoverage of isolated groups is accompanied by only I I of the extending adjacent interacting silanol groups. Activity of silica toward adsorption of water vapor relies completely on the formation of hydrogenbonding both to single and paired OHS. Spectroscopic evidence indicates that in the beginning stages of adsorption, water draws oxygen down [58] making respectively one ortwo H bonds withisolated or adjacent vicinal silanols. The interaction via one silanol group is very weak[59,60] if cornpared with that involving adjacent pairof OH groups. At laterstages, hydrogen-bonded water clusters are formed beforeall the single hydroxyls are taken up. This is explained on the basis of energetic considerations. The calorimetric heat of a single hydroxylwater bond is 6 kcal/mol which is lessthan the heatof liquefaction water, kcal/mol, therefore wateris energetically favored in the liquid compared to the bonding by a single hydroxyl. On the other hand, multiple bonding of water to the surface OHSis accompanied by higher heat of adsorption, in which case the surface occupied to a large degree before the clustering occurs. The clustered molecules arernore translationally androtationally restricted in their motionthanthose weakly bound to the silanols, and therefore have a lower entropy. Thus, at low

coverages, water adsorption on silanols is favored by entropy, whereas at higher coverages, water clustering is favored by energy stabilization in clusters. The differences in the strength of adsorbed water on single and paired indicate that the pyrogenic silica will behave differently than silica gel or precipitate. Zettlemoyer [62] showed that the flame-hydrolyzed Cabosil silica has hydrophilic/hydrophobic site ratio of about 1:4, whereas the dehydrated HiSil (precipitated) hasaratio of1:1.7. Also,theprecipitated silicais shown to be hilic than the Cabosil as measured by their respective surface area )/A(Nz),of 1.09 !nd 0.38 For silica samples with adsorbed water t least 50 to 70 A in thickness, the adsorbed films undergo a phase transition with heats lie below those for the ice-water transition [64]. The lower heats in the case of HiSil compared with Cabosil are interpreted in terms of an alteration in structure and propertiesof adsorbed waterby the silica surface. This is an effect whose magnitude decreases with increasing distance from surface. the The strength of the interaction and the distance to which it extends from the surface increase withan increase in the hydrophilic characterof the surface, that forHiSil is greater than for Cabosil. The iSil surface, which is highly hydroxylated, affords an opportunity for each adsorbing H 2 0 molecule to form multiple hydrogen bonds to surface silanols. Since hydrogen bonds to water molecules arehighly directional, the position and orientation of the molecules in the first monolayer are strongly determined by the silica surface. This layer also influences the structure of the next layer, and it can be visualized as the persistence of a surface-depende~t structure to a considerable distance from the surface itself. The Cabosil surface, on the other hand, is not fully hydroxylated, and contains numerous isolated silanol groups, so water molecules adsorbed on thesesites retain a greater degree of freedom and the structurillg influence of the surface in this case would be expected to be less pronounced than in case of the HiSil sample. The covalent characterof the Si Si bond on the surfaceof silica that has not been subjectedthermaltreatmentmakes these sites hardly reactive toward wateradsorption.asurementsona pristine silica surface, i.e., one which had never been in contact with a high relative humidity or with aqueous, media, have shown that the process is dynamic and depends on temperature, partial pressureof water vapor, and the contact time Activity of silo thermal treatments toward water is thoroughly covered

he reaction of the surface hydroxyls may can substitute the hydrogen atom of the the metal, or metalloid, atom

that the adsorbate molecule he reaction with this hydro-

n different sizes; TiC14 and BC13 may react bifunctionally, whereas Al(Me)3, and essentially HMDS do not; Ca(Me), is a monome se, while Al(Me)3 is a dimer; 61, only react at an appreciable rate in the gas phase with silica at temperatures of the orderof 350°C. Thisimplies that the properties of the adsorbate are the determining factors of the extent of interaction with silica. The ability of a reactant to react bifunctionally is important in determining the initial reactivity of the H-bonded silanols relative to those that are isolated, and its size determinesthenumber ofinaccessible or blocked silanols 1661. However, it can be said that a difference in the reactivity of isolated and bonded silanols may be of significance in determining the strength of interacti with a specific HS agent. It is supposed that Al(Me)3 preferentially reacts with isolated OH groups [67], since the lone pairs of electrons on their oxygen atoms are readily available to interact with molecules having the electron acceptor property. The oxygen atoms in the hydrogen-bonded hydroxyls, in contrast, have their lone-pair electrons participating in the hydrogen bond,and such oxygen atoms will necessarily be less reactive toward electron acceptor molecules. ~otwithstanding, some authors believe that the reaction of silica with A1(Me)3 proceeds mainly via the hydrogen-bonded sites, and even the strained siloxane bridges formed due to heating at 700°C are more reactive than the isolated OH groups [68]. Alcohols have been shown to undergo condensation reaction with surfacesilano1 groups on silica. The reaction may involveboth isolated and H-bonded groups depending on the sizeof the alcohol. Methanol was found to react withboth isolated and interacting silanols [67]. With increase in the size of the alcohol the free hydroxylsaremore susceptible due to steric effects[69] and theextent of reaction shows some dependence on the structure of the alcohol. Thus, primary alcohols are found to react to a somewhat greater extent than secondary alcohols, and the product is thermally more stable. ~nsaturationin the 2-position modifies reactivity, while beyond the 2-position the effect appears negligible. The surface of silica adsorbing alcohols exhibits good hydrolytic and thermal stability [69]. The no nun if or^ distribution of surface hydroxyls on silica is the reason Eo activity in the adsorption processes since the reactivity of a single isolated group differs fromthat of apairand essentially fromagroupperturbe extended mutual hydrogen bonding. This fact can be seen clearly in the process of wateradsorptionthatcanproceed either weakly throughone sing1 toan isolated OH group,orstrongly via twobondstoapair of adjacent,oups. Although the functionality of the adsorbate potentially affects the adsorption process, the distribution of groups appears to play an important role. monofunctional adsorbate requires the presence of a single OH group; however, the extentof adsorption will be higher when theOH groups are distributed so that no clustering takes place, otherwise steric factors, depending on the size of the adsorbate, will intervene to limit the extent of adsorption. One canexpect that the strongest mode of adsorption takes placein the caseof a monofunctional adsorbate and a partially hydroxylated silica. On the other hand, highly a dehydroxylated surface mayrepresent the other extremeof reactivity in the adsorption process. The extent of hydroxylation of silica,as discussed before, is the consequenceof the origin and history of the specific type of silica.

The usual picture on the surface of silica is the presence of regions composed of siloxane links interrupted by sites exposing silanol groups. The numberand type of the latter on a“ ~ i r g i n sample ’~ are related to the class and history of the silica. The term virgin refers to the after-production sample, that is, the sample synthesized and subjected to no further treatment. process that makes silica cast off the nature of itssurface is amodification process, resulting in amodified surface. Although the different modes of preparation lead to different extents of surface hydroxylation, the preparation method does not seem to be included among the different ways that are considered as a sourceof modification. Whatever the mode of preparation, it results in a certain degree of hydroxylation which is inherent in the details of the processes takingpart in theoverall preparation method. Thermal treatments that cause surface dehydroxylation can be considered as a source of modificatioll, since it results in a surface accommodating fewer hydroxyls than it has to according to the method of preparation. It has to be noted, however, that this kind of modification is temporary in the major part as it has longbeen known that rehydroxylation can be effected, even though dehydroxylation is drastic, with liquid water either at room temperature or at slightly higher thermal level. In contrast to the normal natureof silica surface, the modified surface exposes, instead of the normal silanol groups, Si-0-M groups where M is any species except H and may be simple or complex. The simplest picture of a modified silica surface is that of the surface covered repeatedly and regularly with such links in which M is a single species that reacts in the same way with surface OHS. Of this kind of modified silica is that in which the lnodification takes place via interaction with an alcohol that results in a surface covered, albeit incompletely, with Siwhere R is alkyl group derived from the alcohol. The initial heterogeneity of the surface may lead, even with a noncomplex modifier, to heterogeneously modified silica. The situation becomes more complex when the modifier of the type that can react via more than one groupin its skeleton. (Figure 6) shows many possibilities in which (3-amino~ropy1)triethoxysilane can modify the surface of silica A modifier mayreact differently with the different functionalities on the surface, leading to different forms of modified silica. Boron trichloride (BC13) may react with isolated or vicinal OHS in addition to bridges leading toa surfacecomposed o i-0-BC12, (=Sil, and SiCl. Thereplacement of Si-0 h Si-C1 can be also achieved through the reaction of silica with SOC12,CC14,C12,SiC14, and SiHC13 The BC13-modified silica may result in chemically adsorbed BNHz groups through the reaction with NH3 attemperatures less than 100°C according to the reaction:

which is complete after 30 min. At temperatures greater than 200”C, secondary amines are increasingly formed through a consecutive reaction mechanism forming a borazine-like structure (B3N3H6) on thesilica surface. At 300°C the NH3 reacts solely with the surface boron species to form secondary amine groups

Different possibilities in which (3-aminopropyl) triethoxysilane (APTS)may modify the surface of silica. (Adapted from Ref. 70.)

Si

0

B(Cl),]

7NH3

Si

0

(NH2)B+ NH

4NH4Cl

A fast chemisorption of NH3 was also shown [73] to take place at room temperature on SiC14- and CC14-modified silica (composed of SiCl), forming SiNH2 groupsand NH4C1. However,aftersublimation of the NH4C1, the GSiNH2 formed could revert to the =Sic1 group via reaction with HCl at 600°C [35]. A surface dominated by =SiNH2 species can be obtained by a total substitution of the surfacesilanols on silica pretreated at 800°C through the reaction 300°C at in an NE13 flow [SO]. An esterified silica surface is obtainable through a modification step involving the reaction with an alcohol. The pyrolysis of the methoxy groups of methylated Aerosil, obtained from the reaction with methanol [81], produced Si-0 SiH2 on the surface, and prolonged outgassing caused the eliminationof hydrogen from SiH2 with the formation of Si radicals whose presence is confirmed by ESR [81]. This silica permits thedissociative chemisorption of hydrogen to occur readily at temperatures as low as 25"C, which is well known to be improbable in case of normal silica. Esterification of silica occurs more readily with primary and secondary alcohols containing two or more carbon atoms [82]. ~odificationby tertiary alcohols does not proceed readily and furthermore, the modifier decomposes by dehydration at elevated temperatures. The esterified silica surface is preferentially wetted by at er-immiscible organic liquids, andasthe degree of esterification increases it becomes less hydrophilic and more organophilic. It is interesting to note that the adsorption of nonpolar adsorbate is also affected by the silanol content of silica surface; the heat of adsorption of nitrogen at about -196°C decreases as a function of the surface silanol content Heat of adsorption of nitrogen is more sensitive to replacement of hydroxyl groups with

alkoxy groups than it is to the removal of hydroxyl groups through dehydroxylation [84]. The importance of nitrogen as adsorbate stems from being the recommended adsorptive, except with solids of very low surface area, for evaluation of both the surface area and pore size distribution from a single adsorption isotherm [85]. The evaluation of the specific surface area makes use of the famous BET equation [86] that has the linear form: l

P

V(P0

C-1P

P)

where is the volume equivalent to a monolayer coverage, and Po is the vapor pressure of the adsorbate. The constant C has been given the form:

where E lis the average heat of adsorption of the first layer, and EL is the heat of liquefaction of the adsorbate. Itis the value of this C constant, calculated from the slope and the intercept of the linear relation P) versus that has been taken as a measure of the heat adsorption, and its value determines theapplicability of Eq. (4) to surface area determination. The postulates used in deriving the applied equation emphasize neglecting the adsorbate-adsorbate interaction if compared withtheadsorbate-adsorbentinteraction,inthebeginning of adsorption. This implies that a complete monolayer takes place u n i f o ~ ~ lbefore y the multi~ayer formation, suggested by the theory in [86], begins. These conditions are accompanied by a high valueof the C constant, greater than50 an approximation, and the decrease in the value of C is taken as an indication of the intervention of the adsor~ate-adsorbate interaction duringthe monolayer formation. In this case the evaluated area suffers from uncertainty that increases with the drastic decrease in the value of C. With modified silica surface, the replacementof the silanols, which are the high-energy nitrogen adsorption sites [83] leads to decrease in the heat of adsorption of nitrogen, for which the obtained value of area has to be taken with caution. Itis helpful in this respect that the casesof very lowvalues of C correspond to type I11 or adsorption isothermsof nitrogen 1851 which are convex to the axis over a large range enclosing the monolayer formation region of PIP0 (0.05he data in these cases are not amenable to the BET method for evaluation of surface area.On the other hand, the decrease, upon modification, of high-energy adsorption sites was found advantageous, being accompanied by more uniform surface superior for chromatography [87]. major problem for good performance in gas-solid cllromatography is the surface heterogeneity of the solid adsorbent. When ordinary silica gels are used as column packing material, the resulting solute usually asyl~metricalwith the rear edge of the peak having tailing ause of the variability of adsorption energy for thedifferent adsorption sites, the solute molecules remain on the active or high-energy sites for a longer time than the average,giving rise to peak tailing. This tailing decreases the effectiveness of the colunnn and makesquantitativemeasurementsmore difficult. A better performance is obtainedwithmoreuniform silica obtained by substitution of more inert groups for the active surface silanols. Effective modifiers in this respect

ilic

are alcohols, chlorosilanes, and hexamethyldisilazine [88-901 that reducethe adsorption of molecules such as H20, HC1, and C6H6. Furthermore, because of its high rigidity and slight tendency to swell, silica gel functionalized with organofunctional groups was shown to be highly selective in the exchange process and therefore a good candidate to be used in high-performance liquid-exchange chromatography [91]. Silica modified gel with pyridinium ion, Si(CH& N+C5H5C1-, was shown to selectively adsorb Cu(I1) as CuC142- fromethanol solution, and the dimeric Cu2C162- fromacetonesolution [92]. For Co(TI), the adsorbed species was the tetrahedral C O C ~ complex ~ ~ - from both solvents; however, the adsorption was much higher from acetone than from ethanol solution. The ability of organofunctional groups to absorb metal ions from nonaqueous solutions has been revealed by using a modified silica with different organofunctionality for preconcentration and separation of metal ions dissolved in organic solvents [93-95]. In the process of supporting a catalyst, the first step involves a bond formation between the catalyst precursor and the surface of the support. This leads to a modified surface which, in the case of silica, exposes ESi links, where M may be MO,Cr, or other elements which differ in their affinities toward bonding tosilica. The calcination step, which in most cases comprises heatingup to 500°C results in the breaking of a great number of such links, leaving only the strongly held species. This thermal treatment, in addition, leads to a considerable amount of dehydroxylation. Removal of OH groups upon condensation occurs by random interactions of adjacent hydroxyls from therichest areas, and up to800"C, the density of isolated OH groups increases by virtue of condensation of silanol groups pairs from a group composed of an odd number. A. distinguished state of dehydroxylation exists when a pair of OH groups are presentin an isolated manner in the sense that they form bonds with each other only, being not a part of an extending donlain of cluste groups. This pair of OH groups results in two IR bands at 3720 cm" due to the protonacceptor group, and at 3520-3540 cm" due to proton donor group196,971. However, this pair of OH groups is not present on the silica heated to 500°C [32,66]. The essence a normal surface silica is the probable presence such a pair of silanol groups as long as the heating t e ~ p e r a t u r e does not exceed and their absence in case the s ~ m p ~heated e at 500°C or h i g ~ e r The . detection of the above-mentioned IR bands on the surfaceof silica gel heated 500°C has been reported in the case of ~o-nlodifiedsilica [50 Fi 7 The results are interpreted in terms of the weak bonding in ESi-0which are present on the surface as (=Si -0)2 -MO, and which up0 leave 2(ESi 0)in their positions. The bulk waterthat may evolvedue to heating at 500°C may easily hydrolyze thesesites, resulting in 2( ESi 0 are necessarily quite near to each otherso as to be H-bonded groups. The behavior of terminal group is judged from the defined character of the band at 3720 cm" (12% MO)and attributed to the environmental nearby these silanols. The surface is dominated either by an isolated OH group, or a siloxane group ehydroxylation of groups that primarily escaped the interaction The consecutive increasein the intensity of the considered bands was shown to be accompanied by a decreasein the intensity of the

4000

3200

Wavenumbers

spectra of silica loaded with different Ref. 50.)

of MO.(From

ilica

band at 3750 cm", due tofreely vibrating isolated groups. When thenewly formed OH groups are positioned near to an isolated group, the latter suffers from perturbation via hydrogenbonding, which leads to adecrease in its IR band. Therefore a modified surface results, which exposesdifferent surface functionalities than should be expected on a surface heatedat 500°C. Thesilanol pairs detected in this case are considered as groups that have been protected from condensation through the involvement in a weak interaction with the M O species. Although no studies have been conducted to test the properties of this modified surface, it has to be less hydrophobic than normal silica heated at the same temperature. In all cases, modification of silica surface was due to one ortwo demands, either developing anew surface of different activity, or probing thefunctionality that may present. The nonuniform distribution of surface silanols is the working factor in both cases. The difference in activities of isolated, vicinal, or geminal groups has motivated the work concernedby masking or even depleting one kind, keeping, as a dominant type, one or two of the other types. Modification in relation to characterization has benefited from the different modes of interaction exhibited by different types of OHS toward modifiers of different activities. The vast amount of literature onmodified silicas reveals the highlevel of interest in the surfaceof silica. The silica surface plays acrucial role in determining its activity in the process of adsorption. The appearance of this surfaceshowsastrongdependenceonthe method through which the underlying matrix has been formed. Numerous factors significantly affect the formation of the three-dimensional network in the bulk, which leads to a wide scope of variations that canbe obtained through the control of a specified group of parameters. It follows that one can obtain many kinds of silica of variable properties fulfilling a specific number of requirements. This gives silica a particularflexibility, making itof usein several fields. far as the surface is concerned,a very broad area of objectives coversthevastrange ofspecificity displayed by the surface of silica in different media. A surface exhibiting variable pictures whenused undercontrolledconditionsfurnishesmanyadvantageous properties, making it useful for many purposes. Silica with its fascinating nature is one of those few materials lending themselves to fields of different disciplines. The adsorption process, by its definition, finds the surface of silica as an excellent candidatethatbothpermitsitsoccurrenceand highlights thebasic principles involved in it. Silica and adsorption are strongly related to each other, and the interesting properties of silica surface appear to be the key to this relation.

1. A. R. Ruffa. Phys. Rev. Lett. 25:650 (1970); Non-Crystalline Solids 13:37 (1973/ 1974). 2. C. L. Mckinnis and J. W. Sutton. J. Am. Ceram. Soc. 42:194 (1959);24250 (1959). C. G. Dodd and G. L. Glen. J. Appl. Phys. 39:5377 (1968). 4. K. Fajans and N. J. Kreidl. J. Am. Ceram. Soc. 105 (1948). 5. Brunauer, D. L. Kantro, and C. H. Weise. Canad. J. Chem. 34:1483 (1956). W. D. Harkins and G. Jura. J. Am.Chem. Soc. 66:1362 (1944).

7.

K. Her, The Chemistry

Silica, Cornel1 University Press, Ithaca, N. Y., 1955

8. J. H. De Boer and J. M. Vleeskins. Proc. K. Ned. Akad. Wet. Ser. B: Palaento]. Geol. Phys. Chem. B61:85 (1958). 9. J. B. Peri and L. Hensley. J. Phys. Chern. 72:2926 (1968). 10. D. Sindorf and G. E. Maciel, J, Am. Chem. Soc. 105:1487 (1983). G. C. Pirnentel and A. L. McClellan, The ~ y ~ r o Bond, ~ e n F. Freeman Co., San Francisco, CA, 1960. 12. Barby, in C~~racterizfftion Surjaces (C. D.Parfitt and K. S. W. Sing, Academic Press, New York, 1976, p. 402. 13. Weiser, Chemistry, 2nd ed., Wiley, New York, 1953, p. 307. 14. R. K. Iler, The C h e ~ i s t r y 2nd ed., Wiley,New York, Chichester, isbane, Toronto, 1979. 15. System DANA, 7th ed., Vol. 3, Silica ~ i ~ e r a l s , Frondel, Wiley, New York, 1962, p. 154. 16. G. D. Ulrich. Combust. Sci. Technol. 4:47 (1971). 17. annemacher. J. Colloid Interface Sci. 125:61 (1988). 18. A. P. Legrand, Y. Chevallier, and J. C. Morawski. Colloids 19, H. A. Classens, L. J. M. van de Yen, J. W. de Haan, and C. A. Cramers, J. High esolut. Chromatogr. Chromatogr. Commun. 6:433 (1983). 20. eiderer er, K. Albert, E. Bayer, L. J. M. van de Ven, J. W. de Maan, and C. A. ramers. Phys. Chern. 944189 (1990). 21. J. Fripiat. ACS Syrnp. Ser. 194:165 (1982). 22. P. M. Carrott and K. S. Sing. Adsorption Sci. Technol. 1:31 (1984). 23. Flernming. Physik Chem. 41:427 (1902). 24. A. P. Brady, A. G. Brown, and H. Huff. J. Colloid Interface Sci. 8:264 (1953). K. Iler. Phys. Chem. 57:604 (1953). 25. 26. K. R. Anderson, L. S. Dent-Glasser, and D. N. Smith, in S o l u ~ Silicates l~ (J. lcone, ed.), ACS Syrnposium Series 194, 1982, p. 115. Shafer, D. D. Awschalom, and J. Warnock. Appl. Phys. 61:5438 (1987). 27. Coltrain, S. M. Melpolder, and J. M. Salva, in ~ l t r a s t r u c t ~ r e 28. ~ a t e r i ~(D. l ~ sR.Uhlrnann and D. R. Ulrich, eds.), Wiley, NewYork, 1992, p. 69. 29. D.Avnir and D.Farin. Nature 308:261 (1984). 30. A. C. de Pradel and B. Irnelik. C. R. Acad. Sci. 242122 (1956). 31. L.Hench and D. Ulrich, ~ l t r a , s t ~ u c t Processing ~re Ceramics Glasses and C o ~ ~ o s i tWiley, e ~ ~ , New York, 1984, p. Morrow and A. McFarlan. Langrnuir 7: 1695 (1991). 32. 33. R. Cillespie. Chern. Ed. 40:295 (1965). C h e ~ i s t r yof ON Group, Prentice-Hall, Englewood Cliffs,N. J., 34. L. P. Clapp, 1967. 35. Peri. Phys. Chem. 70:2937 (1966). Dugger, H. Stanton, B. N. Irby, B. L. McConnell, W. W. Curnmings, and 36. R. W. Maatrnan. J. Phys. Chem. 68:757 (1964). 37. ~arshall, G.L. Ridgewell, C. M. Rochester,and J. Simpson.Chern. Ind. (London) 19:775 (19741.

38. M. L. Hair and W. Hertl. J. Phys. Chem. 74:91 (1970). 39. T. Hiemstra, W. H. van Riel~sdijk,and G. H.Bolt. J. Colloid Interface Sci. 133:91 (1989). 40. P. W,Schindler, B. Fiirst, R. Dick, and P. V. Wolf. J. Colloid Interface Sci.5.5:469 (1976). 41. T, Hiemstra, J. C. M. De Wit, and W. H. van Riemsdijk. J. Colloid Interface Sci. 133:105 (1989). 42. G. M. S. El Shafei. J. Colloid Interface Sci. 182:249 (1996). 43 J. H. Stanton and R. W.Maatman. J. Colloid Sci. 18:132 (1963). 44. A. M. Petrov. Tr. Khim. i Khirn. Technol. 1:229 (1958). 45 D. N. Strazhesko and G. F. Yankovskaya. Ukr. Khim. Zh. 2.5:471 (1959). 46. S. S. Bhatnagar, K. N. Mathur, and P. L. Kapur. Indian J. Phys. 3:53 (1928). 47. M. T. Rogers and R. van der Vennen. J. Am. Chem. Soc. 7.5:1751 (1953). 48. C. M. Frensh and J. P. Howard. Trans. Faraday Soc. .52:712 (1956). 49 R. J. Faber and M. T. Rogers. J. Am. Chem. Soc. 81:1849 (1959). 50. C. M. S. El Shafei and M. M. Mohamed. Colloid Surf. 94267 (1995). 51. M. M. Mohamed and G. M.S. El Shafei. Spectrochemica Acta A .51:1525 (1995). ~ n gDekker, , New York, 52. M. D. Joeslen and L. Schaad. Hydrogen ~ o ~ ~Marcel 1974. 53. B. A. Morrow and I. A. Cody. J. Phys.Chern.79:761(1975);80:1995(1976); 80:2761(1976). 54. G. M. S. El Shafei and M. Mohamed. J. Colloid Interface Sci. 17.5:518(1995). 55. D. C. Jones and L. Outridge. J. Chem. Soc. 1574 (1930). 56. R. W. Maatrnan, A. Geertsema, H. Erhage, G. Baas, and M. Du Mez. J. Phys. Chem. 72:97 (1968). 57. J. A.Anderson and C. H. Rochester. J. Chem.Soc. Faraday Trans. I 85:3505 (1 989). 58. K. Klier and A. C. Zettlemoyer. J. Colloid Interface Sei. 58216 (1979). 59. H. Knozinger, in The ~ ~ ~ ~ Bond o g (P. e Schuster, n G. Zundel, and C. Sandrofy, eds.), North Holland, Amsterdam, 1976. 60. J. Texter, K. Klier, and A. C. Zettlemoyer.Prog.Surf.MembraneSci.12:327 (1978). 61. D. R. Bassett. Thesis, Lehigh University (1967). 62. A. C. Zettlemoyer. J. Colloid Interface Sci. 28:343 (1968). 63. P. B. Barraclough and P. G. Hall. J. Chem. Soc. Faraday Trans. I 74:1360 (1978). 64. M. N. Plooster and S. N. Gitlin. J. Phys. Chem. 7.5:3322 (1971). 65. V. R. Deitz and N. H. Turner. J. Phys. Chem. 7.59718 (1971). 66. B. A. Morrow and A. J. McFarlan. J. Non-Crystalline Solids 220:61 (1990). 67. D. J. C. Yates, G. W. Dembinski, W. R. Krell, and J. J. Elliott. J. Phys. Chern. 73:911 (1969). 68. J. Kunawicz, P. Jones, and J. A. Kockey. Trans. Faraday Soc. 67:848 (1971). 69. R. J. Azrak and C. L. Angell. J. Phys. Chem. 773048 (1973). 70. G. S. Caravajal, D. E. Leyden, G. R. Quinting, and C. E. Maciel. Anal. Chem. 60:1776 (1988). 71. K. Possemiers, K. C. Vrancken, P. van der Voort, and E. F. Vansant. J. Chem. Soc. Faraday Trans. I 91:2173 (1995).

72. 73. 74. 75. 76. 77. 78. 79 80. 81. 82.

83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96. 97

G. A. Blomfield and L. H. Little. Canad. J. Chem. 51:1771 (1973). P. Fink and I. Plotzki. Wiss. Tech. Univ. Friedrich Schiller 29:815 (1980). M. P. McDaniel. J. Phys. Chem. 85:532 (1981). M. L. Hair and W. Hertl. J. Phys. Chem. 73:2373 (1969). H. P.Boehm and M. Schneider. Z. Anorg. Allg. Chem. 301:326 (1959). P. vanderVoort, I. G. D’Hamers, and E. F. Vansant. J. Chem.Soc. Faraday Trans. 86:3751 (1990). P, van der Voort, 1. G. D’Hamers, K. C. Vrancken, and E. F. Vansant. J. Chem. Soc. Faraday Trans. I 87:3 174(1 99 Possemiers, P. van der Voort, and E. F. Vansant. J. Chem. Soc.Faraday Trans. I 92:679 (1996). Tech.IJniv.FriedrichSchiller P.Fink, 1. Plotzki, and G. Rudakoff.Wiss. 39:2 17 (1 990). C. Morterra and M. J. D. Low. Phys. Chem. 73:327 (1969). C. C. Ballard, E. C. Broge. R. K. Her, D. S. St. John, and J. R. McWhorter. J. Phys. Chem. 65:20 (1961). J. W. Whalen. Phys. Chem. 72: 1557 (1967). W. K. Lowen and E. C. Broge, J. Phys. Chem. 65:16 (1961). Reportingphysisorption data for gaslsolidsystems.PureAppl.Chem.57:603 (1985). Brunauer, P. H. Emmett, and E. Teller. J. Am. Chem. 60:309 (1938). J. Sore11 and R. Rowen Jr. Anal. Chem. 42: 1712 (1970). A.V, Kiselev, in (M. V. Swaay, ed.), Butterworth, London, 1962, p. 34. 2440 Bohemen, H. Langer, R. H. Perrett, and J. H. Purnell. Chem. (1 960). J. H. Purnell, Wiley, New York, 1962, pp. 108-109. M.Grimpel and K. Unger. Chromatographia 17:200 (1983). Y. Gushikem and M. S. Iamamoto. J. Colloid Interface Sci. 134:275 (1990). R.F. Sturgeon, S. S. Berman, S. N.Willie, and J. A. H. Desaulniers. Anal. Chem. 53:2337 (198 P. Sutthivaiyakit and A. Kettrup. Anal. Chim. Acta 169:331 (1985). J. C. Moreira and Y. Gushikexn. Anal. Chim. Acta 176:263 (1985). A. Morrow, I. A. Cody, and L. S. M. Lee. J. Phys. Chem. 80:2761 (1976). P. Hoffmann and E. Knoezinger. Surf. Sci. 188:181 (1987).

V Department of Chemistry, M. Moscow State University, Moscow, Russia

Lomonosov

I. Silica Chemistry Surface Adsorption 63 and chemistry SurfaceA. silica of adsorbents Determination of surface silanol groupconcentrationon silica72 C. Molecular interactions at adsorption of various compounds on silica surface D. Peculiarity of adsorption water on silica E. Evaluation of energy of hydrogen bonds of silanol groups on silica F. Chemical ~odification silica of surface Silica adsorbents deposited with carbon layers H. Silica grafted with fullerene molecules I.modification Adsorption silica of J. Application of hydroxylated and modified silica

63 74 79 86 87 93 95 97 100

Conclusion

113

References

114

Many papers and monographs (e.g., Refs 1-5) have been dedicated to the investigation of adsorption properties and surface chemistry of silica. Silica is one of the most important adsorbents used in chemical technology and chromatography, for immobilization of enzymes, a catalyst support, well filler for polymer l ~ a t ~ r i aand l s raw materials foroptical products and so forth. Through the study of surface chemistryof silica, intermolecular interaction adsorption on silica well the mechanism of adsorption on silica has become of great importance. Many

laboratories and groups around the worldinvestigate properties of silica and a lot of information on properties of silica has been collected. Some of the mainresults on surface chemistryand adsorption propertiesof silica obtained in the laboratory of adsorption and gas chromatography founded by Professor A. V. Kiselev (Chemistry Department of M. Lomonosov, Moscow State University) are presented. For the investigation of silica surface chemistrydifferent physicochemical methods are used, such as adsorption, adsorption calorimetry, chromatography, and optical and nuclear magnetic resonance (NMR) spectroscopy. Infrared (IR) spectroscopy has provedto be one of the most informative methodsof investigation of silica surface chemistry, which makes it possible to determine both qualitative and quantitative compositions of functional groups on silica surfaces and inside silica particles, as well observe the interactions of functional groups with adsorbed molecules and their participation in chemical reactions. It was found that surface samplesof different forms of silica such as Aerosilhighly dispersed silica (pyrogenic silica), Aerosilogels, prepared from aerosils by hydrothermal treatment, silica gels, and porous glass in a limiting hydroxylated state, havesimilar properties. So the results of investigation of surface propertiesof different silicas are comparable. In silica there are single or "free"(3750cm") and hydrogen-bonded (3550 cm") surface silanol (hydroxyl)groups and intraglobular silanol groups (3650 cm") (Figs and 2) which may be detected only after deuterium exchange reaction of vapor D 2 0with surface silanol groups [6,7]. The quantity of free and hydrogenbonded silanol groups may be changed by heating silica samples or by chemical reaction with modifiers. The 3650 cm" band is attributed to intraglobular silanol spectra of silica containing adsorbed water in the region 40003000 cm" in all cases have overlapping bandsof both water molecules and silanol groups hydrogen-bonded with each other and with water molecules.So it is difficult only from the spectral data in this region to distinguish silanol groups from hydroxyl groups of water molecules. In the region 8 0 0 0 ~ 0 0 0cmm1there are overtone bands of stretching vibrations of silanol groups and water (2uOH) as well as combination bands of water and silanol groups (uOH S) which do not overlap with the first two bands. So these bands can be used for detection of the presence or absence of water in silica Spectra of silica in this region can be obtained using apile of pressed thin plates of Aerosil. The transparency of such a pile in the ~ 0 0 0 ~ 0 cm-' 0 0 region is quite enough to register good IR spectra with a double-beam spectrometer. Figure 3 shows IR spectra of liquid water (l) consisting of two bands: overtones of stretching frequency of vibrations of interacting water hydroxyl groups 7000 cm") and a combination band of stretching and deformation frequency of water hydroxyl groups S200 cm-') as well spectra Aerosil 163 m2/g) standing in air with adsorbed water(2). In the spectrum there are broad bandsin the region 7000 cm" consisting of overlapping bands of surface silanol groups (2usiOH), surface silanol groups perturbed by adsorbed water molecules, and the adsorbed water molecules themselves (2uOH), band S300 cm-' (uOH S) of adsorbed water, as well as a combination band 4550 cm" (vsiOH S) of silanol groups.

Infrared spectra of Aerosils with different specific surface area (I 180, I1 136, 111 100, and IV m2/g) evacuated at 200"C, before ( l ) and after (2) deuteriumexchangereaction at ambienttemperature(a)andafterevacuation at 400°C (From Ref. 7.)

Evacuation of samples uptoTorratambienttemperature even after lh result in large decrease of band 5300 cm-' (Fig. 3) and after 9 h this band practically disappears from the spectrum. Aerosil with adsorbed waterwas investigated by IR spectra in the region of the deformation frequency of water 8H20[lo].The adsorbed water band at 1620 cm-' disappeared from the spectrum of nonporous silica at evacuation at room temperature. It means that adsorbed water is removed from the surface of nonporous silica at ambient tern erature as was determined in [9,lO]. The evacuation of Aerosil at 150°C up to 10" Torr did not result in visible changes in the spectrum. This means that at evacuation of silica up to 150-200"C, dehydration of the surface took place, but not dehydroxylation. Thus if the sample includes water in any state, in the spectrum the band S300 c111-l should exist. Figure 4 shows the comparison of two spectra ofsilica: (a) Aerosil with S 163 m2/g and(b) Aerosilogel withS SO m2/g preparedby hydrothermal treatment Aerosil which contained a remarkable concentrationof intra-

J.

Distinguishing of bands of surface hydrogen-bonded silanol groups(dotted lines) in infrared spectra of Aerosils with different specific surface area (I 180, 11 136, I11 100 m2/g)after evacuation at 2OO"C, (1) before and (2) after deuterium exchange reaction at ambient temperature with D 2 0 vapor. (From Ref. 7.)

globular silanol groups [6]. The comparison of spectra of these silica containing adsorbed water shows thatintensity of the 5300 cm" band is larger for Aerosil at about equal sample thickness of tablets: 0.33 and 0.26 g/cm2 respectively. Such a difference is due to a distinction in specific surface area of samples. After evacuation at 150°C for 2 h (curves 2) the combination band of molecular water disappeared. Deuterium exchange reaction at room temperature makes it possible to replace all accessible hydroxyl groups of silica by groups In spectra 3a and 3b (Fig. 4) there are the bands of groups SiQD for both samples but of different intensity. In the Aerosil spectrum the band 5430 cm-' (overtone of band SiQD) is more intense than for Aerosilogel because of the difference in specific surface area of samples. In the spectrum of region ~000-6000 cm", band (2usioa) disappears almost completely; at the same time in the Aerosilogel spectrum thereis a band of intraglobular hydroxyl groups inaccessible fordeuteriumexchangereactions at ambient temperature.

Infrared spectra of liquid water (1) and Aerosil with adsorbed water before (2) and after evacuation (3-7).

Thus the bands 3660-3650 cm" and 7150 cm"' are due to intraglobularsilanol groups but not water moleculesexisting inside silica globules. Dehydration of pure silica occurs at ambient temperature, but the rate of this process depends on the porosity of the adsorbent. For removal of adsorbed water from fine porous silica, a long time or elevated temperature are needed. The energy of interaction of bonded and intraglobular silanol groups with each other can be characterized qualitatively by a shift of frequency of the ~ a x i m u mof their bands. ~uantitatively,theenergy of interaction of silanol groupscan be evaluated by using the changesin integral intensity of bands of these groups, taking into consideration the correlation between integral intensity and hydrogen bond energy 1 Thus to determine the energetic characteristics of different silanol groups of silica it isnecessary to evaluate the integral intensity of bands of surfaces free and hydrogen bonded as well as intraglobular hydroxyl groups of silica.

where aOH is concentration of hydroxyl groups in samples (~mol/m2), and are intensity of initial and transmitted radiation with frequency U , B is a measure for limiting values of aps, the integral intensity of the band, S is specific surface area (m2/g) and plate mass per unit area (g/cm2).

cm"

Infrared spectra of Aerosil (a) and Aerosilogel (b): (1)sampleswith adsorbed water before evacuation, (2) after evacuation at 150°C for h, and (3) after deuterium exchange reactionat ambient temperature and evacuation 150°C.

The true integral intensity of a band of free silanol groups on a silica surface, can be evaluated from spectra of silica plates of different thickness, prepared from Aerosil samples havingdifferent aoH, and S. For determination of aOH the correlation between aoH and temperature of evacuation for different samples of Aerosils was used for which the aoH were determined by different methods. a result the averageaOH values were accepted for Aerosil samples evacuated at 700, 800, 900, and 1000°C respectively, 2.0, 1.5, 1.1, and 0.7 p/mol/m2, or 1.2, 0.9, 0.6, and 0.4 groups/m2. The calculated true integral band intensity of free silanol groups on a silica surface, AOH, is 9.8 pmol" cm The bands of surface silanol groups of silica evacuated at 200 and 450°C can be distinguished by graphical subtraction of overlapping bands of intraglobular silano1 groups from the spectra of Aerosil evacuatedat 200 and 450°C before and after a deuterium exchange reaction of D 2 0 with silanol groups. The band area of hydrogen-bondedsurface silanol groups can be found by graphical subtraction of the free silanol group band from the spectra of Aerosil evacuated at 200°C. Then taking into consideration that, under heating from 200 up to 450"C, the silica sample loses 3.7 pm01 OH/m2 [7],it is possible to evaluate integral intensity of these hydroxyl group bands. For two samples of Aerosil with

163 m2/g the average integral intensity of hydrogen-bonded surface silanol groups is B 43.7 pmol" cm ~121. For determination of spectral and energetic characteristics of intraglobular silano1 group band silicas with a large quantity of these groups are needed. Silica samples under hydrothermal treatment transform into wide-pore silica adsorbents. Such adsorbents have large a quantity of intraglobular hydroxyl groups. Fordetermination of the integral intensity of theintraglobularhydroxylgroupband, Aerosilogels with specific surface area 109, 50, and 13 m2/g, prepared by hydrothermal treatment of Aerosil were used 1121. Areas of the band 3650 cm"' were determined from infrared spectra of these Aerosilogels after deuterium exchange reaction with vapor D 2 0 at ambient temperature and evacuation at 150°C. For calculation of integral intensity it is necessary to know the concentratio11 of intraglobular silanol groups, which can be determined by loss of sample weight after calcination at 1000°C in air. At this temperature, il1traglobular and most of the surface silanol groups are removed. In 131 it was shown that after evacuation of silica at 750-8OO"C, there are no intraglobular silanol groups in the spectra. Small concentrations of residual surface hydroxyl groups at 1000°C can be evaluated from the dependence of surface concentration of hydroxyl groups on calcination temperature 171. Thequantity of intraglobular silanol groups in silicas can be determined by subtraction of the amount of surface hydroxyl groups from the loss of mass sample at 1000°C calcination. The integral intensity of intraglobular silanol groups was calculated from

S

where cOHis the concentration of intraglobular groups 0 (pmol/g), is plate massperunitarea (g/cm2), andthe integral is area o band 3650 c111-l of intraglobular silanol groups(cm-').The limits of integration were from 4000 to 3000 cm-'.Calculated integral intensity of intraglobular silanol groups is pmol" cm 1121. Now it is possible to evaluate the concentration of intraglobular silanol groups in Aerosil with S 163 m2/g using the area of the 3650 cm"' band and integral intensity. In this case, to determinetheconcentration of intraglobular silanol groups of Aerosil by loss of weight at calcination is very difficult because the main change of weight will be determined on the whole by removal of surface silanol groups, and the precision of the experiment does not make it possible to evaluate the change of mass, owing to removal of intraglobular silanol groups. For two samples of Aerosil evacuated at 200°C the average concentration of illtraglobular silanol groups was 0.144 mm01 OH/g, that is, about 10% of surface silanol groups. The dependence of intraglobular silanol group concentration on. specific surface area of silica is shown in Fig. Concentration of these silanol groups of silicas prepared by hydrothermal treatment and evacuated at tem~eraturesnot higher than 150-200°C reaches a constant value with decreasing specific surface area of the sample. It means that the degree of polycondensation of silicic acid is approximatelythesamewhen large particles of hydrothermalsamplesareformed. ncreasing the surface area, i.e., decreasing the diameter of globules, results in a

Concentration of intraglobular silanol groups of silica samples with different specific surface area. (From Ref. 12.)

decrease of shared intraglobular silanol groups of silica and in the limit, the concentration of these silanol groups is zero. For silica samples with specific surface area S 200-300 m2/gand higher, theconcentration of intraglobular silanol groups is negligible and concentration of surface hydroxyl groupsof silica samples can be determined by simple calcination, by weight loss. For determinat~on of the concentrationof surface silanol groups of samples with smaller specific surface area as well as silica adsorbents with chemically modified surfaces, other methods have to be used. In Aerosil evacuatedat 450°C the average concentrationof intraglobu~arsilanol groups 0.049 mm01 OH/g, i.e., more than half these groups at such heating are removed. Heating of Aerosil at 750-800°C results in removing all intraglobular silanol groups 121. Evacuation at heating from 200 to 450°C results in not only removal of some intraglobular silanol groups and a decrease of intensity of the 3660 cm"' band, but also shift of frequency to larger wavelength at 20 cm". The asymmetrical band 3660 cm"' at heating of Aerosil decreases irregularly. This means that there is the distributio~of intraglobular silanol groups on energy interactions similar as for hydrogen-bonded surface silanol groups of silica. If we compare the changes of concentration and specific surface area of different silica (Aerosil) silanol groups with temperature of evacuation it is possible to find the interesting regularity presented in Fig. 6. The most essential changes in the structure of sample occur at temperatures above 450°C when free hydroxyl groups are removed. The removal of hydrogen-bonded silanol groups does not result in visible changes in structure of particles, their density or specific surface area. So rehydroxylation of silica, heated up to 450"C, makes it possible to return to the initial state, but rehydroxylation of silica heated at higher temperatures results in formation of other samples with smaller specific surface area. Intraglobular silanol groups are formed during the silicic acid polycondensation process at hydrotherl~al treatment ofsilica or in its synthesis. At formation of three-dimensionalpolymersfrom silicic acid not all hydroxylgroupscanform water molecules. As a result, amorphous silica with large quantity of intraglob-

150

Concentration changes of various silanol groups of Aerosil with S 163 m2/gat heating of samplefrom 200 to 1000°Cas well asspecificsurfacearea changes.

ular silanol group is formed, and violation of structure with formation offine channels (micropores) occurs. Formation of ultrapores is probably determined by the structure of the initial silica being the framework of a new structure. clarify the influence of initial silica structure on concentrationof intramolecular silanol groups and on the variety of ultraporosity, samples of Aerosilogel with S 47 m2/g, prepared by hydrothermal treatment of Aerosil A175 and silica gel with S 123 rn2/g, prepared by hydrothermal treatment of initial silica gel with s 220 m-/g, were investigated by an infrared method 214.1, Study of the deuterium exchange reaction of silanol groups of these samples by IR spectroscopy shows that both silica gel and Aerosilogel prepared by hydrothermal treatment contain a considerable quantityof intraglobular silanol groups which are relatively more available for deuterium exchange with D 2 0 for silica gel. Exchange reaction occurs noticeably even at ambient temperature and most of the intraglobular silanol groups exchange at 200°C. Deuterium exchange of D20 with intraglobular silanol groups of Aerosilogel at ambient temperature hardly occurs and only at very slow passes at 200°C This means that Aerosilogel has extremely narrow ultrapores, these being gaps between polysilicic acid chains not fully reacted with each other. In infrared spectra of porous glasses the intraglobular or intraskeleton silanol groups are manifested also. At removal of the borate phase by alkali, very fine canals are produced.Onlysmallwater molecules canpenetratethese canals. Depending on the conditions of alkalization and the composition of boron-silica glass, porous glasses with different specific surface areas and containing various intraskeleton silanol groups are formed 151. Concentration of intraglobular silanol groups can be determined by the NMR method 161if beforehand all surface silanol groups are replaced by OD groups using the deuterium exchangereaction. By this method, concentrationof intraglobular silanol groups in samples of Aerosilogel with S 50 m2/g was determined [17].

Theconcentrations of intraglobular silanol groupsdetermined by the IVMR method and by weight at calcination coincide quite well [13,17]. Thus the combinationof the infrared method with deuterium exchange reaction and reaction of dehydroxylation makes it possible to determine the maxima of absorption bands of three main types of silica hydroxyl groups and also determine their concentrations, but onlyin silica samples from which it is possible to prepare rather transparent plates in the region 4000-2200 cm".

Adsorption properties of adsorbents and supports aswell as fillers of polymers are mainly determined by concentration of surface silanol groups. To determine the concentration of silanol groups fromIR spectra by integral intensity of band silanol groups in many casesisdifficult becauseit is too complicated and sometimes impossible to prepare transparent adsorbentplates in the infrared region spectrum. To increase transparency of adsorbent plates, sometimes these plates are placed in carbontetrachloride orin oil. But determination.of concentration of surface hydroxyl groups from infrared spectra obtained in such a way is rather difficult. In Refs 18 and 19 the method of determination of surface silanol group concentration by deuterium exchange with mass-spectrometric controlis presented. It is accepted that the isotopic composition of water in thegasphaseunderthe adsorbent is equal to theisotopic composition of surfwe groups capableof deuterium exchange reaction, i.e., it is considered that the coefficient of separation of isotopic water molecules onthis adsorbent is equal to 1. Probably such an approximation canbe accepted for many cases because the coefficient of separation differs from by less than the measurement error by this method, which is several percent. The values of concentration of surface silanol groups of nonporous and highly porous silica determined by the deuterium exchange method and other chemical methods are similar [7]. Determination of isotope composition of surface silanol groups may be simplifiedbyusing infraredspectranot of thesamples themselves but of suitable auxiliary adsorbents having high transmittance, for exampleAerosil. An ampoule holdingasampleunderinvestigationcontainingaknownportion of D20 is connected toan optical cellwith a plate of an auxiliary adsorbent (Aerosil), thespectrum ofwhichwas detected preliminarily. After isotopic equilibration between samples under investigation, the vapor phase and Aerosil plate infrared spectrum of this plate is detected again, It possible to consider that isotopic compositions of surface silanol groups of the Aerosil plate being in contact with the adsorbent under investigation and in the vapor phase are the same as the isotopic composition of surface silanol groups of adsorbent. This isotopic cornposition can be determined by the initial spectrum of Aerosil and the spectrum after isotopic equilibration with the adsorbent under investigation [20]. Thus the deuteriumexchangemethodwithmass-spectrometric(MS)controlmay be replaced by the same method with IR spectroscopic control [20]. This method can be applied for determination of surface concentration of any groups

-NHz, NH, capable of deuterium exchange reaction with D2Q on any adsorbents. The quantityof hydroxyl groups on adsorbent surface aoH can be determined by

where ([Hl/[D])a~~~rben~m([Hl/[Rl)Ae~~~il @OD initial I>OD equilibrate)/ DODequilibrate, mDz0 and are mass of D20 taking part in the deuterium exchange reaction and mass of adsorbent, ROD log is the optical density of 2760 cm" band free OD groups on the Aerosil surface. If the specific surface area of adsorbent S is known, then the concentration of surface hydroxyl groups can be calculated:

This method was used for determination of concentration of silanol groups on surfaces of two adsorbents used in liquid chromatography, silica gel LiChrosorb SI 60 and silanizedsilicagel LiChrosorb RP 2. The average surface concentrations of silanol groups for LiChrosorb SI 60 and LiChrosorb RP 2 were 8.0 and 4.5 pmol/m* f 5% respectively. Surface concentration of silanol groups of the first sample is thesameasfor all limiting hydroxylated silica[7,21], andfor modified silicagel LiChrosorb RP 2theconcentration of hydroxylgroups is much smaller. Thismeans that not all hydroxylgroupsareremovedfromthe surface, but only free hydroxyl groups; hydrogen-bonded surface silanol groups remain [7,8]. Other auxiliary adsorbentshavinghydroxylgroups available fordeuterium exchange reaction may be polymer film framework structures [22], for example, polyester-melamine film [17]. In contrast withAerosil, this film is a poor adsorbent of water and so there is no need of evacuation at high temperature for removalof adsorbed water. Many polymers, prepared by polycondensation reaction and polyaddition, have remaining functional groups with "active" hydrogen which do not take part in the reaction of polymer formation owing to steric hindrance. These grou (hydroxyl groups) are responsible for wide bands in the region 3700-3000 cm-'of infrared spectra owingto formation of intra- and intramolecular hydrogen bonds that may take part in deuterium exchange reaction with vapor D2Q, which results in corresponding changesof spectra: the band in the region 3700-3000 cm-' disappears and the bandin 2700-2200 cm-' appears. The deuterium exchange rate very high, so the kinetics of deuterium exchangein the film will bed e t e r ~ i n e don the whole by the diffusion process. The surface concentration of hydroxyl groups of LiChrosorb Si 100 was determined with polyester-melamine film as auxiliary adsorbent [17]. After evacuation of LiChrosorb Si 100 at 150°C the average concentrationof surface silanol groups was TZOH f0.5 QH/nm2 which is similar to the values for all hydroxylated silica [2 l]. Thus the deuterium exchange method with spectroscopic control has turned out to be asimplemethod of determination of surfweconcentration of hydroxyl groups of adsorbent any nature.

At adsorption of nonionogeniccompounds on hydroxylated silica surfacetwo types of main interactions are manifested: van der Waals being universal for any compounds and interaction with formation of hydrogen bonds between the surface silanol groups and adsorbed electron donor molecules. Hydrogen bonding in many cases brings the essential contribution to the adsorption energy and so determines the adsorption process. Many studies have been dedicated to investigation of hydrogen bonds of silanol groups on adsorption. One of the brightest manifestations of the formation of hydrogen bonds, is changes in infrared spectra of interacting proton donor groups. Athydrogenbonding between hydroxylgroups of acid and base molecules, a decrease of stretching vibrationfrequencies of the OH group observed,The most sensitive spectral characteristic of hydrogenbondingnevertheless is the changes of integral intensity of the stretching band of acid hydroxyl groups [23]. Since on the silica surfacethereare both free and hydrogen-bonded silanol groups, it is important to investigate the role of each kind of groups in adsorption of various compounds. The investigation of the participation of free and bonded silanol groups in adsorption was made in Refs 24-27. In Ref. 24, the very small difference between the isotherms of diethyl ether adsorption onsilica evacuated at 150 and 400"C, i.e., on silica having free and bonded surface silanol groups and on silica with only free silanol groups, means that the main adsorption species of hydroxylated silica are free silanol groups [24]. Infrared study of adsorption of many classes of organic cornpounds supports this statement [24-271. In Fig. 7 the spectra of Aerosil evacuated at 200 and 400°C before and after adsorption of triethylamine, dibutyl and diethyl ethers, tetrahydrofurane, cyclopentanone,acetone, acetonitrile, ethyl acetate, andnitromethaneat different vaporpressuresarepresented. As a result of adsorption of compoundson Aerosilevacuated at 200"C, in thespectratheband 3750 cm-' offree silanol groups decreases in intensity and shifts to a lower frequency and a broader band of perturbed silanol groups is seen before adsorption as free silanol groups as well as bands of groups -CH, appear.The frequency of thebandmaximum of perturbedhydroxylgroupschangeswithincreasingadsorbate pressure. If one accepts that only free hydroxyl groups take part in adsorption 1241 and the band of bonded silanol groups has to be unchanged, then it is possible to make a graphical subtraction of the band of bonded silanol groups from the common broad-bandperturbed silanol groups. By suchasubtractionit is possible to distinguish theonlyband of silanol groupsperturbed by adsorbedmolecules which before adsorption were free. Maxima of these distinguished bands (dotted lines) are at the same wavenumber (Fig. 7). On the silica surface the only free silanol groups exist on evacuation at 400°C and therefore on adsorption of the c o ~ p o ~ l nin d ,the spectrum the only hydroxyl group band perturbed by adsorbed molecules appears. Figure 7 shows that maxima of silanol group bands perturbed by adsorbed molecules on silica evacuated at 200 and 400°C coincide [24]. A similar picture is observed on adsorption of alcohols [25,26].

400°C

0.5

0

0

3800

3200

3800

3200

3200

3800

3200

Infrared spectraofAerosilbeforeand after adsorption ofvarious compound vapors at different surface coverage, evacuated at 200°C (a) and 400°C (b) (drop lines are determined by subtraction of the bonded surface silanol group band from the general band of perturbed silanol groups). Ref. 27.)

Thus it is important to know not only the concentration of hydroxyl (silanol) groups on the silica surface but the relative quantity of free and bonded silanol groups, i.e., qualitative composition of silanol groups. In Ref. 28 the influence of the only residual free silanol groups of amino silica gel, but not residual bonded surface silanol groups, on retention of electron donor compounds in gas chromatography was demonstrated. To investigate interaction of free silanol groups on adsorption of various compounds by the infrared method it is more convenient to use samples evacuated at 400450°C or at higher temperature when on the surface only free silanol groups exist and the band of free silanol groupsbecomessymmetrical. In Aerosilsample was evacuated at 800°C andtheinfraredspectrumhasonlya narrow, almost symmetrical, band of stretching vibration of surface free silanol groups; it was very convenient to follow the changesof this band on adsorption of various compounds. The determinationof integral intensity of surface free silanol groups has been presented. For calculation of integral intensity of these groups perturbed by adsorbed molecules, the following equation was used:

where AoH is integral intensity of theband offree silanol groups, is area under the curve recalculated in optical density of silanol groups band perturbedby adsorbed molecules, is area under the curveof the band of free silanol before adsorption and S& is the area under the curveof the band of residual free silanol groups after adsorption. In Fig. 8 spectra of Aerosil evacuated at 800°C before and after adsorption of various organic compounds as well as the spectrum of Aerosil in liquid CC rom these spectraby Eq. (5) integral intensity AOH...M and Avow, band of stretching vibration of free silanol groups of Aerosil perturbedby adsorbed ecules were calculated. n Fig. 9 the correlations between energetic and spectral characteristics of hydrogen bonds at interactionof surface free silanol groups with organic bases at adsorption from gas phase as well as of phenol with various organicbases in solution alitative and quantitative correlations between therl~odynamic andspectroscopic characteristics of formation of hydrogen bonds in solution and on adsorption of various molecules on the surface of pure silica are similar and may be presented by the same equations. The main deviation from relationships found rved at hydrogen bonding of pyridine with hydroxyl groups. he valuesof ASiOH...Mand AvOH, as well as heats of f o r ~ a t i o lof l the hydrogen bond of surface silanol gorups ofsilica ..Mon adsorption ofseries of compoundsfromargontotriethylamine (see 9) gradually increase. The relative contribution of hydrogen energy bondiith silanol groupsto thecommon energy of interaction with the surface on adsorption for this series of compounds also increases. For argon and saturated hydrocarbons this contribution ison1 a few percent and for triethylamine it is about 50% of the common energy. interaction of saturated hydrocarbons withfree silanol groups theintegral intensity

-L"."""-

N

N

Infrared spectra of Aerosil evacuated at 800°C before and after (2,3) adsorption of various compounds at different surface coverage. (From Ref. 11

band OH increases about 1.2 times but, on interaction with triethylamine,intensity increases 25 times 1l]. Correlation between the energyof hydrogenbondformation(kJ/mol) and changes of integral intensity both for solution and adsorption can be expressed by

This equation maybe used for evaluation of the energyof hydrogen bonding using spectral data. The measurement of integral intensity in the adsorption system is a rather complicated problem. So for evaluation of integral intensity changes it is possible to use (AoH)~] -0.59

0.147(AuOH)4

Heats of compound adsorption on dihydroxylated silica surfaces are usually near their heat of condensation. So the common heatof adsorption of a compound on a hydroxylated silica surface at 0.5 can be evaluated by condensation heat and from

Thevalue of can be calculated by the shift of the stretching vibrationband of hydroxyl groups (Au0I-I from Eqs (6) and (7).

As already discussed, on adsorption of electron donor compounds and alcohols on a hydroxylated silica surface, free silanol groups are the most important and hydrogen-bonded surface silanol groups do not take part in adsorption of such compounds [24-273. From spectral data it follows that, for adsorption of water on hydroxylated silica, not only free silanol but also bonded surface silanol groups are very important 1261. For understanding of the mechanism of water adsorption and investigation of the participation of various surface functional groups of silica the thermodynamic characteristics of wateradsorption were determined.Measurement of water adsorption 1291 anda spectral investigation [30] were carried outonthe same sample of silica-Aerosil with S 163 Isotherms of adsorption at different temperatures were determined from the vacuum adsorption line with a calibrated microburette. As was discussedearlier, physical adsorbed water canbe removed by evacuation even at ambient temperature. For investigation of the adsorption of water on a fully hydroxylated silica surface the sample was evacuated at 40 and 150°C. For measurement of water adsorption on a silica surface having commonly only free silanol groups, the sample was evacuated at 450°C. Fully dehydroxylated silica surfaces cannot be achieved even by evacuation at 1000"C, so for the preparation of a silica sample with an almost dehydroxylated surface, Aerosil was evacuated at 900°C. At this temperature there is no remarkable sintering and at the same time the concentration of residual silanol groups is small. Aerosil evacuated at 900°C can be considered as silica with siloxane groups (Si-Si) only on surface.

80

Davydov

Adsorption of water on silica samples evacuated at temperatures higher than 200°C is nonreversible because of chemisorption of the water, resulting in restoration of hydroxyl (silanol) groups on a partly dehydroxylated (depending on temperature of evacuation) silica surface. The degree of rehydroxylation increases with increasing water vapor pressure over silica, time of contact of water vapor with the surface and temperature of the sample. So for measurement of physical adsorption but not chemosorption on a partly dehydroxylated surface of Aerosil it was necessary to reduce the contact time of water vapors and work in region of relatively small relative water pressure PIPo and possibly at low temperature. To fulfill these demands after measurement of one or two adsorption points the sample again was evacuated at a corresponding temperature and only after that the next adsorption points were measured. Adsorption values and vapor pressure were fixed for 1 h. Preliminarily it was found that on the same sample with a fully hydroxylated surface the adsorption equilibrium was attained in that time. So this time was adopted, but the quantity of chemisorbed water was still too small and in all following points was not increased because of the evacuation of the sample at the corresponding temperature. Owing to this, it was possible to consider the isotherms of water adsorption on Aerosil evacuated at 450 and 900°C as close to the equilibrium adsorption isotherms. Differential heats of water adsorption on silica evacuated at 450 and 900°C cannot be determined by a calorimeter, owing to the long experiment time and therefore chemosoption process to some extent. Even a small quantity of chemisorbed water will make a relatively large contribution to the common heat effect, especially at small coverage. So measurements of the isotherm of water adsorption at different temperatures on silica samples with surfaces dehydroxylated to different extents taking in to consideration all above-mentioned precautionary measures, are probably the only method of evaluation of physical adsorption energy and of the contributions of different functional groups to the common energy of water interaction with the silica surface. In Fig. 10 the isotherms of water adsorption on Aerosil evacuated at different temperatures are presented. The isotherms of water adsorption on Aerosil evacuated at 40, 150, and 450°C are reversible (desorption points on samples evacuated at 900°C are not determined). From these isotherms the isosters were determined and presented in Fig. 11. The isosters are near to lines and can be used for calculation of heats of adsorption. The isotherms of water adsorption at 20°C on Aerosil evacuated at 40 and 150°C practically coincide. This means that surface chemistry of samples evacuated at 40 and 150°C are similar and that it is possible to remove physically adsorbed water, i.e., to dehydrate the surface without removing the silanol groups at low rather than elevated temperatures. This is correct only for nonporous and mesoporous silica. In Fig. 12, isosteric heats of water adsorption are presented. Isosteric heats of water adsorption on Aerosils evacuated at 40 and 150°C are similar, which is expected because of the coincidence of the corresponding water adsorption isotherms. The isosteric heats of water adsorption on silica evacuated at 450"C, i.e., carrying only free hydroxyl groups, are much smaller than the heats of adsorption on Aerosil evacuated at 15O"C, i.e., carrying both free and hydrogen-bonded sur-

Adsorption on Silica Surfaces

81

1

I-

2 0

900 O C L?

I

2

p,Wa

FIG. 10 Isotherms of physical adsorption of water vapors at different temperatures on Aerosil preliminarily dehydrated at 40 and 150°C and partly dehydroxylated by evacuation at 450 and 900°C (open symbols are adsorption, closed are desorption). (From Ref. 29.)

face silanol groups. For Aerosil evacuated at 450°C the heats of water adsorption are near the heat of water condensation but exceed it slightly. On fully dehydroxylated samples carrying on the surface mainly siloxane bonds the heats of physical water adsorption at small coverage are smaller than the heat of water condensation and gradually come near to it with increasing surface coverage. Such changes of heats of adsorption with quantity of adsorbed water are usual for adsorption of water on adsorbents with hydrophobic surfaces. So surfaces of silica carrying mainly siloxane bonds may be considered as “hydrophobic.” Such a surface is unstable and in contact with water vapor is gradually hydrolyzed. But it is

0.5

-0.5

0.

0.75

-0.5

0.5 0.25

f0

Isosters of water vapor adsorption on silica, evacuated at different temperatures.

50

Isosteric heats of physical adsorption of water vapors on Aerosil surfaces containillg both free and hydrogen-bondedsilanolgroups(evacuation at 40 and l.SO"C), containing mainly free silanol groups (evacuation at and on highly dehydroxylatedAerosilsurfacescontainingmainlysiloxane bonds (evacuation at 900°C). (From Ref. 29.)

important that siloxanebonds existing on hydroxylated silica surfaces maybe considered as hydrophobic parts of hydroxylated silica adsorbents. Additionalinformationaboutthemechanism of wateradsorptionon silica surfaces may be obtained by simultaneous measurement of water adsorption and registration of the spectrum of silica with adsorbed water [30]. In this case it is possiblequiteaccurately to attribute values of adsorptionto spectral changes. Usually adsorption measurements are carried out on samples of relatively large massandinfrared investigations are carried out on smallthin tablets, i.e., for adsorption and spectral investigations different samples are used. Moreover, cornparison of adsorption and spectral data is complicated by the fact that the tablet of adsorbent is heated by IR spectrometer radiation and the temperatureof the tablet is higher than ambient temperature; but usually this temperature is not measured, and owing to this it is not possible to attributespectra data to adefinite quantity of adsorbed water. Investigation of the mechanismof water adsorptionin the overtone regionof the spectrum is suitable because of the possibility to observethe 5300cm"' band attributed to adsorbed water; at the same time, owing to the small intensity of overtone bands it is necessary to use a thick tablet of adsorbent mass for which it is sufficient to precisely determine the quantity of adsorbed water by a calibrated microburette. Thus to each quantity of adsorbed water can be attributed an infrared spectrum. If, to determine spectra in a monochromatic beam, in such a case strong heating of the tablet is avoided, its temperature willbe near to ambient temperature [30]. Figure 13 shows that the increase of adsorbed water results in an increase of intensity of the combination band of stretching and deformation vibrations of water molecules at 5300 cmm1 and a decrease of intensity of the overtone band of stretching free surface silanol groups at 7326 cm-'. This means that free silanol groups take part in adsorption of water. In Fig. 14 the changes in optical density of the 5300 cm-' and 7326 cm" bands with increasing adsorbed water concentration, on silica surfaces are presented. For samplesevacuated at 25 and 150"C,with theincrease of concentration of adsorbed water at small coverage the optical density of the 5300 cm" band of water increases and the optical density of the 7326 cm" band of free silanol groups decreases. Approximately the same picture is observed for Aerosil evacuated at 450°C. Afterevacuation of samples at 1050"C, on the silica surface there is a slow process of chemisorption of water, but the water adsorption is very small, so it was difficult to register the changes of optical density of the 7326 cm"' band, These results also show that both free andhydrogen-bondedsurface silanol groupstakepartinphysicaladsorption of water,Similar results presented in Fig. 14 were obtained using reflectance spectra for Aerosil samples evacuated at temperatures higher than 500°C [31]. Thus the existence of silica surface bonded silanol groups essentially increases water adsorption capacity of the surface. Free silanol groups on the silica surface take part in water adsorption both in the case where on the surface there are also bonded silanol groups and in the case where free silanol groups exist alone. The decrease of free silanol group concentration results in a decrease of adsorption

Infrared spectra of Aerosil evacuated at different temperatures before and after (2-13) adsorption of increasing amounts of water. (From Ref. 30.)

0.2

W

Correlation of adsorption and spectra data for water adsorption on Aerosil (symbols with points inside are desorption). (From Ref. 30.)

capacity relative to water and in an increase of nonreversible chemisorption of water. Monolayer capacity of water on hydroxylated silica (Aerosil) under investigation formally evaluated by the BET equation is 6.04 pmol/m2. From Fig. 14 one can see that at such coverage not all free silanol groups disappear. Probably this can be explained by the formation of clusters, as was proposed in First, portions of water adsorb on free silanol groups. Then, as vapor pressureincreases, adsorption of water occurs on free silanol groups and on already adsorbed water molecules, with formation of clusters. The heat of water adsorption on silica surfaces with only free silanol groups is not much different from heat of water condensation (Fig. 12) and probabilities of adsorption on adsorbed water molecules and on free silanol groups are Comparable. Figure 14 shows also that approximately the same numberof free silanol groups interact with adsorbed molecules but at different relative pressures when the same quantity of water adsorbs on silica surfaces carrying both free and bonded silanol groups and on silica surfaces with only free silanol groups. It means that adsorbed water molecules on free silanol groups form some more hydrogen bonds with the oxygen of bonded surface silanol groups of silica. In this case the energy of water adsorption will be larger. At higherrelative pressures adsorptionof water occurson both silanol groups and adsorbed water. Thus the existence of silica surface bonded silanol groups probably increases energy of interaction of water with free silanol groups. But more important is the existence of free silanol groups, although the energy of bonding water molecules with them is not so large. if, to remove free silanol groups from silica surface by

reaction with trimethylchlorosilane, the surface bonded silanol groups remain on the surface, then they are accessible for adsorption of water. Nevertheless heats of water adsorption onmodified silica becomemuch smallerthan on the hydroxylated silica surface [32j. Low heats of water adsorption cannotbe explained by screening of bonded surface silanol groups by grafted trimethylsilyl groups since deuterium exchange of bonded surface silanol groups is easily carried out [7,8] and therefore residual bonded silanol groups are accessible for interaction with water. Neverthelessthesurface of theadsorbent becomes hydrophobic similar to the case of removing free and bonded surface silanol groups by heating of silica at high temperature (Fig. 12) when strained siloxane Si bonds are formed. These siloxane bonds are hydrophobic from the point of view of physical adsorption of water, but these bonds can be broken gradually by contact with water molecules.This also means thatthe energyof interaction of water molecules with bonded silanol groups only is not very high 1321. Thus free silanol groups are the most important species of adsorption of water on silica. The peculiarity of adsorption of water on silica in comparisonwith adsorption of othercompounds is probablyconnectedwiththe possibility to form hydrogen bonds not only with surface silanol groups of silica but also with each other. Experimental data show that alcohols interact with free silanol groups in the same way as electron donor molecules [26,27].

Correlation between energetic and spectral characteristics of theformation of hydrogen bonds in solutions is the same as for adsorption systemsand is probably universal, which makes it possible to evaluate energetic characteristics of adsorption and the contributions of hydrogen bond energy to the common energy of adsorption by using only spectral data. Energy of hydrogen bonds of surface and intraglobular silanol groups of silica can possibly also be evaluated by use of these correlations from spectral data. In Table 1 the energyof interaction between silanol groups on the surface and in the volume of globules of silica obtained using integral in~ensity are presented. From this table one cansee that surface bondedsilanol groups have greaterintegral intensity, strongerhydrogenbonds,agreater shift of vibrationfrequency and smaller distance between silanol groups in comparisol~with intraglobular silanol groups of silica. So at heating, surface silanol groups are removedfirst (temperatL~re from 150-200 to 400450°C). Intraglobular silanol groups also are partly removed from silica at thesetemepraturesbecause of the existing distancedistribution between the silanol groups, but at 400450°C intraglobular silanol groups still exist in rather great quantity inside silica (Fig. Surface bonded silanol groups are removed almost completely under these conditions. From data of Ref. 33 and spectral data it is possible to evaluate the average distance between silanol groups which is presented in Table The distancebetween free silanol groups has to be larger than 0.310 mm.

SpectralCharacteristicsandEnergy Silanol Croups Silica

Surface free Surface bonded Intraglobular

3750 3550 3650

200

of HydrogenBonds of Interacting

9.8 43.7 0.283 35.5 0.293

0.310

11.7 9.2

vSioH frequency of maximum band of silanol groups; AvsioH shift frequency; integral intensity of band; ~ j s i o ~ . . . heat ~ of hydrogen bond formation; l? distance between oxygen OH. .O. Ref. 12.

Surface silanol groups of silica can be involved in chemical reactions with various organic and inorganic compounds which make it possible to graft to the silica surface different functionalgroups and by this method to changeadsorption properties of surfaces of silica adsorbents. In monographs and reviews many reactions are presented, by means of which on hydroxylated silica surfaces it is possible to create relativelyhigh concentrations ofnew functional groups. The most promising chemical modification reactions of silica surfaces are those with chlorosilanes, silanols, disilazanes, and esters of silicic acid. Also other types of chemical reactions are used for preparation of adsorbents for immobilization of enzymes [37]. The reaction of chlorosilanes with surface silanol groups of silica occurs at a highrate at hightemperatures Butmore suitably, this reactioncan be achieved in organic solvents by using amines as donor electrons. The influence of electron donors on the rate of reaction of modification silanols with surface silanol groups of silica wasinvestigated in detail It was shown that even a donor such as water accelerates this reaction. But the presenceof water at modification of silica by di- and trichlorosilanes may result in polycondensation of silanes. Usually, as a donor of electrons the amines, e.g., pyridine or triethylimine [43-45] are used. The mechanism of reaction of chlorosilanes was investigated by infrared spectroscopy (e.g., Refs and S), by means of which it was possible to determine not only the appearanceof grafted functional groups on adsorbent surface but tofind which silanol groups take partin ~odificationreactions. In these reactions the most important are free surface silanol groups. The bonded surface silanol groups can only partly interact with dichlorosilanes. So it is difficult to obtain a concentration grafted modifying groups greaterthan that of free silanol groups. If branched or big volume groups are grafted, then the concentration may onlybe smaller. In the case where a polycondensation process is used in chemical modification by di- or trichlorosilanes, theconcentration of graftedmodifyinggroupscan be more thantheconcentration of free silanol groups of hydroxylated silica. Detailed

investigations of modification reactions carried out later show that the concentration of grafted groups was about 4 ,urnol/m2 [34]. Chemical modification reactions may be controlled by the presenceof carbon on theadsorbent [46]. Infraredspectroscopycan beused notonlyfor qualitative composition of grafted groups but also for quantitative evaluationof carbon content on the adsorbent surface. In Ref. 43, for chemical modification of silica the following chlorosilanes were used: diphenyldichlorosilane, methylphenyldichlorosilalles, methyl-~-phenyl-ethyldichlorosilane, and chlorosilaneswithunsaturated andsaturatedhydrocarbon groups--methylvinyldichlorosilalle, dimethyldichlorosilane, and octadecyltrichlorosilane. Chemical modification of silica was carried out on silica gels: LiChrosorb SI 60, LiChrosorb SI 100, LiChrospher SI 100, LiChrospherSI 500, and LiChrospher SI 1000 (Merk), and Silacorb 600 (Lahema). The degree of coverage by grafted modifying groups on silica surfaces can be determined by the carbon content of the modified adsorbent [46]. For LiChrosorb SI 100 modified by different chlorosilanes the carbon content was determined by combustion in a rnicrocalorimeter. In Table 2 the results on determination of modification degree of silica surface are presented. Table2 shows that degree of modification of adsorbent is rather high and probably is near to complete coverage by corresponding functional groups. In accordance with the area occupiedby grafted functional groups thelargest content of these groups is obtained with modification by dimethyldichlorosilane and the smallest one with modification by diphenyldichlorosilane. The average area of benzene in a dense monolayer on silica is about 49 A2 [47]. The average area per grafted Si(C6H5)2 group inmodified samples is about 8 1 A2, i.e., is near to the van der Waals area of two benzene molecules in a dense monolayer. If one takes into consideration that phenyl groups in the modifying layer are arranged not parallel to the surface but at angles to each other, it is possible to consider that all surfaces of the adsorbent are coyered by phenyl groups. Table 2 shows that the area per Si(CH& group is 39 A2. This means that the

Carbon Content in Silica Gel LiChrosorb SI 100 ( S 285 m2/g) Modified by DifferentDichlorosilanes (C), Average Concentration of Modifying Groups per Mass of Adsorbent and PerUnit of Surface (a), well Area of Grafted Modifying Groups and Number of Carbon Atoms in Modifying Groups

Si(C6HS)2 12 si(cH3)(cH2)2(c6H5) 9 si(cH3)(c6H5) 7 Si(CH3)CH CH2 3 Si(CH3)2 2 Ref.

7.69 8.32 6.47 3.51 2.70

0.527 0.770 0.770 56 0.975 45 1.125

2.04 3.04 2.96 3.66 4.2 1

81

39

surface of thesample modifiedby dimethyldichlorosilane is almostcompletely covered byo methylgroups, since, on average,one free silanol groupoccupies about 33 A2 onthehydroxylated silica surface(evacuated at 200°C)[7]. It is possible to consider that practically all free silanol groups took part in the modification reaction with the surface of such samples [S]. Bonded surface silanol groups remaining on the surface after modification are under the modifyinglayer and probably areaccessible only to small molecules such as water. It is possible that a very small number of free silanol groups may remain on the surface as a defect of coverage and may be detected by a c~romatographic separation process. Samples modified by diphenyldichlorosilanes and methyl-~-phenylethyldichlorosilane have the largest concentration of carbon atoms on the surface (Table 2). The area of the Si(C6H5)2 grafted groupis only about two times larger than the Si(CH3)2 group, but the number of carbon atoms of the first is six times larger than the second. So it is possible to propose that the adsorption potential of silica modifiedby diphenyldichlorosilane willbe higher in comparison with a sample modifiedby dimethyldichlorosilane. The comparison of area per grafted group, the number of carbon atomsin each and their structure makes it possible to predict that the adsorption potential of silica modified by methyl-~-phenylethyldichlorosilane, diphenyldichlorosilanes and will be similar to or higher than that of samples modified by dimethyldichlorosilane and carrying on their surfaces only methyl groups. very suitable method of qualitative and quantitative control of the modification reaction is infrared spectroscopy, making it possibleto observe the graftingof modifying groups and to evaluate their surface concentration (Fig. 15). To decrease scattering by the pressed silica gel plate, the spectra of samples are determined in carbon tetrachloride. In carbontetrachloride the band of free silanol groups at 3750 cm-' of untreated silica shifts by 58 cm to a lower frequency [l l] and in the spectrum there is a broad band at 3700-3000 cm" consisting of overlapping bands of free silanol groups perturbed byCC14 and of surface bonded silanol groups.From Fig. 15 onecan see that after chemicalmodifications by chlorosilanes on the surface ofsilicagel the bonded silanol groups still remain and the bands of corresponding groups appear. To make a quantitative analysis of graftedfunctionalgroupsfrominfrared spectra only is quite difficult, since it is necessary to know the extinctioncoefficient ut reproducibility of modification reactions or degree of coverage after modification without exact quantitativeanalysis of content grafted functional groups on surface is quite simple. In Fig. 16 the infrared spectraof silica with different specific surface areas after chemical modificationby diphenyldiclllorosilane are presented. The degreeof coverage of the surface by phenyl groups after modification canbe evaluated from the band at 3083 cm-'. From the spectra presented in Fig. 16 it is possible to evaluate the productlira is the extinctioncoefficient, and is the concentrationof grafted groups to the adsorbent surface) and to compare it for different samples for evaluation of the reproducibility of the modification reactionand for determination of the degree of surface coverage. Calculations showthat Ka for samples withsimilar

3200 cm”

Infrared spectra of untreated silica gel LiChrosorb SI 100-1 and modified by various dichlorosilanes (2-6) (samples beforehand were evacuated at 200°C to Torr and placed in CC14), (2) diphenyldichlorosilane, (3) methyl~henyldichlorosilane, (4) dimethyldichlorosilane, methylvinyldic~lorosilane,and (6) methyl-@phenylethyldichlorosilane.(From Ref. 43.)

specific surface areaand particle diameter aswell identical conditions of reaction are similar, although Ka decreases a little with increase of specific surface area on untreated silica gel. It is probably connected with steric difficulties at ~ o d i ~ c a t i o n of part of narrow pores. In calculating Ka it isnecessary to take into consideration that the specific surface areaof modifiedsamples decreases because of increasing of adsorbent, owing to the attachment of modifying groups at practically the same area of sample. This decrease will be larger for adsorbents with higher specific surface area at the same degree of modification. If the K value of grafted groups is known, then it is possible to evaluate the concentration of modifying groups on the silica surface by infrared spectra. The concentrationof grafted diphenylsilyl groups canbe calculated from spectra because the value of a for LiChrosorb SI 100 was determined earlier from the content of carbon on the modified surface and the coefficient of extinction K was determined, which is the same for all samples.

cm"

Infrared spectra of silica gels modified bydiphenyldichlorosilanes (samples beforehand were evacuated at 200°C to Torr and placed in CC14): (I) LiChrosorb SI 60(500 m2/g), LiChrosorb SI 300 (286 m2/g), LiChrospher SI 500 (50 m2/g), and (4) LiChrospher SI 1000 (20 m2/g). (From Ref. 44.)

In Table 3 the concentrations of diphenylsilyl groups on surfaces of different silica samples, takinginto consideration changesof specific surface areaof modified adsorbentsarepresented.Thedatashowthatconcentrations of diphenylsilyl groups on the surface for all adsorbents are nearly the same and are near to full coverage. tenfold change in specific surface area has a relatively small influence on concentrationof grafted diphenylsilyl groups per square meter of silica samples. One technique in chemical modification of silica is covering of the surface of silica by a polymer layer both by adsorption of the polymer from solution and by p o l y ~ e r i ~ a t i oofn monomers on the surface (e.g., Ref. 38). It is possible to one or more ends of polymeric chains of fragments on the silica surface. In Ref. 48 the

~oncexltrationof Crafted ~ ~ ~ h e n y lCroups s ~ l y ~ on Samples with SpecificSurfaceArea and DifferentDiameter of Particles dp Sample

Silasorb 600 1.6 LiChrosorb SI 60417 1.7 L~chrosorbSI 60 417 LiChrosorb SI 100 .9 LiChrosorb SI l 100 49 LiChrospher SI 500

45

5 5 5 265

256

7 10

process of copolymerization of N-vinylpyrrolidone on silica surfaces was investigated by infrared spectroscopy. The preparation of silica with a grafted copolymerof ~-vinylpyrrolidone chains occurs in two stages: first, grafting to the silica surface methylvinylsilyl groups by reaction of vinyltriethosysilane or methylvinyldichlorosilane with silanol groups, and then copolymerization of attached methylvinylsilyl groups with N-vinylpyrrolidone. Figure 17 shows spectra of various silica gels after reaction with =CH, in the presence of (C2H&N and after reaction with e in the presence of a catalyst. After the silanization reaction the bands of -"CH3 and -CH=CH2 appear in the spectra. Intensities of the bands, taking into consideration the thicknessof plates are in accordance with thespecific surface areaof adsorbents. After the polymerization reaction bands of viny practically disappear from the spectra and there are intense bands of groups of hydrocarbon chains and pyrrolidonerings. The absence in spe at 3020 and 3060 cm-l may indicate that the reaction of grafted vinylsilyl groups in copolymerization is complete. The appearance on the silica surface of pyrrolidone rings can be detected from the bandat 3350 cm-' of overtone vibrationof C ples with large specific surface area s 300 m2/g the intensity of groups and bondedsilanol groups so large that in the region samples became nontransparent. 3800

3000

em"

l '

Infrared spectra of silicagels after reaction with methylvinyldichlorosilanes in the presence of Et3N and following reaction of polymerization with Nvinylpyrrolidon (2) (samples were evacuated beforehand at 200°C to lo-' Torr and placed in CC14).

One of the forms of modification of the silica surface is the covering of this surface by carbon layers (49-54). If it would be possible to deposit on the surfaces of silica pores the carbon layers with formation of graphite planes, then such an adsorbent would be ideal from the point of view of chromatography owing to uniformity of surface, hydrolytic resistance, and possibility to regulatepore size adsorbents. ~nfortunately,upto the present, such anadsorbenthasnot been prepared. Silica adsorbents with carbonlayers have properties similar to those of hydrophobic chemically modified silica with high adsorption potential, but the efficiency of columns with such adsorbents usually is not too high. In some papers [55-5'73 it was demonstrated that silica adsorbents with carbon layers can be prepared by pyrolysis of CH4 and CC14 mixtures. The carbon layer deposition reaction occursat relatively low temperature, so the porous structureof silica does not change, while the methods of deposition of pyrocarbon on silica surfaces at 850°C and higher may result in noticeable structural changes. Infrared spectroscopy in this case is a suitable method for the study of changes of the surface chemistry of silica. For spectral investigation, Aerosil with l65 m2/g was usedfrom which transparent plates wareprepared by pressing. The Aerosil plate, in a holder, was placed in a reactor where the deposition of a carbon layer on the surface of the Aerosil was carried out under different conditions and the infrared spectraof this plate then recorded. Deposition of the carbon layer was carried out on a pure surface and a surface with grafted trimethylsilyl groups [58]. The reaction of trimethylchlorosilane was realized in flow CH4 at 400-420°C. In Fig. infrared spectra of untreated Aerosil and(2-6) after deposition of a carbon layer in a flowing CC14 and CH4 mixture at 400°C are presented. After heating of the Aerosil in the mixture of CC14 and CH4 for1 h there are no changes in the sample spectrum but the transparency decreases with time owing to formation of the carbon layer. To accelerate the processof deposition of the carbonlayer in these conditions, it is possible to modify the silica surface by reaction with tri~ethylchlorosilane from the gas phase.At 400°C this reaction occurs with a high rate. In Fig. 18b the spectra of Aerosil (1) before and (2) after reaction of silica surface with trimethylchlorosilane at 400°C in flow CH4 arepresented. In h the reaction of free silanol groups withtrimethylchlorosilane finishes and in spectrum 2 bands of -CH3 groups appear and band of free silanol groups at 3750 cm" disappears. After passing the CC14 and CH4 mixture across the surface of the modified samplefor 1 h the plate transparency decreases. Intensity of bands of -CH3 groups considerably decreases. This means that grafted trimethylsilyl groups take part in the process of formation of the carbon layer. In 4 h the Aerosil plate becomes almost nontransparent (6). Thus it is possible to accelerate the formation of carbon layers on silica surfaces by apreliminary silanization reaction. The acceleration of the formation rate of the carbon layer from the CC14 and CH4 mixture on the modifiedsilica surface is probably connected with formation of higherconcentrations of hydrocarbongroupsnearthesurface in comparison with the case when unmodified silica is placed in CH4 at atmospheric pressure.

Infrared spectra of Aerosil: (a) (1) heated in CH4 at 400"C, h; after passing a mixture of CH4 and CC&at 400"C, 1 h; the same, h; (4) the same, h; the same, 4 h; the same, 5 h; (1) heated in CH4 at 4OO0C, h; after reaction with(CH3)3Sic1 at 400"C, h; after passing mixture CH4 and CC14 at 400"C, 1 h; (4) the same, h; (5) the same, 3 h; the same, 4 h. (From Ref. 58.)

At the adsorption of CH4 on the silica surface is much smaller than the concentration of grafted methyl groups. From the spectra in Fig. 18a, b one cansee that at formationof the carbon layer on the silica surface the structures with hydrocarbon groups do not form, or their concentration is too small. Except for the concentration of hydrocarbon groups near surface, the rate of deposition of the carbon layer obviously depends on the flow rate of the gas mixture and temperature of the sample surface. Noticeable commencement of the carbon layer deposition reaction from the CC14 and CH4 mixture probably occurs at 400°C and the rate of this reaction quickly increases at higher temperatures. The peculiarity of this process is the increase of specific surface area of adsorbent after nucleation of the carbon phase on the silica surface. This means that after deposition of carbon in such conditions there are no dense uniform graphite carbonlayers on the silica surface since during deposition of the graphite layers the specific surface area of adsorbent should remainpractically unchanged. This result is supported by the C constant of the BET equation for nitrogen adsorption on untreatedsilica and silica with deposited carbon.For initial silica samples, C is about 160, which is usual for silica, for silica treated by trimethylchlorosilane at 400°C in flow CH4 is about 20, for silica treated by CC14 and CH4mixture in an hour C about 60 and with deposited different quantity of carbon in different time C equals several hundred, which is smaller than the value of about 1000 for graphitized carbon black During pyrolysis CC14 and CH4 mixture the particles of carbon are probablydeposited between silica particles as a new carbon phase that results in an increase of specific surface area of prepared adsorbent.

To avoid formation of this new carbon phase between silica particles it is probably necessary to remove hydrocarbon from gas phase, i.e., to replace CH4 by CO2 as the carrier gas. In this case the pyrolysis of CC14 at 400°C on silanized silica surfaces results in the formation of only carbon film on silica surface [60]. The specific surface area of silanized silica gel changes very little and the C constant of the BET equation for nitrogen adsorptionincreases from about 50 for untreated silanized silica gel (LiChrosorb SI-60) to 80-90 for samples treated for different times at 400°C. It means that decomposition of CC14 occurs on the surface of adsorbent and the forming carbon layer is not porous, but at the same time it is not a graphite layer, since C for this adsorbent is much smallerthan forgraphitized carbon black lOOOj. Reposition of carbon layers on hydroxylated silica surfaces can also be realized by treatment of silicaat 400°C with a vapor mixture of CC14 and chlorosilanes [e.g., C12Si(CH3)2] inCO2 as thecarrier gas. In this case, at first the silanization reaction occurs and then formation of the carbon layer. Comparison of gas chromatographic data on retention, heats of adsorption of different class compounds aswell the contributionsof various functional groupsof molecules to these characteristics shows that the adsorption potentialof silica with a carbon layer higher than for untreated silica but smaller than carbon adsorbents such as Carbopack or graphitized carbon black. At the same time the influenceof residual polar (silanolj groups ofsilica surface on retention of polar compounds is remarkable [61],

The discovery of a new carbon modification-fullerenes [62] and the development of a method for its preparation on macroscopic scale E631 makes it possible to use carbon clusters (molecules) for modification of silica surfaces. By bonding fullerene molecules to the silica surface it possible to prepare the carbon-containing adsorbents with a smaller adsorption potential in comparison with carbon adsorbents, which has some practical uses. InRef. 64,silicamodified withbonded fullerene c60 by reaction of surface silanol groups with chlorosilane derivatives of fullerene is described. The most simple derivatization reactionof fullerenes is that with primary amines [65]. This reaction was used for immobilization of fullerene to silica surfaces [66]. The first step of immobilization of fullerene is modification of silica by y-arninopropyltriethoxysilane and then reaction of fullerene with grafted y-arninopropylsilyl groups. Of course for immobilization of fullerenes on silica it is possible to use aminosilicas, these being commercial products. Silicamodified by fullerenes has some peculiarity from the point ofviewof changes of the specific surface area of the prepared adsorbent. Usually, chemical modification of silica results in a decreaseof specific surface area of adsorbents if a porous surface layer is not formed. The bonding ofrelatively large “spherical” fullerene c60 molecules forms additional surfaces owing to these molecules that should result in an increase of specific surface area of silica adsorbents [66]. Silica adsorbents with differing content of immobilized fullerene were prepared by changing of solvent composition and temperature.

In Fig. 19a the results of measurements of specific surface area of adsorbents with different amounts of bonded fullerene are shown. After reaction of the silica with ~-al~inopropyltrietho~ysilane, the specific surface area of adsorbent becomes smaller in comparison with untreated silica gel. Usually at chemical modification andafterimmobilization of fullerene the specific surfacearea of adsorbent increases; this increase is approximately proportional to the contentof immobilized fullerene. So it is possible to evaluate the specific surface area of i~mobilizedfullerene. The specific surface area of fullerene molecules bonded to silica calculated from the correlationbetween fullerene surface and contentof immobilized fullerene is about 725 m2/gfromdataonnitrogenadsorptionat 77 K (Fig. 19b). The geometrical surface and specific surface area of isolated fullerne molecules can be calculated by using the van der Waals diameterof the molecule, 1.0 nm From geometric dimensions of fullerene molecules the specific surface area should be 2620 m2/g. The experimental value is 3.6 times smaller. This may be owing to the use for determination of surface area as adsorbate nitrogen molecules with 0.162 nm2 in a dense monolayer. In this case the meanlinear dimension of the adsorbed nitrogen molecule is about 0.4 nm, which differs not too much from the diameter of fullerene molecules. Another reason is probably connected with the impossibility of nitrogen molecules to interact with all surfaces of spacer-bonded fullerene molecules owing to contact of the bonded fullerene molecules with each other or with the surface of the silica adsorbent. If one proposes the uniform distribution of arninopropyl groups with concentration 1.9 pmol/m2 for one of our samples (usual concentration of amino groups of aminosilica) then the average distance between amino groupswill be about 0.834 nm. This meansthat the arrangementof can be only in the manner presented in Fig. 20 and concentration of bonded fullerene molecules is 0.53 pmol/m2, i.e., only about 28% of amino groups can be involved in reaction of modi~cationby full-

M n

0-

(a) Changes ofspecific surface area ofsilicagel LiChrosorb SI 100 before and (2) after chemical modification by ~-aminopropyltriethoxysilane,and (3) after il~l~obilization of various amounts of correlation between the increasing of surface area of silica geland amountof immobilized fullerene,(From Ref. 66.)

Model of distribution of immobilized fullerene moleculeson the surfaceof gel.

erene CGO.Experimentally determined concentration of bonded is 0.54 pmol/ m2. From Fig. 20 one can see that only small molecules such asN2, H20, etc. can penetrate between the attached fullerene molecules at the highest concentration of fullerene on the silica surface. One important problem is the determination of the surface concentration of immobilized fullerene on silica. As it has turned out, the adsorption of dyes from water solution is the most simple method of determination of this concentration. The adsorption properties of silica with bonded fullerene are determined by the concentration of immobilized fullerene. The greater the concentration of immobilized fullerene, the larger the specific surface area of adsorbent the higher adsorption of dye (Fig. 21) but the adsorptionof Rhodamine B well as Rhodamine relative to immobilized fullerene concentrationfor all adsorbents is practically common and the limiting adsorption of dyes on silica with bonded C60 isvery near toone (Fig. 22). This meansthat one molecule of dyeadsorbs on onemolecule of bonded fullerene and the isotherm of dye adsorption can be used for determination of the concentration of inmobilized fullerene on silica,

Changing the adsorption properties of silica surface can be realized not only by chemical modification withgrafting some groups to the surface, but by adsorption modification when bifunctional molecules form strong adsorption bonds with functional groups of the silica surface. As result the strong adsorption modifying layer is formed on the silica surface [68-711. In ion-pair and ligand-exchange chromatography the pretreatment of the column is possible to consider as adsorption l~odifi~ation of adsorbent surface170,711. Although the mechanismof separation in

80

L

0

Adsorption isotherms of Rhodamine B from water solution on silica gels: KSK-2 and (9) LiClxosorb Si withvarious (1,2) LiChros~herSI concentrations of immobilized fullerene: (2) (6) (7) 121, and (9) 232 p,mol/g. (From Ref. 66.)

8 c12

Relative adsorption of Rhodamines B and 6 1s withimmobilizedfullerene. (1) Rhodamine (From Ref. 66.)

from water solutions on and(2)Rhodamine 6C.

ion-pair chromatographyis not completely clear, the role of adsorption modifiersis probablyconnectedwithstrongadsorption ofpositively or negatively charged groups with hydrocarbon chains from water solution on the hydrophobic modified silica surface, which takes part in separation of acids, salts, and bases. In the caseof ligand-exchange chromatography, example, ~-alkylderivatives of a-amino acids adsorb onhydrophobized surfacesof silicaowing to interactionof the hydrophobic partof the molecule with the hydrophobiclayer of adsorbent, and then form ion complexes(usually Cu2+);such an adsorbent is capable of separating racemic mixtures of amino acids on optical isomers completely [71]. For separation of carbohydrates, at present the most often used silica gels are those with bonded ~-aminopropylsilyl groups[72]. On chemically modified adsorbents with grafted -NH2 groups the separation of carbohydrates occurs as a result of differences in energyof interaction of their moleculeswith electrondonor amino groups, But adsorption modifiers can be used also to form on the silica surface the layer of amino groups, as adsorption modifiers have tobe molecules which, at one end molecules can be retained strongly on the silica surface owing to relatively strong adsorption bond and at the other end these molecules have to be amino groups turned tosolution. was mentioned, amines form strong hydrogen bonds with surface silanol groups of silica. So convenient molecules for formation on silica surfaces of strong electron-donor groups turned to solution are diamines. In Fig. 23 the models of a silica surface with grafted-NH2 groups by chemical modification and with NH and "H2 groups of adsorbed molecules from solution owingto stronghydrogen bonds of diamines withsilanol groups of surface are presented. Ethylenediamine and piperazine are strong bases. Their molecules are attached by hydrogen bonds of one amino group to SiOH and other amino groups can form hydrogen bonds with molecules having electron-acceptor groups. Silanol groups are moreacidic than hydroxyl groupsof carbohydrates for example. So diamine molecules form stronger bonds with the silica surface than carbohydrates in solution, whichmakes it possible to separate carbohydrates on silica surfaces by an adsorption modification [68,69]. The concentrationof diamine in solution necessary for the formationof a dense monolayer on the silica surface is determined by the equilibrium adsorption con-

Models of ~ o d i f y i nlayer ~ of silica surface after chemical modification ~-a~inopropyltriethoxysilane (a), and after adsorption l~odification piperazine and et~ylenediamine(c). (From Ref. 68.)

stant in solution. The larger the adsorption constant, the smaller the concentration of diamine in solution needed. Of two diamines used, ethylenediamine and piperazine, themore suitable asanadsorptionmodifier piperazine being a solid compoundmoderatelysolubleinwater. Incontrastwithethylenediamine,the moleculeof piperazine isrelativelyrigid and does not change conformation on adsorption. So one of two amino groups turn to solution.

Silica is an important material which is widely usedin industry, technology,science, protection of the environment, etc. We mention a few applications of silica in more detail with regard to silica samples inscience and chromatography. One example of pyrogenic silica is Aerosil (or Cabosil), which is used as a model sample for investigation of silica surface chemistryby infrared spectroscopy, since from this silica it is very easy to prepare transparent plates in the infrared region of the spectrum. Samples of this silica have been used from the earliest studies up to thepresent The adsorption properties of pyrogenic silica in the hydroxylated state are similar to properties of other samplesof hydroxylated silica so Aerosil has turnedout tobe very useful for investigation of the mechanism of adsorption on silica, chemical modification of silica, transformation of silica under heating, etc. For the same reason Aerosil (Cabosil) was used as supports of catalyst (metal) for investigation of the mechanism of catalytic reaction by infrared spectroscopy (e.g., Refs 73-75). As an example, in Figs 24 and 25 the spectra of Cabosil with deposited platinum before and after adsorption of hydrogen and benzene are presented Spectral investigation shows that under 400 Torr of hydrogen, in the

5

h,

Infrared spectra of Cabosil with deposited platinum: (1) after adsorption of hydrogen (a) and deuterium (b) at -400 Torr at ambient temperature, (2) after evacuation of gas phaseat ambient temperatureand (3) at -400 Torr of hydrogen or deuterium at -300°C. (From Ref. 76.)

Infrared spectra of Cabosil with deposited platinurn: (1) after evacuation, (2) after adding benzene vapor(-40 Torr), after evacuationof benzene vapor, (4) after adding hydrogen(“400 Torr), after evacuationof hydrogen, after adding hydrogen, (7) after evacuation of gas phase, (8) after adding hydrogen, and(9) after evacuation of gas phase. operations atambient temperature. (From Ref. 76.)

spectrum of Cabosil-supported platinum, bands, at 2128 cm-’ or 1527 cm-l for deuterium appear. After evacuationof the sampleat room temperature these bands disappear. These bands are attributed to bands of hydrogen and deuterium bonded with platinum, since, at ambient temperature, adsorptionof these gases on silica is negligible. At 300°C the bends at 2128 and 1527 cm“ were practically unobservable even at 400 Torr of hydrogen or deuterium respectively (Fig. 24). Onadsorption of aromatichydrocarbons on silica-supportedplatinum it is possible to observe intermediate compounds able easily to accept and to return hydrogen. Figure 25 shows the spectrum (1) silica-supported platinum, (2)

sample with adsorbed benzeneat room temperature. In spectrum 2 there are >CH bands of aromatic rings. After evacuation of benzene at room temperature, physically adsorbed benzeneis removed and in the spectrum only a broad band at 3040 cm" of very small intensity is observed (3). After adding hydrogen 400 Torr), in the spectrunn the intense bands of -CH2 at 2935 and 2859 cm" appear (4). After evacuation of hydrogen at room temperature these bands practically disappear (5). But addition of a new portion of hydrogen 400 Torr) results in the appearance in the spectrum of these bands with almost the same intensity (6). At evacuation these disappear again (7). Repeating these cycles results in the same pictures, (8) and (9). This means that on adsorption of benzene on the platinum surface a Chemisorption complex is formed which is able to accept and to lose hydrogen very easily [73-751. Similar pictures are observed chemisorption of toluene, p-xylene, and mesitelene on silica-supported platinurn 1761. odified silica is widely used as an adsorbent forliquid chromatography. A lot of modifying groups canbe attached to silica surfaces for separation any organic compounds from simple to complicated biologically active substances and polymer molecules. Many separations can be made on silica with bonded alkyl groups and with octadecylsilyl groups in particular (ODS silica gel). Alkyl chains attached to the surface, owing to conformational mobility, may change the thickness of the modifying layer depending on eluentcomposition.Attachment of diphenylsilyl groups to silica surfaces makes it possibleto endow the adsorbent with hydrophobic properties of the surface the same as ODs silicagel, and at the same time achieve a thickness of modified layer independent of eluent composition. Silica with grafted diphenylsilyl groups was used for separationof cardiac glycosides and steroid hormones and forinvestigation of mechanism of their adsorption and separation [45, 69,77,78]. In Figs and 27 the separationsof some cardiacglycosides, which are the most important and effective biologically active compounds for treatment of abnormal circulation of the blood due to high blood pressure, atherosclerosis, and infarction

1.

Separation of (l) convallatoxin and (2) desglucoheiroto~inon silicagel with bonded diphenylsilyl groups from 3:7 ethanol-water used as the eluent (temperature: 40°C). (From Ref. 78.)

of the miocardia, are presented.In reversed-phase liquid Chromatography the main intermolecular interactions determining retention are the van der Waals interaction of molecules with hydrophobic groups attached to the silica surface and the strong intermolecular interactions of the eluent molecules with each other. In the case of chromatography of cardiac glycosides, the molecules of which contain aglycone and glycone (sugar part), the intermolecular interactions are rather complicated. First of all, it is possible to note the effect of the geometry of molecules on retention. Figure 26 shows the separation of desglucoheirotoxin and convallatoxin [78]. These molecules are identical from the pointof view of their general formulae but the same aglycone, strophantidin, is connected with a-l-rhamnose in convallatoxin and p-D-gulomethylose in desglucoheirotoxin. Such a difference results in various configurations of the whole molecules. a result the more flat desglucoheirotoxin moleculeis retained more strongly than the convallatoxin molecules, and they can be separated on silica with bonded diphenylsilyl groups. evaluate the role of glycone and aglycone in retention of cardiac glycosides on silica with bonded diphenylsilyl groups some artificial mixtures were prepared chrol~atograms,which are presented in Fig. 27. The glycoside mixtures were chosen in such a manner that the glyconeswere similar in size in each compound group. Thisfigure illustrates that the aglycone structure, the number of hydrophilic and hydrophobic groups and the position of these groups in the aglycones have a great influence on the retention of cardiac glycosides. For example, in Fig. 27a all compounds may be subdivided into three types of glycosides: glycoside with the most hydrophilic oubagenin-aglycone, glycosides with the strophantidin aglycone and glycoside with the oleandrigenin aglycone. The effect of properties of aglycone on retention of glycosides is demonstrated most clearly in the separation of lanatosides and C having the same glycone (Fig. 27d). Themosthydrophobicaglyconein this mixture is digitoxigenin. Therefore lanatoside has the longest retentiontime. Introduction of one hydroxyl group in the CI6position decreasesby more than a factorof two the retention time of lanatoside B, while introduction of hydroxyl groupsto position C12decreases the retention time of lanatoside C by a factor of five in comparison to lanatoside Figure 27d demonstrates also that glycosides with the same aglyconeare separated according to the hydrophilic and hydrophobic properties of the glycone. Separations of some estrogens onsilica gelwith bonded diphenylsilyl groups are presented in Fig. 28 [45,69]. The order of elution from the columnis determined by hydrophilic and hydrophobic properties of these steroid hormones. ~ydroxylatedsilica gels also can be used for separation of cardiac glycosides in thin-layer chromatography (TLC) using benzene-ethanol eluent [79]. The mechanismof separation in this case will be different. The main interaction is that of hydroxyl groups of cardiac glycosides with surface silanol groups ofsilica and interaction of hydrocarbon groups with eluent molecules. In Table4, comparison of the contribution of different functional groups of cardiac glycosides to adsorption on silica with bonded di~henylsilyl groups and adsorption on hydroxylated silica are presented. The data of Table 4 showthat in reversed-phase liquid chromatography, i.e., on separation of cardiac glycosides on hydrophobized silica the only hydrocarbon

Ho

OH

(b)

CHs

4

I

1

C

0'1 k

20

I0

f-

,

30 Y,minoms

I

J

30

7,min

FIG. 27 Separation of cardiac glycosides mixtures on silica gel with bonded diphenylsilyl groups. Eluent: ethanol-water in the following proportions: (a) 7:13, (b) 3:7, (c) and (d) 2:3. Temperature: (a) and (d) 50°C, (b) 40"C, (c) 60°C. Peaks: (a) (1) G-strophanthin, (2) convallatoxin, (3) desglucoheirotoxin, (4) erysimin, ( 5 ) cymarin, (6) oleandrin; (b) (1) corelborin-n, (2) olitoriside, (3) K-strophanton-B; (c) (1) K-strophanthoside, (2) digoxin, (3) digitoxin, (4) acetyldigitoxin; (d) (1) desacetyllanatoside C, (2) lanatoside C, (3) lanatoside B, (4) lanatoside A. (From Ref. 78.)

Separation of estrogens hormones on silica gel with bonded diphenylsilyl Croups at 50°C. Eluent: ethanol-water (a) 7:13, (b) Peaks: (a) (1) estriol, (2) estradiol, (3) ethynylestradiol,and (4) estron;(b) (1) estradiol, (2) mestranol, (3) estradiol valerianate, (4) estradiol benzoate, (5) estradiol dipropionate. (From Ref. 69.)

Contributions of Various Groups of Cardiac Glycosides to Their Retention on Silica from Benzene-Ethanol Mixture (3:1) at Ambient Temperature 1nk' (on TLC data) andonSilicawithBonded ~iphenylsilylGroups from Ethanol-Water Mixture (35:65) at 50°C Ink' (on HPLC data)

-0.096 0.052 0.115 0.218 0.593 0.309 0.285 0.762 0.600 0.612 0.343 0.403

0.193 -0.136 -0.659 -0.278 -0.49 1 -1.533 -0.809

1

-0.720

Ref. 79.

groups of their molecules increasetheretention (Ink' 0) and all hydrophilic groups decrease the retention (Ink' 0) but in normal-phase chromatography, on separation on hydroxylated silica the hydrophilicgroupsincreasetheretention owing to hydrogen bonding with silanol groups and hydrocarbon groups decrease the retention. Silica with bonded poly-~-vinylpyrrolidoneis used for separationof proteins. In this case the modifying layer protects proteins from strong adsorption on hydroxylated silica. The grafting of N-vinylpyrrolidone groupsis usedfor coatingof silica capillary walls on separation of proteins by capillary zone electrophoresis ~-vinylpyrrolidone groupsgrafted to silica surfaces include both hydrocarbon and electron-donor fragments. So such adsorbents can be used for separation of polyaromatic hydrocarbons and proton-donor compounds. In Fig. 29, separation of some aromatic hydrocarbons on silica with bonded poly-~-vinylpyrrolidone from heptane is presented. The order elution corresponds to the molecular mass of the compounds. In Fig. 30a, b, the separations of very polar compounds (acids) are presented. Gallic acid sometimes is used as a nonadsorbedcompoundfordetermination of the void volume of columnsin reversed-phase liquid chromatography,owingtostronginteraction of sucha hydrophilic molecule withaqueous eluents. This molecule practically doesnot adsorb on the hydrophobic surface of modified silica. But on silica with bonded N-vinylpyrrolidone groups, gallic acid has the highest retention for compounds in mixtures under investigation (Fig. 30a). In Fig. 30b, the separation of some derivatives of benzoic acid as an example of separation of hydrophilic compounds on silica with bonded ~-vinylpyrroiidone are presented. The elution order is deter-

5

min

Separation of some polyaromatic compounds on silica gel Silasorb 300 by N-vinylpyrrolidone. Eluent: heptane. Temperature: 40°C. (1) toluene, fluorene, (3) anthracene, (4) triphenylene, perylene.

mined not only by interaction of carboxyl groups of separating molecules with pyrrolidone groups bonded to the silica surface but by substituted groups with eluent molecules. After the discovery of fullerenes and the development of the method of their preparation. on macroscopic scale [62,63] the problems of separation, purification, and analysis of fullerenes arise. It was found that silicas with different bonded modifiers are very useful adsorbents for separation and analysis of fullerenes by HPLC [81-861. Separation of fullerenes on silica gel itself ispossible, but selectivity is not high and it is necessary to have a very effective column to separate and A more useful adsorbent for separation issilica with bonded diphenylsilyl groups.Theseparation is based onthestronger interaction of fullerenes with

I Separation of some aromatic acids on silica gel Silacorb 300 modified by ~-vinylpyrrolidone.Eluent: phosphate buffer (pH 2, I 1S). Temperature: 40°C. (a) (1) benzoic acid, (2) 3,4-dihydrobenzoic acid, (3) gallic acid; (b) phenylalanin, (2) p-an~inobenzoicacid, (3) p-nitrobenzoic acid, (4) benzoic acid, (5) p-oxybenzoic acid, (6) p-chlorobenzoic acid, (7) p-toluic acid.

phenyl groups on silica surfaces than with n-hydrocarbon molecules used as eluent. For the separation of fullerenes, silica with bonded octadecylsilyl groups is widely used. Adding alcohol to the eluent makes it possibleto separate oxide of fullerene from fullerene. For separation of fullerenes it is possible to use hydrogen bonding between fullerene molecules havingto some extent aromatic propertiesand bonded to silica surface hydroxyl groups. A column with the silica gel LiChrosorb Diol is very useful for analysis not only of fullerene itself but also oxide derivatives of fullerenes [8 1,821. At present many typesof stationary phaseson bases of modified silicahave been prepared for effective separations not only of fullerenes but also their derivatives [64, 8 1-86]. naryphasessuchastripodal arrange~ents three 2,4-dinicovalently attached through a Clo spacer a m to silica er" [83] as well as silica with bonded groups 2-(1-pyrenyl)ethyl, 3(1-pyrenyl)propyl, 2-nitrophenylethyl, 3-(~-nitrophenoxy)propyl[84], tetraphenylporphyrin [85], and dinitroanilinopro~yl [86] can be used for separation of fullerenes not only on analytical but on semipreparative and preparative scales.

40

Separation of fullerenesonsilicagelswithbondeddiphenylsilylgroups (a), with bonded octadecylsilyl groups (b),and silica gel LiChrosorb Diol (c). (From Ref. 81.)

For separationof fullerenes on preparative scale, activated carbon is used (e.g., Ref. 87). Silica with a deposited carbon layer can also be used for separation and purification of fullerene in semipreparative scale In Fig. 31a, b, c the separationsof fullerenes on different modified silica gelsare presented. In Fig. 32a, b the separations of oxide derivatives of fullerenes on the silica gel LiChrosorb Diol are demonstrated. Figure 33 shows the semipreparative separation of and on silicagel with deposited carbon layer. Fullerenes can be separated on silica with bonded fullerene (Fig. 34) (64,881. Application of adsorption modification of silica can be demonstrated on separation of carbohydrates. In this case hydrogen bondingbetween the hydroxyl groups of carbohydrate molecules and amino groups bonded to the surface silica adsorbents is used. In a number of publications the separationsand analysis of carbohydrates are carried out on silicamodifiedby ~-amonopropyltriethoxysilane(e.g., Ref. 72). Adsorption modification for the formingof amino group layers for this purpose can also be used. Figure 35 shows the separationof sugar mixture on hydroxylated silica gel, and on silica gel with pieperazine adsorbed from eluent, as well as on silica gel with bonded ~-amonopropylsilyl groups [68,69]. From this figure it can be seen that on the hydroxylated silica surface this mixture of carbohydrates practically cannot be separated from the acetone-water eluent because the difference in intermolecular interaction of carbohydrates withsilanol groups of adsorbents from water solution is too small. Moreover, acetone and water molecules of the eluent are competitors, forming quite strong hydrogen bonds with silanol groups. After addition of piperazine to the eluent, satisfactory separation of the sugar mixture takes place. The elution order of the carbohydrates monosaccharides, then disaccharides, and finally trisaccharides, in accordance with the increasing numberof

6

0

~eparationofoxidederivativesoffullerenes(a) c 6 0 and (b) onsilica gel LiChrosorb Diol. c 6 0 ; (2) C 6 0 0 ; (3) C6005 (4) c6003;(5) c6004 (6) C6005; (1) C60; (2) (3) C7,0; (4) C7002; C7003.(From Ref. 82.)

proton-donor groups in their molecules. Similar separation of the same mixture takes place on silica with bonded amino groups. The effect of molecular geometryon retention of carbohydrates of silica adsorption modified by piperazine is important. In Fig. 36 the separationsof some pentose and hexose stereoisomers are presented. By using the adsorption modifier piperazine it is possible to completely separate these isomers on hydroxylated silica gel. The elution order of mannose, glucose, galactose on the one hand, and lyxose, xylose, arabinose on the other, is determined by the similarity in the geometry of

Se~ipreparativeseparation of fullerenes deposited carbon layer. (From Ref. 60.)

and

on silica gel with

L

fullerene

Separation offullerenes (1) (From Ref. 88.)

and (2)

on silica gel withimmobilized

30

Separation of carbohydrate mixture on hydroxylated silica gel at 20°C from 4:l acetone-water eluent on hydroxylated silica gel (a), on silica gel adsorbed from eluent piperazine and on silicagel with bonded aminopropylsilyl groups (c). Peaks: (1) ribose; (2) xylose; (3) fructose; (4) glucose; (5) sucrose; (6) cellobiose; (7) melezitose; (8) raffinose. (From Ref. 68.)

their structure, Inspite of the fact that the numberof hydroxyl groupsin all pentose is the same (as well as for hexose also) the different geometrical arrangement of hydroxylgroups in their molecules makes it possible toseparatetheisomers completely. Adsorption modifiersin liquid chromatography in some cases makes it possible to replace compounds used for chemical modification. In water eluents, adsorbents with bonded amino groups have basic media which result in removal of bonded NHz groups from the silica surface. The column in this case loses capacity with time. By using anadsorptionmodifierthesurface of theadsorbent is always covered by an adsorbed layer of modifier if its concentration in the eluent is sufficient and the column works stably. Some kindof adsorption modification takes place during deposition of fullerene on silica surfaces by evaporation of solvent from thefullerene solution added to the silicagel. Theadsorptionandchromatographicproperties ofsilica essentially change. On silica with deposited fullerene it is possible to separate ethers (Fig. 37) [61j,

2

4

0

l0

30

min

0

"0. 30

r, min

Separation ofhexose (a)andpentose(b)isomersonsilica gelwith adsorbed piperazine from eluent. Eluent: (a) and (b) 9:l acetone-waer, both including 0.44 mg/ml piperazine. Peaks: (a) (1) talose; (2) mannose; (3) glucose; (4) galactose; (b) ribose; (2) lyxose; (3) xylose; (4) arabinose. (From Ref. 68.)

In recent years, for investigation ofsilica surface chemistry the method of cross-polarization ~agic-angle-spinningnuclearmagneticresonance (CP NMR has been used. In Refs 89 and 90 this method makes it possible distinguish to single and geminal silanol groups, the concentration of which is relatively small compared with all surface silanol groups of silica. The review of investigations of silica by this method is presented in Ref. 5. Unfortunately there are few data on differences in interactions with adsorbed molecules and in chemical modification reactions between single free and geminal silanol groups. But even if these differences exist, and they should exist, the main relationships of adsorption properties already determined cannot be changed noticeably because the concentration of geminal silanol groups is not so high. Moreover, by infrared spectroscopy it is difficult to find the difference in frequency of single and geminal silanol groups and only by such a sensitive method as"Si (CP MAS) has it been possible to distinguish them and therefore differences in properties of single and geminal silano1 groups should not be so noticeable.

" I " 8

(a) Separation of some etherson silica gel with deposited fullerene (extract mainly a mixture -85% and 15%). (b) For comparison chromatogram diethyl ether on initial silica gel. Ref. 61.)

The investi~ationof adsorption of col~poundsfrom ~ulticomponentsolutions including aqueous solutions on hydroxylated and modified silica adsorbents is of great interest for technology, chromatography, and e~vironmental protection. The main results of these investigations are presented in Ref. 9 l .

l . L. A. 2. A. New 3.

Little. Spectra Species (with supplementary chapters by Kiselev and I. Lygin), Academic Press, London, New York, 1966. ~ h y , Press, Kiselev and Ya. 1. Yasbin. ~ a s - ~ d s o r ~Ct ~i or o~ ~ u t o g r ~Plenum York, 1969. Unger. Silica, Its Pro~ertie~s and Use S ~ ~ p o in rt L~~uid Elsevier, Amsterdam, 1979. 4. Her. The C h e ~ i s t r y Silica, A Wiley-Interscience Publishers, New York, 1979. 5. P. Legrand, (ed.), The S ~ r f a c e~ r o ~ e r t i e s Silicas, Wiley Sons Ltd, New York, 1998. 6. Ua. Davydov and A. Kiselev. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 37:2593 (1963).

7. V. Ya. Davydov, A. V. Kiselev,and L. T.Zhuravlev. Trans. Faraday Soc. 60:2254 (1964). 8. V. Ya. Davydov, L. T. Zhuravlev, and A. V. Kiselev. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 38:2047 (1964). 9. V. Ya. Davydov, A. V. Kiselev, V. A. Lokutsievskii, and V. Lygin. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 47:809 (1973). Ar~ltiunyan,A. V.Kiselev, and T.I. Titova. Dokl. Acad. Nauk SSSR 10. B. 251:1148 (1980). 11. G. Curthoys, V. Ya. Davydov, A. V. Kiselev,S. A. Kiselev, and B. Kuznetsov. J. Coll. Interface Sci. 48:58 (1974). Kiselev, and S. A. Kiselev. Kollooid Zhurn. (Colloid 12, V. Ya. Davydov, A. Russian) 41:227 (1979). 13. N.V. Akshinskaia, V. Ya. Davydov, A. V. Kiselev, and Yu. S. Nikitin. Kolloid Zhur. (Colloid Russian) 28:3 (1966). 14. R.L. Gorelik, Ya. Davydov, L. T. Zhuravlev, G. Curthoys, A. V. Kiselev, and Yu. S. Nikitin. Kolloid Zhurn. (Colloid Russian) 35:456 (1973). 15. R. L. Gorelik, V. Ya. Davydov, S. P. Zhdanov, L. T. Zhuravlev, C. Curthoys, A. V.Kiselev, and Yu. S. Nikitin.Kolloid Zhurn. (Colloid Russian)35:1152 (1973). Basic ~ r i n c ~ l e s Vol. 7, Springer, Berlin, etc., 1972, Pfeifer. 16. 53. Ya. Davydov, A. V. Kiselev, H. Pfeifer and Yunger. Zhur. Fiz. Khim. (J. 17. Phys. Chem., Russian) 57:2535 (1983). 18. L. T. Zhuravlev, A. V. Kiselev, V. P. Naidina, and A. L. Poliakov. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 37:2258 (1963). in ~ x ~ e r i ~ e n teth a l hod in ~ ~ s o r ~ t iand on ~ o l e c ~ l ~ r 19. L. T. Zhuravlev, C ~ r o ~ a t o g r a(V. ~ ~P. y Dreving and A. V. Kiselev, eds.), MSU, Moscow, 1973, pp. 250-272. Davydov, A. V.Kiselev, S. A. Kiselev, and V.0.-V. Polotnyuk. J. 20. V. Interface Sci. 74378 (1980). 21 L. T. Zhuravlev. Langmuir 3:316 (1987). Kiselev. Kolloid Zhurn. (Colloid 22 V. Ya.Davydov, T. V. Erexneeva, and A. Russian), 43:144 (1981). 23. A. V.Togansen. Teor.Eksperirn.Khim.(Theor.Exp.Chem.,Russian)7:302 (1971);Intensity InfraredSpectra and HydrogenBonds, Moscow, Moscow State University, 1969. 24. V. Ya. Davydov,A. V.Kiselev, and V. I. Lygin.Kolloid Zhurn. (Colloid Russian) 25: 152 (1963). 25. V. Ya. Davydov, A. V. Kiselev, and V. I. Lygin. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 37:469 (1963). 26. V. Ya. Davydov and A. Kiselev. Kolloid Zhurn. (Colloid Russian) 30:353 (1968). 27. V. Ya. Davydov, A. V. Kiselev, and B. V. Kuznetsov. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 39:2058 (1965). 28, V. Ya. Davydov, T. M. Roschina, N. M. Khrustaleva, and A.A. Mandrugin, Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 67:2428 (1993).

29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46.

47. 48. 49. 50. 51. 52. 53. 54.

Ya. Davydov and A. Kiselev. Kolloid Zhurn. (Colloid J., Russian) 42:348 (1980). Ya. Davydov, A. Kiselev, V. A. Lokutsievskii, and I. Lygin. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 48:2267 (1974). C. Zettlemoyer and K. Klier. Surface C ~ e ~ i s tof r yOxides. Disc. Faraday Soc. 1971, No. 52, pp. 46-49. Yu. Babkin and A. Kiselev. Zhur. Fiz.Khim. (J. Phys.Chem.,Russian) 36:2448 (1965). C . S. Pimentel and A. L. McClellan, The ~ ~ d r o g eBond n (L. Pauling ed.), San Francisco, London, 1960. L. Boksanyi, 0. Liardon, and E. Sr. Kovats, in Advances and Interface Science, 1976, Vol. 6, pp. 95-137. A. Tertykh and L. A. Beliakova. C ~ e ~ ~Reactions i c a ~ with ~ a r t i c ~ a t of i oSilica ~ Surface, Naukova Dumka, Kiev, 1991. A. Dabrowski and V. A. Tertykh (eds.),Adsorption on New ~norganic 99, Elsevier, Amsterdam, Sorbents Studies in Surface Science and Catalysis, 1996. Beresin, B. K. Antonov, and K. Martinek(eds.), ~ ~ ~ o ~ i IEnzymes. ized ~ o d e r nState and Perspectives 1, MSU, Moscow, 1976. Silica 98. Extended Abstracts Oral and Poster ~resentations,Mulhous, France, 1-4 September,1998. A. Kiselev, B. V. Kuznetsov, and S. N. Lanin, J. Coll. Interface Sci. 69:148 (1979). K. Unger, K. Berg, and E. Gallei. Kolloid Z. Z. Polymer. 234:1108 (1969). K. Berg and K. Unger. Kolloid Z. Z. Polymer. 246:682 (1971). Ya. Davydov and I. Nazarova. VestnikMosk.Univ.(Russian)Series2, Chemistry 29:580 (1988). Davydov, M. Elizalde Gonzalez, A. Kiselev, and K. Lenda. Chromatographia l#: 13 (1981). Ya. Davydov, M. Elizalde Conzalez,and A. Kiselev, J. Chromatogr. 204293 (1981). A. P. Arzamastsev, V. Ya. Davydov, M. Elizalde Gonzalez,A. Kiselev, and R. Rodionova. Khirn.-Farm. Zhur. (Chern.-Pharm. J., Russian) 17:494 (1983). W. Wemetsberger, M. Kellerman, and Ricken. C h r o ~ a t o g r a p ~ i10:726 a (1 977). V. Ya. Davydov, A. V. Kiselev, and B. Kuznetsov. Zhur. Fiz. Khim. (J. Phys. Chem., Russian) 6#:1 (1970). L. B. Zubakova, N. Borisova, S. K. Koroleva, V. Ya. Davydov, and C . N. Filatova. Prikl. Khirn. (J. Appl. Chem., Russian) 60:1491 (1987). N. K. Bebris, R. G.Vorobieva, A. V. Kiselev,Yu. S. Nikitin, L. Tarasova, I. Frolov, and Ya. Yashin, J. Chromatogr. ll7:257 (1976). W. Colin and G. Cwiochon. J. Chromatogr. 137:19 (1977). J. H. Knox, B. Kaur, and G. R. M i l ~ ~ a rJ.d .Chromatogr. 352:25 (1986). K. K. Unger. Anal. Chem. 55:361A (1983). Plzak, F.P. Dousek, and J. Jansta. J. Chromatogr. 147:137 (1978). C . G. Guiochon and H. Colin. Chromatogr. Rev. 4:2(1978).

55. E. P. Smirnov, N. A. Bogomolov, S. I. Koltsov, and V. B. Aleskovskii. Zhur. Prikl. Khim. (J. Appl. Chem., Russian) 52:2590 (1979). 56. P. A. Tesner, E. V. Denisevich, and I. Yu. Kulikov. Kinetika i kataliz (J. Kinet. Catalysis, Russian) 25:990 (1984). 57. E. V. Denisevich and P. A. Tesner. Neftekhimia (Petrochemistry, Russian) 21:552 98 1). Ya. Davydov and V. I. Nazarova.VestnikMosk.Univ.(Russian)Series2, 58. Chemistry 29:580 (1988). 59. V. Ya. Davydov and V. I. Nazarova,VestnikMosk.Univ.(Russian)Series2, Chemistry 32:127 (1990). 60. V. Ya. Davydov and G. N. Filatova.Vestnik Mosk. Univ.(Russian)Series2, Chemistry 37:357 (1996). 61. V. Ya. Davydov, T. M.Roschina, G. N. Filatova, and N. M.Khrustaleva. Vestnik Mosk. Univ. (Russian) Series 2, Chemistry 36:518 (1995). 62. H. W. Kroto, J. R. Heath, S. C. O'Brien, R. F. Curl, and R. E. Smally. Nature 318: 162 986). 63. W. Kratschmer,L. D, Lamb, K. Fostinopoulos, and D. R. Huffman. Nature 347:354 (1990). 64. Y. Saito, H. Ohta, H. Terasaki, H. Nagashima, and K. Jinno. J. HighResol. Chromatogr. 18:569 (1995); Separation of Fullerenes by Liquid Chromatography (K. Jinno, ed), RSC, UK, 1998. 65. F. Wudl, A. Hirsch, K. C. Khemani, Suzuki, P.-M. Allemand, A.Koch, and Srdanov,in ~ ~ l l e r e n e sSynthesis, : Properties and C h e ~ i s t r y Large Carbon Hammond and V. J. Kuck, eds.), The American Chemical Soc., Clusters Washington, D.C., 1991, p. 161. 66. V. Ya. Davydov, G. N. Filatova, and T. D.Khokhlova. Mol. Mat. 10:125 (1998). and 67. M. S. Dresselhaus, G. Dresselhous, and P. C. Eklund, in Science Carbon ~anotubes,Academic Press, San Diego, 1995, p. 63. 68. M. Boumahraz, V. Ya. Davydov, and A.V.Kiselev. Chromatographia 15:751 (1982). y Proceedings the Advances in Li4uid 69. V. Ya. Davydov, in C h r o ~ ~ t o g r a p h'84, Chro~atography,Szeged, Hungary (H. Kalasz and L. S. Ettre, eds.), Akademiai Kiado, Budapest, 1986, pp. 351-373. 70. L. R. Snyder and J. Kirkland. ~ntrod~ction to ~ o ~ e Liquid r n Cl~~o~ato~ru~hy, 2nd ed., Wiley-Interscience, New York, 1979. 71. V. A.Davankov,A. S. Bochkov,A.A.Kurganov, P. Rourneliotis, and K. K. Unger. Chrornatographia 13:677 (1980). 72. M. Verzele, G. Simoens, and F. Van Darnme. Chromatographia 23:292 (1987). 73. N. Sheppard and J. W. Ward. J. Catalysis 15:50 (1969). 74. B. A. Morrow and N. Sheppard. Proc. Roy. Soc. A. 311:391 (1969). 75. N. Sheppard and C. De La Cruz in Advances in Catalysis (D. D. Eley, W. Haag, B. Gates, and H. Knozinger, eds), Academic Press, San Diego, Vol. 41, 1996, pp. 1-112; Vol. 42, 1998, pp. 181-312. 76. V. Ya.Davydov, in~ x p e r i ~ e n t~u l e t h ino Adsorption ~ ~ and C h r o ~ ~ ~ t o ~ r(Yu. aphy Nikitin and R. S. Petrova,eds.),MoscowUniversity,Moscow,1990, 193-211.

77. V. Ya. Davydov. J. Chromatogr. 356: 123 (1986). 78. V. Ya. Davydov, M. Elizalde Gonzalez, and A. V. Kiselev.J. Chromatogr. 248:49 (1982). 79. Ya. Davydov and G. N. Filatova. VestnikMosk.Univ.(Russian)Series2, Chemistry 37:128 (1996). C. Vidal-Madjar, Sebille, and J. C. Diez-Masa. J. Chromatogr. A. 80. R. J. 730:2139 (1996). 81. V. Ya. Davydov,in Adsor~tionon and ~ o d ~ rnorganic e d or bent^^ (A. DabrowskiandV. A. Tertykh, eds.)StudiesinSurfaceScience and Catalysis, Vol. 99, Elsevier, Amsterdam, 1996, pp. 899-914. 82. Ya. Davydov, 0. M.Knipovich,M. V. Korobov, N. B. Tamm,and G. N. Filatova. Vestnik Mosk. Univ. (Russian) Series 2, Chemistry 38:238 (1997). 83. C. J. Welch and W. H. Pirkle. J. Chromatogr. 609239 (1992). 84. Kimata, K. Hosoya, T. Araki, and N. Tanaka. J. Org. Chern. 58:282 (1993). H. Francis. 85. C. E. Kibbey, M. R. Savina, B. Parseghina, M. E. Meyerhoff, and Anal. Chem. 65:3717 (1993). R. Ruoff, D. C. Lorents, and R. Malhotra, in F ~ l l ~ r e ~Recent ~es. 86. D. S. Advances in the Chemistry and Physics Fullerenes and R e l ~ ~ e d ~ a t e(K. ~ i M. als Kadish and R. Ruoff, eds.), The Electrochemical Soc., ~ e n n i ~ g t oN. n , J., 1994, pp. 191-199. 87. W. A. Scrivens and M. Tour. Org. Chem. 575032 (1992). Ya. Davydov, L. Feltl, G. N. Filatova, T. D. Khokhlova, Pacakova, and 88. Stulik, in Fuller~nes.Recent Advances in the C h e ~ i s t r yand Ph-vsics ~ ~ l l e r e ~ e s elated ater rials (K. M. Kadish and R. S. Ruoff, eds), The Electrochemical Soc., ~ennington, 1997, Vol. 4, p. 567. 89. G. E. Maciel and D. W. Sindorf. J. Am. Chem. Soc. 102:7606 (1980). 90. D. W. Sindorf and G. E. Maciel. J. Am. Chem. Soc. 103:4263 (1981). ion ~ o d ~ e d ~ o r ~ e n (A. ts Ya. Davydov,in ~ ~ ~ ~ o r ~ tNew 91. Dabrowski and V. Tertykh, eds.),StudiesinSurfaceScience and Catalysis, Vol. 99, Elsevier, Amsterdam, 1996, pp. 673-703.

Centre de Recherche sur la Matikre Universith d'Orlhans, Orleans, France

Tntroduction TI. Curvature and Confinement A. A pinch of stereology and mathematical morphology connectivity and Curvature 128 in silicas From C. two-dimensiollal three-di~ensional to ~dsorption,disorder, and c o n f i ~ ~ m e n t Adsorption, disorder, curvatureand 111. Roughness, Scale Invariance, and Scale invariance: self-similarit Thenanoscaletexture of silicas and their surfaceroughness Experinlental methods and 154 evidence Conclusion eferences

119 123 123 134 139 141 143 146 150 163 163

This chapter is devoted to a simple question: does the of the surface of silica influenceits adsorption properties? A this is a general questionwhich. is not limited to silicas. Any porous medium be described as interface exploring the embedding th.ree-dimensional space partitioning it into two subspaces: matter and void. The morphological complexity of that interface, from subnanometric to macroscopic scale, is virtually unlimited. It is the result of all the elemen"Current ~ ~ ~ l i a t i oLaboratoire n: de Physico-Chirnie, Structurale et ~a~rornoleculaire, Ecole SupQieure de Physique et Chirnie Industrielles, Paris, France.

tary processes involved in the material synthesis. Considering the variety of processes involved in amorphous silica synthesis (polymerization, growth, aggregation, sintering, and drying, to quote a few), some of them occurring in thermodynamic equilibrium conditions and some in out-of-equilibrium conditions, it is no wonder that Silicaland offers a particularly rich variety of landscapes, being rivaled in that only by porous carbons. This morphological complexity is usually classified in three categories corresponding to three structural levels. At atomic and molecular level,when interpreting the molecular size-dependenceof monolayer or multilayer adsorption, it is common to think in terms of At larger scales, when discussing adsorption capacities, capillary condensation,or adsorption~esorptionhysteresis, the concepts invoked are those of and At even larger scales, when considering adsorption kinetics or hysteresis phenomena again, the notion of may be introduced.Although this clear and simple hierarchical classification may be relevant (Fig. la), it is in no way guaranteed that it applies to all situations. Even simple concepts such as pore and pore rnay be difficult to apply in disordered porous materials. Silica aerogels, for instance (Fig. lb) [l], are probably an extreme case of such a difficulty, buttheproblem is raised for all silicas outsidethe crystalline ones. Similarly, distinguishingthepocketsdue to surfaceroughnessfromthose whichrnaybe called "pores" may become a purely semantic matter in some cases (Fig. IC) [2]. Fortunately, in the overwhelrning majority of cases, the problem of describing a complex interface rnay be put more generally by introducing parameters or properties suchascurvature,correlation length, or scale invariance, whichwillbe developed further in the following. ince the purpose of this chapter is to discuss shape and nlorphology,it is naturalto refer to ~nfortunately, images ofsilica suitable for quantitative analysis, at the right l ~ a g n i ~ c a t i ofor n discussing adsorption or condensation, are extremelydifficult to obtain." Even more difficult (actually, impose, for the time being) is to obtain three-dimensional (3D) representations. No imaging tech-nique (magnetic resonance microimaging, x-ray microtomograetc.) so far reaches the necessary resolution. Most of the time, the nanoscale and mesoscale morphological information we have corning, via models, from scattering tech-niques or from the adsorption data themselves. In spite of this, we stantial part of this chapter to the as-quantitative-as-possible mages or 3D structures. The reason is that, with the development of computer modeling methods of all kinds, the trend is clearly towards the comparison of experimentaladsorption data withcomputeddataobtained by MonteCarloormoleculardynamicssimulations in modelporousstructures morphologically as similar as possible to the real material (Fig. 2). Thus, faithful representation of the real material willbe essential, as willbe the quantitative deter~inationof the important morphological parameters. "This is not the case for other mineral or biological porous media like rocks, or bones, where the interest shifts towards mechanical or hydrodynamic transport properties and where the relevant length scales are well above the micrometric level.

(a) 2D section through a hypothetical porous material where the notions of surface roughness, pore size, pore shape, and pore network would have obvious meanings. sketch of the structure of a silica aerogel. (Adapted from Ref. What is the meaning of pore size and pore shape in such a case? (c) 2D section through a porous calcium silicate hydrate. (From Ref. 2.) What is the difference between surface roughness and pore size distributio~in such a case? Before ill~~strating this approach, we will spend some time recalling some basic morphological notions and their relation with the surface morphology of one or another familyofsilicas. Although fractal concepts have beenextensivelyused since the early 1980s to analyze adsorption on silicas and other porous solids, they will not represent the main focus of this chapter." We deliberately chose to diversify our tools in a pragmatic spirit. We also gave up covering the immense literature where adsorption data were correlated in some way or another to

sdace energy

chemistry molecular dynamics

The scheme which is likely to become the general scheme for studying the adsorption properties of disordered porous materials in the near future. Realistic 2D and representations of the material at nano- and mesoscales are the cornerstones of the process.

morpbological feature (pore size, fractal dimension, etc.). If an emphasis of some kind has to be put forward in this chapter, it is definitely on methodology and generic phenomena.

Curvature is a key parameter in any growth, condensation, or coalescence process in solution or in the gas phase, since it controls, in association with the surface energy and via the Kelvin equation, thesolubility of the vapor pressureof a solid or liquid surfaceinequilibriumconditions,* In its simplestform,fora spherical meniscus, a particle or a droplet of radius r, the Kelvin equation reads [4] R T ln(a/ao)

2

where ais the activity in the gasor solution phase, y the surface energyof the solid, and its molar volume; a0 is the activity over a flat surface, R is the universal gas constant for anideal gas, and Tis the absolute temperature. Incrystalline particles, the surface energy dependson the typeof crystal face and this leads to well-defined equilibrium shapes.In amorphous materials, may be assumed to be constant, like in liquids. Among other things, the Kelvin equation predicts that small particles will have a larger solubility and vapor pressure than large particles, which is the basis for the well-known Ostwald ripening processin which small particles tend to disappear to feed the growth of the larger ones. This is potentially an intrinsic part of any amorphous silica synthesis, whether in solution for gels and precipitated silicas or in the gas phase for high-temperature silicas. However, silicas are not simple collections of independent “elementary” spherical particles. They are, at least, connected and partially sintered assemblies of spheres with a broad distribution of convex or concave local curvatures, which means that mass transfer during synthesis has taken a much more subtle form than in the above description. In addition, this distribution of local curvatures introduces strong local bias in multilayer adsorption and condensation. Furthermore, curvature may be the signature of connectivity, which is importa~t for diffusion. It is therefore interesting to spend some time examining the ways curvature and connectivity may be defined and measured, locally and globally, and then to look at the situation insilicas.

A any metric (obtained from a measure) or topological (obtained from a numbering) parameter is useful to describe a 3D structure or a 2D image. The problem is different if we are interested in identifying those parameters which are suitable to establish the links between 2D images and the real 3D structure, whichis the aim of stereology. ath he ma tical morphology permits us to identify the global parameters which are suitable for that. In order tobe suitable, they have to *The Kelvin equation is usually considered for the vapor pressure above a liquid liquid droplet, but it applies also to solid surfaces.

meniscus or a

satisfy a number of criteria, the so-called Hadwiger conditions: (l) they have t o be invariant by translation or by rotation (not to depend on the position and axis orientation in the material); (2) they have to satisfy the homogeneity condition (measurements performed at different magni~cationsshould yield the same result); (3) they have to be continuous functions (a small change in the size of the piece of material analyzed should notgive a large fluctuation of the parameter); and finally (4), they have to satisfy the additivity condition (which speaks for itself). Four parameters satisfy theaboveconditions:the porosity, the specific surfacearea, the integral of mean curvature, and the integral of Gaussian curvature. The three first are metric parameters. They can be measured on 2D sections of the porous material. The integral of Gaussian curvature is of a different nature. It is a topological parameter, which can only be determined on a 3D representation of the interface.

Porosity may be defined either as the ratio of void volume over total sample volume, Vv or, alternatively, as the void volume per unit mass, Vm,which is the common practice in physical chemistry. Interestingly, provided the sample is isotropic, may be measured directly on sections through the sample: it is the ratio of void area over total area. Even more simply, it may be obtained from the average length of the line segments (“chords”)on the image.A chord is defined as a line segment which belongs totally to the phase of interest (void or

Chords are obtained by tracing random straight lines (“rays”) through an image or a 3D structure. There are “pore chords” and “mass chords.” The average chordlengthpermits us to calculatetheporosity and thespecificsurface area, whereasthechordlength distribution characterizesthetype of disorderinthe solid. (From Ref. 28.)

matter) and which has bothits ends on the interface. Chords may be generated by tracing rays through the 3D structure or the image (Fig. 3). If (Zm) the average chord lengthin silica and ( l p )the average chord lengthin the porespace, then vv

( Z ~ ) / ( ( l ~( )L ) )

The surfacearea, Sv is the interface areaperunittotalvolume. Alternatively, according to common practice in physical chemistry, the specific surface area mayalso be defined per unit mass,Sm.S, may also be calculated on sections, from the average chord lengths: S,

4/(Vp)

The integral mean curvature, Adv, and the integral Gaussian cursature, Gv. The intersection of an interface with plane containing the normal to that interface generatescurvature lines (Fig. 421). Atanyplace on theabcissa S, there are two planes suchthat the curvature radii and R2(s)are maxi mu^ and mini mu^, respectively. The mean curvature is defined

1 pG(S>

1/R2(S)l

(4)

whereas the total or Gaussian curvature is defined [l/R~(S)~[l/R2(S)I l/[R1(s)Rz(s)I The curvatureradii are countedpositive if they are within the mediumof interest (pore space for instance) and negative otherwise. Thus, in the case ofnecks

Minimum and maximum curvature arcs of a surface. The portion of surface shown has positive mean and total (Gaussian) curvatures. (b) A saddle-like surface, with negative Gaussian curvature.

between partially sintered particles, with asaddle-like surface at the junction, the Gaussian curvature is always negative (Fig. 4b). The integral of mean curvature per unit volume, of a volume V of a biphasic porous materialis calculated by integration of the mean curvature over all the interface elements contained in this volume:

It may be shown that this integral is equal to

kfv

27tNA

where iVA is the so-called per unit area on the 2D image. Thus, the integral of mean curvature mayalso be determined on a 2D section of the material. Similarly, the integral Gaussian curvature perunit volume, Gv, is obtained by integration of and is straightforwardly related to the connectivity number of the interface per unit volume in 3D, (8) F

In order to understand what the connectivity numbers NA and iVv are, it is necessary tointroduceanothernotion:the G, of a surface. G is the number ofways a surface may be cut by a simple curve without losing its connectivity. For instance, in 3D, the genus a sphereor anyspheroidal surface is zero, whereas the genusof a torus is 1 (Fig. On a 2D image, the genusof a circle is zero. If the image of a section of unit area of the porous materials contains independent(disconnected)spots Si belonging to the pore space,

On a spheroidal surface, one single cutby a curve enough to separate the surface in two pieces. Its genus is zero. On a torus, one cut is not enough. Two are needed to obtain two pieces. Its genus is 1. How many cuts are necessary to separate the intricate solid network of Fig. l b or the multiply connected porous network of Fig. 16 into two pieces?

the connectivity number of the interface on this image, NA, is given by j

1

In the case of more-or-less circular spots, N A is just the number j of independent spots. Asimilar equation holds for but t h e j surfaces to consider are now developing in 3D space In real porous materials like a silica xerogel or aerogel, the pore-matter interface is generally avery complex structure. Inspite of this, porosity, specific surface area, and integral of mean curvature are still easy to determine, since they are obtained from simple measurements on random 2D sections (provided well-resolved images of such sections may be obtained). This is no longer true for the integral of Gaussian curvature and the connectivity numberin 3D, which cannot be obtained from metric parameters on 2 0 sections. The only way to obtain them is to reconstruct the interrace in 3D, for instance, by performing serial sections as close as possible to each other and piling them up, as has been done on several natural media such as embryo organs [7], soils [8], or waste-water flocks [9], for instance. This is not a purely gratuitous game.If a realistic simulation of capillary condensation or mercury intrusion in a porous silica ever has to be done, it will necessarily require such a reconstruction. In porous silicas, due to the small scale of the pores most of the time, it wouldrepresentextremely difficult and tediouswork. Fortunately, as we shall see later, numerical simulation methods may provide an interesting alternative. Assuming that a 3D representation of the interface has been obtained, its analysis may be made much easier by working on its graph. The graph is obtained by progressively narrowing the pore space, starting from theinterface, up to the point where a section through the void space is reduced to a single point, acording to a process reminiscent of arteriosclerosis (in mathematical morphology, this is called skeletization). The graphis a set of vertices connected by edges which have a vertex on each end (Fig. 6). It is characterized by a set of numbers, the so-called Betti which is the number numbers [lo].The most important is the first Betti number, of loops. The first Betti number may be obtained from the following formula: E l , where E and V are the number of edges and vertices, respectively. An interesting property of the first Betti number of an object is that it is equal to that of its complementary space. Thus, /?,(P) for the percolating pore space of a pieceofsilica aerogel or a particle ofsilicagelis equal to Pl(S)for the solid network. Once the graph is obtained, it is quite easy to determine its genus G and connectivity number by applying a numerical exploration algorithm. There is a problem remaining though: thefinite size of the sample. Indeed, any reconstruction is necessarily performed over a limited portion of the material and loops may be missed. It is nevertheless possibleto determine the maximum and minimum values of G by connecting all the pores arriving to the cell surface to a single vertex or by stopping them at the cell surface, respectively (Fig. 7).

Sketch of a 3D porous medium reconstructed from a series of 2D sections (left) and its graph (right). (Adapted from Ref. 83.)

It is not easy to obtain sharp contrastimages of the curvature profile of an “internal” silica surface. This is definitely not possible so far at the atomicscale, since the atomic force microscope is not able to penetrate into the porosity. It is easier at larger scales, either by scanning (SEM)or transmission (TEMI) electron microct visualization of silica particles has the advantage of providing images which can reveal someof the internal structureof particles close to their surface. However, due to their ill-defined nature (neither really 3 0 nor strictly 2 0 cuts) suchimagesare not suitable for quantitative analysis. we

Method to calculate the maximum (left) and minimum (right) valueof the genus of a cell of porous material. In the first case, all the pores arriving at the cell boundary are connected to one single vertex, whereas in the second case, they are stopped at the cell boundary. (Adapted from Ref. 84.)

pointed out above, a quantitative analysis of curvature and connectivity can only be performed on 2D sections or on 31) reconstructions. 2D sections are usually obtained by embedding the silica in a resin and slicing it with an ultramicrotome [12]. The smallest thickness attainable is of the orderof 60 nm, which is still large as compared to the pore size of many silicas. Such slices are suitable for TEM study but, due totheir finite thickness, the information obtainedis not strictly 2D. To our best knowledge, no 3D reconstruction from a set of 21) sections has everbeen performed on silica. Figures 8 and 9 illustrate the mesoscale morphology of a typical porous silica gel particle for chromatography. At thescale of these images, the presence of negative mean and Gaussian curvaturesis quite obvious (peanut-shapedunits). Such saddlelike surfaces (Fig. 4b), which have a negative Gaussian curvature, are the natural consequence of the partial sintering of particles with a positive curvature (Fig. 4b). They are the necessary condition for building a continuous network. A good illustration of this, which may eventually serve as a limiting model for the structure of this type of silica is the so-called cubic bicontinuous phase or, better, the sponge phase of surfactant molecules (Fig. 10) [13,14].* In the perfect cubic phase, which may be used as a model for partially sintered spheres, each phase (each side of the interface) is of equal volumefraction (50%) and the integral of Gaussian curvature is zero. The structureis said to be bicontinuous because eachside of the interface is continuous throughout the whole system. The interface is a succession of sphere portions and saddlesof equal curvature radii (in absolute value). When disorder is introduced by thermal ~uctuationsthe structure retains its high connectivity (the topology of the graph is basically unchanged) but the local size of the blobs of phase l and phase 2 are no longer equal. In silicas, without deliberate ternplating methods 1151, such a highly symmetrical structurehasnoreasonto be formed.Nevertheless, since partial sintering of particles? is an intrinsic step in many synthesis methods, in solution or in the gas phase, one may expect a somewhat similar, though not so well balanced, sequence of negative and positive curvatures of the surface. The overall symmetry has no reason to be that of a cubiclattice, since it is determined by the presinterillg packing of the silica particles, which is itself determined either by the rules of random space filling in the more dense systems or by the process of diffusion limited aggregation of particles. In the first case, one expects random (glassy) close-packed arrangements (Fig. 11) 116-181, whereas in thesecond case, fractal structures oflower average coordination are expected (Fig. 12) [19,20]. This will be developed further in Section 111. Whensuchmore-or-lessconnected assemblies aresubmitted to sintering, the contact regions between the spheres gain negative local curvatures and strings of peanut-shaped units are generated. The connectivity numbers, genus,

*The analogy should be kept at a purely geometrical level and should not be pushed too far, since the driving forces for generating this type of molecular assembly (interface elasticity, vanishing surface tension) are very different from those producing a connected network of partially sintered elementary silica particles (finite surface tension, diffusion). ?At this stage, we will not define further what a “particle” is. we shall see later, a particle may itself be a complex assembly of smaller units.

SEM micrographs of silica gel particle (LS 60, Merck) surface and subsurface, at two different l~agni~cations, showing that the texture of the particle is that of partially sintered network of blobs with alternating positive and negative curvatures.

ow-magnification TEM micrograph of ultrathin section of the same sample as inFig. 8. Note thedifferenceinscale,inspiteofthemorphological similarity, which means that the connected blobs seen in Fig.8 are in fact connected assemblies of smaller blobs. (Courtesy S. Bonnamy [l l].)

(a) The cubic bicontinuous phaseof surfactant solutions, witha vanishing integral of Gaussian curvature; (b) the so-called sponge phase, built according to the same principle, but with some disorder due to fluctuations.

A random close packing of nlonosized hard spheres. The average coordination number of the particlesis high and the porosity is low (0.38). (From Ref. 18.)

and first Betti number of these networks of strings are the direct reflection of the parent packing. Random close-packed structures generate dense networksof tight loops with a rather well-defined average curvature, whereas the more open fractal networks generate curvaturesat all scales, providing collapsingof the structure can be avoided. From that pointof view, sintering by surface diffusion preserves much more the initial topology of the medium than sintering byviscousflow,which induces drastic shrinkage (Fig. 13) [21,22]. Another interesting case of curvature, akin to the bicontinuous sponge phase of surfactant solutions described above,is the structure obtained by spinodal decomposition of melts. Phase separationby the spinodal mechanism involves the growth of concentration fluctuations leading finally to two separated continuous phases, as long as the volume fraction of each phase is not below the percolation threshold [23]. By dissolving selectively one of the phases, ahighly connected porous network may be obtained. Porous Vycor silica (trademark Corning) a typical example 11241. It is obtained by spinodal decomposition of Bz03-Si02 glasses, followed by acid leaching of the boron-rich phase. The residual solid is a monolithic porous silica (porous glass rods, for instance) with a porosity of the order of 0.30 and a typical pore “radius” of 3.5 nm, as inferred from analysis of the nitrogen desorption data in terms of cylindrical pores, using the Kelvin equation [25-271. A transmission electron micrograph of an ultrathin section of Vycor silica is shown in Fig. 14 [27]. In spite of the relatively low contrast of the images, theexistence of both positive and negative curvatures (in 2D) is clearly vislble.

The 2D projection of a fractal particle madeby diffusion-limited aggregation of 6000 monosized hard spheres. The averagecoordination is much lower than in Fig. 10. The porosity higher and, as in all fractal structures, depends on thesize of the system. (From Ref. 21.)

Once such imagesare digitized and binarized (pixels are either in the amorphous silica lattice or in the voidspace), several calculations canbe performed, as outlined in Section ITA,and compared to experiments. first parameter easy to calculate is porosity, Vv, which isjust the number of void pixelsover thetotal numberof pixels intheimage. expectedfrom stereology, thesurfacefraction of porespace measured on the digitized images (0.31) corresponds closely to the porosity derived from adsorption data at saturation (0.30) [27]. A second parameter is the average pore size. was pointed out in the Introduction, defining thesize of a pore is not obvious in disordered porous solids. In adsorption experiments, the pore sizeis usually inferred from the Kelvin equation assuming cylindrical a shape, unless some specific information on the pore shapeis available from other sources of information (from the crystal structure for instance, as in exfoliated graphite or clays) [29,30]. On the other hand, on a binary image, an average or most probable pore

Evolution of the structure of a 2D fractal aggregate upon sintering by surface diffusion (top, from left to right) or by viscous flow (bottom, from right to left). (Adapted from Ref. 22.)

“size” may be calculated from the pore chord length distribution. shown in Fig. 15, in spite of the complex morphology of the void space, the most probable pore chord length is readily obtained by image analysis (7.5 nm) and corresponds closely to the average pore diameter derived from nitrogen desorption data (7 nm) [27]. Finally, the specific surface area per unit volume,S,, may also be computed on the images using (3) for instance and converted to the specific surface area per unit mass, Sm,knowingthedensity of amorphous silica. Interestingly, thevalue obtained (77 m2/g) was significantly lower than thenitrogen BET surfacearea (103 m2/g) [27]. discussedbyLevitz et al. [27], this islikely to reflect the existenceof somesurfaceroughness, sensible to theadsorbingmoleculesbut smoothed out in the images. We will come back to this in Section 111.

Analysis of the2D images has obviouslimitations as faras adsorption is concerned. Since a direct 3D reconstruction of the porous silica from a large number of images of ultrathin sections would be extremely tediousand of the utmost difficulty, alternative techniques were developed, which reconstruct a model medium from one single random 2D section. Since, strictly speaking, the Gauss curvature and connectivity can only be measured in 3D, all these methods necessarily rely on some

silica glass; graph in

Transmission electron micrographof an ultrathin sector of Vycor 7930 superposition of the digitized pore network on the original micro(From Ref. 27.)

i

10

20

30

40

50

pore diameter; chord length [nm]

Dotted line:porediameter distribution of porous Vycorsilicaglass obtainedfromnitrogen desorption usingKelvin equation for cylindricalpores. Continuous line: pore chord length distribution calculated from the digitized image of Fig. 14b. (Adapted from Ref. 27.)

approximations and assumptions, the most important one being that the mediumis statistically isotropic. This is quite reasonable for amorphous silicas, but is not guaranteed for other materials, such as rocks. The porous mediumis defined by phase function is one inside the pore network and zero in the solid phase. In such a description, the porosity is nothing else than the average value of Vv

(10)

where represents a volume average. The morphology of the medium is characterized by the two-point bulk autocorrelation function, CV(u), which gives the probability of being in the pore space at position U, knowing that we are in the pore space at

u)l[Z(r)l)

(1 1)

Since the medium is assumed to be isotropic, vectors and U may be replaced by their scalar norms, r and Volume averages may be replaced by surface averages and Cv(u) may be calculated on the binary image of a thin section. The general strategy to reconstructa 3D model of themedium is to produceacorrelated random Gaussian field which is separated into two subspacesby applying a threshold in such a way as to obtain theright autocorrelation function. This method can

be applied either in a discrete manner on a lattice of sufficiently small cell size and the porous medium is then composed of small cubes [31,32], or in an off-lattice manner, using continuous functionals [33], The latter method may be advantageous in several cases, for instance, when large systems have to be handled or when a realistic distribution of surface chemical groups is desired (see Section 11.E). It should be noted that using the bulk autocorrelation as a central criterion of validity implies a good correspondence between the experimental small-angle xray or neutron-scattering spectrum of the real material and the computed spectrum of the reconstruction, since the small-anglescattering intensity is directly related to the Fourier transform of the density fluctuation autocorrelation function, [28]:

where @(U) is the density distribution. Thescattered intensity is thus related to the second derivative of the bulk autocorrelation function. In a biphasic matrix with a sharp interface, all these density fluctuations are located at the interface. key question in any reconstruction procedure is: How large? More precisely, the problem is to know what is the minimumsize of the mediumto treatin order to capture all its essential properties. This is often referred to as the problem of definition of theelementaryrepresentativevolume (ERV). For instance, it is quite obvious by looking at Fig. 14 that it would be a waste of effort (or at least a luxury of effort) to treat the whole image. portion of it would certainly be enough to measuretheaveragechord lengths, to haveagoodsampling of all curvatures, and to correctly determine the connectivity on the reconstructed medium, but a portion of what size? There is no general answer to this question but there are cases where a safe answer canbe given. Figure 14 is precisely one of them, The scattering spectrum is not a monotonous curve but exhibits a well-defined peak for a scattering angle corresponding to a distance of 27 nm [28]. Thus, inspite of the amorphous structureof the matrix and the disorderof the pore network, the system is in fact strongly correlated. Knowing that you are on the interface at some point of co-ordinate there is a fair probability to be again on the interface at x This correlation length may be considered as the size of a kind of “unit cell” or elementary representative volume, equalto the average pore chord length plus the averagesilica matrix chord length. In this case, the values of these parameters are approximately 7 and 17 nm, respectively. Statistically, the whole picture in Fig. 14is butthereproduction, by translationoperations,of this elementary representative volume.Its statistical properties contain all the useful textural information about the material (porosity, average pore size, specific surface area, etc.) and a 311) reconstruction of the size of a few unit cells would be a safe choice, However, the situationis not always so simple. Many porous materials exhibit a continuous hierarchyof textural features (Fig. ICis an example) or, on the contrary, a discontinuous hierarchy of textures, which leads us to consider much larger pieces ofmaterial or to takeaverages at a given scalebefore movingup to the next one. 3D reconstructionsfrom 2D sections have not yetbeenachieved far for porous silica gels, mainly due to the difficulty of obtaining good-contrast TEM

micrographs of thin sections of these materials with a resolution significantly better than the size of the pores, but they havebeen done for sedimentary siliceous rocks (sandstones and for Vycor silica glass An off-lattice reconstruction of the latter material is shown in Fig. 16. Also shown in Fig. 17 is a cross-section of this structure which, when visually compared with the actual digitized version of the TEM micrograph, evidences a striking analogy but also some faint differences with the real medium. There is little doubt that this type of approach will provide a powerful toolto model further adsorption, condensation, transport, and separation inporous silicas. Inparticular,molecularmodeling of adsorptioninsuch 3D structures takes automatically into account the structural disorder and the distribution of local interface morphologies (cf. Section 1I.E).

Three-dimensional off-lattice reconstruction of porous Vycor silica glass

17 Comparison between a cross-section in the reconstruction of porous Vycor silica (left)and the digitized versionof a TEM thin section (right). (From Ref. 33.) I

When molecules are in contact with a porous silica or withanyotherporous material, they are submitted to two effects. The first is an attraction to the wall or, more specifically, to some sites on the wall. This leads to adsorption, strictly speaking. The secondeffect simply that they are in finite box or, in other words, that they are confined. It is usually considered that this leads to bulk first-order phase transition, which is capillary condensation. In some cases, in very narrow pores called micropores according to IUPAC nomenclature, the wall and the box effect are mixed and adsorption takes some special form which, in phenomenological terms, appears as adsorption at “vanishingly OW" relative vapor pressures, i.e., as vertical step along the ordinateaxis in the adsorption isothermwhen the amount adsorbed is plotted versus PIPo It is not easy to separate clearly the effect of confinement from that of adsorption. A clear reason is that there are always some attractive interactions with the solid surface, at least the London dispersionforces. Another reasonis that thesilica surface, and that of many other finely divided or porous materials, is most of the time mor~hologicallydisordered at the mesoscale. In those conditions, considering simple pore shapes such cylinders or spheres with a well-defined diameter is generallyacrudeapproximation.Thevariousparameters related to curvature, chord length, and correlation briefly introduced in Sections and may be useful tools to go beyond such a description by enabling us to consider complex shapes with a distribution of local situations. In anycase, whatever the descriptive tools, what is real though is that molecules in the void space of a silica are in a confined environment with variable wall-to-wall (chord) distances. discussed by Evans the questions raised by this situation are of the followingtype: To what

extent can notionsof bulk phase equilibriabe employed in confined systems? Down to which wall-to-wall distance is the concept of a meniscus valid? Do the condensation and freezing transitions persist in narrow pores? Are therewell-defined critical points for a confined fluid? What is the microscopic structure of a confined gas or liquid? How does it depend on the morphology (curvature) of the porewalls and on their separation? And finally, at a larger scale, does the behavior of molecules in a single pore have any relevance for gas adsorption on real porous solids containing network of interconnected pores? So far, the most detailed information on the effect of confinement in disordered media comes from molecular simulation studies, the only ones able to take into account the effect of morphological disorder of the walls. Analytical theories are restricted to simple shapes, but important phenomena have nevertheless been predicted [36,3’7]. For instance, the critical temperature for thevapor-liquid transition in a slit-shaped pore oran infinite cylinder with no external surfaceis lower than in the bulk. Below this critical capillary temperature adsorption is characterized by a disco~tinuity,the capillary condensation. Furthermore, adsorption and desorption branches do not coincide, so that a hysteresis loop is predicted. However, this hysteresis is not related to any pore blockingin constrictions or “necks”, since the theory considers a single and morphologically simple pore. Nor is it related to variations in wetting conditions (contact angle). Actually, during adsorption, at pressures higher than the hysteresis closure point, the adsorbed film remains in a metastable state, being unable to overcome the free energy barrier for condensation. The system does not have fluctuations of sufficient amplitude to find the equilibrium state [38]. On the contrary, during evaporation, the fluid always at equilibrium. At temperatures higher than the discontinuity and the hysteresis loop disappear and the adsorption/desorption isotherm exhibit monotonic evolution. The smaller the cylinder diameter, the lower the critical temperature (for a given adsorbate). For diameterssmaller than 1.5 nm, atransition is nolonger predicted.Recentexperimentsonthe ~ C ~ 4 l - t y pmesoporous e silicas (which are so far the only amorphous mesoporous silicas with a simple pore geometry) have confirmed the previous scheme [39]. Matrixdisorderadds some significant new features to these confi~ement effects.” For instance, inMonte Carlo studiesof the phase diagramof an off-lattice molecular modelof a Lennard-Jones fluid confined in a rigid matrix of spheres with parameters representative of a silica xerogel, disorder was shown to induce a much smaller density change during condensation than in an ordered system [40]. The hystereticbehavior is recovered,butthe vapor-liquid coexistence curveinthe adsorption branch is further narrowed as a consequenceof both the wetting behavior of the fluid in the matrix and the disorder in the matrix (Fig.18). In addition to a main vapor-liquid transition analogous to capillary condensation, the simulations provide evidence for a second transition associated with the wetting behavior of the fluid in the more dense regions of the matrix. In such conditions, the coexisting phases exhibit high degree of spatial inhomogeneity and disorder, as illu*The purpose this chapter is not to cover extensively the molecular simulation literature. broader survey simulation studies can be in Chapter 6 this book.

much

f0

Typical adsorption~esorptionisotherms for a 12-6 Lennard-Jones fluid in a disordered sphere matrix. The dotted lines mark the limitsof the hysteresis loop. The existence of a pre-condensation transition is clearly visible on the adsorption branch. (Adapted from Ref. 40.)

strated in Fig. 19. It should be noted that the strength of the fluid-matrix interaction plays an importantrole in these phenomena. Itis only in the weak interaction limit that they are observed. Strong interactions (those corresponding to methane in a silica gel, for instance) may create field suf~ciently strong to suppress all fluid-phasetransitionsinthe void space, leavingonlyacontinuousadsorption process. This may provide the link to the more classical picture of adsorption in rnicropores.

was pointed out in the Introduction, all the morphological information about porosity, pore size distribution and pore space connectivity is contained in the way the solid-pore interface explores the embedding space. In this respect, the local curvature of that interface is a key parameter in adsorption phenomena, since it determines the integral of all the interactions between the adsorbate molecule and the atoms of the solid in its environment (Fig. 20). In addition, positive mean curvatures (Fig. 20a) generate pockets in which confinement effects may appear. Thus, the influence of local curvature is expected to be particularly sensitive in physical adsorption.

Snapshots of Monte Carlo simulations for vaporand liquid states from the adsorption isotherm of Fig. 18, closeto coexistence. Thesnapshot on the left is close to the saturated vapor state and the one on the right is close to the saturated liquid state. (From Ref. 40.)

A remarkable illustration of the influence ofcurvature on the adsorption properties of disordered silica adsorbent is provided by the recent grand canonicalMonte Carlo (CCMC) simulation of adsorption of xenon in a Vycor-like porous silica matrix by Pellenq et al. [41]. The matrix was obtained by applying the off-lattice functional used to create Fig. 16 [35] to a box containing thesilicon and oxygen of 15 l 5 15 unit cells of cristoballite. This allowed portions of the cristoballite lattice to be cut out and a set of morphologically and topologically equivalent

Sketch of three possible situations from an adsorbate molecule, with different mean, and Gauss, curvatures: in (a), 0 and 0; in c 0 and 0; in (c), H($) 0 and 0 (curvatures are counted positive when the center of curvature is in pore space). The first situation is obviously the most favorable to adsorbate-adsorbate interactions.

numerical 3D porous samples tobe generated, asillustrated in Fig. 21 1421. In order to reduce the computationalcost, a homothetic contraction of a factor of approximately 2.5 of the simulation. box dimensions, which preserves the morphology and topology of the pore network, was applied. This pseudo-Vycor porous silica exhibits the expected correlation peak,shifted by a factor of 2.5 with respect to the real material. The average density and porosity of the numerical samples were 1.49 g/ cm3 and 0.32, respectively, in good agreement with real values (1.50 g/cm3 and 0.30). A chemically realistic surfacewasthenobtained by removing all silicon atoms in an incomplete tetrahedral environment and saturating all oxygen dangling bonds with hydrogen. Each hydrogen atom was placed in the pore void in the direction perpendicular to the interface at 0.1 nm from the closest unsaturated oxygen. The potential function used in the adsorption simulation was similar to the potentialused to model xenon adsorptionin silicalite, a purely siliceous crystalline zeolite 1431. It was the sum of a dispersion interaction term with a repulsive short-range contribution and an induction term. The independent variables were, as usualin GCMC simulations of adsorption, the temperature(195 K), the pressure of the bulk gas in equilibrium with the solid, and the volumeof the simulation cell containing the porous solid. The adsorption isotherm was obtained by calculating the ensemble average the number of adsorbate molecules. The main point of interest as far as interface morphology concerned is the mechanism of adsorption and pore filling, as illustrated in a series of snapshots in Fig. 22. At 195 K, xenononly partially wets the silica surface. In parallel to adsorption of isolated molecules on the surface, adsorption proceeds by formation of condensate pockets in the places of highest concave surface curvature. This leads to an unexpected situation where parts of the pores are filled with liquid condensate, while other regions remain uncoveredor only partially covered with an adsorbate film. Therefore,monolayercapacity-basedmethods(suchasthe approach) to determine thespecific surface area cannotbe used in such nonwetting situations.* Finally, capillary condensationoccurs in thewholeremainingvoid space. The shape of the adsorption isotherms reflects this gradual condensation process. In contrast to what is obtained in a simple infinite cylindrical mesopore [37], the slope of the isotherm has a finite value. Rather than being associated with a distribution of pore sizes, it is merely the consequence the curvature fluctuations of the interface.

Very few solid surface are molecularly flat over large areas. Mica and graphite are, among others, noticeable exceptions which makes them so valuable for the studyof surface forces /45] or the phase diagram of two-dimensional adsorbed films [46]. Most of the time, if we except the low-index faces of singlecrystals, there is evidence that the surface of solids exhibits roughness at length scales larger than a few not the case for other adsorbates such as krypton or nitrogen, which wet the silica surface

interatomic distances. This roughnessmay be a purelyinterfacial property orit may be the boundary manifestation of the complex internal texture of the solid. It may come from the growth of the solid, from fracture or wear, or from chemicalreaction with the environment, including dissolution. It may appear at all scales, on macroscopic pieces of material or on the surface of finely divided particles. Many of the rough surfaces generated by the above processesexhibit structural features at many different scales (Fig. 23 [47]; Fig. 24 1481). Interestingly, in spite of their randomly disordered appearance, there aresimple statistical rules relating the number and size of the features at one length scale to the number and size of the features at anotherlength scale. Furthermore, these relationships appear tobe scale independent, within certain limits. This scale i ~ v a ~ i a is~ coften e described in terms of fractal concepts 149-511. Before examining to what extent the surface roughness of silica particles is of this type and bow it influences adsorption, we shall briefly recall some definitions and properties about scale invariance.

Scale invariance may be qualitatively defined as the general property of an object which keeps the same morphology (i.e., looks the same) under a change of scale. The simplest form of scale invariance, self-similarity, may be rationalized by considering scale changes such that

Diffusion front in 3D of a material in gray, into another material B (not shown), frombottom to top. separation surface very rough and self-similar. In addition, there are “islands”of in B and “lakes” left behindby B in (From Ref. 47

The surfaceof a 2D deposit of circular particles with uniformly distributed diameters which arrive ona surface and roll to a position in which they touchat least two other particles. Note the disordered roughness of the surface, with features of many different sizes, that this simple process is able to generate. (From Ref. 48.)

Thus, if a scale change of a factor h is applied to a length element A, of the object along the axis, a statistically similar object is obtained provided the same factor is applied to the elements Ay and A, along the y and axes, respectively (Fig. 25). S t ~ ~ i ~ tsimilarity ic~Z means that it is only by averaging for different A, elements that scale invariance is obtained. For instance, asimple similarity operation is zooming with a camera. The calcium silicate interface in Fig. IC and thediffusion front in Fig. 23 are examples of self-similar surfaces. Self-affinity is a generalizationof self-similarity, by allowing fordifferent scaling factors along the variousaxes. It may be defined by considering the followingscale changes:

Scaling an element A, by a factorh, along x will not change the morphologyof the object provided the scaling factors alongy and are h, and h,, respectively, which are different from h, (Fig. 26). Thus, in order to keep the morphology statistically

A regular (or deterministic) self-similar profile.

unchanged, it is necessary to deform the object it is zoomed (i.e., using different zooming factors along and z). Most of the time, for surfaces, only one direction is different from the others, usually thedirection perpendicular to the average plane of the surface. For instance, for a fracture surface, the directions in the fracture propagation plane are usually equivalent, so that h, h, h. In order to obtain a group structure for the mathematical transformations of Eqs. (14), it is necessary to relate the scaling factors by homogeneous functions

(a) A regular (or deterministic) self-affine profile; dom) self-affine profile.

a statistical (or ran-

where H is the self-affinity or Hurst exponent. In self-similarity, all scaling factors are identical and H 1. In thecase of self-affinitywith only onescaling factor (h,)different from the two others,the self-affinity propertymay be expressed byOne single relationship. Whatever the direction in the plane perpendicular to (Az(hA,))

This equation permits us to estimate the statistical evolution. oftheroughness amplitude Az corresponding to a “horizontal” length elementAx,when the length of is changed:

(&(Ax))

(17 )

In other words, the roughness amplitudeis a power law function of the distance in the average surface plane. Interestingly, the “apparent roughness” lAz/Axl), i.e., the ratio of the roughness amplitudeto the “horizontal” distance overwhich it was estimated, also follows a power law: (18) Thus, according to Eq. the apparent roughness of a self-affine surface with H 1 tends to zero as its lateral extension tends to infinity. A self-affine surface is asymptotically flat. The self-affine model of surfaces has been remarkably well validated for natural and laboratory fracture surfaces overseveral orders of magnitude In particular, the surfaceof cracks in soda-lime-silica glass has been shown to be self-affine from a few angstroms to more than 100 nm [54]. The use of atomic force microscopy (AF~), which allows for thedirect determination of the surface topography, has been of invaluable help in that respect. Self-affine surfaces are characterized by their Hurst or roughness exponent, H , which takes values between zero and one. The closer to one H is, the smoother (correlated) the surface appears. Thecloser to zero, the rougher (anticorrelated) the surface looks. What about self-similar objects? Their Hurst exponent is trivial (H and doesnot tell anything about their shape and their self-similarity. The usual way to characterize them is to look at how the number N of boxes necessary to cover them (or a part of them) is changing when a scale factor h is applied either to the box size or to the size of the part of the object to be covered. Thus, if the box size is changed by a scale factor h, is expected to change as

h-” where L), the fractal dimension, is usually a noninteger. On the other hand, if the region of the object to be covered is changed by a scale factor A, is expected to change as NNA”

9b)

so that, in general o(

(~l/h)~

(20)

For a surface, D 3. The larger D,the more convoluted the surfaceis. It may be shown that self-affine surfaces are characterized by a fractal dimension equal to (3 [49].

Why is the surface of silica expected to be rough at the molecular scale? As briefly mentioned in the introduction to this section, the main reason is that many of the elementary physical processes involved in silica synthesis (diffusion, dissolution, precipitation, polymerization, aggregation, drying and shrinkage, etc.) may often lead to scale-invariant structuresat scales which go down to the size of the elementary structural unit involved. Thus, if the elementary unit involved in a polymerization process is a monomeric silicic acid unit, roughness may eventually appear down to the scale of a single silica tetrahedron. On the other hand, if the process which operates the agglomerationof already-formed larger particles, the boundary of that agglomeratemay be expected to be rough atscales larger than the size of the particles (this does notpreclude the surface of those particles frombeing themselves rough at a smaller scale). Synthetic amorphous silicas are all prepared by hydrolysis of a monomeric precursor (sodium silicate, silicon tetrachloride, tetramethoxysilane,hexamethyldisiloxane, etc.) and condensation of the hydrolysis product. They may be classified into three categories, according to classical nomenclature*: 1.

are prepared by hydrolysis and condensation (polymerization) of sodium silicate in acidic conditions. This leads first to a hydrogel which is either washed and dried, or first aged at elevated temperature and then dried. Theheattreatment induces the coalescence of the hydrogel particles. If drying is performed in subcritical conditions (xerogels), the capillary forces induce shrinkage and densification of the assemblies, but they may also lead to flaws if the stresses become too large. If drying is performed in hypercritical conditions (aerogels), capillary forces are avoided, leading to more open textures. Other precursors such as tetraet~oxysilane (TEOS) or tetramethoxysilane (TMOS), for instance, may follow the same general route as sodium silicate. Hydrolysis and polymerization may also be conducted in emulsions of aqueousdroplets in an immiscible liquid.This allows us toprepare agglomerates of the same size as the emulsion droplets, which may be further sintered according to a hierarchical process. The xerogels shown in Figs. 8 and 9 are typical examples. Another subfamily which should be added here is the recently developed micellar template silicas obtained by performing hydrolysis andcondensationina liquid crystalline solution of surfactant molecules 151. are also prepared from aqueous sodium silicate solutions, but in alkalineconditions.This leads to aggregates of partiallysintered

*For an extensive description of synthesis methods and a detailed discussion of subcategories, the reader is referred to the classical references or more recent summaries

micrograph of precipitated silica aggregates on a carbon grid and a typical silhouette.

spheroidal particles (Fig. 27) which can be reasonably well described in terms of self-similarity, like the aggregate shown in Fig. 12 [58]. are obtained by high-temperature decomposition of a silicon halide or another precursor such hexamethyldisiloxane as (HMDS). Recompositionmay be achievedina flame (fumed silicas), in an arc or a plasma discharge, or in a high-temperature vapor. These processes generate very spherical particles which, while sticking to each other, may remain well individuali~ed(arc, plasma, or vapor silicas, Fig. 28) or which may lose their individuality and form deeply sintered chainlike fractal aggregates (fumedsilicas, Fig. 29). Aerosil from Degussa, Cabosil from Cabot Corporation, orHDK silica from Waclser Chemie are the archetypal examples of the latter type. Transmissionelectronmicroscopy [l 1,601 provides evidence that all finely divided silicas, including those which are considered as nonporous, such as the high-temperature flame, arc, or plasma silicas, consist of dense packings of very small primary units of diameter around 1 nm. An exampleis shown in Fig. 30 for a xerogel, but the same type of nanotexture has been evidenced in flame or plasma silicas [59]. Most of the time, those primary particles do not make up the surface

micrograph of a thin section through a resin-embedded low surface area (20 m2/g) silica fume forming in the high-temperature vapor phase of ferrosilicon alloy furnaces. Note the quasi-independence of the silica spheres. (From Ref. 85.) area for adsorption. They form compact and sintered packings with a high coordination (like in Fig. so that the porosity (if any) of such packings not accessible to gas molecules. In high-t~~perature precipitatioll or silicas, the spherical particles or the spheroidal unitsin the chainlike aggregates,which are the most visible features in TEM (Figs. 27-29), would have such a structure.In gels, depending on the details of the condensation process during synthesis (aging, heating), the primary units mayformsecondarydense particles of intermediate size,which

micrograph of a flame silica aggregate (Aerosil 200) and a typical silhouette.

thennselves form larger packings with poresdirectly related to thepacking geometry of thesecondary particles (Fig. 31). The well-defined pore size distributions achieved by synthetic-silica manufacturers,withaveragepore sizes going frorn less than 10 nrn to more several hundreds of nm, is the best proof that this secondary coalescence process is well mastered. However, careful examination of the images (Fig. 30) shows that this coalescence scheme should merely be considered as a guide for thought and that reality is probably more disordered than what our schemes suggest. Comingback tothequestion of thepossible origin of roughness on amorphous silica surfaces, it clear that, whatever the size of the secondary particles, the existence of the primary condensation structural units is a possible source of roughnessassoon asrandomdeposition, diffusion-limited aggre~ationor dissolution-precipitatioll processesareinvolved [50,51]. Ontheotherhand,the growth of roughness is opposed by onestrongthermodynamic force: surface tension. Inthat respect, thehigh-temperature silicas areexpected to havea much smoother surface than the silicas prepared at low temperature. Similarly, the gels which have been aged in hydrothermal conditions are expected to have lowersurfaceroughness than those which have been driedwithoutaging treatment.

High-magnification TEM micrograph of an ultrathin section of the same silica gel as in Figs. 8 and 9. This material hasa sharppore size distribution, obtained by nitrogen adsorption and water thermoporometry, centered at 6 nm in diameter (the thickness of the section used to take this micrograph is more than ten times the average pore diameter). (From Ref. 11

Direct measurements of the roughness offinely divided silica particles even more, of the wall surface of pores in porous silicas are far less obvious than for macroscopic surfaces like that of a piece of glass. Most of the information on this roughness comes from adsorption or scattering experiments. Scattering (light, x-rays, neutrons) is a powerful method to probe the surface roughness of finely divided or porous solids and, more generally, their interface morphology, includingscale invariance properties[28,61-631. In a scattering experiment, the scattered intensity is measured as a function of the modulus of the scattering vector, q: (43t/h)sin(6j2)

(21)

where 6 is the scattering angle and h the wavelength of the incident radiation. Generallyspeaking, scattering probesthedensity fluctuations in themedium which, in a biphasic medium, are locatedat the interface. However, several regions may be identified in the scattering spectrum of a porous or finely divided material. In the very large q domain (domain of interatomic distances), the eventually spec-

Primary particles

Drying

nm)

Secondary particles

aging Hydrothermal

Hydrogel

Sketch of the coalescence process used pore sizes. (Adapted from Barby, Ref. 60.)

Xerogel

to prepare gelswithwell-defined

trum contains the Bragg peaks due to diffraction by the lattice if the material is crystalline. In the large domain (small distances), the spectrum probes thewall of the individual pores or of the surface of the individual particles composing the medium (Fig. 32). This is the so-called Porod domain in which, if the surface is smooth, the Porod law is observed:

If the surface is self-similarly rough,

where D is the fractal dimensionof the surface. Equation (23) may be considered as a generalization of Eq. (22). If the surface is self-affine, one has

Thus,roughsurfaces are generallyexpected to yield exponents between -3 and At smaller (larger distances) the scattering spectrum starts probing correlations of the interface morphology as scales corresponding to the distribution of pores or particles in space (Fig. 32). Thismayeventually also display scale

Cartoon illustrating the different scales probedin a scattering experiment. For small values of the reciprocal scattering vector modulus, l/g (large scattering angles), the radiation probes the surface of particles or the wall of pores. For large values of l/y, the scattered radiation probes longer-range correlations between particles or pores.

invariance, in which case, if this distribution is statistically self-similar, one expects

where D is now the fractal dimension of the mass distribution, which should be smaller than 3. On the other hand,if there is well-defined correlation lengthin this distribution, asin Vycor silica glass, the spectrum exhibits a correlation peakwhich reflects the pseudoperiodical distribution of pores or particles. Surfaceroughnessmay also be probed by adsorption of molecules. It is intuitively obvious that small molecules cover roughsurfacemore closely than larger ones, which leave cavities under them. Surprisingly, it is only little less than twentyyearsago that this simple idea, which is intuitively quite obvious, was put on a quantitative basis by D. Avnir and P. Pfeiferby considering the case of self-similar roughness [64] and applying it to many different materials The number N J r ) of molecules ofsize r needed to cover (or to “tile”) se~f-si~ilar surface of fractal d i ~ e n s i o nD by monolayer is expected to scale as

which is equivalent to Eq. (l9a) and which, for practical analysis of adsorption data, is converted in terms of the molecular area to

Experimentally, Nm is themonolayercoverageobtained fromthe Emmet-Teller (BET) equation.* Thus,by plotting monolayer content versus molecular area in logarithmic coordinates, one expects a straight line of slope -r>/2. Several conditions have, nevertheless, to be fulfilled for this. In particular, the shape of the molecules hasto be ideally spherical and the fractal exponent should notbe too high, since self-similar surfaces of high fractal dimensions generate constrictions which prevent the molecules from approaching the surface [73]. The expected behavior islesssimple on self-affine surfaces, since the relative roughness of a self-affine surface depends on the scale at which it is observed [Eq. (18)]. Indeed, simple tiling simulations show that log N,%versus log plots are no longer linear (Fig. 33). It is only in the asymptotic limit of small molecules that the expected exponent is recovered. What are the results of scattering and tiling experiments on silicas?? The general trend from scattering experiments is the one which is expected: high temperature silicas are those which have the smoothest surface. The slope of the log (scattered intensity) versus log plots is always close to -4, as expected from Porod law for a smooth surface with D 2. Thus, Hurd et al. found a slopeof -4 for Cabosil M5 samples [74]. Legrand et al. obtained -3.9 (D 2.1) for Aerosil200 [SS]. However, one Cabosil sample (EH-5) yielded a slope of -3.46 (D 2.54), which suggests a rather rough surface [75]. Interestingly, heating this sample wellbelow its glass transition temperature (250 K below Tg)led, after 3 h, to D 2. This is a beautiful example of the action of surface tension on a softened material, but the relatively low temperature at which this smoothing takes place is also an indication of the extremely small size and fragile nature of the surface asperities which form the surface roughness (single tetrahedra or very short chains?). Even xerogels for chromatographywere found tohave smooth or nearly smooth surfaces. Schmidt et al. investigated x-ray and neutron small-anglescattering from a series of silica gels for chromatography with average pore radiiof 6, 20, 50, 100, and 250 mm, respectively (Si-60, Si-200, Si-500, Si-1000, and Si-2500, respectively) [76] (Fig. 34). With only one exception (Si-60), the slopes were between -3.9 and -3.8 (D between 2.1 and 2.2) for distances the order 1 to 10 nm. For xerogels from another manufacturer (Si-40, Si-60, Si-100, and Si-4000), Drake et al. measured slopesof -4 (D 2) for distances of the order of to nm (Fig. 35) with the exception of Si-100, which showed departure from Porod behavior at length scales smaller than 1 nm [77,78]. It is interesting to note that the latter results could be put on one single master curve by rescaling the scattering curves by a factor directly related to the average pore size (Fig. 3 9 , showing that this family of samples is

*This implies that the BET equation is still valid on self-similar surfaces. Although this is definitely not the case if one considers the entire isotherm from 0 to close to it is reasonably valid in the monolayer region [66-681. On the other hand, several models of multilayer adsorption on self-similar surfaces have been developed outside the framework of the BET theory, based either on hierarchical capillary condensation 168,693 or on FHH-type theories TA detailed review of literature data may be found in Ref. 3. Once more, we will only insist here on methodology and main trends.

Va

2

r

me

l~orphologicallyhomogeneous, the only difference being the size of the building blocks. Porous silica glasses turn out to exhibit also some weak departure from Porod behavior. Thus, a slope of -3.8 (D 2.2) was obtained by Hohr et al. [79] for Controlled Pore Glass (CPG) samples with nominal pore sizes going from 7.5 to 200 nm, in a range of length scales going from about l to 7 nm for the small pores samples to more than 30 nm for the large pores samples (Fig. 34). Similar results (slope -3.7, D 2.3) were found for porous Vycor silica glass [27, 801. If a general picture hasto be extracted from thescattering results on amorphous silica surfaces, it is that of smooth or moderately rough surfaces, at length scales going approximately from one to at least a few nm. An illuminating sample case, which helps us to decide what this “moderate roughness” could represent, is given by the analysis of Pellenq et al. on adsorption and scattering by porous Vycor silica [41]. In order to accelerate their simulation runs referred to in Section Pellenq et al.used a grid-interpolation procedure,in which the simulation box volume (Fig. 21) was split into a collection of lo6 cubic voxels of size close to 0.1 nm. The Xesilica adsorption potential energy was calculated at the corners of each elementary cube. A grid of lo6 energy points was obtainedin that way for each of the numerical silica samples. cut through such a grid is showni n Fig. 36. The interest of such a representation is that it shows the pore wall surface as it is “seen” by the xenon atoms. This surfaceis clearly rough at the molecular scale and its roughness comes merely from “cutting” the periodic atomic lattice of the matrix at random angles. Such grids can also be used to obtain theporosity and specific surface area by considering the negativeenergy grid voxels. Whilst the porosity is not seriously affected by taking this roughness into account, the surface area is now in good agreement with nitrogen BET values (the value calculated on the smooth TEM images used to reconstruct the 3D simulationcells was too small, see Section 1I.B). The mostinteresting point is that the decayof the simulatedscattered intensity is characterized by a power law with exponent equalto -3.6, in good agreement with experiments (experimental value:-3.7). Thus everything happens as if the pore walls of this material were fractal, with a surface fractal dimension of 2.4. What the results of Pellenq et al. actually show is that cutting an ordered atomic lattice with a highly curved but smooth interface, generated by a thermodynamic process (spinodal decomposition), enough to generate a surface with a roughness amplitude of the order of l nm, which presents all the symptoms of a morphologically scale-invariant surface. Cutting a vitreous lattice instead of a crystalline lattice would probably make only marginal differences.

(a) Monolayer coverage of circles on a self-affine curve of fractal dimension L) 1.2 and roughness exponent H 0.8. (b) log (monolayer content vs. log (radius)plotsforthreedifferentroughnessexponents, 0.8, H 0.5 and 0.2, from bottom to top, respectively. The expected slopes (-D -1.2, -1.5, and -1.8, respectively) are only recovered in the small “molecular” size limit. The curve appears smooth (-D 1) for “molecules” largerthan afinite value, which is the maximum roughness amplitude of the curve segment which is probed.

1

loe 1O’O

r

t

IO8

Id

lo4

0

Small-angle x-ray scattering curves for (a) a series of porous silica glasses withdifferentporesizes,Si-60,Si-200,Si-500,Si-1000, and Si-2500,from top to bottom. (From Ref. 76.) A series of CPG-10 porous silica glasses with different nominal pore sizes,7 3 , 17, 50, 70, and 200 nm, from top to bottom. (FromRef. 79.)

The analysis of adsorption data in terms of scale-invariant roughness is much more difficult than that of scattering data. The reason is that the shape of the adsorbate molecules [73] and the nature of their interactions with the surface [78] may seriously perturb the purely geometrical aspect of tiling. In addition, roughness necessarily introduces strong local curvatures which, as seen in Section II.E, may induce phase transitions which, in turn, pervert the meaningof the monolayer coverage values derived from the isotherms. Finally, the scale range in which power laws of thetype of Eq. (27)between and G areobserved is usually very restricted, typically from 0.3 or 0.4 nm (the size of a water nitrogen molecule) to 0.7 nm (the sizeof a butane molecule) [65,81]. Using concepts such as selfsimilarity or self-affinity in suchasmalldomain is at the limitof geometrical owever, being aware of such limitations, fractaldimensions may nevertheless be used as operational parameters to quantify surface roughness

lo6 lo6

lo4

l

(a) Small-angle x-ray scattering curves for xerogel samples Si-40, Si-60, and Si-100. Rescaled scattering curves for Si-40, Si-60, and Si-4000. (Both parts from Refs. 77 and 78.)

1821. Their main valueis to put comparativestudies of the interactionof the various members of silica family with series of molecules on a rational (though complex) basis. Rather than entering into the highly debated interpretation of the wealth of fractaldimensionsobtainedfromadsorptionexperiments on silicas andother adsorbents [Sl], we would like to close this chapter by discussing what is probably the only exampleof a casein which the adsorption data couldbe directly related to the visualization of the interface morphology. Once more, it is the case of porous Vycor silica, which so far unique case because it is the only silica on which direct information on the surface morphology has been obtained (271. Figure 3’7 illustrates the evolution of the specific surface area (monolayer coverage molecular cross-section) as the molecular cross-section area increases (from nitrogen to butane). Although thelog-log plot does notexhibit the linearity expected for a self-

A cutthrough. an energygriddescribingthe porous Vycor silica glass. (From Ref. 41.)

adsorption ofxenon on

2.2

Specific surfaceareavs.moleculararea log-log plotforadsorption of nitrogen, methane, ethane, propane, and butane on porous Vycor silica glass. The evolution towards the asymptotic value expected for a smooth surface, and effectively measuredon the TEM images, is typical of self-affine surfaces.(From Ref. 27.)

similarly rough surface, the evolution clearly points to a significant roughness, in agreementwiththe scattering dataand withthesimulation results (Fig. 36). Interestingly, thecurvereaches“asymptotically” (for butane)thegeometrical valuecalculatedfromachord analysis onthe digitized TEM micrographs (Section 1I.E). Thus, with molecules of the size (and the interactions) of butane, the roughness of the silica surface has been completely smoothed out. What this suggests is that, if a fractal model has to be chosen for the pore wall roughness of Vycor silica (Fig. 36), a self-affine model would probably be more realistic than a self-similar model.

Adsorption on amorphous silicas is definitely not two-dimensional physics. Even when the surface is smooth at themolecular scale, curvature or confinement introduce specific effects. The direction for progress in our understanding of the interplay between surface morphology and adsorption is clear. Whenever it is possible, an integrated approach should be used, combining quantitative image analysis on thin sections and 3D reconstructions with molecular modeling and experimental study of the adsorption process. Ideally, this morphological approach should integrate as much information (either computational or experimental) as possible on the surface chemistry and the surface roughness.

I am immensely indebted to Pierre Levitz and Roland Pellenq. This chapter, especially Section 11, is largely based on their own thoughts. Sylvie Bonnamy is also gratefully acknowledged for her invaluable helpin the TEM investigation of silica gels. Finally, I would like to thank Isabelle Cousin for allowing meso kindly to use several figures of her thesis.

1. A. Bourret, Europhys. Lett. 6731 (1988). 2. Van Damme. Ciments, Bktons, Pliitres et Cbaux 15:782 (1990). 3. F. Ehrburger, in The Surface Properties Silicas (A. P. Legrand, ed.), Wiley, Chichester, 1998, pp. 83-143. W. Adamson, Physical C~emistry Surfaces, 5th ed., Wiley Interscience, New 4. York, 1990, Chapters 111, VII, and IX. 5. Serra, Image Analysis and ~ a t h e m a t i c ~ l ~ o r p ~ oAcademic l o g y , Press, London, 1982. 6. M. Coster and L. Chermant, Pre‘cis d’Analyse d’lmages, Presses du CNRS,Paris, 1989, pp. 13-35. 7. H. M. Gerard, P. Bouchet,P.Jacquemin, P. Lelinaire, and J. L.Mallet.Acta Sterol 187 (1992). 8. I. Cousin, P. Levitz, and Bruand. Eur. Sci. 47:439 (1996).

9. F. Zartarian, C. Mustin, G. Villemin, T. Ait-Ettager, A. Thill, J. Y. Bottero, J. L, Mallet, and D. Snidaro, Langmuir 23:35 (1997). 10. L. K. Barrett and C. Yust. Metallography 3: (1970). 11. A. Raoof, thesis, Universitk de Marne la Vallke, France (1997). 12. Ehret, L. Crouisier, and J. P. Ebherhart. J. Non-Cryst. Solids 86:72 (1986). 13, Ryde, Andersson, K. Larsson, Blum, T. Landh, S. Lidin,and B.W. Ninham, The ~anguage Shape, Elsevier 1997. Porte, in Soft ~ f f t t ePhysics r (C.Williams and M. Daoud, eds.),SpringerVerlag, Berlin, 1999, pp. 15. C. Kresge, M. E. Leonowicz, W. J. Roth, J. C. Vartuli, and J. Beck. Nature 359:710 (1992). 16. B. J. Alder and T. E. Wainwright. J. Chem. Phys. 33:1439 (1960). 17. 1. Snook, W. Van Megen, and P. Pusey, Phys. Rev. A 43:6900 (1991). 18. P. Levitz, Phys. Chem. 97:3813 (1993). 19. Meakin, Adv. Colloid Interface Sei. 28:249 (1988). 20. P. Meakin,in The Fractal Ap~roach to ~ e t e r o g e n e o ~ sC ~ e m i s t ~ y :S u r f a ~ e ~ s Chichester, 1989, pp. 131-160. Colloids, P o l y ~ e r s Wiley, , 21. R. Julien and R. Botet, in Aggregation and Fractal Aggregates, World Scientific, Singapore, 1987. 22. N. Olivi-Tran, R. Thouy, and R. Jullien, J. Phys. I, France 6557 (1996). 23. J. W. Cahn. J. Chem. Phys. 42:93 (1964). 24. H. P. Hood and M. E. Nordberg, US.Patent 2,106,744: H. P. Hood and M. E. Nordberg, U S . Patent No. 2,286,275. 25. V. Enustun and M. Enuysal. J. Pure Appl. Sci. 3:81 (1970). 26. G. Mason. Proc. R. Soc. London, Ser. A 415:453 (1988). 27. P. Levitz, Ehret, S. K. Sinha, and J. M. Drake. J. Chem. Phys. 95:6151 (1991). 28. P. Levitz and Tchoubar. J. Physique I 2:771 (1992). 29. J. Gregg and K. W. Sing, ~ d s o r ~ t i o n , S u ~Area ~ a c e Porosity, Academic Press, London, 1982. 30. F. Rouquerol, Rouquerol, and Sing, Adsorption by Powdersand Porous Solids, Academic Press, London, 1999. 31. J. A. Quiblier. J. Colloid Interface Sci. 98:84 (1984). 32. P. Adler, Porous Media: Geometry T r a ~ s ~ o r tButte~orth-Heinemann, s, Stoneham, 1992, pp. 503-538. 33. P. Levitz. Adv. Colloid Interface Sci. 76-77:71 (1998). 34. P. M.Adler, C. C. Jacquin, and J. A. Quiblier. Int. J. Multiphase Flow 26:691 (1990). 35. P.Levitz,in Dynamics in Small C o n ~ n ~ nSystems, g MaterialResearchSociety, Boston, in press. 36. R. Evans. J. Phys.: Condens. Matter 2:8989 (1990). 37. P. C. Ball and R. Evans. Langrnuir 5:714 (1989). 38. F. Celestini. Phys. Lett. A, 228:84 (1997). 39. K. Morishige, H. Fujii, M. Uga, and D. Kinakawa. Langmuir 23:3494 (1997). 40. K. Page and P. A. Monson. Phys. Rev. E 54: 6557 (1996). 41. R. J. M.Pellenq, S. Rodts, V. Pasquier, Delville, and P. Levitz. Adsorption, in press 999).

42. R. J. M. Pellenq and P. Levitz, unpublished. 43. R. J. M. Pellenq and D. Nicholson. J. Phys. Chern. 98:13339 (1994). 44. R. J. Pellenq, private communication. 45 N. Israelachvili, ~ntermolecularand Surface Forces, 2nd ed., Academic Press, London, 1992, pp. 168-174. 46. Y. Larher. J. Chem. Soc. Faraday Trans. I 70:320 (1974). Gouyet, in The Fractal A p ~ ~ o ~toc h 47. B. Sapoval, M. Rosso, and J.F. ~ e t e r o s e n e o ~Chemistry: s Surfaces, Colloids, Polymers, Wiley, Chichester, 1989, pp. 227-245. 48. Csahok and T. Vicsek. Phys. Rev. A 464577 (1992). 49. B. Mandelbrot, The Fractal GeometryofNature, Freeman, SanFrancisco, 1983. 50. Avnir,ed., The Fractal A~proach to ~eteroseneousChemistry: Colloids, Polymers, Wiley, Chichester, 1989. 51. A.-L. Barabasi and H. E. Stanley, Fractal Concepts in Surface Growth, Cambridge University Press, Cambridge, 1995. 52. E. Bouchaud, G. Lapasset, and J. Plan& Europhys. Lett. 13:73 (1990). 53. K. J. Maloy, A. Hansen, E. L. Hinrichsen, and S. Roux. Phys. Rev. Lett. 68:213 (1 992). 54. F. Creuzet, and E. Bouchaud. Phys. Rev. Lett. 74: (1998). 55. R. K. Iler, The Chemistry o ~ s i l i c aJohn , Wiley, New York, 1979. 56. K. K.Unger, Porous Silica, Elsevier, 1979. 57. A. P. Legrand, in The Surface Properties of Silicas (A. P. Legrand, ed.), Wiley, Chichester, 1998, pp. 5-14. 58. A. P. Legrand, H. Hommel, A. Tuel, A. Vidal,H. Balard, E. Papirer, P. Levitz,M. Czernichowski, R. Erre, H. VanDamme, J. P. Callas, J.F. Hemidy, J. C. Lavalley, Barres, A. Burneau, and Y. Grillet. Adv. Colloid Interf. Sci. 33:91 (1990). 59. F. Rouquerol, J. Rouquerol, and E;. Sing, in Adsorption by ~ o ~ d and e ~Porous s S o l i ~ sAcademic , Press, London, 1999, pp. 288-3 11, 60. D. Barby, in C~aracteri~atio~ ofpowder Surf~ces (G. D. Parfitt and K. S. W. Sing, eds.), Academic Press, London, 1976, pp. 353-425. 61. Teixeira, inOn G r o ~ t h a n ~ F o(H. r mE. Stanley and N.Ostrowski, eds.), NATO AS1 Ser., Applied Sciences, 1986, pp. 145-162. 62. A. J. Hurd, D. W. Schaefer, and A. M. Glines. J. Appl. Cryst. 21364 (1988). s 63. P. W. Schmidt, in The Fractal Approach to ~ e t e r o s e n e o ~Chemi'~try: C ~ l l o i ~Polymers s, (D. Avnir, ed.), Wiley, Chichester, 1989, pp. 67-79. 64. P. Pfeifer and D. Avnir. J. Chem. Phys. 79:3558 (1983). 65. D, Avnir, D. Farin, and P. Pfeifer. J. Chem. Phys. 79:3566 (1983). 66. J. J. Fripiat, L. Gatineau, and H. Van Damme. Langmuir 2562 (1986). 67. P. Levitz, H. Van Damme, and J. J. Fripiat. Langmuir 4781 (1988). 68. P. Pfeifer. Springer Ser. Surf. Sci. 10:263 (1988). 69. A. Neimark. Physica A 191:258 (1992). 70. M. Jaroniec. Langmuir 11:2316 (1995). 71. A. Neimark. Phys. Rev. B 50: 15435 (1995). 72. F. Ehrburger-Dolle. Langmuir 13: l 189 (1997).

73. H. VanDamme, P. Levitz, F. Bergaya, J. F. Alcover,L. Gatineau, and J.J. Fripiat. J. Chem. Phys. 85:616 (1986). 74. A. Hurd. Mater. Res. Symp. Proc. 172:3 (1990). 75. A. J. Hurd, D. Schaefer, and J. E. Martin. Phys. Rev. A 35:2361 (1987). 76. P. W. Schmidt, A. Hohr, H. B. Neumann, H. Kaiser, D. Avnir, and S. Lin. J. Chem. Phys. 941479 (1990). 77. J. M. Drake, P. Levitz, and S. Sinha. Mater. Res. Symp. Ser. 73:305 (1986). 78. J. M. Drake, P. Levitz, and J. Klafter. New J. Chern. 1477 (1990). 79. A. Hohr, H. B. Newmann, P. W. Schmidt, P. Pfeifer, and D. Avnir. Phys. Rev. B 38: 1462 988). 80. D. W. Schaefer, A.J. Hurd, andA. M. Glines, in~ a n ~ ~ u o~ t u~ a t i o ~ ~ s Growth (H. E. Stanley and N. Ostrowsky, eds.), Kluwer, Dordrecht, 1988, p. 62. 81. D. Avnir, D. Farin, and P. Pfeifer. New. J. Chem. 16:439 (1992). 82. D. Avnir, 0. Biham, D. Lidar, and 0. Malcal. Science 278:6557 (1996). 83. C. Lin and M. H. Cohen. J. Appl. Phys. 53:4152 (1982). 84. P. Kaufmann, F. Dullier, I. MacDonald, and C. Sirnpson.ActaStereol.2:145 (1 982). 85. M. Edwards-Lajnef, thesis, Universiti: d’Orleans, France, 1993.

I)epartment of Adsorption and Planar Chromatography7 Maria Curie Sklodowska University, Lublin, Poland

Introduction I. A. Porosity ofsilicagel Investigation of theporosity

167 167 168

TI. Porosity ofSilica Gels from Isobars A.Temperature-programmeddesorption

169 169

111. Porosity of Variousby Silicas T G Method Nuclear A. membranes model a as of porous solids Porosity silica gels of with bonded alkyl grafts Porositysilica C. mixed of samples I). Pore size distribution of mesoporous silica molecular sieves 193 E. Porosity of thermally treated gels silica IV. Conclusions References

182 182 186 192 96

200 l

In contrast to crystalline particles of neutral silica, amorphous silica is characterized by its porosity, introducinga large surfacesarea inside its particles. Amorphous silica gels exhibit a large diversity in structural properties depending on their preparation. The most popular and documented is the sol-gel method of preparation of silica gels. In this process Si(OH)4 molecules condense to form a siloxane network. The precursor solutions are usually soluble alkali-metal silicates

or alkoxysilanes. Hydrolysis of those components and condensation of Si(OH)4 entities may be controlled in various ways, producing silica with variety of structural characteristics [1-10]. A lot of work has been directed toward the modelingof silica texture during gelling and drying processes. Polycondensation occursbetween oligomeric silicaspecies, resulting in larger particles (colloidal, aggregates, networks,sediments).Large particles of spherical shapewhichareformedduring condensation determine the specific surface, and their packing density determines mainlythevolume andporedimensions.Thevariation of physicalproperties between various silicasis caused by changing the pH, addition of various substances as inorganic salts and surfactants to sol--gel reaction mixture, temperature, aging time, and drying process.

operties of the silica matrix, four parameters are sufficient:specific surface area, specific pore volume, pore size or pore area distributions dV(S,), dS(R,), and particle size. A wide variety of physical methods are used for estimation of the above-mentioned parameters. Theyare based on different pore models. Usually the pore shape is not known and the pores are very irregular. Hence, the choice of the model is made rather arbitrary. Two methods are most commonly used: adsorption and ion of the adsorption~esorptionisothermsthemodel of Halenda (BJH) is usually used, assuming cylindrical pores l]. Because during desorption first the core of the pore is evacuated, the problem of determination of the thickness of multilayer remaining on the walls of pores should be solved. Both the pore radiusand surface film thickness are related to the relative pressure by means of the Kelvin and the Halsey equations 12,131. The problem of the surface multilayer thickness does not exist in the case of mercury porosimetry. Pore diameters and pore volumes are calculated directly from intrusion~xtrusioncurves illustrating the dependence of the volume filled by mercury at pressure applied. The relation between pore radius and external pressure gives the Washbourn equation [12]. With adsorption and porosimetry the sample mustbe dried before analysis and an assumption should be made concerning pore shape in order to relate the applied Pr the topore size. above-mentioned methods depend on interfacial curvature effects and the movement of a phase boundary through the capillaries in the solid. Similar effects are registered calorimetrically in the case of thermoporometry for movementof the solid-liquid front when temperature decreasesbelow the freezing point of the liquid [14,15]. The depressionin the freezing temperature is related to the poredimensions and this feature can indeed be used to determine pore shape and size. Small-angle x-rayscattering (SAXS) and small-angle neutron scattering (SANS) are techniques which give us statistical information about pore structures [16-20]. The main assumption in interpreting small-angle x-ray scattering data is that the porous solid contains regions in which there occurs a constant electron density mixed withregions in which theelectrondensity is zero. Analogously,forthe

SANS method one assumes regions of finite scattering length and regions without scattering. The literature on these methods is expanding rapidly and many recent papers illustrate the utility of these techniques. The results obtained from SAXS and SANS, i.e., specific surface area and the average poresize, have been shown to be in reasonable agreement with those using a standard nitrogen adsorption technique. Moreover, small-angle scattering can be used to detect blocked or closed pores inaccessible to adsorptives. Anothergroup of methods consists of themeasurements of adsorption of macromolecules (fractions of polymers) from the liquid mixtures in dynamic conditions 121,221 (size exclusion chromatography) or in static conditions [23]. In this last case the pore structural analysis is based on the excess adsorption isotherms and the relation between extent of adsorption of different-sized molecules and the pore dimensions. Chromatography appearsto be an attractive method for porosity measurements. However, results may depend on choice of relations between solute size, pore dimensions, and elution volume. Recently, in the field of microscopy, NMR and other spectroscopic methods have been increasingly considered as alternative means of investigation (see, e.g., Refs. 24,25).Several studies employing both simulation and densityfunctional theory indicate that none of the conventional methods of pore characterization is entirely satisfactory [26,27].

Temperature-programmed desorption is a convenient method in estimation of the surface properties of solids. This technique is often employed in the investigation of gas-solid i~teractions. It provides information on surface chemistry and the state of bonded molecules, as well as the desorption kinetics and mechanisms like reaction order, activation energy, etc. Thus, this method is commonly used in the investigations of catalysts and other solids with energetically heterogeneous surfaces. The measurements are usually restricted to adsorbed mono- or multilayers. The recording of the weight loss against temperature for adsorbedspecies starting from their boiling point gives the isobars of desorption. In this case the process is equilibrated, i.e., the solid at a given temperature is in contact with the saturated vapors of adsorbate. For samples containing excess of liquid adsorbate, when the pores aretotally filled the desorptionis divided into several stages. At thefirst stage, when the temperatureis lower than theboiling pointof the liquid, evaporation only takes place. At the boiling point the excess liquid out of the solid is evaporated. Above boiling point starts the desorption from pores of different dimensions and internal surfaces of the solid. Thus the temperature desorption processanalogous to desorption under isothermal conditions when we control the presure. Temperature of evacuationfrompores is dependent on their dimensions.The lower the radii of the pores, the higher the temperature required for their evacuation. Similar relations are observed in the caseof evaporation of liquid wetting the porous solid if the rate of desorption (evaporation) is measured at a constant

temperature against time [28]. The measurements were made in open chambers under normal pressure. Aswas observedduringthe first stage of evaporation, therate is constantandcorrespondstotheevaporation of thebulk liquid. Duringthesecondstagethe liquid presentintheporesevaporates.Thethird stage of the process corresponds to the remaining molecules of liquid on the surface. In both cases the pure adsorptive is removed at a slow and changing rate and a gravimetric technique used to follow the variation of the amount condensed or adsorbed at constant temperature orwith increase in temperature. ~ e n e r a lConditions the ~ ~ ~ e r i ~ eAt nthe t beginning . of the 1990s the thermogravimetric(TG)technique was appliedfor investigation of thermal desorption of liquids from porous materials [29-341. This chapter is a review of some of the capabilities of T G analysis for characterizing mesopores using examples taken mainly from workby the author’s group onsilica gels. The object of this review is to outline the technique, which can be applied in the characterization of both the structure of porous silica and the other phenomena occurring atsilica surface. The T G experiments were performed with a Derivatograph Hungary). This apparatus possesses several heating programs which make it possible to carry out the desorption in more-or-less equilibrated conditions. The desorption curves (isobars) were measured under atmospheric pressure. The samples of porous solid were prepared by wetting with an excess of liquid and placed in platinum crucibles of conical or labyrinth type, with which the Derivatograph is also equipped [35]. The constructionof these measurement chambers guarantees the creation a selfgenerated atmosphere of saturated vapor of liquid over the sample. Temperature desorption was realized using the so-called quasi-isothermal program [36,37]. This special heating mode permits the maintenance of conditions inside the measurement chamber as close as possible to equilibrium during the transformation of the sample. In the case of quasi-isothermal programs the heating rate is not constant and changes during the experiment, If evaporation of liquid is slowthe fixed weight-loss level(in our case 0.5 mg/min)regulatingthe speedof theprogram is not exceeded. Asa result the linear increase of temperaturewithin this measuring range is realized. At a certain temperature, corresponding to the temperature of conversion, the above-mentioned level is exceeded and isothermal conditions are establishedforthetimeduring which transformationsconnectedwith weight change take placein the sample. In fact, in the case of our experiments even within thequasi-isothermalrange of measurements,thetemperature slightly increases according to the variation of the dimensions of the pores evacuated. Thus, the desorptionprocesscan be faster or slower depending on the assigned heating rates. The time required for a single measurement increases proportionally to the reduction in temperature gradient. Figure l shows the typical desorption curve for a liquid which wets the adsorbent perfectly and interacts with its surface only physically [38]. The results have beenexpressed asconventional T G curves, weight loss againsttemperature,

F

1500

ui

150

0 min

l \

Temperature, "C

Desorption curvesof n-butanol from silica gel Si-100; (b) temperature vs. time (solid line, quasi-isothermal program; dotted line, linear program). (Reprinted from Ref. 38 with permission from Elsevier Science.)

A m =f(T). In Fig. 1 the dependence of temperature against time for the same desorption processis also shown. For illustrative purpoes, the dotted curves represent the results obtained by applying dynamic, linear heating program with continuouslyincreasingtemperature. In this casethe characteristic points on the desorption curve connected with the textural properties of the solid disappear. Two main conclusions can be drawn from Fig. 1. The first, and most obvious, is that the totaltime of analysis depends on the heating mode. The secondis that the linear heating modeis inadequate in the investigations of the porosity. Segment I of this curve represents the bulk liquid outside the pore structure of the adsorbent. Intensive evaporationat this stage of the process takes placeat the boiling pointof the liquid (perpendicular segment). When thefirst stage is completed, the temperature increases at heating rate lower than that during the initial heating program and starts the desorption from pores. Segment 11 corresponds to the capillarycondensed liquid within the pores together with the adsorbed film on the walls of pores and is, therefore, measure of the total pore volume. The steepest part of the curve above the boiling point of the liquid corresponds to the desorption from pores of the greatest part of the total pore volume. The plot A m =f(T) may be converted into plot of volume loss versus pore radius R,, A using the Kelvin equation: ln-P

2Y v;, ---cos6 IiKRT

Equation (1) gives the relationship between the saturated vapor pressure of liquid over the curved (inside the pores) and the flat liquid surface, p and respectively. RK is the radius of the liquid meniscus, y is the surface tension of liquid, 6 is the contact angle between the liquid and the wall of pore (for perfect

wetting it is assumed to be zero), is the molar volume, R is the universal gas constant foran ideal gas, and Tis the absolutetemperature. When 0 0 the Kelvin radius RK is equal to the core/pore radiusRP. In T G experiments the valueof is constant and equal to external atmospheric pressure. The parameters y, and VM are functions of temperature. It follows from Eq. (1) that the sharpemptying of the pore of given diameter occurs under the experimental conditions applied when the pressure of liquid vapor becomes equal to atmospheric pressure. For large pore dimensions the evacuation of liquid from the pores occurs at temperatures only little higher than the boiling point of thebulk liquid. Whentheporedimensionsdecrease highertemperature is required for their emptying. ~ifferentiatingthe dependence the appropriate pore size distribution curves may be calculated. It is interesting to compare the desorption isobars measured with theTG technique with adsorption~esorptionisotherms for the same samplesof silica gel. For illustrative purposes T G desorption curvesof benzene forsilica gel Si-60and Si-500 are shown in Fig. Fig. shows appropriate adsorption~esorptionisotherms of nitrogen at -195°C. Figure demonstrates that in the case of silica Si-500 above 100°C the total amount of the adsorbate present in the pores and bonded physically with the surface is practically desorbed. Comparing the curves in parts and of Fig. 2 one can observe that there exists great similarity of their shapealthoughthey were measuredusingvarioustechniquesunder different experimental conditions. It is main reason for the assumption that isobars, similarly to isotherms, represent valuable data for derivation of the parameters characterizing the porosity of solids. general rule one can observe that the pores get smaller the desorption curve becomes more extended along the temperature axis. For nonporous solids, segment 11disappears. The best illustration of these effects is desorption of liquid

'

b

m n

Desorption curves of benzene from silica gels Si-500 and Si-60 (2) measured using the TG technique under quasi-isothermal conditions, (b) adsorption-desorption isotherms of nitrogen at -195°C on silica gel Si-500 and Si-60 (2). Opencircles and squares, adsorption; solidcircles and squares, desorption. (Reprinted from Ref. 38 with permission from Elsevier Science.)

for adsorbents of different mesoporosity. Figure 3 shows differential curves Am/ for the series of silica gels and nonporous silica glass (desorption of benzene) [39]. The peaks of these curves represent the mostly intensive evaporation corresponding to emptying of those pores with a greatest share of the total pore volume. Thesolid line represents the differential curve for glass, i.e., corresponds to evaporation of benzene at its boiling point. From our previous studies of desorption isobars measured with theTG technique arise that two main factors should be taken into account for obtaining meaningful results, i.e., the wetting liquid and appropriate heating program. Wetting It is necessary to choose an appropriate liquid (adsorbate) for desorption.The best results areobtained in thecase of liquids for which the dependence of theporeradius R, on temperature T,according tothe Kelvin equation, guarantees the highest accuracy. In order to demonstrate the effects of the typeof adsorbate on desorption curves and derived PSD the investigations were performed for the homologousseries of alcohols from n-propanolto n-hexanol and three members of the hydrocarbon series from n-hexaneto n-octane. Desorptionof higher members may be accompanied by loss of surface hydroxyl groups of silica gel. Alcohols were desorbed from Si-100 silica geland hydrocarbonswere desorbed from Si-60 silica gel [32]. For example, thermodesorption curves (massloss against temperature) for three systems are shown in Fig. 4. A feature of the curves for the alcohols is the relatively steep segment corresponding to the desorptionof adsor-

1 50

in

Te~perature,"C

Differential curves A m / for: nonporous glass; (2) (3) Si-60; (4) Si-40. (Reprinted from Ref. 39 with permission from Elsevier Science.)

900

750

250

160

180

70

90

600

50

Temperature,

Thermogravimetric curves of desorption for the different systems investigated: (a) n-propanol/Si-100; n-hexanol/Si-100; (c) n-hexane/Si-60. Part I of these curves correspondsto desorption of the bulk liquid. Part represents desorption of liquidsfromthepores and their internal surface.(ReprintedfromRef. 32 with permission from Elsevier Science.) bate from the pores above their normal boiling temperature; for n-hexane and for the other hydrocarbons this segment is more extended. It is understandably attractive to compare the results of the TG method with the results of the most common methods by which we characterize porous solids, i.e., physical adsorption/desorption of nitrogen and mercury porosimetry. After transformation of the TC experimental curves into the dependence of volume loss on core radius in the manner described earlier appropriate were calculated d V / d R K (see Figs. 5 and 6). In the samefigures from nitrogen method and mercury porosimetry are also shown.

'F 0.5

0.3

0.1

0.7

0.5

0.3

0.1 4

4

Radius,

5 Pore/core size distribution curvesforsilicagel Si-l 00: solidline, TG method; dotted line, nitrogen method; histogram, mercury porosimetry. (a) n-propanol; n-butanol; (c) n-pentanol; (d) a-hexanol. (Reprinted from Ref. 32 with permission from Elsevier Science.)

Theadsorption-desorptionisotherms of nitrogen at -195°Cwere measured with an automatedSorptomatic 1800 apparatus (Carlo Erba, Italy). Surface areas, SE&, were calculated from the linear form of the BET equation over the linear range of relative pressure between about 0!.5 and 0.4, taking the crosssectional area of the nitrogen molecule to be 16.2 The pore size distribution curves were calculatedfromthedesorptionisotherm by the Barrett-JoynerHalenda (BJH) method l] with corrections of the pore radii with respect to the surface film thickness t , where t is calculated using the Halsey equation:

and cr is the average thicknessof a single molecular layer of nitrogen equal to 0.354 nm. Total pore volumes were estimated from the desorption isotherms at 0.98. Mercuryporosimetryanalyses were generatedoverapressure

'P

t i

0.2

0.

0.0

0.2

0.1

0.0

0.2

0.1

0.0

Radius,

size distribution curvesforsilica gel Si-100: solidline, TG method; dotted line, nitrogen method; histogram, mercury porosimetry. (a) n-hexane; (b)n-heptane;(c)n-octane.(ReprintedfromRef. 32 withpermissionfrom Elsevier Science.)

range of 20-4000 bar using a porosimeter 4000 (Carlo Erba, Italy). Pore size dis-

P=

-2y cos

RP

cylindrical model of pores was assumed. An arbitrary mercury contact angleof 141" and a surface tension y of 0.48 N/mwere used to calculate the PSD.

As seen in the case of alcohols (Fig. 5) the location of the peak Rizk of the distribution curveis shifted towards smaller radii the length of the hydrocarbon chain increases. Correspondingly, the difference (R:& REzk) increases (see also Table 1). stated earlier after the evacuation of the core of pore at given temperature a monolayer or multilayer film of liquid adsorbate remains on the pore walls. The difference (R:& R;feCak) may be assumed to be the mean thickness of this film. The mean thickness is not constant but decreases the temperature increases. This thinningof the surfacefilm was observed in our previous experiment for benzene and n-butanol on different silica gels [30]. From Fig. 5, it is seen that the difference (Rp”ezak andthusthethickness of the film, increases asthe length of the hydrocarbon chain in thealcohol molecule increases. This effect suggests that the alcohol molecules are perpendicularly oriented on the silica surface. The formation of monolayer of alcohol molecules is possible by hydrogen bonding to thesurface hydroxyl groups, with hydrocarbon chains oriented towards the bulk liquid. This structure could produce awell-defined boundary between the surface layer and the bulk liquid filling the core of the pore. The regular changes in the film thickness indicate that at the peak of the distribution for each alcohol, a monolayer of alcohol is present on the surface. This could be accidental if we take into account that the thickness of the film depends on the temperature. During desorption of amyl alcohol and n-hexanol the temperature appears to be relatively unstable (see Fig. 4b). Temperature fluctuations for higher alcohols influence the shape of the distribution curves and consequently the location of the distribution peak In other words, the equilibrationof the desorption processis determined by both the heating mode and the character of the adsorbate for the same heating program. The difference in the location of the distribution peak in the case of alcohols is connected with this last factor. We shall now discuss the results for hydrocarbons and silica gel Si-60 (Fig. 6). Again the peaks of the distribution curves calculated from theTG experiment are located at lower radii than those from the nitrogen method. However, in this case the locationof the peaks fordifferent hydrocarbons is identical, and independent of the length of the hydrocarbon molecule (Table1). This means that the thickness of the surfacefilm for these adsorbatesis the same. Thisis plausible if we assume aflat parallel orientation of hydrocarbon molecules to the adsorbent surface. Also, in the case of hydrocarbons, where the heat of evaporation is similar and much lower than foralcohols, the fluctuations of temperature for higher members of the homologous series do not appear; the T G curve is smooth even for n-octane. Figures 5 and 6 also contain the pore size distributions obtained from mercury porosimetry. Since mercury porosimetry and the nitrogen method both lead to pore size distribution, we wouldexpect their distributioncurvesto be close together; in fact they are far apart. In Table are collected the values of total pore volumes, Vp,obtained by the different methods. For silica gel Si-100,TG and mercury-derived pore volumesare very similar (see Table Total pore volumes in the case of the TG method were calculated from theamount of liquid desorbed between its boiling temperature and the end of the process. Appropriate corrections for the changes in liquid density withtemperature were made.Simultaneously,itshould be noted that mercury

Parameters Characterizing the Porous Structure of Silica Gels Obtained by Using Various Methods

Si- 100 n-Propanol 316 322 0.92 0.89 0.98 6649 64 316 322n-Butanol 0.92 0.92 0.98 6649 63 n-Pentanol 316 322 0.92 66 0.92 0.98 316 322n-Hexanol 0.92 0.81 0.98 49 49

Si-60 27 34 n-Hexane 0.33"0.70 0.72 n-Heptane 23 28 34n-Octane

27

49

34

'Low values are probably the result pressures used. Ref. 32.

unavailability

part

the pores for mercury at the

porosimetry underestimates thespecific surface areaand the pore volume for the Si60 sample owing to the unavai~abilityof a portion of the pores to mercury. The portion of pores with R 19 A cannot be entered by mercury at 4000 bar. The nitrogen method leads to a relatively high total pore volume and pore radiusat the peak in the case of Si-100. The comparison of the surface areas and pore size distributions obtained by nitrogen adsorption~esorption and mercury porosimetrygiven in the literature for a range of porous oxides also showed a varying degreeof agreement between these two techniques [40]. c. ~ ~ Anappropriate ~ gfurnaceheating ~ program ~ should ~ be . used. Theheatingprogrammust be neither too slow nor too fast. If it is too slow, desorption willbe too slow and air can mix with the vapor. When the pressure of liquid vapors is lower than the external pressure, evaporation from all pores at each measuring temperature is possible. If the heating program is too fast the tem~eratureof the sample lags behind the measured temperature. As a result the pore radius calculated by the Kelvin equation is smaller than the actual radius. The heating modesubstantially influences the timeof the desorption process and consequently the shapeof the pore/core size distribution curve. Figure 7 shows T C curvesforn-butanol,representingthemass loss againsttemperatureandthe increase in temperature of the sample against time for two quasi-isothermal proAs seen in the case of the QI-3 program, which is characterized by a fast linear (within linear range of heating)increase in temperature (15 Kjmin)and large inflows of heat within the quasi-isothermal ranges the temperature is not stable

I

I

90

Temperature,

7 Desorption curves (weight loss against temperature and time of desorption against temperature) of n-butanol from Si-100 silica gel for different quasi-isothermal programs: QI-1 (dashed line); QI-3 (solid line). Shaded areas illustrate the time of desorption of thesame amount of adsorbate for different heating programs. (Reprinted from Ref. 32 with permission from Elsevier Science.)

and shows relatively high periodical fluctuations. The time of desorption for the same mass of sample is much shorter, 17 min in this case, in comparison with 28 min for the QI-l program (1.5 K./min-'). The conditionsof desorption connected with the heating program mean that the process is rnore-or-less equilibrated. Figure 8 shows the TG curves for desorption of n-butanol from silica gel Si-l00 at various heating programs used. The experiments were performed using three quasi-isothermal programsQI-l, QL-2, and QL-3 (with different heating coefficients 1.5, 5.5, and 15 K/min, respectively) and one dynamic program D-4 at a heating rate 3 K/min [33]. As is visible the shape of these experimental curves dependent on the heating rate applied. Step I1 of the experimental curve becomes more extended when the velocity of desorption decreases. Applying the procedure of the calculations described earlier the appropriate distributionscurves AVIAR, versus R, forvariousprograms were calculated. Figure 9 shows the obtained distribution curves. One can observe that the conditions of desorption influence the shape of these curves. As is visible, ther exist some optimal conditions for which the pore size distribution isclose to that obtained from the nitrogen method. The pore size distribution curve for the investigatedsilica gel calculated on the basis of adsorp-

110

120

Temperature, "C

Desorption curves of n-butanol from silica gel Si-l00 at different heating programs. Curves 1-4 represent programs QI-1, QI-2, 41-3, and D-4, respectively. (Reprinted from Ref. 33 with permission from Elsevier Science.)

0.0 10 Radius,

Core size distribution curves for silica gel Si-l00 calculated on the basis of TC data at different heating programs. Curves 1-4 represent programs QI-l, 41-2, QI-3, and D-4, respectively. Dashed line represents PSI3 from the nitrogen method. (Reprinted from Ref. 33 with permission from Elsevier Science.)

tion-desorption isotherms of nitrogen is represented by a dashed line. The same symmetrical distribution forSi-100 silica gel waspresented elsewhere (see, e.g., Ref. 41). Comparing the distribution curves presented in Fig. 9 one can conclude that the linear program is not suitable for our investigations. Very fast heating with continuously increasing temperature causes the time of evaporation for a given range of temperature tobe insufficient for total emptyingof the pores. The remaining part of the liquid is desorbed at hi eraturestogetherwiththe liquid evacuated from narrower pores and tceir internal surface. The curve presented in Fig. 9 shows a small peakat 45 A, which may be ascribed to pores having the \ighest share in the total pore volume, i.e., pores with a diameter equalto about 60 A. However, the temperatureregistered at this small peak is higher than the equilibrium temperature corresponding to desorption from given a group of pores. As a result the calculated distribution curve is deformed and shifted in the direction of smaller pore radii. Most of the liquid is desorbed when the temperature of the sample exceeds 140°C. Moreover, the shape of the curve is unrealistic, since this silica exhibits a monomodal pore size distribution. Much more realistic shapes of distribution curves were obtained in the case of quasi-isothermal programs. The slow heating program QI- gives 1 the distribution curve showing the shape similar to the nitrogen pore size distribution curve. However, this curve is not symmetrical and contains an elevated segment at wider pore radii. A slow heating mode causes impure vapors of adsorbate to be present over the sampleand evaporation fromall pores takes place. This is the main reason for additional weight loss of adsorbate at lower temperatures. A slightly higher heating rate (program QI-2) leads to a symmetrical curve. For these desorption conditions the surface film effect diminishes. Fast heating (program QI-3)results in an elevated segmentof the curveat higher temperatureswhich is connected, according to our earlier suggestions, with the desorption of the surface film remaining on the walls of pores at lower temperatures. Simultaneously with the changes of the shape of the distribution curves, the location of their maxima also changes. In the case of program QI-1 the pore/ core radius at the peak is close to the pore radius calculated from nitrogen data. Appropriate differences of the pore/coreradii at the peak of the distribution curves for nitrogen and TG methods are collected in column of Table 2. The shift representsthethickness of theadsorbed film at the corresponding temperature. Because usually the calculated mean pore radius is a function of themethod applied, thethickness calculatedasa difference (Rr2ak is not really valid and not very informative. This may be useful for comparative purposes, e.g., in the case of the same adsorbate and different adsorbents. Column 3 of Table 2 contains total pore volumes calculated from step I1 of the desorption curves. As can be seen, in the case of the linear heating program the pore volumeis higher than that fromnitrogen adsorption~esorption data. For quasi-isothermal programs values are similar and a little lower in comparison to values. From Table 2 it follows that there is a general decrease of values with the increase of the heating rate of QI programs. Moreover, for obtaining valuable results in T C experiments the following factors should be taken into account:

Parameters Characterizing the Pore Structure of Silica Gel Si-100 Estimated at Different Desorption Conditions

QI- 1 QI-2 41-3 0-4

0.1 0.4 0.9 2.4

0.97 0.94 0.92 1.17

1.10

Ref.

High-purity liquids should be used because impurities maysubstantially change the boiling point of a liquid and consequently cause errors in calculations of the pore radii from the Kelvin equation. Physically adsorbed water shouldbe removed from the hydroxylatedsilica surface. A measuring crucible of appropriate construction should be used. The best results are obtained by usingaconical or labyrinth type of crucible, which makes possible the formation of a self-generated atmosphere over the sample.

High-energy heavy ions produce stable defects in organic foils and in most dielectrics in the form of microchannels (tracks). Depending on the atomic numberof a bombarding particle, its energy and ion trajectory angles, one can regulate the size of the tracks and their spatial distribution. The main feature of a “nuclear membrane” structure is a high uniformity of pore size, The shape, volume, and dimensions of thetracksarethe subject of intensive theoretical and experimental investigations. The maininterest is how to determine the pore dimensions for membranes which are wet in liquid solvents, i.e., under the conditionsin which they are usually used. The organic membranesswell in contact with most solvents. Recently, thermoporometryhas been appliedin investigations of polycarbonate membranes aged in water,ethanol,anddodecane 1421. Apartfromthe swellingeffect theorganic foils bombarded by heavy ions may be considered as a model of porous solids. Poly(ethyleneterephtha1ate) (PETP) foil 0.1 mm thick was irradiatd by multiply charged I2’Xe8’ ions withan energy of 124 MeV accelerated in the U-400 cyclotron in the Laboratory of Nuclear Reactions, JINR, Dubna, Russia f431. The ion beam was swept over the surface of the foil, producing a uniform distribution of cylindrically shaped microchannels.The dose (equalto thedensity of the microchannels) was 1.3 10’ cmv2. At this dose the overlapping of channels produced by particular ions could beneglected[44]. The channel diameter was a natural one as produced by penetrating ions-no etching process was applied.

Benzene and acetone of thepurestgrade(POCh,Poland) wereused asthe wetting liquids in the TG experiment.The foil wasaged in these solvents for different periods of time: 2 h and 7 days. The rolled foil (about 0.5 g) was placed in a platinum crucible of the labyrinth type, and then an excess of the solvent was added. The sampleswere outgassed before the experimentto facilitate the penetration of pores by the solvent. In Fig. 10 the desorption curves for benzene and acetonefromthesamples of PETP foil areshown.Segment I of these curves represents the bulk liquid outside the pores. Intensive evaporation at this stage of the desorption takes place at the boiling point of the liquid (perpendicular line). The step denoted as segment IT of the desorption curve corresponds to the capillary-condensed liquid within the pores. Segment I11 represents theliquid which takes part in the swelling process and is squeezed between the polymer chains. The boiling point for each liquid, in the case of the investigated foil, is reached twice. The first boiling point corresponds to the liquid outside the poresand the second, at a higher temperature, corresponds to the liquid filling the pores. The te~perature at which the second boiling point is reached depends on the pore dimensions. When the pore dimensions decrease, a higher temperature required for its emptying. The relationship between the temperature of desorption and the radiusof the pore is determined by the Kelvin equation. Under the conditionsof the TG experiment, emptying of the pores of a given diameter occurs when the pressure of the liquid vapors becomes equalto the atmosphericpressure. The pore radii for benzene and 100 90

60

Temperature,

Thermal desorption curves (weight loss against temperature) of acetone (curve 1) and benzene (curve 2) from porous PETP foil: segment I, desorption of liquid outside the pores; segment desorption of liquids from the pores. The continuous line denotes an aging time 2 h; the broken line denotes an aging time of seven days. (Reprinted from Ref. 44 with permission from Elsevier Science.)

acetone at different aging times estimated from the T G experiment using the Kelvin equation are given in Table 3. Figure 11 shows the dependence of the pore radius on the energy of "'Xe ions taken from scattering data (filled diamonds) [443. Circles represent the pore radii calculated by the Kelvin equation from the benzene desorption curves, and the triangles represent the pore radii from the T G experiment with acetone as the adsorbate. As can be seen, the values from the samples aged for seven days (empty symbols) are little higher in comparison with the values estimated in the small-angle scattering experiment. At a shorter agingtime (2 h) the values (filled symbols) become smaller and very similar to those obtained by using the SANS method [44]. This means that the solvents cause a textural modification of the sample.Whentheaging time increases, theporediameters and porevolumes also increase. This observation maybe explained by swelling ofthe organic skeleton of the polymer. As was mentioned earlier, the decrease in the desorption curve following segment I1 is connected with the evaporation of liquid squeezed in the polymer skeleton. The height of segment I1 is a measure of the total pore volume Vp. The values for both investigated liquids and two different wet storage times are given in Table 3. Because the desorption of liquids from the pores takes place pratically at a constant temperature, one can assume the pore tracks to be of uniform size. Our observation is not consistent with the results of Quinson et al. [42] obtained by ther~oporometryfor a polycarbonate symmetrical membrane. Their investigations of the liquid-solid transfor~ationof a capillary condensate inside the pores of this membrane show a pore size distribution. The results of the TG analysis of the investigated PETP foil are also compared with results from the conventional nitrogen method. Figure 12a shows adsorption-desorption isothermsof nitrogen at 77 K. The value was estimated from the desorption isotherm atp / p s 0.98 (interpolation for 0.9641 p/p, 0.9966). In Fig. 12btheBETplotforthesame systemis shown. As may be seen,owing to very small adsorptionof nitrogen on the investigatedfoil,

Parameters Characterizing the Porous Structure

Benzene Acetone

2h 7 days 2h 7 days

Nitrogen "Theinternalsurface of theporesinthePETP S VP/RP. %ET value from nitrogen adsorption. Ref. 43.

8.9 0.34" 14.3 0.45" 9.3 15.8 15.3

Investigated PETP Foil

46 56 48 57 51

0.34" 0.46" 0.6b

foil calculatedaccording to therelationship

5

0

Plot of pore radiusagainstionenergy: scattering data [44] calculated from the desorption curves for the PETP foil aged in acetone om the desorption curves for the PETP foil aged in acetone for radius calculated from the desorption data of benzene, 2 h and sevendays,respectively.(ReprintedfromRef,43withpermissionfromElsevier Science.)

E Kn

0.00

1.o

Adsorption-desorption isotherm of nitrogen on PETP foil at 77 K ion; 0,desorption. (b) BET plot BET b/p,)/ erimentalvalues;solidlinerepresentsleast-squaresfit to the space between the broken lines represents the confidence interval for thestraight-linefit.(ReprintedfromRef. 43withpermissionfrom Elsevier Science.)

a large experimental error occurs within the range (0.05, 0.5). The straight line in Fig. 12b represents the least-squares fit to the experimental data. The linearity of the BET plot within this range ofp/ps may be controversial. For theabovementioned reasons, theS, value is not precise and may even involve a 50%er5or. Hence the hydraulic radiusof the pores estimated from the nitrogen data of 51 is surprisingly, and rather by accident, close to the R, value from the scattering and TG data. Generally, one can say that the nitrogen method is not suitable for the characterization of the porosity of the investigated foil. The results of this method are given for illustrative purposes only. On the basis of presented data one canconclude that the TG method is useful in the investigations of the porosity of organic membranes producedby the treatment of high energy heavy ions. Moreover, theT G method makes it possible to examine the evolution of the structure during the aging process of swelling materials. Good consistency exists between parameters characterizing the pores of the investigated PETP foil obtained by the SANS and T G methods.

Silica gels of RP type with bonded alkyl chains are widelyused as supports in liquid-solid chromatography. The dense nonpolar grafts showsome interesting phenomena and structural changes which depend on temperature and the presence of molecules of different types. The significance of these phenomena in chromatography well recognized now and many papers are devoted to this problem [45481. Following the suggestion of Morel and Serpinet 149,501 and Claudy et al. [5l] we have also detected a phase transition for the C-l8 phase by using the annihilation method 1521. The main conclusion which can be drawn from these investigations is that the thickness of the bonded phase changesat around 20°C. Below this temperature the alkyl phase shows higher compactness and in this connection is thinner. Some changes in the structure of the alkyl graft below and above the transition temperature occur during its contact with liquid adsorbates. It should be pointed out that only some surface silanols take part in the reaction with the modifying reagent. For steric reasons the surface density of alkyl chains cannot surpass 4.2-4.3pmol/m2, which corresponds to approximately half-coverage of the silica surface [53]. The coating of the modified silica surface by free normal paraffins of similar chain lengths, which insert themselves between the bonded chains, leads to a much denser mixed layer [53-551. The aim of the investigations presented below was to investigate the influence of the presence of a nonpolar surface phase on the parameters characterizing the porous structure obtained by using the nitrogen and T G methods (561. We chose to study unmodified silica gel and two chemically modified silica gels for reversephase chromatography. Silica gels Si-100, Lichrosorb RP-8, and LichrosorbRP-l8 from Merck (Germany) were used as the adsorbents. The silicaswere dried by prolong~dheating at 150°C to remove the physically adsorbed water. The initial characterization of these materials, including the determination of the specific surface and pore size distribution, was performed on the basis of adsorption-desorption isothermsof nitrogen. The adsorption-desorption isothermsof nitrogen measured at 77 K for three silicas are presented in Fig. 13.

4

m P

600

N

>z

600

V

c 600

0 0.6

PJP, Nitrogenadsorption-desorptionisotherms at 77 IS forsilica els:(a) (b)LichrosorbRP-8, and (c) LichrosorbRP-18. 0, adsorption; desorption. (Reprinted from Ref. 56 with permission from Elsevier Science.)

Parameters characterizing the pore structureof the investigated adsorbents,i.e., specific surface areas, and pore volumes, Vp,were calculated in the standard manner. The specific surface areas, have been determined using the BET methodfromnitrogenadsorption at 7'7 K (following outga!sing at 150°C and takingthe cross-sectional area of an Nz molecule 16.2 A).Pore radii were derived from the desorption branches of the isotherms. It should be noted that the pore filling and emptying occurred over narrow range of p/p, 0.7-0.9. The and Vp values obtained are collected in Table 4. Thermogravimetric curves representing the weight loss against the temperature for desorption of benzene and n-butanol from the RP-18 sample are presented in Fig. 14, while Fig. 15 shows the desorption curves benzene from all three silica

Parameters Characterizing the Porous Structure of Silica Gels Obtained by Using the Nitrogen and Therl~ogravimetricMethods

0.97

62

Benzene

Si-100

1.10

320 n-Butanol

0.98

63

BenzeneRP-8

0.73

65

77

214

n-Buta11ol

0.83 73

0.75

47

RPBenzene 18 65

168 62

tz-Butanol

0.5 0.55

0.50

Source: Ref.

400

20

80

100

120

140

Temperature, "C

Desorptioncurvesofbenzene(curve l j and n-butanol(curve from Lichrosorb RP-18. The parts of these curves above the circles corresponds to desorption of the bulk liquid and that below the circles represents desorption of liquid from the pores and their internal surface. (Reprinted from Ref. 56 with permission from Elsevier Science. j

1000

800

i

200

0 20

40

60

80120 100

140

Temperature,

Desorption curves of benzene from silica gels: curve 1, Si-100; curve 2, Lichrosorb RP-8; curve 3, Lichrosorb RP-18. (Reprinted from Ref. 56 with perrnission from Elsevier Science.)

gels. Those parts of the curvesin Figs. 14 and 15 above thecircles correspond to the bulk liquid outside the porous structure of the silica. Core size distribution curves forsilica gel Si- 100,RP-8, and RP-18 are shown in Fig. 16a-c respectively. Dotted lines representthepore size distributioncurve obtained from the nitrogen method. It is seen that in the case of silica gel Si-l00 (Fig. 16a) the pore/core radius at the peak of the distribution curve fromTG data andthenitrogen dataforboth benzene andn-butanolare close to each other. The difference S;:k) d represents the thickness of the surface film at the appropriate temperature. For RP-8 and RP-l8 silicagels with n-butanol and nitrogen, the peaks are, again, close together, i.e., the separation d is almost the same whether the alkyl phase is absent (see Fig. 16a), or is present in short-chain (RP-8, Fig. 16b) or longchain (RP-18, Fig. 16c) form. Chemically modified silicas of RP type, particularly with a densely packed C- 18 phase, are of hydrophobic character. In polarsolvents, like water, poor wetting of the surface is observed. Decreasing the polarity of the wetting liquid may activate the bonded phase to swell and therefore the chain to stretch, thereby providing space for its molecules to penetrate between the alkyl chains. Thus, for this surface phasemodel,adsorption of n-butanol even ontheinorganicmaterial (silica) between radicals would be possible. However, sometimes it is assumed for molecules of alcohols that the hydrocarbon radical resides between the upright chains,

oworek

0.6

0.6

6

10

Porelcore size distribution curves for silica gels: (a) Si-100, (b) Lichrosorb RP-8, and Lichrosorb RP-l8 obtainedby: (1) TG method(desorption of nbutanol); (2) TG method (desorption of benzene); (3) nitrogen method. (Reprinted from Ref. 56 with permission from Elsevier Science.)

but its hydroxyl group points towards the bulk liquid [48], In this connection, duringemptyingtheporesthepolarmolecules of n-butanol squeezedbetween alkyl chains can be easily removed. A relatively nonpolar molecule suchas benzene will be forced more deeply into the alkyl phase than a polar molecule such as nbutanol. The solvation of the alkyl chain cannot be described as an adsorptive process because bonded groups do not present a classical surface. Thus, depending on the thickness of the grafted phase the solvent molecules intercalate themselves between alkyl chains entirely or only partially. The stationary phaseof the dynamic solvated layer presents a structure determined by many factors. This structure for the ODs

phase is strongly dependent on temperature. At lower temperatures ODs groups are more organized, losing mobility. Hence the interchain access of solvent molecules is restricted. At higher temperatures the accessibility of interchain space is easier 71. The results of our experiments indicate that in nitrogen adsorption aswell as nbutanol desorption experiments the alkyl chains shrink together and resemble a liquid film covering the silica surface. Thus the cores of pores would have asimilar diameter. It should be noted that,forthe silicas investigated, thetotalpore volumes, ‘vp, estimated by various methods are similar. Withbenzeneandnitrogen,however, when the alkyl phase is absent,the separation d is almost the same as with n-butanol and nitrogen, but when the alkyl phase is eresent, is considerably enhanc$d, and more so for the longer chain, i.e., 18 Aforthe octadecylsilyl and 12 Aforthe octylsilyl phase.This suggests a model where benzene molecules penetrate the bonded film completely and interact withthemethylenegroups of thechain,forminga mixedalkylbenzene layer with chains oriented steeply toward the surface. The presence of an alkyl phase on the silica surfacecausesa dimil~utionof the pore diameter estimated from the T G method in comparison to the nitrogen method (see Table 5). Thelongerthe alkyl chain is, the biggeris thedecrease of the radius. Theoretically, the attachment of a chemisorbed layer to a porous silica leads to diminution of the mean pore diameter by twice the thickpess of the layer d. For the C-l8 alkyl chain, one can assume the value dl 24 A 148,581. In the case of the RP-l8system the difference abetweenothe poreradii estimated by the nitrogen and T G methods, A 47 A A, is similar to this value. It should be pointed out that under the conditions of the low-temperature experiment (nitrogen method) the surface phase containing shrunk alkyl chains also has a thickness which cannot beprecisely determinedonthe basis of the data presented. Our results indicate the different states of bonded phases under the conditions of gas adsorption experiment in the absence of liquid adsorbate (nitrogen method) in comparisonwith wet samplesundertheconditions of thethermaldesorption experiment. As mentioned above, a largely stable and thick surface filmis formed in the case of benzene andthe C-18 bondedphase.Atthesametimeonthe distribution curve for this system, derived from T G data, an elevated segment on the left side observed. Deformation of the distribution curve is probably connected with desor-ption of liquid accumulated between alkyl chains and adsorbed on the silica surface. Desorption of this part of the liquid requiresahigher temperature. The above presented results show that in liquid benzene the alkyl layer is thickened, probablyby interaction of the benzene molecules with the CH2 groups of the alkyl chain thence a reduction in pore width is observed, The effectis small or absent with polar n-butanol. Thus, though the T G method leads only to an approximate core size distribution, the information it can provide may perhaps be more relevant in a chromatography context than the more precise pore size distribution obtained from the standard low-temperature desorption of nitrogen method, where interaction is less intense.

Mixed samples Si-40jSi-l00 were prepared gravimetrically at the wjw proportion l:3, 1:1, and 3:l. Silicaswere driedbeforeexperiment by prolongedheating at l8O"C. Thermal desorption curves for all investigated silicas are reported in Fig. 17. The curves were normalized to mass of 1 g of dry adsorbent. For the same samples the adsorption-desorption isotherms of nitrogen at 77 were also measured. Porejcore size distribution curves V / h R = f ( R ) for investigated silicas were calculated in the manner described earlier [31]. Numerical values of the parameters characterizing the pore structure of the investigated samples obtained by two different methods are collected in Table 5. The data derived from nitrogen isotherms show a decreaseof specific surface area and increase of porevolume as conta~inationofsilica Si-l00 in thesample increases. Pore size distributionsfor these samplesareshown in Fig. 18. Solid lines in Fig. 18 represent PSD calculated on the basis of T G data. The mixed samples of silica gelcontain mesopores in a wide size range. It should be noted that the PSDs for Si-40 and Si-l00 silica gel do not overlap, i.e., these silicas do not contain pores of the same dimensions. Thus, analysis of the desorption data for mixed silica samples is a test of efficiency of TG method. expected in the case of investigated silicas two inflection points on the TG curves are present. Depending on the amount of a given silica in the sample an increase or decrease of the appropriate segment on the desorption curve is observed. Pore size distributions derived from these data are of bimodal type. Taking into account the amount of liquid adsorbate desorbed above its boiling point, the total pore volumes were calculated. is seen in Table 5 the total pore volumes determined by using the T G and nitrogen methods for the same sample of

l

140 Temperature, "C

Desorption curves of benzene for mixedsamplesSi-40/Si-100and w/w ratios of: 1:3; (2) 1:l; (3) 3:l. (ReprintedfromRef. 38 withpermissionfrom Elsevier Science.)

0.0 10

Radius,

Pore/core distribution curves for mixed samples Si-40/Si-100. Labeling as in Fig. 17. Points represent pore size distribution from nitrogen method for mixed sample with w/w ratio 1:1, (Reprinted from Ref. 38 with permission from Elsevier Science.) silica are close together. The results confirm the high efficiency of T G method in a wide range of mesopores.

r Twomesoporous silica molecularsievesweresynthesizedby two previously described methods [59,60]. Onemesoporous silica designated as EUSI 144was preparedaccording totheprocedure ofBeck etal. at Pennsylvania State University. A hexagonal mesoporoussilica (EUHMS) was prepared by the method of Tanev and Pinnavaia [60,61], which uses a neutral templating route unlike the cationic template route [59,62] used above. Figures 19 and 20 show the nitrogen

ParametersCharacterizingPorousStructure Using Nitrogen and TG Methods

Si-40 Si-4O/Si-100 (l 3) Si-4O/Si-100 (1 1) Si-4O/Si-100 (3: 1) Si-100 Ref.

18

19/54 18/54 18/54 56

0.44 0.91 0.8 1 0.56 0.97

ofSilicaGelsObtained

21 842 20/6 448 1 2 1525 /66 20/62 64

0.52 0.94 0.89 0.56 1.10

705 320

by

adsorption~esorptionisotherms of the silicas [63]. The appearance of a capillary hysteresis between and 0.45p/pSin the isotherm clearly indicates the presence of mesoporosity. Figure 20 shows the desorption curvesof benzene from the mesoporous silica molecular sieves, Figure 21 shows pore size distribution for mesoporous silica. Similar distributions were obtained for Si-l44 silica. A comparison of the results by both methods shows avery good agreement, and the deviation is within experimentalerror (Table 6). The results collected in column of Table 6 obtained by N2 adsorption using two different types of instruments, one from Pennsylvania State University and one from Maria Curie Sklodowska University, are also very similar. These results clearly indicate that thetemperatureprogrammeddesorption method is suitable in determining the pore size distribution in materials containing narrow mesopores, especially those that are based on silica.

m cn

l

~dsorption~esorption isotherms of nitrogen on a mesoporous silica prepared using a cationic and neutral template and calcined at 300°C for h. Open circles and squares, adsorption; solid circles and squares, desorption. (Reprinted from Ref. 63 with permission from Kluwer Academic Publishers.)

Ther 100

l

6

Temperature, "C

Thermal desorption of benzene from a mesoporous silica made by the cationic(a)and neutral (b)template route and calcined at 300°C for 4 h. (Reprinted from Ref. 63 with permission from Kluwer Academic Publishers.)

Parameters of Mesoporous Silicas byN2 Desorption and TPD Methods

Cationic template Neutral template

7

"Pennsylvania State University. bMaria Curie-Sklodowska University. Ref. 63.

0.73b 54b

14.0a

5.4b 15.7b

13.3b 5.gb

1.o

0.5

1.5

1.o

0.5

10

0.1

Radius, nm

Pore size distribution obtainedusing TPD method (a) and nitrogen method (b) of mesoporous silica made by neutral template route and calcined at 300°C for 4 h. (Reprinted from Ref. 63 withpermissionfromKluwerAcademic Publishers.)

Recently we reported the effect of heat treatment on the textural and adsorption characteristic of silica gelsusually used in gas chromatography Silica gelsafter pretreatment athigher temperature appear as model porous interfaces for studying the influence of surface irregularities on the desorption processin T G experiments. Comparison of the desorption curves fromT G experiments and PSDs obtained by using two different techniques for thermally modified silica gels was the aim of the investigations presented below, Five samples of silica gel Si-l00 from Merck, Germany, after heat treatment at 200, 400, 600, 800, and 1000°C wereused in the experiment. The physical characterization datafor these silica sampleshave been previouslyreported Information concerning the porosityof four of those silicas is collected in Table 7. can be seen, temperatures below 800°C do not influence substantially the specific surface area of silica gel and its pore structure. owever, concentrations of the surface silanols depend strongly on temperature. In column 2 of Table 7 the

er

ro~ity

number of OH groups per mm2 (aoH) for four of the investigated samples are collected. These values were calculated from the relation 2

lo3

W

where Wdenotes the percentage water content of silica geldetermined from theloss of silica mass due to its calcination at 1000°C [65]. Let us consider the desorption of pure liquids from silica surfaces. Thermodesorptioncurves of carbon tetrachloride and acetonefor investigated silica samples are shown in Figs. 22 and respectively [66]. The lower parts of the desorption curves forsilicas with different concentration of surface silanols are almost identical, which may indicate that the chemistry of the surface does not influence the thickness of the surface film. This effect is observed for all the silicas studied, irrespective of thetype of the adsorbate. The identity of the ends of desorption curves is especially surprising in the case of acetone for which surface silanols are the main adsorption centers [67]. A cornparison of the PSDs from the nitrogen method and T G experiment for various liquid adsorbates is presented in Fig. 24. In the case of carbon tetrachloride the distributions obtained by the different methods are close together. For acetone systems the PSDs derived from T G data are shifted toward smaller radii in comparison to the nitrogen method. similar effect isusually observedin the caseof silica geland strongly polar componentslike water or methanol [68]. The difference of the peak locationof pore size distribution curves maybe ascribed to the presenceof strongly bonded liquid film remaining on the wallsof poresafter their emptying. It is noteworthy that the difference RE:,) is similar for both silica samples (see Table 7). In other words, the thickness of the surface filmis independent of the concentration of surface

F" m-

0 50

70

75

80

Temperature, "C

Thermogravimetric curvesof desorption of acetone for various silica gels: (1) Si-100(200), (2) Si-100(400), (3) Si-100(600), and (4) Si-lOO(800). (Reprinted from Ref. 66 with permission from the Royal Society of Chemistry.)

2400 2000

ui

1200 800

0 70

75

"C

Thermogravimetric curves of desorption of carbon tetrachloride for various silica gels: (1) Si-100(200), (2) Si-100(400), (3) Si-100(600), and (4) Si-lOO(800). (Reprinted from Ref. 66 with permission from the Royal Society of Chemistry.)

0.4

0.0

0.4

4

8

2

4

6

8

PSDs obtained by using various methods: points, nitrogen method; solid lines, thermogravimetry. (a) Si-100(200), Si-lOO(8OO)~esorptionof CC14, (c) Si100(200), (d) Si-lOO(2OO)~esorptionof acetone.(ReprintedfromRef. 66 with permission from the Royal Society of Chemistry.)

er

Parameters Characterizing the Porous Structure of Investigated Silica Gels method Nitrogen

method vp

Silica

vp

(oH/nm2> (m2/g) Adsorbate (crn'/g) (A)

315 Si-l00 (200°C)

3.86

Si-l00 (400°C) Si-l00 (600'C)

5.83

l 30 2.21

230 Si-l00 (800°C)

1.01

(cm3/g) 6 6 f 3Acetone

319

1.08

1.05 0.88

CC14 rt 3 Acetone CC14 66 f 3 Acetone CC14 f 3 Acetone CC14

(A)

0.91 52 10 0.98 64 f 3 0.90 51 f 10 0.92 65 3 0.92 52 f 10 0.84 3 0.76 5 0 f 10 0.75 6 7 f 3

Ref.

silanols. This effect correlates reasonably with the similarity of the shape of TG curves for various silica gel samples (see Figs. 22 and 23). A comparison of the total pore volumes estimatedby the different methods can be seen in Table It appears that in the case of the TG experiment the total pore volumes Vp are similar for acetone and carbon tetrachloride. These valuescorrelate very well with the values of from the nitrogen method. Corre~tionof Desorption Curves for instru~entaiResponse It should be noted that the instrumental response in the case of definite pore size is not a delta function. The evaporation rate highest at the boiling point, but this rate is nonzero also at lower temperatures. For a definite boiling temperature we loss versustemperature obtainthe instru~entalreaction P(T),andthemass observedexperimentally is thusaconvolution of instrumentalreaction with the mass loss distribution expected in an ideal R(T) (Le., not distorted by the The real pore size distribution can be derivedfrom R(T),not from

T):

where P(T is normalized to unit area. The function is found experimentally, as well as P(T),which can be taken from the experiment with nonporous material. The functions and were approximated by respective historgrams, thus instead of integration we have the summation:

Equation (2) represents a set of equations which allow us to determine the value, to reconstruct the real desorption curve.

Figure shows for illustrative purposes the initial and corrected curves Am/ versus LT for the carbon tetrachloride/Si-l00(200) system, is seen, the difference between these curvesis very small and occurs within thelower temperatures which correspond to desorption from large mesopores. Summing up, one can state that the chemical characterof silicasurfaces does not influence the shape of thermal desorption curves for given adsorbate. Thus, PSDs derived from these curves and the surfacefilm effects for various samples are very similar.

The T C technique is nondestructive and can providedetails of the porous structure of wet silica gels containing organicsolvent. Temperature-programmed desorption (TPD) may be usedfor rigid materials but also for those whose texture may swell in liquid medium. From the comparative analysis of experimental data obtained using various methods one can conclude that the TPD method is accurate and sensitive enough for characterization of the porosity of silica gels within the mesopores. It also possible to distinguish whether given sample contains only mesopores or also rnicropores. Especially valuable results may be derivedfromdesorption data of various liquids for silica adsorbents with chemically modified surfaces. Thus,althoughthe TPD methodleadsonlyto core size distribution,the information it can provide may be more relevant in chromatography context 60 50

ui

10 0 90

85

80

100 Temperature, "C

Arn/AT vs. for desorption of from silica gel Si-lOO(200); solid line, initial curve; broken. line, corrected curve. (Reprinted from Ref. 65 with permission from the Royal Society of Chemistry.)

than the more precise pore size distribution obtained from the standard low-temperaturedesorption of nitrogen, where interactions of theadsorbateare less intense.

1. K. K. Unger, in Porous Silica, Elsevier, Amsterdam, 1979. 2. R. K. Iler, in The Chemistry of Silica, John Wiley Sons, New York, 1979. 3. C. J. Brinker and G. W. Scherer, in Sol-Gel Science, Academic Press, New York, 1989. 4. E. F. Vansant, P. Van Der Voort, and K. C. Vrancken, in C?zaract~rizationand Chemical ~ o d ~ c a t i oofnthe Silica Surface, Elsevier, Amsterdam, 1995. 5. R. Yu. Sheinfain and I. E. Neimark. Kinet. Katal. 8:433 (1967). 6. A. V. Kiselev, Yu. S. Nikitin, and E. B. Oganesyan. Kolloidn. Zh. 31:525 (1969). 7. T. Murakata, Sato, T. Ohgamara, T. Watanabe, and T. Suzuki. J. Mater. Sci. 27: 1567 (1992). 8. L. L. Hench and J. K. West. Chem. Rev. (1990). 9. R. Deshpande, D. V. Hua, D. H. Smith, and C. J. Brinker. Non-Cryst. Solids 144:32 (1992). 10. R.K. Iler,in Surface and Colloid Science (E.MatijeviE,ed.),Vol.6,Wiley, London, 1973. 11. E. P. Barrett, L. G. Joyner, and P. H. Halenda. J. Amer. Chem. Soc. 73:373 (1951). J. Greggand K. S. W. Sing, in Adsorption, Surface and Porosity, Academic 12. Press, London, 1982. 13. G. D. Halsey. J. Chem. Phys. 16:931 (1948). 14. M. Brun, A. Lallemand, J. F. Quinson, and C. Eyraud. Thermochimica Acta 2159 (1977). 15. J. F. Quinson and M. Brun, in Characterization of Porous Solids (K. K. Unger, J. Rouquerol, K. S. W.Sing, andH. Kral, eds.),Elsevier,Amsterdam,1988, p. 307. 16. 0 . Glutter and 0 . Kratky, in SrnaZl Angle X-ray Scattering, Academic Press, 1982. C. Dore and A. N.North, in C~ar~cterisation of Porous Solids (F. Rodriguez17. Reinoso, J. Rouquerol, K. S. W.Sing,and K. K. Unger,eds.)Elsevier, Amsterdam, 1991, p. 245. 18. P. G. Hall and R. T. Williams, J. Colloid Interf. Sci. 104:151 (1985). 19. P. T. Branton, P. G. Hall, A. Mengel, and R. T.Williams, in Characterization of Porous Solids Rouquerol, F. Rodriguez-Reinoso, K. S. W. Sing, and K. K. Unger, eds.), Elsevier, Amsterdam, 1994, p. 247. D. F. Ramsay and R. G. Avery, in Ch~racterizationof Porous Solids (F. 20. Rodriguez-Reinoso, J. Rouquerol, K. S. W.Sing, and K.K. Unger,eds.) Elsevier, Amsterdam, 1991, p. 235. Yau, J.J. Kirkland, andD.D. Bly, in Modern S i z ~ - ~ ~ c l u ' ~ i o n 21. W. C h r o ~ u t o g r a p ~Wiley-Interscience, y, New York, 1979. Kirkland, in Int~oduction to Modern Liquid 22. L. R. Snyder and J. Chromatograp~y,Wiley-Interscience, New York, 1979.

k

23. U. A. Eltekov. Pure Appl. Chem. 61:1987 (1989). 24. W. L. Earl, W. Kim, and D. M. Smith, in Churucterization of Porous Solids (J. Rouquerol, F. Rodriguez-Reinoso, K. S. W.Sing, andK.K. Unger,eds.), Elsevier, Amsterdam, 1994, p. 301. 25. C. A. Koh, J. A. Zollweg, and K. E. Gubbins, in Cha~acterizationo f P o ~ o u sSolids Rouquerol, F. Rodriguez-Reinoso, K. S. W, Sing, and K. K. Unger, eds.), Elsevier, Amsterdam, 1994, p. 61. 26. J. P. Olivier, W. B. Conklin, and M. Szombathely, in Churacterization Porous Sol~ds (J. Rouquerol, F. Rodriguez-Reinoso, K. S. W. Sing, and K. K. Unger, eds.), Elsevier, Amsterdam, 1994, p. 81. 27 C. A. Jessop, S. M. Riddiford, N. A. Seaton, J. P. R. B. Walton, and N. Quirke, in Charucterizution Porous Solids (F, Rodriguez-Reinoso, J. Rouquerol, K. S. W. Sing, and K. K. Unger, eds.), Elsevier, Amsterdam, 1991, p. 123. 28. IC. Birdi, D. T.Vu, S. I. Andersen, A. Winter, H. Topsoe, and S. Christensen in Characterizution Porous Solids (F. Rodriguez-Reinoso, J . Rouquerol, K. S. W. Sing, and K. K. Unger, eds.), Elsevier, Amsterdam, 1991, p. 15 1. 29. J. Goworek and W. Stefaniak. Colloids Surfaces 57:161 (1991). 30. J. Goworek and W. Stefaniak. Colloids Surfaces 60:341 (1991). 31. J. Goworek and W, Stefaniak. Colloids Surfttces 62:135 (1992). 32. Goworek and W. Stefaniak. Colloids Surfaces 69:23 (1992). 33. Goworek and W. Stefaniak. Mat. Chem. Phys. 32:244 (1992). 34. J. Goworek and W. Stefaniak. J. Thermal Anal. 45:999 (1995). 35. F. Paulik and J. Paulik. J. Thermal Anal. 5:253 (1973). 36. F. Paulik, Paulik, and M. Arnold. J. Thermal Anal. 32301 (1987). 37. F. Paulik, in in T ~ e r ~ u l John Wiley Sons, Chichester, New York, Brisbane, Toronto, Singapore, 1995. 38. J. Goworek and W. Stefaniak,in C~zaracterizution Porous Solids (J. Rouquerol, F. Rodriguez-Reinoso, K. S. W.Sing, and K. K. Unger,eds.), Elsevier, Amsterdam, 1994, p. 401. 39. J. Goworek and W. Stefaniak. Thermochimica Acta 286: 199 (1996). 40. L. A. de Wit and J. J. F. Scholten. J. Catal. 36:36 (1975). 41. P. Lewitz, J. H. Drake, and J. Klafter. J. Chem. 89:5224 (1988). 42. F. Quinson, N. Nameri, and B. Bariou, in Churacterizatio~ Porous Solids (F. Rodriguez-Reinoso, J. Rouquerol, K. S. W.Sing,and K. K. Unger,eds.), Elsevier, Amsterdam, 1991, p. 209. 43 J. Goworek and W, Stefaniak. Colloids Surfaces 82:71 (1994). 44. D. Albrecht, P. Armbruster, R. Spohr, M. Roth,K. Schaupert, and H. Stuhrmann. Appl. Phys. A 37:37 (1985). 45 G. E. Berendsen and L. de Galan. J. Chromatogr. 196:21 (1980). 46. P. Roumeliotis and K. K. Unger. Chromatogr. 149:211 (1978). 47 C. Horvath and W. Melander. J. Chromatogr. Sci. 15:393 (1977). 48. G. E. Berendsen and L. de Galan. J. Liq. Chromatogr. 1:561 (1978). 49 D. Morel and J. Serpinet. J. Chromatogr. 200:95 (1980). 50. D. More1 and J. Serpinet. J. Chromatogr. 248:231 (1982). 51. P. Claudy, J. M. Letoffe, C. Gaget, D. Morel, and J. Serpinet. J. Chromatogr. 329:331(1985).

52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63 64. 65. 66.

67. 68.

Wawryszczuk, M. Lewandowski, W. Gorniak, J. Goworek, and T. Goworek. Chromatographia 25:721 (1988). Riedo, M. Czencz, 0. Liardon, and E. Kovats. Helv.Chim.Acta62:1912 (1978). G. Korosi and E. Kovats. Colloids Surfaces 2:315 (1981). T. C. Schunk. Chromatogr. 656:289 (1993). Goworek and W. Stefaniak. Colloids Surfaces 80:251 (1993). Goworek, F.Nooitedacht, M. Rijkhof, and H. Poppe. Chromatogr. 352:399 (1986). B. Buszewski and Suprynowicz. Chromatographia 24:573 (1987). Beck, J. C. Vartuli,W. J. Roth, M. E. Leonowicz, C. Kresge, K. D. Schmitt, C.T. W. Chu, D. H. Olson,E. W. Sheppard, S. B.McCullen, J. B. Higgins, and J. L. Schlenker. J. Am. Chem. Soc. 224: 10834 (1992). P. T. Tanev and Pinnavaia. Science 267:865 (1995). P. T. Tanev and Pinnavaia, in Advances in Porous ~ u t e r i ~ l s Komarneni, M. Smith and J. S. Beck, eds.) Mat. Res. Soc. Symp. Proc. (Pittsburgh, PA, 1995), Vol. 371, p. 63. C. T.Kresge, M. E. Lenowicz, W. J. Roth, J. C. Vartuli, and J. Beck. Nature 359:710 992). Komarneni, V. C. Menon, R. Pidugu, J. Goworek, and W. Stefaniak. Porous Mat. 3:115 (1996). Goworek and A. Nieradka. J. Colloid Interf. Sci. 280:371 (1996). G. Linsen,in Physical ChemicalAspects Adsorbents and C~talysts, Academic Press, London/New York, 19’70. Goworek, W. Stefaniak, and T.Goworek, in Churucterization Porous Solids W(B. McEnaney, T. Mays, J. Rouquerol, F. Rodriguez-Reinoso, K. S. W. Sing, and K. K. Unger, eds.), Royal Soc. Chem., 1997, p. 481. Goworek and R. Kusak. Colloid Polym. Sci. 267:539 (1989). Goworek. Am. Lab. 5:16 (1995).

This Page Intentionally Left Blank

Institut de Chimie des Surfaceset Interfaces, Mulhouse, France Rhcjne Poulenc Industrialisation, CR Carrikres, Saint Fons, France

Introduction I. A. Principle of IGC B. Advantages of IGC over other adsorption techniques C.Limitations of IGC

206 206 206 207

11, IGC in Infinite Dilution Conditions (IGC-ID) A. Dispersive component of surface energy B. Surfacenanoroughness C. Specific component of surface energy D, Acid-base properties of the silica surface

207 207 208 210 21 l

III. IGC in Finite Concentration Conditions (IGC-FC) A.Adsorptionisotherms B. Adsorption energy distribution (surface energetic heterogeneity)

214 214 214

TV. Application of IGC-ID to a Variety of Silica Surfaces A.Nonspecificinteractionpotential B. Surfacenanoroughness C. Specific interaction potential of unmodified silicas D. Acid-base properties of silica

216 216 218 219 220

V. ~pplicationof IGC-FC to Silica Surfaces

220

VI. Application of IGC to the Study of Modified Silica Surfaces A. Heat-treated silica examples B. Silica grafted with alkyl chains

224 224 224

VII. Computer Modeling and

23 l

IGC

apirer et

VIII. Conclusion

238

References

238

Knowledge of the physical interaction potential of silica is required to understand and monitor the behavior of silica, particularly for important applications such as chromatographic processes thickening and thixotropic agent for liquids, fillers for elastomers, e.g., silicones, and, more recently, in rubbers used in the tire industry Thestudy of adsorption, using different adsorption techniques such volumetric and gravimetric methods, but also inverse chron~atograp~y (IGC) has been applied for some years, for the testing of the interaction potential of various solid surfaces: powders [3], fibers [4], and polymers [S]. IGC, in infinite dilution conditions (IGC-ID) or in finite concentration conditions (IGC-FC) is unique and efficient method to determine thermodynamic data from which series of fundamental and practical information may be derived. Changing the nature and size of the solute (molecular probe) injected in the GC column that contains the solid of interest will deliver information on the surface nanomorphology [6], the surface energy acid-base properties [8], and energetic heterogeneity of silica surfaces

The principle of IGC is rather simple: Fill a column for chromatography with thesolid of interest, select molecularprobes,injectthemeitherin very small amounts (infinite dilution condition, IGC-ID) or in measurable amounts (finite concentrationcondition,IGC-FC),anddeterminetheretention time (timeit takes for the probe to cross the column when pushed by carrier or analyze the GC peak shape. Obviously, both retentiontime and peak shape will depend on the solid probe interaction, i.e., on the characteristics of both thesolid’s surface and of the probe. The known propertiesand expected interaction of the probewill then allow us to identify the interaction potentialand surface characteristicsof the solid itself. Numerous other adsorption processes and applications of IGC are discussed in books and review papers [IO-131.

IGC presents number of advantages over classical adsorption techniques. The necessary equipment exists inalrnost every laboratory, since commercial GC fitted with. sensitive detector will be satisfactory. There is no need for outgassing facilities, since this will be achieved by conditioning the sample inside the column and flushing it with inert carrier often He and N2 at the desired temperature. Indeed, the possibility of an easy change of the measurement temperature constitutes major advantage of IGC. Furthermore, simultaneous determinations may be performed just by injecting mix of probe molecules. In situ treatments (thermal, chemical) may be applied to the samples without any reexposure to atmo-

sphere. Finally, the IGC data are superior accuracy.

usually obtained in a shorter time and with

There are physical limitations: the solutes should be volatile and the analyzed solid stable at the chosen measurement temperature. The preparation of the column needs special care. The GC support, the solid of interest, should be in a suitable form: silica powders having avery small particle size require agglomerationso as to meet common GC support dimensions (a few hundreds of microns). Problems that are typical of chromatography processes, nonlinearity, nonideality, and kinetic or diffusion parameters, may arise: solutions to these problems are discussed in the literature

At first approximation, the surface energy, i.e., the potential to interact through physical or reversible interaction, may be considered as a sumof two components: the dispersive componentandthe specific component London, dispersive, nonspecific, or universal interactions originate fromtherapidelectrondensity change, even in a nonpolar symmetricalmolecule presenting no macroscopic dipolar moment, such as methane, generating instantaneous dipoles. Those dipoles will interact with neighboring ones or induce instantaneous dipolesin close-proximity molecules. The London energy of interaction between substances and 2 is given by the following expression:

The dispersive interaction energy dependson the numberof interacting sites the polarizabilities (a),and ionization potentials (I) of the interacting partners. We may thusexpect that each variation of these parameterswill interfere in the adsorption phenomena and will be detectable by IGC. In other words, IGC will be very sensitive to minutesurfacepropertyvariationsupontreatments ofsilicas, for instance. The London interaction capacity of a silica surface may be determined using the methodof Dorris and Gray 161. For this method, a homologousseries of n-alkanes is injected in very small amounts into the column containing silica and standard variations of free energy of adsorption (AG:) are calculated from the net retention times (or volumes V,) according to

where C is a constant depending on the choice of a bidimensional reference state R is the universal gas constant for an ideal gas, and the absolute temperature at which the measurement is performed. As seen in Fig. AG: of alkanes varies linearly with the number of carbon atoms of the injected n-alkane probes.

25 -AGOa

l

l

8

Variation of the standard free energy of adsorption (AG:) of n-alkanes with theirnumber of carbon atoms, on a pyrogenicsilica (HDK T30, from WackerChemie)having a specificsurfaceareaof300 m2/g. Measurementperformed at 110°C.

The slopes of these straight lines vary for different (crystalline, pyrogenic, or precipitated) silica samples. The steeper the slope, the higher is the intensity of interaction between the solid surface andthen-alkane. From theslope of the straight line, it is possible to compute an incremental value, AGFH2 corresponding to the free adsorption energy variation of a CH2 group on the silica surface:

A G : ~ ~ RTln(vr,+l/v,) nowing AGEH2,

is then computed from 2

y ~ = ~ [ - ]

In this equation, all is known the area of an adsorbed CH2 group; Avogadro’snumber; yCH,: thesurfaceenergy of a solid entirely made ofCH2 groups, i.e., polyethylene) or measurable (AGFH2), except It is important to underline that physically meaningful y: values are only obtained when the examined surface is flat on the molecular level and, furthermore, when it is energetically homogeneous. Such conditions arescarcely encountered with actual solid surfaces. Therefore, the recorded y: should be analyzed with suspicion and only considered as indicative.

flat at themolecular level. Crystalline silicas, such esent identifiable surfaces with structural defects; other samples are porous. This meansthat, compared to the size of the probe usedto explore their surface, they may appear rough. In practice, one will observe size exclusion effects depending on the shape and bulkiness of a given probe. Various methods may be employed to describe the surfaceirregularities, e.g., determination

of surface fractality 18,191 and measurements of size exclusion effects through ICC [20]. The surface of pyrogenic silica, having a specific area of about 130 m2/g, has a fractality parameter close to 2. It is concluded that this sample may be considered as rather smooth at the molecular scale. Precipitated silicas have fractality parameters between 2and 3. Thismeans that their surfaceaspect is notplanar. Obviously, this will influence the interaction capacity of, say, branched or bulky alkanes. Consequently, a nanomorphology index may be simply defined by comparing the adsorption behaviorsof linear and branched alkane isomers, on the flat reference (pyrogenic)silica surface and on the oneof the examined sample. Yet the resulting valuewill depend on the choice of the alkane isomers. More recently, Brendlk and Papirer [21] proposed the use of topology indices (xt) of solutes, based on Wiener indices [22j, that perform satisfactorily to establish quantitative structure-property (boiling points, etc.) relationships. Xt indicesof some cyclic and branchedalkaneprobes thatare currentlyemployed in ICC arereported in Table 1. On a roughsilica surface 123,241, bulky alkanes stay apart from andbelow the alkane line depicted in 2. The decrement (--AGa corresponds to the contribution of the morphological factor to the loss of interaction energy between cyclic or branched alkanes on a rough silica surface. The nanomorphology index (Iln) can then be expressed by

Topology Indices of Some Common IGC Probes Used for Testing the Surface Morphology of Silica Probes Pentane Hexane Heptane Octane Nonane Decane Cyclohexane Cycloheptane Cyclooctane Methyl 2-butane Methyl 2-pentane Methyl 2-hexane Dimethyl 2,5-hexane Trimethyl 2,2,4-pentane Tetramethyl 2,2,3,3-butane aFor figures.

Xt

C5 C6 C7 C8 C9 C10 Cyclohex. Cyclohept. Cyclooct. MeBut. MePent. MeHex. TMPent. TeMBu.

5O .O 6.00 7.00 8.00 9.00 10.00 6.15 7.32 8.32 4.84

6.84 7.68 7.40 7.09

1nfVn)

. p

8 Xt

Determination of the decrementof free energy of adsorption of a branched probe, (iso-octane)on a crystallinesilica (H-~agadiite).Measurementmade at 110°C.

where AGF is the decrement of the free energy of adsorption caused by the size exclusion effect, evaluated by thedeparture of therepresentativepoint of the sterically hindered alkane probe (branched or cyclic alkane) from the reference linear alkane straight line. is, in fact, equal to the apparent ratio of the adsorption area of the sterically hindered alkane probe compared to a hypothetical nalkane having the same morphological index

Specific interactions include all types of interactions except London interactions. “Specific” means also that the types of interactions susceptible to exchange are specific to the two partners in contact. For instance, a nonpolar partner, such as polyethylene, willbe able to undergoonly dispersive interactionswith silica, whereasapolar solute, like acetone, will exchangewiththe silica surface both specific and dispersive interactions. For ICC purposes, the evaluationof the respective contributions of dispersive and specific interactions is based on the comparison of nonpolar (n-alkanes) and polar probe adsorption behaviors. The specific interaction capacity(AGip) is evaluated by subtracting, from the global variation of free energy of interaction of the polar probe (AG,), the contribution of the dispersive interactions. In practice, this is estimated by first plotting the “n-alkane reference line” relating the free energy AG, of adsorption of n-alkanes to a given molecular descriptor (vapor pressure, enthalpy of vaporization, xi).The representative points of polar probes, which are more interactive with the polar surfdcethan alkanes, are in principle located above the “alkaneline” and the specific interaction parameters are determined as shown in Fig. 3. This approach supposes that the polar probe has access to the same area of adsorptionthanthen-alkanehavinganequivalent size or parameter.This assumption will be fulfilled only if the surface is flat at the molecular level. On a

25

10

Determination of the increment of free enthalpy of adsorption of polar probem (acetone) ona pyrogenic silica (A130 from Degussa). Measurement madeat 80°C.

rough surface, only an “apparent may be observed, taking into account both the incrementof free energy dueto thespecific interactions and to the decrement of free energy related to the size exclusion effect.

‘Various forces (dipolar, H bond, acid-base, metallic, magnetic, ionic, hydrophobic) may intervene when a polar surface interacts with a polar solute. In nonaqueous media, whenever possible, acid-base type interactions become prevalent [25] compared to dispersive or pure polarinteractions. The solutes used to assesss the acidbase character,in the Lewis acidity or basicity definition, are acids, like chloroform, bases, like ether, or amphoteric solutes, such as acetone. Amphoteric means that this molecule possesses both acid- and baselike properties. In principle, it seems possible to evaluate the capacity of silica to exchange acid-base properties just by studying the adsorption of selected probes. This becomes only possibleif the acidbase characteristics of the solutes are themselves well defined.Up tonow, there has been no theoretical way to calculate these characteristics. Hence, only empiricalor semiempirical acid-basescales are available: the hard-softacid-basescale of Pearson [26], the E-C relation of Drago [27], thedonor acceptornumbers of Gutmann [28], and the solvatochromic parameters 1291. The hur~-soft acid-base scale Pearson [26]is based on molecularorbital theories. So far, theseconceptshavenot beenused for IGC given the largely unknown and complex nature of actual solid surfaces. j o ~ r - ~ u ~ u ~semieter e ~ ~ i r i c ea l ~ ~ ~ t i ~ n [27]is based on a relationship predicting acid-base reaction enthalpies, in the gas phase or in poorly solvating solvents:

The acid and the base (B) are characterized by two values: an E value that describes the ability of A and I3 to participate in electrostatic bonding (hardness) and value C (softness) that indicatestheirtendency to formcovalentlinks. Considering iodine as reference substance, for which EA; and making a series of calorimetric measurements, mixing iodine with the substance of interest, allows, by computer fittingof Eq. C and E values to be calculated forsome tens of molecules. The Drago approach has been used with success, starting from wettability l~easurements,for the evaluation of surface properties of oxides, yet not through IGC; but there is no peculiar reason for that. One major drawbackof the Drago scale comes from the fact that a molecule is defined either an acid or The ~ e t h of o ~ base.Inreality,most molecules exhibitamphotericproperties. ~ a ~ l and e t Taft E301 is called the “solvatochromic method,” since the measurementsare usually performed using UV or visible spectrometry.Peakshiftsare detected when the solid is mixed with solvatochromic probe molecules of known characteristics.Spange et al. E291 appliedthismethod to series of chemically modified silica samples and related the H-bonding accepting property to the isoelectric points determined by zeta potential nleasurements. Based on linear solvation energy relationships, thesilica hydrogen bond acidityand basicity well its polarizability(Londoninteractions)canalso be evaluated by IGC [31]. The ~ u t ~ a nelectron n acceptor ( A N ) and donor numbers ( D N ) 1281, that respectively stand for acid and base properties of molecule or substance, are determined e~perimentally. is defined the enthalpy of formation of /l molecular adduct of given molecule with reference Lewis acid (SbC15). DN is measured by the NMR shift of P when the given molecule ismixed with solution of oxotriethylphosphine (Et3PO) taken a reference base, the shiftbeing normalized by takingthevalue 0 forthesolvent(1,2-dichloroethane) and 100 for SbC15. Knowing the of a “pure” acid and DN of a “pure” base, one may simply calculate the enthalpy of interaction using

For thecalculation of the enthalpy of interaction between twoamphoretic molecules, having respectively (AN)1 and (AN)2 acceptor numbers and N)2 as donor numbers, weE321 suggested the application of

This relation hassince been often used in the literature for the evaluationof the acid-base properties of varioussolids.Theacceptornumbers and donor numbers (DN) of the most common IGC probes are gathered in Table 2. An example of the determination of the acid-base properties of a silica sample is displayed in Fig. 4. This figure results from linearization of (8). a at ever the method chosen, rather than focusing on the absolute physical meaning of the determined A N and DN values of the solid surface, designated only relative variations of acid-base properties should be considered. Practically, the determinationof the acid-base properties of solid surfaces supposes the ability of the polar probes to cross the chromatographic column in finite and

ilie

Acceptor Number (AN) and Donor Number (DN) of Common IGC Probes Probes

AN (Yo)

DN (kcal/mol)

0.7 5.4 4.7 2.5 l .5 1.4 0.5

0.0

Benzene

0

Chloroform 1 Acetonitrile 17.0 Acetone Ethyl acetate 19.2 Ether 20.0

14.

measurable time interval. This condition is usually not fulfilledby strong basic probes,such as ether or THF, that remainquite irreversibly adsorbedonthe acid silica surface under common IGC conditions. This constitutes an important limitation of IGC-ID forthedetermination of thermodynamic quantities, but another major limitation is related to the influence of the silica surface heterogeneity. Indeed, the thermodynamic data determined using IGC-ID have only a true meaning whenthey are measured on an energetically homogeneoussurface on which the residence times of the probe, on an adsorption site, will be independent of its location on the solid surface. On an energetically heterogeneous surface, this residence time will change from one point to another depending on the energy of interaction of the visited sites. Moreover,themobility of themolecule, in the adsorbed state,will amplify the contributionof the sites having the highest energies because they are able to act as potentialwells and capture the migrating molecule

Determ~natio~ at 110°Cof the acid-base properties of a silica surface (A130 from Degussa).

f331. The evaluation of the total interaction capacityof a solid surface with agiven molecular probe will require the determination of this capacity, in a large domain of surface coverages. In other words, only IGC in finite concentration conditions may deliver the wanted information.

The methods allowing usto calculate adsorption isotherms from chromatographic data are well documented in the literature [14], in particular “frontal analysis’’ and “elution by characteristic point” (ECP). However, care should be taken to identify and possibly correct difficulties arising from unwanted peak deformation due to concentration dependent behavior and also to nonlinear, nonideal, and kinetically controlled chromatographic processes. Only the ECP method willbe presented, since it offers the advantage of allowing the determination of a whole part of the adsorption isotherm just by evaluating a single chromatographic peak. Using this method, the first derivative of the adsorption isotherm can be readily calculated, starting from the retention times and signal the height of characteristic points taken on the diffuse descending rear front of the chromatogram according to thesimplified Conder equation

where is the number of absorbed molecules, P the pressure of the probe at the output of the column, L the column length, ti the net retention timeof a characteristic point on the rear diffuse profile of the chromatogram, the corresponding net retention volume, the James-Martin coefficient taking into account the compressibility of the gas dueto thepressure drop inside the chromatographic column, D the output flow rate,m the massof adsorbent, R the universal gas constant for an ideal gas, and the absolute temperature atwhich the measurement is performed.

Lamellar crystalline silicas possess macroscopic surface heterogeneity, since they present identifiable surfaces: the flat basal and theperipheral surfaces. Quartz exhibits different crystallographic planes having different chemical characteristics and hence different adsorption potentials. Amorphous nonporous silica surfaces are also covered by hydroxyl and siloxane groups, but the organization of these groups may differ from one silica sample to the other. Pyrogenic silicas (of moderate surface area) according to several investigations [19,20] have a “statistically” homogeneous surface, whereas precipitated silicas always present much heterogeneity: islands of hydroxyl groups, partially condensed polysilicic acid chains, chemical impurities, andratherrough surfaces. All these are reasonsforthe appearance of a multitude of different adsorption sites. In recent years [34,35],

convenient methods have been developed for the evaluation of the distribution of theadsorption sites. All approachesdescribed in theliteraturearebasedona physical model that supposes that an energetically heterogeneous surface, with continuous distribution of adsorption energies, may be described, in the simplest way, as a superposition of a series of homogeneous adsorption patches. Hence, the amount of adsorbed molecules (probes) is given by

where N ( P m , is the number of molecules adsorbed at pressure P, and temof measurement is the number of molecules needed for the forperature mation of monolayer, is the local isotherm, the adsorption energy of a site, and x(&) is the distribution function (DF)of the sites seen by the probe. The range of adsorption energiesis included between minimal (cmin)and maximal values. Various methods have been proposed to solve Eq. (10). The most popular one supposes discrete distributiol~of monoenergetic sites, based on a discretization formof Eq. (10). Considering a numbery1 of measurements, one endsup with a linear system of equations. To control the inherent instability of such systems, many authors, like Szombathely et al. and Jagiello [37j propose the use of regularizationparameter.Otherapproaches call on local adsorptioniostherm approximations reviewedby Nederlofetal. The oldest andthesimplest approximation of the local isotherm is the condensation approximation The condensation approximation supposes that thesites of adsorption of given energy are unoccupied below a characteristic pressure and entirely occupied above it. The distributionfunctionforthecondensationapproximation(DFCA) is directly related to the first derivative of the isotherm corrected for the multilayer adsorption, according to

where and are respectively the amount of adsorbed probe and the equilibrium pressure of the probe corrected for the multilayer adsorption, is the amount of adsorbate corresponding to the monolayer formation, R the universal gas constant for an ideal and the absolute temperature at which the measurement is performed. This approximation is all the better as the temperature of l~easurementapproaches the absolutezero. In the usual IGC measurement conditions, at room temperature and above room temperature, this approximation fails completely and it becomes necessary to use other approximated forms of the local isotherm. Among them, for Langmuir local isotherms, the extended Rudzinslci et al. method [40] allows the computation of the actual distribution function (DFRJ) from a limited development of the even derivatives of the DFCA, according to with (RT)2Jb~~X~ ~(c)

1 and b2j -----y”-l-(

(2j

l)!

An exampleof a distribution functionof the adsorption energies of propane-2-01 adsorbed on a pyrogenic silica sample is given in Fig. 5. A bimodal distribution function of the energiesof adsorption of propane-2-01is observed. The population, at lower adsorption energies, possiblycorrespondstosurface siloxane-rich domains, whereas the one at the higher energies is probably related to zones rich in silanols able to strongly interact by H bonding with isopropanol. Much care should be taken when interpreting the results, since each selected probe provides only its own appreciation of the solid's energetic heterogeneity.

Nonspecific interaction potentials of numerous silica samples (pyrogenic, precipitated, or crystalline) have been measured 1411, but usually at a given temperature only. in the temperature range where no significant alterations of surface chemistry occur, vary linearly with temperature. Table 3 141-461 displays thermal susceptibilities (t?y,d/t?t) and values, calculated at0°Cand 110°C either by interpolation or by extrapolation, from literature and from our own data, obtained on various silica samples. The value, 72 mJ/m2, extrapolated at room temperature, obtained with A130 compares fairly well with the one obtained by Kessaissia et al. who applied a different method (liquid contact angle orwettability nleasurements). Since a macroscopic method (wettability) and a molecular method (ICC) deliver the same value for y,d, one may concludethat this silica sample exhibits a rather homogeneousand smooth molecular surface at least at the scale of the probing molecule, thus corresponding to the applicability requirements of the Dorris and Gray method. The

x (pmol/((kJ/mol)) 0.10

0.05

10

16

28

Distribution function of the adsorption energies of propane-2-01 ona pyrogenic silica sample (HDK 3 from Wacker-Chemie).

y: andThermalSusceptibilities

DeterminedforVariousSilica

Samples Samples HDK HDK

A1

Wacker Wacker Wacker Degussa Degussa Degussa Degussa Degussa Degussa Degussa Degussa Rh6ne Poulenc

FIC. RP Lichroprep Merck Nippon Nipsil Silica H-~agadiite Lab

values of of the examined samples, determined at 1lO"C, vary between 30 and mJ/m2. The lowest value is observed for the precipitated silica (Lichroprep from Merck) and the highest for the crystalline silica (H-magadiite). Moreover, the y t of pyrogenic silicas vary only in a narrow range from to mJ/m2 when measured at 110°C: a proof of the low variability of the surface structure of this type of silicas. Nevertheless, one notices some differences between Degussa and Wacker-Chemie samples, possibly due to variations of the preparation conditions (flame composition and temperature) that may induce variations of the number of high-energy adsorption sites on pyrogenic silicas. The silanol density, determined by acid-base neutralization, is higher for the Degussasilicas S i O ~ / n m 2than ) for the Wacker silicas (-1.8 SiOH/nm2). Furthermore, the higherdispersive interaction capacity of the latter may also be related to a higher content of stressed siloxane groups. The thermal susceptibilities of vary in a large interval, especially for precipitated silica samples, as do the values themselves. The highest value is observed for H-magadiite. Two questions arise now: Why does the temperature susceptibility vary from one sample to another and is this susceptibility related to the surface structure of the silica sample? The answer will be partially given by size exclusion IGC, which provides information about the surface morphology at the molecular scale (surface nanoroughness).

AGM and corresponding Im values measured with three branched alkane probes, according to the procedure described before, are listed in Table The A200 silica egussa) and the Z175 sample (RhEjne Poulenc) do not exhibit size exclusion effects, leading to negligible AGMvalues and to IM values close to 1. Similar results were obtained for other pyrogenic silica samples, indicating their smoothness at a molecular scale. Conversely, on NipsilLP, significant size exclusion effects are observed that become very important for H-magadiite. Brendlt: and Papirer showed that AGM and do not vary with the measurement temperature, on A130. The same was observed with another pyrogenic(HDK N20) silica (Fig. As seen, all alkanes, linear and branched, fall on the same lines. The experiments were repeated, at 100" and 130"C, using a silica (H-magadiite) that is known to exhibit size exclusion effects (Fig. 7). The a-alkane~ycloal~ane lines get closer with increasing temperature of measurement, indicating a decrease of the size exclusion effects. However, AGM and parameters are temperature dependent on a rough silica surface. This observation indicates, firstly, that the morphology indices allow us to describe correctly the interaction capacity of linear, branched and cyclic alkane probes with smoothsilica surfaces. Secondly, it is seen that the differences in theadsorptionbehaviors of linear, branched, or cyclic alkanestend to diminishwithincreasingmeasurementtemperature. Inother words, on a rough surface, the energy of adsorption of a linear alkane decreases more rapidly than thoseof branched alkanes having the same morphological index. On a rough surface, a reasonable hypothesis is that the ease of insertion of alkanes intothe slotlike defects, located on thelateral surfaces of the HMagadiite crystal, decreaseswithtemperature,due to an increase of the steric dynamic radius of the adsorbing molecule. This steric hindrance will lead to a progressive exclusion effect of the n-alkane probes with increasing temperatures. However, branched or cyclic alkanes that are excluded from these slotlike structures are not affected as muchby the samesteric effect. Consequently, both typesof alkanestend to exhibit thesameadsorptionbehavior at highertemperature. Finally, this temperature dependence would explain that thesilica samples having the highest surface roughness exhibit the highest ysf.

T

Morphology Indices Determined DMHex

Samples A200 2175 Nipsil LP H-Magadiite

ay:/at

0.25 0.10 0.55 1.63

A G ~

-0.15 0.20 1.09 6.61

for Various Silica Samples TMPent A G ~

1.05 0.94 0.78 0.12

-0.1 1 -0.26 0.73 7.03

TeMBu ACI"

1.03 1.09 0.70 0.11

1.13 6.92

0.13

RT ln(Vn)

n-alkanes Cl Methyl2-alkanes 0 DMHex TeMBu

20

16 12 8 4

5

7

8

10

9

6 Variation of the free energy of adsorption of n-alkanes, branched alkanes, and cyclic alkanes on H-Magadiite.

In recent review on the acid-base properties of fillers and fibers that are used for the reinforcement of polymers, Belgacem adn Gandini [41] do not cite any paper concerning the determination of the specific interaction potential of the silica surface. our knowledge, there exist only a few studies This originates from the experimental difficulty of eluting polar probes, even at higher tempera-

130°C

7

8

Topology indices

Variation of the free energy of adsorption of n-alkanes, branched alkanes, and cyclic alkanes on H-Magadiite.

7.5 2.61 3.61 5.43 4.77

tures. Information is available only on somepyrogenic and precipitated silica samples. The specific interaction energies (Isp), determined in the way presented before and measured on different silica samples are reported in Table 5. The precipitated silica sample exhibits the highest even for acidic probes. This is related to its high functionality and certainly also to its higher mineral impurities content. For pyrogenic silicas, Isptends to decrease with increasingspecific surface area. The lowest values observed for the Cabosil-HS5T silica possibly result from the thermal treatment applied to this sample [46].

From the values, and knowing the acceptor (AN) and donor (DN) numbers of the injected polar probes, the acceptor and donor parameters (ISA, KD) of these silicas may be computed (Table 6). The precipitated silica has the strongest acidity because of its higher content of superficial silanol groups [SO], and also the highest basicity possibly due to the presence of mineral impurities (carbonates, etc.). All pyrogenic silica samples presentclose KA andKD values when taking into account the measurement errors. In fact, IGC at infinite dilution conditions is unable to deliver a complete description of the silica surface, since the very low number of injected probe molecules will preferentially be retained on the most energetic sites.

Only limited information on the surface heterogeneity of silica samples usingIGCFC available in the literature 151-531. Since most authors adopt different resolution methods for theintegral equation (9) and use different probes and experimental conditions, it is quite impossible to compare the published results. We shall therefore describe only results from our laboratory. From isotherms obtained by IGC, the specific surface areas and the BET constants may be readily computed (Table 7).

Specific Energies of Interaction

Determined at 110°C (kJ/mol)

Probes

A200 Kt 50

A1A130

CHZC12 14.1 CHC& 8.7 MeCN Acetone EthAc Ether 19.7

6.22.86 3.21

7.4 31.0

4.85

26.1

Cabosil-HS5

7.7 13.2 8.7

13.2 8.7

11.7

11.7

75

Main Characteristics of Fumed Silica Samples. Acceptor (KA) and Donor (KD) Parameters, Measured at 1 Using IGC-ID Silica A12.4 30 A150 2.4 A200 2.4 Cabosil-HS5 2175

SiOH/nn12

S (rn2/g)

130 150 200 310 175

KA

KD

Y

0.94 0.989 l .21 0.67 0.997 l .7 0.80 0.992 1.1 0.76 0.985 0.58 l .2 0.994 2.6

6.0

Given the doubt about the true values of the molecular sections of the adsorbates, one may consider neverthelessthat IGC leads to a correct estimation of the surface area of silica, taking as a reference the value obtained from nitrogen volumetric measurements, performedat 77 K. It is surprising that the T30silica exhibits a lower BET value than does the S13 sample using the Si2 probe. This is attributable to the surface roughness of the former that does not authorize the formation of a denseand well-organized adsorbed monolayerthat canbe easily formed on the flat S3 3 silica surface. Propane-2-01, the more polar probe, leads to a much higher value of the constant, testifying its ability to exchangestronginteractionswiththesurface silanol groups. The distribution functions of the energies of adsorption of n-heptane, benzene, and propane-2-01 on the S13 were computed using Balard's method 1541 and are displayed in Figure 8. The shapes of the distribution functions are most sensitive to the nature of the interactions that take place between the probe and the silica surface. On the one hand, propane-2-01, which is able to exchange strong interactions with the silanol groups, leads to the mostwidely extended distribution function exhibiting aclearcut bimodal structure. On the other hand, heptane, which is able to undergo only weak and nonspecific interaction, is quite insensitive to thelocal change of chemical surface f~lnctionality and leads to a quasi monomodal distribution function. What generates the bimodal structure observed for the propane-2-01 distribution function? The comparison between the adsorption energy distribution functions recorded on the bare silica surface and on a partially silylated (treated with tri-

SpecificSurfaceArea and BET Constant (forHexamethyldisiloxane (Si2) and Propane-2-01)

29 HDK S13 20 HDK T30 194HDK S13 Molecular section

122 Si2 270 Si2 139 Propane-2-01 Si2: 78

2

127 283 127 0 2

molecular section of propane-2-01: 35 A

apirer et al.

22

0.20

Benzene Propane 2-01

0.15 0.10

28

2210

16

40 Adsorption energy (Wmole)

Distribution functionsof the adsorption energies of n-heptane, benzene, and propanol on the sample.

~ethylchlorosilane)sample will give a first answer. Figure 9 shows the distribution function on S13 and on itspartially silylated derivative (19% of the total amountof silanol groups initially present being silylated). It is seen that the partial silylation leads to an increase of the intensity of the low-energy peak, whereas the numberof sites having the highest adsorption energies diminishes significantly. Hence, it is reasonable to attribute thefirst peak to surface domains containing mainly siloxane bridges, whereas the second is related to surface domains that are rich in silanol

0.12

I

10

16

3422 28 Adsorption energy (kJ/mole)

Distribution functions of the adsorption energies of propane-2-01 probe, measured at on bare and trimethyl silylated (S13TMS81)silica samples.

groups, having the highestenergy of interaction with propane-2-01 through hydrogen bonding. Si2, a bulky molecule, exchanges onlyweak specific interactions with the silica surface and will lead to distribution functions having a different shape. Figure 10 shows the distribution functions of the Si2 adsorption energies measured on S13 and the T30. It is seen that S13 is almost energetically “homogeneous” from the point of view of the Si2 adsorption, whereasthis is not the case for T30 silica that exhibits mainly a trimodal distribution function, with a major component appearing for the same energy valuethan on the S13. Taking into account the fact that Si2 exchangesmainly nonspecific interactions the variation of the sole surface functionality cannotexplainsuchapolymodality. Moreover, Si2 is a bulky molecule whose adsorption behavior will be sensitive to the surface geometry of the adsorbent. Consequently, the peak corresponding to the high adsorption energy (32 kJ/mo1) and the shoulder (at 30 kJ/mol), rnay be attributed to sites in which the probe may be more-or-less inserted: respectively, interaggregate and inter-primary-particle junctions. The main peakat 21 kJ/mol is related to the flat silica surface, the one that forms the majority of the S13 silica surface. Then, according to this assumption, thelow-energy peak rnay be attributed to the interaction of the probe with theentries of the micropores: locationsthat do not permit a good fit of the molecule with the solid surface and therefore lower its energy of interaction. Thisexampledemonstrates that ICC-FC is, according to thenature of the selected probe, very sensitive to the solid’s surface heterogeneity stemming from chemical and geometrical differences. Only a systematic study will lead to closer insight into the silica surface structure, aswill be demonstrated when studying the influence of a controlled silylation process on the surface property alterations of silica samples.

Oa20

10

I

\

28

Adsorptionenergy distribution functions of propane-2-01,measured at 53”C, on S13 and 13TMS81 silica samples.

The following exampleswill demolistrate how the couplingof an accurate determination of the surface energeticcharacteristics of treated silicas, through IGC techniques, and controlled surface modifications may contribute to a better physicochemical understanding of the silica surface, as well as the elucidation of the consequences of the modification upon the silica surface properties.

Heating is a most common procedure to remove adsorbed water molecules, or when going to higher temperatures, to activate the silica surface. The change of the surface functionality leads necessarily to a variation of the physical interaction potential of the silica surface. This variation is complex,amaximum of is recorded in the temperature range between 400 and 500°C [57,58], depending on the origin of silica. Three treatment temperature domains may be distinguished. The first, below 150"C, mainly corresponds to the departure of physisorbed water molecules, then between 150°C and 450°C, three-membered highly stressed polar siloxane cycles are formed leading to an increase of the values. Above these cycles rapidly transform into four-membered [59] nonpolar cycles: a transformation accompanied by a decrease of The third domain is attributed to a surface relaxation phenomenon allowed by the increasing mobility of the surface atoms [21]. The behavior of the crystalline H-magadiite is quite different [60]. An important drop of around 2OO"C, is correlatedwiththepartial crystalline structure destruction shown by x-ray diffraction. Above 400"C, the residue of the heated crystalline silica follows the trends evidenced with other amorphous silicas.

The esterification of the silanol groups with alcoholswill change the specific interaction capacity by lowering the density acidic silanol groups. Conversely, the relative density of siloxane groups, having a base character, will increase as longas no secondary siloxane opening reaction does take place. This is documented by the results of Wang et al. [48,49] shown in Table 8. The esterification of the silica surface leads indeedto a decreaseof the silica acid character and a relative increase of its base character. But, this effect is relatively weak in the case of methylation with methanol dueto the formationof new silanol groups resulting from a secondary reaction of methanol, leading to a siloxane ring opening. Esterification of the silica surface with hexadecanol,which is less reactive than methanol, strongly decreases the silica acidity surface, whereas its basicity does not change significantly. When using diols, one may expecteither an esterification involving a single alcohol group or a diesterification supposing that the silica surface structure will allow it. Such a surface should be rather smooth at the molecular scale and pre-

KA and KD Constants of Some Unmodified and Alkyl-Grafted Silica Samples Determined at "C, Using IGC-ID

samples Silica 2.6 RP 1 0.49 RPl-C1 RP1-Cl6 A1 30 A130-C1 A130-C16

Grafts

KA KD/KAKD

None Methyl 2.4 Hexadecyl None 1.4 Methyl Hexadecyl

0.47 0.26 0.38 0.35 0.05

492.7 49 49 49 49l .9

Ref. 5.7 9.2 3.22 4.0 38.0

senting adequately (randomly) distributed silanols. The diol chemisorption is simply achieved by heating together thesilica and the alcohol at200"C, for 2 h and extracting with chloroform the reagent excess [61,62].Using IGC-ID, the enthalpies of adsorption of alcohol probes were calculated from the variation, with the temperature of measurement, of free energies of adsorption. An alternation of AH values with the parity of the number of carbons of the diol grafts was evidenced on a pyrogenic silica (A130). This was explained by assuming that chains adopt a configuration on the flat pyrogenic silica surface. I3C NMR spectroscopy [63] showed that preferential diesterification was favoredfor odd diols, whereas even diol was chemisorbed by only one of the two terminal hydroxyl groups. For a precipitated silica, no difference between odd even grafted diols was observed. This is accounted for by the significant higher density of silanol groups on the one hand and by the molecular-scale rough surface, which does not allow an optimum interaction of the whole diol molecule with the precipitated silica surface, on the other. SilylatedSilicas A pyrogenic silica (S13 from Wacker-Chemie) was progressively silylated by reacting it with trimethylchlorosilane (TMCS). Thesilyation was performed by impregnating the sample with SO% (w/w) of awater-methanol1/1 mix andspraying variable amounts of silane in order to control the TMS coverage ratio [64]. The treated silica samples were dried at 300°C for 2 h. The characteristics of the modified silicas are reported in Table 9. Taking into account the experimental errors, the total number of surface groups (SiOH TMS) per nm2 remains quite constant: a proof of the absence of any secondary chemical reaction [6S]. The increment of adsorption free energy variation, per methylene group, the dispersive component of surface energy and the acid (KA) and base (KD) parameterswere determined on the set of controlled silylated S 13 silica samples (Table 10). The nonspecific capacity of interaction (y:') of the silylated silica samples decreases only slightly. The substitution of the low polarizable silanol groups by TMS groups, which present a similar interaction capacity, does not significantly change the ability of the silylated surface to interact through London interactions.

T

MainMolecularCharacteristics

Silica samplesa S13 .83 S13T81 0.49 S 13T73 0.56 .34 S13T61 0.96 S 1.09 3T44 l S13T41 S137-32 1.58

of S13TMSx and T30 TMSx SilicaSamples

zb

ng(N2)

(TMS/nm2)

(SiOH/nm2)

l 1.4821 l 24 1.12 41 0.81 47 0.75 52 0.59 68

0.0 1.97 .90 l 2.08 1.90 .09 1.981.23 .23 2.17

(TMS/nm2) 0.0 .83 0.49 0.56 0.96 l l 1.58

l

0

"The silica samples arequoted according to the following system: symbol of the initial silica followed by T (for trimethylsilyl graft) and x, which gives the percentage of residual silanol groups, determined by acid-base analysis with regard to the initial surface density of the unmodified silica. For example, 3T32 means a silica S13 silylated using TMCS, having a residual density of silanol groups equal to 32%. bThe relative coverage ratio? of the surface by TMS groups were calculated taking a TMS molecular section equal to 43 A

The KA and KD parameters were determined with basic (ether and THF) and acidic (chloroform and dichloro~ethane)probes (Fig. 11). remains practically constant, whereas KA reaches a plateau value, between 21% and 47%, and then diminishes sharply beyond a coverage ratioof 50%. It is difficult to admit that the acidity of the residual silanol groups could changeso suddenly. It may be possible that, on these samples characterizedby a high surface density of TMS groups,steric effects play an important role. Linear polydimethylsiloxane (PDMS)oligomers, such as hexamethyldisiloxane (Si2), octamethyltrisiloxane (Si3), decamethyltetrasiloxane (Si4) and dodecamethylpentasiloxane (Si5), which are relatively bulky molecules, areinteresting probes for

1" Main IGC-ID Characteristics of Controlled TMS Silylated Silicas Measured at 150°C AGrH2

Samples S13 S13T81 2.28 S13T73 2.17 2.19 S13T61 S 13T44 2.17 S13T41 2.13 S1 2.09 3T32

(mJ/m2) (kJ/mol) 0 21 24 41 47 52 68

2.52 0.73 0.69 0.67 0.57 0.52

KIA 25 21 19 19 19 18 17

KID

1.69 0.83 1.74 .76 0.68 1 2.69

1.4 1.3 1.3 l 1.2 1.3 .94 2.1 1.2 1.4

1.78 l

1.6

1.2

0.8

0.4

0

10

20

50

30

11 Variation, with the surface coverage ratio, of the acidity (KA) and basicity para~etersof silylated silicas.

Figure 12 displays testing the surface topology of the silylated silica samples the standard variations of free energies of adsorption of the PDMS oligomers on the silylated versus their number of monomer units. On the unmodified silica, only the first three oligomers could be eluted and the corresponding ““oligomer straight line,” obtained when plotting AG; versus the number of monomer units of the injected oligomers, stays well above all other straight lines corresponding to the silylated silica samples, that present lower capacity of interaction. But, an unexpected result appears, since silylated silica samples may nowbe classified in

-AGao (kJhole)

30

6

20

OS13

(0)

S13T81 (21) S13T61(41) S 13T44 (47)

10

3

Variation of the AG: of linear PDMS oligomerswiththenumber of monomer units nUMof the oligomer probes for the silylated silica samples (measured at The coverage ratios of the silica surface by the groups are indicated in parentheses.

two well-separated groups. The gap between these groups corresponds to a very slight change of the surface coverage by the groups, occurring at a coverage ratio of around This jump is evidenced when plotting the variation of the increment of free energy of adsorption per monomer unit versus the coverage ratio, as depictedin Fig. 13. After a slight decrease of AG!M when going from the initial unmodified silica to a silica having a coverage ratioof about which probably corresponds to the disappearance of the adsorption sites having the highest capacity of interaction, a jump of the surface properties is clearly observed around a surfacecoverageratio about SO%, aspreviouslyobservedforthe variation of the KA parameter. This evolution may be explained with a simple surface coverage model A random distribution of the TMS grafts on a surface made of a hexagonal bidimensional networkof sites, as earlier proposed by Sindorf and Maciel Figure 13 gives two examples of such model silylated surfaces for two coverage ratios equal to and It is seen that for a coverage ratio of only some holes remainin the TMS layer, whereas at the TMS groups form some islands scattered on. the still-free silica surface. In a chromatographic process, the progress of a molecule along the chromatographic columnis mainly ensuredby the carrier gas flow and by the staticdiffusion in the gas phase inside the porous solid particles. But, the adsorbed molecule may also migrate on the solid's surface itself. If the surface is energetically homogeneous, this phenomenon will have no influence on the measured retention time. Conversely, on a heterogeneous surface, the sites havingthehighest energyof interaction will act as potential wells, capturing the surface migrating molecule. Those sites will mainly contribute to the total durationof the residence time of the molecule, knowingthat this time is proportional to the exponentialof the energy of interaction of the adsorbed molecule. On the initial silica, the adsorbed molecule can freely migrate between its first site of adsorption and the site of desorption.

10

9

Low TMS

Percolation Threshol~ I

50

Variation of the increment of

80

90

100

with the coverage ratio, measured at

TMS grafts, introduced on the surface, will act geometrical barriers and will necessarily restrict the mobility of the molecule in the adsorbed phase. At low coverageratio,theadsorbedmoleculepercolates between the TMS grafts, but beyond critical value corresponding to the percolation threshold, the molecule may migrate only towards very a limited number of adsorptionsites. The threshold between mobile and localized adsorption processes corresponds to jump in the surface properties observed for the variations of the KA parameters and of the increment, AGYM, of free energy per monomer unit. Table l 1 displays the evolution of the specific surface area and BET constants of pyrogenic silica samples with increasing TMS surface coverage For the S13TMSx silica series, the specific areas, measured with Si2, decrease almost linearly with increasing coverage ratios and would approach zero for theoretical coverage ratio of 100%: a proof that, under these experimental conditions, the Si2 probe interacts only with thefree residual silica surface and provides convenient method to estimate it. The same trend is observed in the case of propane-2-01. But, the specific surface areas do not vary linearly with increasing coverage ratios. sudden decreaseis observed, corresponding probablyto percolation threshold for 50% surface coverage ratio. The small size of this probe and its strong capacity for specific interaction through hydrogen bonding, compared to Si2, is responsible forits sensitivity to the change of the surface topology. The BET constant for the Si2 probe decreases sharply up to coverage ratio of about 25% and then di~inishesslowly beyond this limit. The propane-2-01probefollows exactly the same trend. Its ability to exchange strong interactions with the bare silica surface leads to a very high BET constant value for the unmodifiedS13 silica sample. The variation of the BET constants, for both probes, indicates that ICCFC evidences sudden change in surface properties for coverage ratio of about 25%. The distribution functions of the adsorption energies of the Si2 probe on samplesaredisplayed in Fig. 14. This figure indicates significant change of the shapes of the distribution functions with increasing coverage ratios:

Main IGC-FC Characteristics, Measured at 150"C, of TMS Silylated Silica Samples Propane-2-01 (53°C) (45°C) Si2

SBET Sample 195 S13 46 99 S13T81 97 S13T73 25 92 S13T61 67 S 13T4417 66 S13T41 S 13T32

(m2/g)

0 21 24 41 47 52 68

CBET

139 123 104 103 70 17 63

29 16 14 14 12 11

56 41

17

40 e

10

16

22

Distribution functions of the adsorption energies of Si2 (normalized to a surface area of 130 m2/g), measured at 53°C on 3TMSx silylated silicas.

for the initial silica and that having coverage ratio of 21 monomodal distribution functions are obtained. For coverage ratio of and above this value, bimodalshapeswiththeappearance of apeak in thelow-adsorption-energy domain (around 16.5 kJ/mol) are observed. An explanation is offered using our simplified model. One has not onlyto consider noninteractive TMS sites and interactive sites corresponding to the uncovered or free silica surface previously, but to split these interactive sites into two categories: the sterically hindered sites, i.e., sites located between two TMS groups, and the free sites of adsorption according to the upper scheme of Fig. 15. For a coverage ratio of250/0, the sum of the number of hindered sites and TMS groups equals the number of free accessible sites. Beyond critical TMS surface concentration, the free accessible sites will only appear islands, scattered on continuous layer formed by bothTMSgroupsand the sterically hindered sites. Therefore, coveragevalue of 25% will correspond to another percolationthresholdallowing free mobility of theprobe,on longdistance range.Thiscannot be detectedinthe ICC-ID conditions,because of the tooshort residence time of the molecule on the surface. However, population of molecules, used in IGC-FC, will possibly be sensitive to this long-range mobility restriction. It is now interesting to verify if thedistributionfunctions of propane-2-01 confirm this analysis. The distribution functions of propane-2-01 (Fig. 16) are quite different from those ofSi2. Propane-2-01hashigh polarity and is smaller in size. Thepeak, centered on 17 kJ/mol,attributedtosiloxane bridges, is only weakly shifted whenincreasingthecoverageratio,possiblydue to steric hindranceincrease when increasing the surface density of TMS grafts. Conversely, the part of the function situated beyond 20 kJ/mol (in thehigh-adsorption-energydomain) is

1

2 Tvpes of adsorution sites

1 Free access Strong interaction Steric hindrance Weak interaction 12.5 Continuous free silica surface

Discontinuous free silica surface Discontinuous TMS surface coverage

First percolation threshold Sterically hinderedsites Free accessible sites

Different types of adsorption sites (free accessible and sterically hindered) and the evolution of their relative surface density with the surface coverage by the groups.

strongly affected by the controlled silylation process. expected by the blocking of the silanols, their relative contribution decreases, but also theshape of the functions changes when crossing the critical coverage ratio of 25%. Beyond this limit, a fine structure appears. The surface density variation of the residual silanols is not sufficient to generate a silanol peak in each domain delimited by the TMS groups. It seems reasonable to attribute this discrete structure to discrete distributions of the hydroxyl groups in each domain. The IGC-FCresults, using propane2-01 in place of Si2, confirms the existence of the first percolation threshold that takes place around 25% andcorroboratesthe interest of complementary approaches using both IGC-ID and IGC-FC for the exploration of the surface topology of silylated surface. This first critical coverage ratiois not observed with a precpitated silica, which is rugous at the molecular scale (T30 from Wacker-Chernie), due to the presence of crevices and micropore entries permitting the persistence of the long-range connectivity between the TMS beyond this coverage ratio on the silylated surface.

Silica and poly(dimethy1siloxane) interactions have already been examined using quantum chemical modeling[55,56]. Molecular modeling mayalso be considered as an ideal complement to IGCinvestigations, supposing that theinteractions between the probe moleculesthemselves are totally negligible compared to those established

28

40

Adsorption energy (kJ/mole) Distribution functions of the adsorption energies of propane-2-01, normalized to the surface area (130 m*/g), measured at 53°C on S13TMSx silylated silicas.

with the silica surface. In this instance, the aim is to demonstrate that basic molecular modeling techniques, providedby commercial software packages[67], choosingadapted force fields [68-701, may correctly describe and “sample”the adsorption phenomenaof different organic molecules on amorphoussilica surfaces, and allow us to reproduce the progressive coverage of the same surface by the probe molecules, up to the formation of the first monolayer, and, finally, give some new insights in the IGC processes. The first step is to establish a model, representative of a fumed silica surface, at the molecular scale. Our silica model is contained in a cube having a length of approximately 3.0 nm and a depth of about 1.5 nm. The surface is rendered amorphous through the use of simulated annealing of the crystal structure of B-cristoballite. The silanol groups are positioned randomly, with respect to their global amount, determined experimentally (three OH groups per nm2), their chemical nature (isolated, bonded, vicinal, and geminal), and respective contents determined by IR spectroscopy [71]. Furthermore, water molecules may be placed on that surface, so as to interact preferentially with the sites assumed to be the most reactive: the bonded silanol groups. These moleculesare then “relaxed,” allowing them to organize themselves into microaggregates on the surface element (Fig. 17). Thesurfacesampling strategy is baseduponacombination of molecular mechanics and dynamic techniques applied in a “consecutive loop” fashion. The probe molecule is first positioned at random on the silica surface and the interaction energy between the two partners is calculated. Then, molecular dynamics at high temperatureis applied to the probe fora given time, whilst the surface is kept

Stick and CPK representation of a fumed silica surface cluster. Adsorbed molecules are water clusters.

fixed, to force it to explore another part of the cluster. The probe is relaxed again and the probe-surface interactions evaluated. The (dynamics relaxation) cycles then reproduced as many times as necessary:20 to S0 cycles are generally required in order to correctly sample the silica cluster. At the end of the simulation, an average value of the adsorption energy of the probe is calculated. The quantumHSAB semiempirical methodof Pearson can be considered as a good complementof the surface samplingstrategy, even if both methods refer to very specific theoretical bases. Two molecular systems, whether acids or bases, with similar hardness, are morelikely to interact than two systems of different hardness. For a given molecule, this parameter can easily be calculated from the valuesof its frontier orbitalsby applying the following equation: Hardness EHOMO. Using this approach allows us to establish a hardness scale for probe molecules. The value of the “hardness” of the silica structural model will allow the prediction of the strongest probe-surface associations, without even getting the two components in contact. Initial calculations concerned the simulation of the adsorption of small organic probeslike a-hexane, l-hexene, and styrene. Table 12 summarizes the results obtained by applying the HSAB method. HSAB calculations show that the LUMO orbital of the silica cluster model is readily accessible (the value of -0.624 e\r indicating an “acidic” character for its surface). It can therefore accept electrons from the HOMO orbital of the probe molecules. The comparison of the respective hardness of the probes allows us to predict that the silica surface will preferentially interact with styrene. Indeed, the HOMO of styrene is the occupied orbital with the highest energy within the consideredgroup of probes.Comparatively,theinteractionswithl-hexene and nhexaneappear less favorable.These results are in entire agreementwiththose deduced from the surface sampling simulations which also predict that the silica

Hardness of the Structural Silica Model and Different Probe Molecules as Determined by the HSAB Method surfacea l-Hexene Styrene Molecular n-Hexane Silica system 3.365 -0.6241.166 -6.241 -9.261

ELUMO (ev)

0.017 10.025 5.60

2.81

Hardness (eV)

11.278

aOnly silanol groups and their neighborhood considered.

surface interacts, on average, much more favorably with styrene than with l-hexene and n-hexane. A more careful examination of the adsorptionof the different probes shows that hexane and hexene tend to adsorb in much the same way, with an approximately zigzag planar conformation at the silica surface. Styrene, on the otherhand,interacts in a very different way: thestrongest interaction occurs when the aromatic ring is nearly parallel to the silica surface and close to isolated silanol groups. The calculated distancebetween the centers of mass of the ring and the oxygen atom of the silanol group lies between 0.33 and 0.35 nm: a distance associated with strong interactions as shown by quantum mechanics calculations (73,741. Surface sampling simulation results and their comparison with experimental data [75] obtained from IGC-FC performed on a pyrogenic silica (A150 from Degussa) are reproduced in Table 13. Table 13 also reveals that, whereas the order of magnitude of the adsorption energies differ strongly when estimated either by surfacesamplingsimulations or by IGC measurements,the relative variations observed between the different probesfollowthesametrend.Onceagain,the adsorption potentialof styrene is much larger than forthe other probes,its adsorption energy being strongly displaced towards the higher-energy values (a difference of approximately 7.0 kJ/mol is observed with respect to the other probes). This specific feature tends to demonstrate that the presenceof an electronic delocalizationontheprobe molecule reinforces itsadsorptionontothe silica surface. Complementary calculations were performed in order to simulate the experimental

Characteristic Parametersof the Adsorptionof the Three Different Probes onto Our Structural Silica Model and Comparison with Relevant Experimental Data Mean interaction Mean energy site (surface energy Molecular sampling) association (kcal/mol) Silica/styrene Silica/ 6.8 1-hexene 7.9 Silicaln-hexane aReferred

-24.2 14.8-17.5 13.7

adsorption D energya (kcal/mol)

(IGC experiments) D energya (kJ/mol) (kJ/mol)

0 6.2 8.3

the value of the silica/styrene interaction energy.

21.6

0

measurements madeby IGC-ID 1761. Figure 18 shows the variationof the energyof interaction of a-pentane and l-pentene (calculated with a surface sampling strategy), versus the numberof simulation adsorption assayson thesilica surface model, carried out on 50 adsorption assays and various starting positions of the probes on the surface. It appears, as expected from IGC-ID, that l-pentene interacts more strongly than n-pentane with the silica surface. The average values of the interaction energies, Ei,of different probes were calculated from the preceding curvesand plotted versus the topology index (Fig. 19a) and conipared with the free energies of adsorption of the same probes, measured by IGC at infinite dilution conditions. Figure 19b displays the evolution of AG: with XI 1211. The ordersof magnitude of the adsorption energiesdiffer when estimated either by simulation or by IGC measurements. Nevertheless, the trends of the relative variations remain the same. The well-known ability of aromatic probes to interact more strongly with the silanol groups present on the silica surface is clearly evidenced. From the departureof the representative pointof the polar probes from the alkane line, the specific interaction parameters were estimated. The ratios of values of aromatic probes over those of alkenes are reported in Table 14. A good agreement is observed between predicted and IGC-determined ratios using nonpolar and polar probes. As discussed above, IGC-FC is able to assess the energetic heterogeneity of a silica surface. Thisheterogeneity is estimatedfromadsorptionisotherms and expressed as the variation of the number of sites corresponding to discrete values of adsorptionenergies of achosenmolecularprobe as thesurfacecoverage increases. It is therefore interesting to examinehowcomputersimulationcan help to understand how the energyof adsorption changes when increasing the

-4

-8

10

50

n-pentane n-pentene

Evolution of the energy of interaction with the silica cluster surface,of npentane and l-penteneversus the number of computer simulationadsorption assays.

Evolution of the adsorption energy of alkane, alkene, andaromatic probes determined by surface sampling simulations(a) and by IGC experiments with the topology index.

coverage ratios of the surface by the probe molecules. For the simulation of the silica’s surface progressive coverage by the molecular probes, up to the monomolecular layer ~ ~ r m a t i othree n , probes were selected having similar molecular section (around 50 one being apolar cyclopentane and the other two polar tetrahydrofurane(THF)and tetramethyl 2,2,4,4-tetrahydrofurane ( T ~ T H ~ ) . The comparison between the two latter should indicate how the steric hindrance of the four methyl groups, that form a cage around the basic oxygen group, may influence the interaction of the ether function with the silanol groups. To perform the “filling” of the silica cluster surface, we sampled the surface starting with one molecule in order to find its most interactive position. Thereafter, it was kept fixed on this most energetic adsorption site. second moleculewas then allowed to find its most interactive residual site following the procedure previouslydescribed. This processwasrepeated until thewholesurfacewascovered by probemolecules (about eight molecules for the selected silica cluster). The variation of the interaction energy Ei with the nu3ber of adsorbed molecules on a sampledsilica cluster surface area of about 400 is depicted in Fig. 20. The upper x-axis scale indicates the estimated surface coverage ratio by the probe molecules. clear-cut difference of behaviors is observed between the apolar probe for which the energy of interaction decreases monotonously with increasing surface coverage ratios, and the two baselike probes forwhich a sharp decrease is

Ratios of Values of Aromatic Probes Over Those of Alkenes Calculated from Computer Simulation and IGC Measurements benzene/&, alkenes IGC measurements Simulation

Isptohene/I,, alkenes

Ei (kcahhole)

F

l6

(~mo~(kJ/mol)

(a) Evolution of the energy of interaction of cyclopentane THF, and tetramethyl THF (TMTHF) with the numberof molecules adsorbed on the silica cluster and with the estimated coverageratio. (b) Experimentaldistribution functions of adsorption energies for heptane and propanol-2 measured on a fumed silica.

observed afterfixing four to six molecules, whichcorresponds to a coverage ratio of about 60%. This jump may be reasonablyattributed to the fact that the first adsorbed molecules are free to select and adsorb on sites having the highest local density of silanol groups, experiencing with them strong hydrogen bonds. When four molecules are adsorbed, that is to say about one molecule per nm2, the most interactive sites are now occupied or shielded by theadsorbedmolecules. Thereafter, the last arriving molecules may interact only with the remaining, less interactive sites. In other words, the steric hindrance due to thepreviously adsorbed molecules will contribute also to the decrease of the energy of interaction. This

latter hypothesis is supported by the stronger decrease observedin the interaction energy Ei in the case of the most bulky molecule(TMTHF) when comparing with the THF probe. Thesefindings are in line with recentresults concerning the adsorption of hexamethyldisiloxane on silica samples that were progressively grafted with trimethylgroups(TMS). As presented earlier, significant variation in surface properties is noticed when thesurfacecoverageapproached 50%, variation that was attributed to sudden changeof populations of adsorption groups, groups Of on free silica surface being predominant before a surface coverageof 50% course, molecule, like that is unable to have polar interactions, will explore the surface differently: the phenomenon observed leads to monotonous decrease of the interaction energyEi, corresponding to the progressive coverageof the sites having the highest polarizability. Simulation supports the fact that an apolarmolecule considers the silica surface as quite homogeneous, leading to quasimonomodal asymmetrical distribution function of the energy of adsorption, depicted in Fig. 20b, for heptane on fumed silica; whereas polar probe, like propanol-2, that exchanges strong hydrogen bonds with the silica surface, leads to bimodal distribution function, proof of the existence of two typesof adsorption sites on the-silica surface predicted by simulation. Combining computer simulation and IGC will definitely lead to rapid progress, as the near future will demonstrate.

IGC is indeed convenient and powerful surfaceanalysis method, in particular for the evaluation of the physical interaction potential of silica. This results from the high adaptability of IGC, combined with the possibility to readily perform measurel~ents atdifferent temperatures using large set of molecular probes chosen according to their known interaction capacities or to their bulkiness. Applying IGC-ID and cornparing the adsorption behavior of linear and branched or cyclic alkanes provides morphological information on the surface roughness of the silica surface. The acid-base character of a silica surface may be estimated by selecting probes of given acid-base characteristics. IGC applies fairly well to the examination of silica surfaces modified either by thermal or chemical treatments, not only for silica as described inthis instance, but for other oxides, fibers and so forth. Whereas experimentation is quite straightforward, the interpretationof the data is all but simpleand needs generally complementary information provided by modern physical methods. The surface of silica has been investigated for a long time, but obviously even such a “simple” surface generates new and fascinating questions and remains an object of studies for the coming years.

1. K. K. Unger, Packings ~ t u t i o n a ~Phases y for C h ~ o ~ u t o ~ ~ a p h i c ~ e c ~ M. Dekker, New York, 1990. International Patent PCT FR 9600463, to RhCine Poulenc, March 28 (1996).

3. E. Papirer and H. Balard, in Adsorption and Che~isorption rnorga~icSorbents (A. Dabrowski and Tertykh, eds.), Elsevier, Amsterdam, 1995, pp. 479-502. 4. A. Voekel. Critical Rev. Anal. Chem. 22:411 (1991). 5. J. Derminot, Physicochi~ie Polymdres etSurfaces par Chromatographieen Phase ANRT, Paris (1981). 6, G. Ligner, M. Sidqi, J. Jagiello, H. Balard, and E. Papirer. Chromatographia 29:35 (l 990). 7. E. Papirer, H. Balard, and A. Vidal. Eur. Polym. J. 24:783 (1988). 8. E. Papirer and H. Balard. J. Adhesion Sci. Technol. 4:357 (1990). 9. J. Jagiello, G. Ligner, and E. Papirer. J. Colloid Interface Sci. 137:128 (1990). 10. K. K. Her, The Chemistry of Silica, Wiley, New York, 1979. 11. H. E. Bergna, The Chemistry of Silica, Adv. Chemistry Series 234, ACS Washington,1990. 12. P. Legrand, The Surjace Properties of Silica, Wiley, Chichester, 1998. 13. P. Mukhopadhyay and H. P. Schreiber. Colloids Surfaces A 100:47 (1995). 14. J.R. Conder and C. L. Young, Physicochemical ~ e a s u r e ~ e n t by s Chro~atography,Wiley, New York, 1979. 15. F. M. Fowkes. J. Phys. Chem. 66:382 (1984). 16. G. M. Dorris and D. G. Gray. J. Colloid Interface Sci. 71:93 (1979). 17. C. Kemball and E. R. Rideal. Proc. Roy. Soc. A 287:53 (1946). 18. F. Ehrburger~Dolle,in TheSurface Properties of Silicas (A. P. Legrand,ed.), Wiley, New York, 1998, pp. 83-137. 19. H. Barthel, M. Heinemann,L.Rosch, and J. Weiss,in Proc. ~ u r o ~ l l e r95, s Mulhouse, France, 1995,pp.157-161. 20. A. P. Legrand, H. Hommel, Tuel, A. Vidal, H. Balard, E. Papirer, P. Levitz, M. Czernichowski, R. Erre, H. VanDamme, J. P. Gallas, J. F. Hemidy, J. C. Lavalley, 0. Barres, A. Burneau, and Y. Grillet.Adv.ColloidInterfaceSci. 33:91(1990). 21. E. Brendli: and E. Papirer. J. Colloid Interface Sci. 194:207 (1997). 22. H. Wiener. J. Physical Chem. 52:425 (1948). 23. H. Hadjar, H. Balard, and E. Papirer. Colloids Surfaces 99:45 (1995). 24. E. Brendli: and E. Papirer. J. Colloid Interface Sci. 194:217 (1997). 25. F. M. Fowkes and M. A. Mostafa. Ind. Engng. Chem. 17:3 (1978). 26. R. G. Pearson. Acc. Chem. Res. 26:250 (1993). 27. R. S. Drago, G. C. Vogel, and E. Needham. J. Chem. Soc. 93:6014 (1971). 28. V. Gutmann, The Donor-Acceptor ~ p p r o a c h t o ~ o l e c ~Interactions, lar Plenum, New York, 1978. 29. S. Spange, A. Reuter, and E. Vilsmeier. Colloid Polym. Sci. 274:59 (1996). 30. M. K. Kamlet, J. L. Abbout, M. H. Abraham, and R. W. Tafy. J. Org. Chern. 48:2877 (1983). 31. S. Spange and A.Reuter,in Proc. rnternationalSilica Conference, Mulhouse, France, 1998, pp. 305-306. 32. C. Saint Flour and E. Papirer. J. Colloid Interface Sci. 91:69 (1983). 33. J. Gregg and K. S. W. Sing, in Adsorption, SuTface Area Porosity, Academic Press, 1967, pp. 1-19.

34. M. Jaroniec andR. Madey, Physical Adsorption Elsevier, Amsterdam, 1988. 35. Rudzinski and D. H. Everett, Adsorption Gases Surfaces, Academic Press, London, 1992. 36. M.V. Szombathely, P. Brauer, and M.Jaronkc. Co~putational 13:17 (1992). 37. Jagiello.Langmuir13:1020(1997). 38. M.N. Nederlof, W. H. Riemsdjik,and K. Colloid Interface Sei. 135:410 (1990). 39. Roginski. C.R. Acad.Sci.USSR45:194(1944). 40. Rudzinski, Jagiello, and Y. Grillet. J. Colloid Interface Sci. 87:478 (1982). 41. M. N. Belgacem and A. Gandini, in rnterfaciu~ Phenonze~a in Chromutogrup~ly(E. Pefferkorn, ed.), Marcel Dekker, New York, 1999, pp. 42-118. 42.A.Khalfi, Papirer, H. Balard, H. Barthel, and M. Heinemann. J. Colloid Interface Sci. 184586 (1996). 43. A. Khalfi. PhD. Thesis, Univ. Haute-Alsace, Mulhouse, France, 1998. 44. E. Brendle and E. Papirer, J. Colloid Interface Sci. 194207 (1997). 45. Joachim, A. Vidal, and E. Papirer, in Chem~ca~ly ~ o d ~ Oxide e d Surfaces (E. Leyden and W. T.Collins, eds.),Gordon Breach, New York, 1990, pp. 361-374. 46. M. A. Zumbrum. Adhesion 181(1994). 47. Kessaissia, E. Papirer, and J. B. Donnet. J. ColloidInterfaceSci.82:526 (1981). 48. A. Vidal, E. Papirer, M. J. Wang, and J. B. Donnet. Chromatographia 23:121 (1 987). 49.E. Papirer, A.Vidal, M. J. Wang, and J. B. Donnet. Chromatographia23:279 (1987). 50. R. Iler, in The Chemistry Silica (R. Iler, ed.) Wiley, New York, 1979, p. 634. 51. M. Pyda, B. J. Stanley, M. Xie, and G. Guiochon. Langmuir 10:1573 (1994). 52.B. Stanely and G. Guiochon. Langmuir10:4278(1994). 53. M.Pyda and G. Guiochon. Langmuir13:1020(1997). 54. H. Balard.LangmL~ir, 13:1260 (1997). E. A. Nikitina, V. D. Khavryutchenkov, E. F. Sheka, H. Barthel, and J. Weis. Surface Rev. Letter. 4879 (1997). A. Nikitina, V. D. Khavryutchenkov, E. F. Sheka, H. Barthel, and Weis. in Proc. Silica 98, Mulhouse, France, 1998, pp. 467-470. 57. G. Ligner, A. Vidal, H. Balard, and E. Papirer. Colloid Interface Sci. 133:200 (1989). 58. E. Papirer, G. Ligner, A. Vidal, H. Balard, and F.Mauss, in C ~ l e ~ i c a l l ~ ~ o d i (E. Leyden and T. Collins, eds.)Gordon Breach, New York, Oxide 1990, pp. 15-26. 59. C. Brinker, R. K. Brown, D. R. Tallant, andR. J. irkp pat rick. Non-Cryst. Solids 120:26 (1990). 60. G. Ligner, A. Vidal, H. Balard, and E. Papirer. J. Colloid Interface Sci. 133:200 989). 61. M. Sidqi, H. Balard, E. Papirer, A. Tuel, H. Hommel,and A. P. Legrand. Chromatographia 27:3 (1989).

62. H. Balard, M. Sidqi, E. Papirer, J. B. Donnet, A. Tuel, H. Hornmel, and A. P. Legrand. Chrornatographia 25:707 (1988). 63. M. Sidqi, H. Balard, E. Papirer, A. Tuel, H. Hommel, and A. P. Legrand. Chrornatographia 27:31 l (1989). IV: 64. H. Balard, E. Papirer, A. Khalfi, H. Barthel, and J. Weis, in From to (N. Auner and J. Weis, eds.), J. Wiley, Weinheim, Germany (in press). 65. H. Balard, E. Papirer, A. Khalfi,and H. Barthel. Composite Interfaces 6:19 (1999). 66. G. W. Sindorf and G. E. Marciel. J. Phys. Chem. 86:5208 (1982). and MolecularModellingSoftwarePackages, 67. Insight Molecular Simulation Inc., San Diego, CA. Goddard, and W. M. Skiff. J. 68. A. K. Rappe, C. J. Casewit, K. S. Colwell, W. Arner. Chern. Soc. 114:10024 (1992). 69. H. Sun and D. Rigby. Spectrochirnica Acta A53:1301 (1997). 70. D. Rigby, H. Sun, and B. E. Eichinger. Polymer Int. 44:311 (1997). 71. C. G. Arrnistead, A. J. Tyler, F. H. Ambleton, S. A. Mitchell, and J. A. Hockey. J. Phys. Chern. 73:(11) 3947 (1969). 72. R. G. Pearson. J. Amer. Chern. Soc. 85:3533 (1963). 73. C. Vergelati, in '95, Mulhouse, 1995, pp. 151-154. 74. T. Suzuki, H. Tarnon, and M. Okazaki. Chern. Lett. 2151 (1994). 75. C. Vergelati, unpublished results. C. Vergelati, H. Balard, and E. Papirer, in Silica '98, Mulhouse, 76. F. 1998, pp. 861-865.

This Page Intentionally Left Blank

Institute of Physics, University of Silesia, Katowice, Poland

Introduction 11. Molecular Dynamics Simulation. A. A “c1osed”pore model “open” B. An Dynamics moleculesof pores in

243 of Molecules Model in Pore

244 245 246 247

III. Molecules in Cylindrical Silica Pores A. Cyclohexane at low densities B. C60/C6H12 mixture C. Sulfurhexafluoride D. Acetone

25 25 1 257 259 265

IV. Liquid-PlasticPhaseTransition V. Concluding

27 l

References

of Cyclohexane in Silica Cavities

274 275

Computer simulation is a useful method to study properties molecules not only in bulk phases but also in porous materials (for review see Ref, 1). Behavior of molecules trapped in a pore depends on the pore shape, features of the solid surface, and nature of surface interactions. These factors canbe relatively easily modified in the simulation of molecular dynamics (MD), which gives possibilities to investigate structural and time-dependent properties of molecules. However, dueto the complexityof the system (porous material with adsorbedmolecules) it is necessary to make certain assumptions about the shape of the pore and its surface, as

well as to model the molecules. In many computer simulations, pores are represented by slits between two parallel walls of structureless surfaces 12-71 or crystalline ones [8-9], whereas in others the pores are in the shape of a cylinder with smooth surfaces [IO-121. In all these studies, molecules are approximated by onecenter particles, and the MD calculations are limited to simulations of the translational motion. In a more realistic model used here the pore walls are irregular, as the solid framework is represented by the atomic structure of amorphous silica obtained from computer simulation. Moreover, amolecule is modeled by a rigid set of interaction centersthat gives the possibilityto describe rotational motion of the molecule in a pore. As adsorbed molecules we selected electrically neutral molecules (cyclohexane and sulfurhexa~uoride,which have only higher multipoles) and polar ones (acetone) that allow us to study the influence of interactions between silica surfaces and the molecules on molecular dynamics inside the pores. Another reason for selecting these molecules is available experimental and theoretical data, including results for molecular dynamics inside the pores of silica glass [13-181. In particular, the studies were stimulated by specific-heat data for cyclohexane in porous silica 13-1 S], Raman measurements of rotational relaxation of molecules in sol-gel glass [16], and measurements of diffusion coefficient of acetone in porous glasses [17,18]. The MD simulation methodis used to study structural and dynamic properties of the molecules inside the pore as a functionof pore size, surface intera~tion, density of adsorbent, and/or temperature. These studies were preceded by the MD sil~u~ations of unbounded systems [19-211, and comparisonof these results provides information on how the restricted geometry affects molecular behavior.

ecause of the limited number of adsorbed molecules which may be used in theMD simulation a pore model must be applied. Two models of the pore are proposed: 1. A “closed”pore in theshape of a cylinder-high symmetry of thepore ~elpfulin studying details of structure and dynamics of adsorbed molecules; in particular, the silica surface is relatively well defined and it possible to study behavior of molecules next to the surface. 2. An“open” pore-a cavityformed by the silica microspheres placed at the corners of a simulation cubicbox allows us to continuously change the density of adsorbed molecules and it is used to study liquid-plastic phase transitions. A basis to create the poresmay be the structure of amorphous silica, and such a structure was achieved using the MD simulation technique. Sixty-foursilicon atoms and 128 oxygen atoms were placed in a cubic box of length 1.44 nm resulting from the density of the system, 2.2 &/cm3. The initial configuration of atoms was a cristobalite structure with small random displacements applied to the lattice positions. Initial velocities were chosen randomly from a uniform distributionin such a way that they were consistentwiththerequiredtemperatureandtheresultant momentum of the whole system was zero.Interaction between two atoms was

described by a modified ~orn-Mayer-Huggins potential proposed by Feuston and Carofalini [22] in the form

where is the distance between atoms i andj, is the formalionic charge, A, and are parameters of the short-range repulsive term, and describes the range of the Coulombic interaction between two ions and qj resulting from screening of the charges. The potential parameters used in the M D simulation are collected in Table l [23]. To calculate forces acting on an atom,neighboring atoms from inside a sphere of radius 0.6375 nm were taken into account, and periodic boundary conditionsas well astheminimumimageconvention [24]were used.The Newtonian equations of motion were solved using a fifth-order predictor-corrector method [25]. In the calculations a timestep of l fs was used. The glass structure was simulatedusingthe melt-quench procedure [22]; the cristobalite structure was melted at 6000 K and then the system was gradually cooled, lowering the temperature by l 0 K every l00 time steps. Additionally, at the intermediate temperaturesof 5000 K, 4000 K, 3000 K, 2000 K, and l000 K the system was allowed to relax for 5 PS, and at the final temperature of 300 K it evolved for 15 PS. Such a procedure leads to “frozen” atom positions at random locations, resulting in different bond lengths and angles between the bonds.

The final structure of the bulk silica was repeated four times along the y , and axes, resulting in a cube of side 5.77 nm or four times along the and y axes and three times along the axis, giving a silica rectangular parallelepiped with dimensions 5.77 nm 5.77 nm 4.32 nm. To create a pore, all atoms from inside a cylinder with a diameter lying along the axis, were removed. We distinguish bridging oxygen atoms (bo), bonded to two adjacent silicon atoms, from nonbridging oxygen atoms (nhO) on the pore surface, attached to only one silicon. randomremoval of atoms may result in less than tetravalent co-ordination of some silicon atomsand trisilanol groupsonthe silica surface. To avoid these nonphysical results the silica surface was corrected by removing all silicon atoms on the pore surface bonded to less than four oxygens and groups consisting of silicon atoms bonded to three nonbridgingoxygens. This procedure gives a cylind-

Parameters of the

Si-0

Potential for Silica Atoms

rical pore with an average diameter and irregular surface. onb bridging oxygen atoms on the surface represent hydrogen groups. Assuming a smooth surface of cylinder one cancalculate the upper limit of surface concentrationof the nonbridging oxygen atoms, and the number of removed atoms allows one to estimate the volume of the pore. The calculations were carried out for pore diameter, ranging from 1.2 to 3.0 nm, and the surface concentration of nonbridging oxygens decreases from 8.5/nm2 to 7.7/nm2. These results agree with the experimental surface concentrations of OH groups for silica gels dried at 115°C [26].

To simulate the liquid-plastic phase transition of cyclohexane in porous silica the molecules were placed at nodes of a face-centered cubic (fcc) lattice, characteristic of the plastic phase. Becauseof the constant volume of a cylindrical pore and periodic structure of the molecules in crystal phase it is impossible to vary continuously the moleculardensity. Moreover, too-short distancesbetween some molecules and the pore surface may lead to large velocities and to the melting of the initial crystal structure. These problems maybe avoided in an inverse system when a silica microsphere is placed at the center of a simulation cubic boxfilled with the molecules, and due to the periodic boundary conditions the system is equivalent to the translated box with silica clusters at its corners (see Fig. 1)[27]. The space

Projection onto xy plane of the simulation box with silica cluster and its images resulting from periodic boundary conditions. The space between silica clusters at the corners of translated box, illustrated by the dashed square, represents a cavity of diameter (From Ref. Copyright 1992American Institute of Physics.)

between the silica microspheres represents a cavity of diameter d which may be estimated by

d=&L-d,

(2)

where L is the side of the cubic boxand d, is the diameter of the silica cluster. In the M D simulations 13 particles (one from the center of the box and its 12 nearest neighbors from thefirst co-ordination zone) from among 108 or 256 molecules were replaced by the rigid silica cluster of 1.2 nm diameter. The microsphere was cut from the bulk system and its size was chosen to avoid too-small distances between the molecules and silica atoms. For 95 (small system) and 243molecules (large system) thecavities have diametersof about 3nm and 5 nm, respectively. In studies of the phase transitions the molecular density was changedand it was necessary to adjust the box length The maximum changein the cavity diameter is less than 10% and this value is smaller than typical distribution of pore size in a sol-gel glass P81*

molecule in a pore interacts not only with like particles but also with atoms of the silica, and the interaction potential can be represented as

where i and j indicate molecules, and a denotes silica atoms. Molecules are modeled by rigid sets of Lennard-Jones (LJ) interaction centers, characterized by the parameters and located on atoms or unit sites (CH2 or CH3 groups), and if necessary fictitious partial charges, which reproduce the experimental dipole moment. Generally the interaction between the ith and $h molecule is defined by uij

where denotes interaction sites in the ith molecule. The whole of Eq. (4) is used SF6) the first for acetone molecules, whereas for nonpolar molecules term only is taken into account. The interaction centers and parameters of intermolecular interactions, together with molecular structure, for cyclohexane 19,271, fullerene 1291, sulfur hexafluoride [20,23,30], and acetone [21,31-331 used in the simulations reported here are collected in Table 2. The forces acting on moleculesand originating fromsilica are representedby the second term in Eq. (3). The interaction between ith molecule and silica atoms may be described analogously to Eq. but the sjthinteraction site should be replaced by the ath silica atom. It must be stressed that available experimental data for molecules in restricted geometries do not allow us to verify the parameters of the adsorbent-solid interaction, and their values remain uncertain. Therefore, in the calculations we used the simplest model of the interactions. For nonpolar mole-

Intermolecular Interaction Sites, the Potential Parameters, and Structure of the Molecules, as well as the Parameters for the Silica Interaction Centers Interaction Molecule site

(nm) 0.3851

Molecular structure

(K) q 61.8

CH,

-CH2 distance CH2 -CH2 angle

Inertia moments:

C

0.3775

30.0

0.270

70.5

0.375 0.296 0.391

52.9 105.5 80.6

0.164 nm 108.43' 194.3

98.5 170.2

Silica

bO nbO Si H

0.270 0.300

230.0 -0.533 230.0

kgm2

340.0 kgm2 CGH,, molecule in the chair form axis is parallel to the C, axis C distances 0.1398 nm 0.1455 nm Inertia moment 1.008 0.1565 nm S F distance Inertia moment 7.809 kgm2 0.12219 nm distance 0.300 0.15072 nm distance -0.424 CH3 1 17.2" 0.062 CH3 C CH, angle Inertia moments: 82.1 kgm2 kgm2 kgm2

All bonds are in (.x,y)plane and y axis parallel to C = O bond distance 0.095 nm -0.629 nbO 116.0" Si 0 H angle 1.283 0.206

cules, interactionswith oxygen atomsareincluded only, andthe potential parameters for silica oxygen are estimated from the ~ir~wood-Muellerformula using the polarizability, the diamagnetic susceptibility, and the van der Waals radius oxygen atom in zeolites The nonbridging oxygen on the surface may re~resenthydroxyl group, and the value CT for the nbO is slightly greater than for the bo-the collision parameter is similar to those used for oxygen in water molecules The small size and small polarizability of silicon atoms justify omission of interaction centers forthese atoms, Charges on silica atoms were calculated using thesemiempiricalmethod PM3 built intothe program for a silica cluster formed by 10 silicon, 28 oxygen, and hydrogen atoms. These calculations give also positions of the hydrogen atoms with respect to oxygen and silicon atoms on the surface, Thehydrogen atoms, attached to the nonbridging oxygens, are represented by charges only in similar manner as was done for water molecules The potential parameters for silica interaction centers are shown also in Table 2.

In the caseof the interactions between like particles as well as themolecule-silica interactions the LJ cross-parameters for different interaction sites are calculated using the Lorentz-Berthelod mixing rules 1241. The simulations concentrate on dynamics of molecules in pores and the silica frame is kept rigid. However, when hydrogen atoms on the pore surface are taken into account explicitly they are allowed to rotate around the Si-0 bonds with constant 0 -H distance and Si 0 -H angle. Orientation of a hydrogenatom is described by azimuthal angle with respect to the Si-0 bond. Initially, molecules are placed in nodes of an fcc or simple cubiclattice, or randomly inside a pore. At the beginning of the simulation, molecules are oriented randomly and molecular orientations are described by quaternions. The chosen translational and angular velocities are consistent with the required temperature. For an “open’’ pore the periodic boundary conditions are applied in three directions, but for the “closed” pore they are reduced to the direction of the pore axis. The forces and torques, acting on a molecules, which result from the LJ interactions, are calculated using the minimum image convention and a spherical cutoff of radius applied to the interaction centers. Moreover, the shifted force potential is used, which makes the force associated withthis potential go to zero outside the cutoffdistance. Value of the cutoff radius depends on the system under consideration, and for cyclohexane 55 nm, for sulfur hexafluoride 0.81 nm, and for acetone 1.35 nm. Spherical truncation is also used to the electrostatic interactions for acetone molecules. However, in that case the cutoff is applied to the center of mass of the molecules and electrostatic potentials are multiplied by a Gaussian switching function [24]. The switching functionfalls from 1 to 0 as either the distance between the center of mass of two molecules, or the distance between the center of mass of a molecule and a silica charge, increases from to For most cases the predictor-corrector method [25] is used to solve equations of motions: fifth-order of the algorithm for the Newtonian equationsand fourth-order for the Eulerianones. However,forrandomly distributed molecules in cylindrical porethe“velocity Verlet”algorithm is applied. In that case initial potentialenergymay bevery large and may result in large velocities and displacements, and this algorithm is convenient to implement constraintswhich prevent the molecules from leaving the pore. When a displacement of a moleculeis too large the moleculeis stopped at the pore surface and returned inside the cylinder. The direction and velocity are found from interactions with the nearest nonbridging oxygen atoms. The constraints are needed during the first 10 to 20 time steps when the molecules are nonuniformly distributed in the pore, and later they are removed. Additionally, in that initial procedure,a very short time step of the order offsisused and every 10-100 steps the time interval is doubled until it reaches the value applied in the main sirnulation run. Such a very short time step and its doubling maybe applied in any case when the initial state of the molecules is far from equilibrium and distances between interacting sites are very small. After that initial procedurethemain part of simulation is started,andthe equations of motion are solved with the same time steps as for bulk systems; 2.5 fs for acetone and 5 fs in the other cases. The first time steps constitute the equilibration run, during which velocities are scaled to obtain the required tern-

perature, and the system can come to an equilibrium state where instantaneous values of the potential energy and pressureoscillate about steady values. The next steps constitute theso-called production stage. When the velocity scaling is still switching on in the production stage the MD simulation is performed at constant temperature, i.e., in the canonical ensemble (NVT), and when no further adjustments are made in that phase the system relax free and the sirnulation is carried out for the microcanonical ensemble (NVE). Numbers of time steps and well the statistical ensemble used in calculations depend on individual case and will be specified later. the simulation advances, running averages for temperature, pressure, kinetic and potentialenergies, mean-square displacement,and others are monitored, andgive us information about stateof the system well proper working of the simulation program. During the production stage the configuration of thesystem, i.e., positions, orientations,and velocitiesof the molecules, is recorded every ten time steps. The molecular positions may be used to calculate single-particle density g(l)(r) and pair radial distribution function For the closed pore, the densityg(')(r) function of the distance, r, from the axis of cylinder is given by

and V(r,r denote the number of molecules and volume, where n(r, r respectively, between two cylinders of radii r and r The definition (5) can also be applied to calculate the molecular density function of distance from the center of the silica cluster for the open pore model; however, in that case one must consider spherical layers of thickness. Time dependence of the positions is used to calculate the mean-square displacements (MSDs). The MSD for an isotropic system, for molecules in open pores, is defined by

and the MSD components for translations along and across cylindrical the pore are given

respectively. Using orientations of molecules we calculated the rotational correlation functions

molecule and PI is the Legendre polywhere u(t) is unit vector bonded with nomial of rank When the vectorU coincide with dipole moment,e.g., for acetone molecule U is parallel to the axis, the correlation function Cl0(t)represents reor-

ynamics ~ i m u l a t i o ~

1

ientations of that moment which may be studied experimentally by the dielectric measurements. The correlation function describes rotational motion of the anisotropic part of Raman tensor for the"Alg mode of cyclohexane (L)3d symmetry) when U lies along axis, or Egmode of sulfur hexafluoride symmetry) when U is parallel to any twofold axis 1401. Other information about molecular motion may be obtained from the translational velocity correlation function (VCF),

as well as from the angular velocity correlation function, Cm(t),which is defined analogously to Eq. (8). In Eqs (6)-(8) the angular brackets indicate average values over all molecules and initial times Attheend of this section onemust discuss theproblem of thenumber of molecules in a pore. has been mentioned, the surface of a model pore has an irregular structure and it is difficult to estimate volume of the pore. As a consequence it is hard to determine the number of molecules to be set in the pore to obtain the desireddensity, Therefore, the following procedure is proposed. Initially, the numberof the molecules in a poreis calculated from theliquid density and pore volume assuming a smoothsurface, and a trial simulation is performed to estimate the pressure inside the pore. For an anisotropic system,e.g., molecules in a cylindrical pore, the pressure is a tensor quantity. Studies of the pressure tensor for particles in a slit pore [7,41,42] showed that the pressure component normal to the pore surfaceis almost constant across the pore, whereas the transverse component behaves similarly to the distribution of molecular density between the pore walls and oscillates around the value of the normal component. Moreover, these studies showed that the components of the pressure tensor can be calculated from the intermolecular interactions using the generalizedviral theorem 141,421, and the pressure in the pore maybe estimated as the average value of the diagonal elements of the pressure tensor determined as follows:

where is the number density and the pore volume. This value is usually higher than the pressure obtained from MD simulation of the unbounded molecules and the number of molecules in the pore mustbe reduced until the pressure in the pore matches the values for the bulk phase. For the systems presented here the maximum reduction in the number of molecules was 6%.

To study the formation of the monolayer and surface phenomena connected with that process we try to simulate filling a cylindrical pore of about 2.5 nm diameter and 4.32 nm length with cyclohexaneby gradual increase of the molecule number. We performed a series ofthe M D simulations for 293.3 K,when number of the

molecules is increasing from 2 to 60. The lowest number of the cyclohexane molecules corresponds to a density of about 0.013 g/cm3 and this value, estimated on basis of volume of the cylindrical pore with smooth surface, is larger than the density of vapor phase. Sixty molecules in the cavity gives a density of 0.396 g/ cm3 which constitutes about 50% of the liquid density. Initially, two cyclohexane molecules are placed on the axis of the cylindrical pore. Then M D sirnulation of cyclohexane molecules is carried out and after that the next two molecules are introduced into the pore. This procedure is repeated until the system consists of 20 molecules, and then the molecule number is increased by four. small numbers of molecules the microcanonical ensemble applied in the MD simulations gives very large fluctuations of temperature. Therefore, the simulations are performed at a constant temperature(N'VT), and 20,000 time steps and 60,000 time steps (300 PS). To improve the statistics the production phasesof the Simulations were performed five times for twoand four moleculesin the pore, four times for six to ten molecules, three times for 12-20 molecules and twice for the other states. The properties of cyclohexane in the cylindrical pore were studied as a function of surface density, i.e., number of the molecules perunit area of smooth surface of the cylinder, whichrangefrom 0.059 to 1.765 The single-particle density of cyclohexane, g(')(r), across the pore, whose examples are shown in Fig. 2 [43], gives us information about the structure of the molecules in the pore. For small densities of cyclohexane all the molecules are adsorbed on thesilica surface, but for 0.8 thecyclohexanesformthesecond layer. Tracingthepositions of molecules with respect to the cylinder axis as the simulation progresses one may find that there are only a few molecules in the second layer and for the highest

0.2

1.0

0.4

r (m) The single-particle density of C6HI2moleculesas a function of distance fromtheporeaxisforthesurfacedensities(inindicated. (From Ref. 43, Copyright Q 1994 Taylor Francis Ltd.)

density thereare four cyclohexaneson average. These observations meanthat most of the molecules are localized near the pore surface and the second layer is not a stable structure, since some of the molecules exchangetheir positions in the contact and second layer. Translational motion of the molecules may be characterized by the velocity correlation functions and mean-square displacements. Because of the pore symmetry the velocity vector for theith molecule at any time, t, can be separated into three components: parallel to the pore axis v/!(t), radial n(r~(t))[v’(t) azimuthal v ~ ( t ) v’(t) v?(& where vfis the velocity perpendicular to the pore axis, and the unit normal to the pore surface is a function of the position ri of the molecule. The VCFs Gi(t), and G$(t), for the corresponding velocity components, are presented in Fig. 3. Almost the same shape for the VCFs G!(t) and G$(t) suggests that translations tangential to the pore surface are hindered in a similar way for anydirection of motion. The tangential translations are much more free than the radial ones for which the VCFs, G:(t), show deep negative parts attributed to a strong cage effect. For the molecules localized near the poresurface, the surface potential causes the radialdeflection of a molecule from its equilibrium position to result in strong forces, reversing the direction of the molecular translation. In other words, radial motion of the molecules is strongly restricted and the molecules are “rolling” on the surface. For larger densities when collisions of a molecule with the other molecules become more frequent the VCFs decay faster. The for translations along the pore behaves similarly to that in the nonrestricted system and it increases almost linearly with time (see Fig. 4). The displacement across the pore cannot be larger than the cylinder diameter, and the perpendicular is a convex function of time approaching a constant value at

1

1

(PS) Velocity correlation functions C!(t) (solid lines), C t ( t ) (dashed lines), and (dotted lines) for the surface densities (in nm“2) indicated. (From. Ref. 43, Copyright 1994 Taylor Francis Ltd.)

25

ro

m-

4

(PS)

Mean-square displacements for translations along, and across the solid and dashed pore, A r i ( t ) , forthesurfacedensities(inindicated. curves represent results obtained from the MD simulations and Eq. (lob), respectively. (From Ref. 43, copyright 1994 Taylor Francis Ltd.)

long times. These results may be explained quantitatively by considering a simple model based on the diffusion equation. In the studied case almostall cyclohexane molecules are moving within the contact layer, and one may neglect the radial motion. Moreover, the VCFs suggest that the diffusion coefficient of a molecule in the monolayer is direction independent. Therefore, translations of a molecule may be treated as diffusion motion on the surface of an infinitely long cylinder of radius R. Assuming that the initial distribution density is uniform,solvingthe diffusion equation, the parallel and perpendicular components of the MSD are given by

The above equations are not valid at very short times where theparticle moving almost freely. Using the diffusioncoefficient, D , estimated fromAri(t) and radiusR obtained from the maximum of the radial density functiong(')(r)one maycalculate Ar:(t). The theoretical MSDs Ar:(t), compare Fig. 4, are in goodagreement with thesimulation results, particularly for 0.8 nm"2. Differences between the

theoretical and simulated data observed for higher densities are result of the molecules moving in the inner part of the pore. Because behavior of those particles is different from that of the molecules in the contact layer the model cannot work properly for high densities. The diffusion coefficients calculated from slopes of the Avi(t), which were used to reproduce the Ar:(t), are presented in Fig. 5. There are also shown values obtained by integration of the VCFs and they match well to the diffusion coefficients calculated from the Up to 0.8 when the molecules begin to form the second layer, one observes rapid decrease in the molecular mobilityand additional molecules introduce much smaller restrictions in the translational motion. To characterize reorientational motion of we calculated the rotational correlation function for unit vector along the axis, and corresponding relaxation time, zZo.Examples of the rotational relaxation function are presented in Fig. 6, and the correlation times as function of the surface densityof cyclohexane are depicted in Fig. 7. For the lowest density, molecular rotations show features typical for free rotor; however, the correlation function C2,(t) tends to zero at long times what is result of hindrances originated from the surface potential. A few additional cyclohexane molecules in the pore restrict strongly molecular rotations and cause drastic change of the correlation function shape. For 0.1 18 nm-2 initial decay of the function is still Gaussian, but at long times C20(t)shows almost exponential decay. The faster relaxation rate of the C20(t) observed for larger is result of the molecules from the second layer, which have more freedom to rnove than the surface molecules. This behavior illustrated by changes of therotationalcorrelationtimewithsurfacedensity of cyclohexane. Initially,

0.0

FIG. 5 The diffusion coefficients calculated from the Ari(t) (full circles) and the VCFs Ct(t) (open circles) as a function of the surface densityps. (From Ref. 43, Copyright 01994 Taylor Francis Ltd.)

U

Rotational correlationfunctionsfor four surfacedensities:0.059 nm"2 (solidline),0.1 18 nm-2(dashedline),0.354 nm-2 (dot-dashedline), and 1.765 nmW2 (dotted line).(From Ref. 43, Copyright 1994 Taylor Francis Ltd.)

2.5

1.5

0.5 1 P,

Rotational relaxation time as a function of the surface density. (From Ref. 43, Copyright 1994 Taylor Francis Ltd.)

p,

hindrances of the rotational motion increase strongly with then they are almost constant, and when the second layer appears, 0.8 nm- they diminish slightly. The above results indicate that the adsorption of cyclohexane on silica surface is nonlocalized, and the adsorbed molecules can move on the surface of the pore. Although the cyclohexane-silica forces are stronger than the cyclohexane-cyclohexane forces, the interactions between the adsorbed moleculessignificantly affect the dynamics of the molecules in the monolayer. When the surface densityexceeds 0.8 molecule per nm2 few of the cyclohexane molecules are pushedout from the contact layer. The molecules form the second layer and their motion is not hindered the motion of the surface molecules.

In the previous section, it has been shown that the parallel and perpendicular to the pore axis can be described by the onediffusion coefficient. However, we considered special case when almost all molecules were localized near the pore surface. In many studies the cavity space was completelyfilled with molecules, and the Einstein formula wasused to estimate diffusion coefficientsof the confined moleculesfrom their fortranslations parallel and perpendicular to the pore walls [4,9,44,45].At short times, but not too short when the particle is moving almost freely, the increase almostlinearly with time and the Einstein formula suggests smaller valueof the diffusion coefficient for motion across the pore than that for motion along the pore. But the displacement perpendicular to the pore walls can never be larger than the poresize (slit width or cylinder diameter) and the MSD at long times must approach constant value. These features of the MSD across the pore were shown by Allen and Masters [46] who sinmlated dynamics of particles between two reflecting walls. Their calculations, based on the diffusion equation, show that the perpendicular to the pore can be described by the diffusion coefficient for an unbounded system, i.e., for translations parallel to the pore walls. In a real pore, surface interactions reduce the value of the diffusion coefficient for motion along the pore; however, onemay suspect that the diffusion constant characterizes translations regardless of the direction of motion. To illustrate the behavior of the of the molecules in restricted volume we performedthe MD simulation of mixture of one fullerene molecule and 112 cyclohexane molecules in the cylindrical pore of2.5 nm diameter and 4.32 nm length. The choice of the system should allow us to study simultaneously translational motion of large molecule with diameter of 1.1 nm) and small one (C6HI2with diameter of 0.6 nm). In particular, long-time behavior of the displacement perpendicular to the pore axis for the two molecules shouldbe different. The number of cyclohexanemoleculescorresponds to the liquid density at 293 K. Initially, the fullerene molecule was placed at the center of the cylindrical pore and the cyclohexanes werelocalized at nodes of the fcc lattice. The simulation was performed in the microcanonical ensemble, with 20,000 and 80, 000 time steps (400 PS). The MSDs along, Ari(t), and across, the pore for fullerene and cyclohexane molecules, calculated from the simulation results, are presented in Fig. 8. For both molecules the parallel to the pore axis are similar-they increase

(PS)

Mean-squaredisplacementsofthefullerene and cyclohexanemolecules. Solid and dashed lines denote results obtained from the MD simulation and Eq. (1 l), respectively. (From Ref. 29, Copyright 01994 Taylor Francis Ltd.)

linearly with time, however, slopesof the functions are different. At the same time, the perpendicular Components seem to be different not only quantitatively but also qualitatively. The results may be interpreted by solving thediffusion equation for particle in an infinite cylinder of radius R.Assuming that the walls of the cylinder reflect the particle and the initial distribution density is uniform, one can calculate the components of the which have the following forms: (1 1 4

AUi(t)

cl,,

In Eq. (1 lb) is the nth root of the equation (x)/dx 0, where Jl(x)is the esse1 function of order 1. The equations 1) are not valid at very short times where inertial effects are dominant and although Eq. (1 b) gives the slopeof A r f ( t ) at zero time equalto it cannotbe used to estimate thediffusion coefficient. The diffusion constant canbe calculated from the slopeof Ari(t) using Eq. (1 la). Its values obtained from simulation results for the cyclohexaneand fullerene molecules in the pore are presented in Table 3. The values are smaller than those obtained for an unbounded system, which illustrate hindrances in molecular motion originating from the surface poten-

T Diffusion Coefficients, D,and Effective Radii, for Cyclohexane and Pullerene Molecules in the Pore and in the Unbounded System

Molecule

D(lo”* m2/s)

(nm)

Pore C6H12 c60

Bulk system C6H12 c 6 0

5.28 1.12

0.95 0.10

9.37 2.56

tial. It must be stressed that thediffusion constant for cyclohexane in the pore is an average value over all molecules. The mobility of cyclohexanes which are moving within a layer in contact with the pore surface is strongly restricted by the silica potential. The motion of the other cyclohexane molecules is perturbed by surface interactions to different degree depending on the distance from the surface and their dynamics is additionally affected by forces originatedfromTherefore, to be able to use Eqs (1 l), which involve the position-independent diffusion coefficient, the considerations presented here are limited to the simplest case when the as well as the diffusion coefficient are pore-averaged quantities. For the fullerene molecule, becausesingle particle motion is involved, the and diffusion coefficient may be treatedaspositionindependent. is surrounded by cyclohexanes and its position is “oscillating” around the pore axis where changes of the silica potential are weak. Therefore, the motionof is determined mostly by interactions with C6HlZmolecules and one may assumethat the fullerene molecule is moving in a constant silica potential. The diffusion equation was solved for the point particle and to estimate Ar:(t) one needs an effective radius of the pore, $teff. The pore radiusis reduced by radius of the molecule, and for fullerene it is additionally reduced by the diameter of the cyclohexane molecule,i.e., by thickness of the contactlayer formed by cyclohexane molecules. The values of $teff are collected in Table 3. Applying the data from Table 3, Eq. (1 lb) gives theoretical MSDs for translations across the pore which are presented in the Fig. 8. Although the modelis very simple, the diffusion coefficients obtained from the parallel to the poreaxis allow us to reproduce the main features of the perpendicular MSDs for C6Hi2and and agreement between the theoretical and simulation results is quite satisfactory. Moreover, seemingly different results for fullerene and cyclohexane molecules can be described within the same model.

In this section we study the influence of the pore size and density of liquid SF6 on the molecular dynamics. We considered cylindrical pores with diametersof 1.2, 1.7,

and 2.3 nm, and length 5.77 nm, and liquid densities 1.5, 1.7, and 1.9 &m3, which correspond to the experimental densities of bulk SF6 [47,48] at temperature 296 K for pressures84, 3 l 1, and 1167 bar, respectively. The number of SF6 molecules in a pore depend on the pore diameter and the densityof adsorbent, and its values for the systems considered here arecollected in Table 4. To study the effect of surface interactions on the behavior of SF6 molecules the calculations were repeated for density 1.5 &m3 using a shallow potential of silica oxygenatoms with the potential well depth &/kB 58 K, the valueused for simulation of solid oxygen [49].The M D programsimulatesa system inthemicrocanonicalensemble (NW)wherethe number of molecules, the volume, and the energy are fixed, and the equilibrium run consists of iVeq 1000 time steps and the production stage 6000 time steps (30 PS). The single-particle density of SF6 molecules across the pore, presented inFig. 9, shows a layered structure of the liquid, and thickness of a layer is about 0.45 nm. The shallow surface potential does not introducequalitative changes in the density profiles; however, fewer molecules are localized nearthesurface and they are weakly attracted by silica. The dependence on liquid density is clearly observed for the pore of 2.3 nm diameter. Distribution of molecules in the inner part of the pore becomes more uniform when decreases, but near the pore surfaceisitalmost unchanged. The last observation suggests that the contact layer forms a stable structure, what was confirmed by tracing the positions ofmolecules across the pore as the simulation progressed. For thecylinder of diameter 2.3 nm in the contact layer there are 55-65% of all molecules, whereas for 1.2 nm all molecules are the surface molecules. Contrary to the surface molecules which are localized at the pore surface, a large fraction of the molecules from the centralpart of pore exchange their positions between the inner layers. Therefore, the molecules can be divided into two groups: the surface molecules and center ones. The time correlation functions for translational and rotational motion inside the pore were calculated for all molecules as well as separately for the surface and center nlolecules. Shapes of the correlation functions aresimilar as those for molecules in liquid phase and hence they are not presented here. However, one must note that confinementof molecules in the pore leads to restrictions of rotations and translations and the restrictions are greater for the surface molecules than for the

Number of Liquid Density (g/cm3)

Molecules as a Function of Pore Diameter

and

Number of SF6 molecules d (mm)

1.2 1.7 2.3 Source: Ref.

1.5 40 81 145

1.7 45 91 165

1.9 51 102 185

.o r (m) The single-particle densitiesof molecules across the pores function of the distance from the axisof the pore. The solidand dashed linesare for the deep and shallow surface potential, respectively, and density l g/cm3. The dotted line is forthedeep potential and density1.9 g/cm3. (From Ref. 23, Copyright 1991 American Institute of Physics.)

center ones. These changesin motion of SF6 molecules are described quantitatively by the diffusion coefficients for translations along the pore, D,rotational correlation times of second rank, and the angular velocity correlation times, The dependencies of the quantities for all molecules on the pore size and density are presented in Fig. 10 where data for 0 were obtained from theMD simulation for bulk liquid The values are averaged over all particles in the pore andresult from contributions of the surface and center molecules. In Fig. it is seen that these contributions are also pore diameter dependent. For the surface molecules changes of D, and are very weak, which indicates that environment of these molecules, and in consequence their interactions with the surface and neighboring particles, are almost poresize independent. In the caseof the center molecules,their number depends on the pore diameter, and this number increases with pore size. Moreover, the dynamicsof the center moleculesis perturbed by surface interactions to a different degree dependingon the distance from thesurface. These two factors explain nonlinear pore size dependence of D, and for the center molecules. However, the changesof D with pore diameterare much stronger than those for and This observation suggeststhat interactions originated from molecules from theimmediateneighborhoodhavemain influence onrotationalmotion,and molecule separated from the silica surface by molecular diameter reorients almost in the same manner in bulk liquid. Contrary to reorientational motion translations involve more neighboring particles, and the correlation range is greater. The depthof the fluid-solid potential hasan influence on values of the dynamics parameters, but it does not change the character of their dependence on the pore

5 g L

p

(m-')

The total diffusion coefficients,D,and rotational, and angular velocity correlation times, tu,as a function of pore size for three densities: 1.5 g/cm3 (circles), 1.7 g/cm3 (squares),and 1.9 g/cm3 (triangles). The solid and dashed lines are for the deep and shallow surfacepotential, respectively, and they are to guide the eye.(From Ref. 23, Copyright 1991 American Institute of Physics.)

size. For the shallow potential, one observes increase in the diffusion coefficient, which is caused by the increase in the diffusivity of both the surface and center molecules. ~imillishingof the well depth of surface potential leads to less pronounced differences in rotational motion of the surface and center molecules: the surface molecules rotate more freely, whereas rotations of the center molecules are interrupted slightly stronger. This behavioris caused by more uniform distribution of the molecules in the pore and confirm our earlier conclusion that rotations of SF6 are mainly affected by the molecules from the first co-ordination zone. In Figs 12 and 13 [50], the rotational correlation functions, C20(t), andcorrelation times, and are compared with the experimental data obtained from pressure Raman measurements of the mode of SF6 molecule in sol-gel pores of average pore diameter2.3 nm [16]. In the experiment thesol-gel sample was placed into a high-pressure cell where gaseous SF6 was introduced, and the pressure was increased from 84 to 1167 bar, which corresponds to the bulk liquid density in the range 1.5 to 1.9 g/cm3. Analysisof the results for the experimental densities shows that agreement between theoretical and experimental data is poor.Although

L

g

/d (m-‘)

Thediffusioncoefficients, D,and rotational, andangular velocity correlation times as a function of pore size for the center (full symbols) and surface molecules (dotted symbols). The lines and symbols have the same meaning as in Fig. (From Ref. Copyright 1991 American Institute Physics.)

changes of the theoretical and experimentalcorrelationtimeswithdensity are similar, their values are different and the difference indicates that the simulated rotational motion of molecules in the pore is more hindered than reorientation of themoleculesinthe real system.Theshallowsurfacepotentialonly slightly diminishesthediscrepancies between the theoretical and experimental results, which suggests that restrictions in the simulated reorientation arise mainly from too-strong intermolecular interactions used in the MD simulations. However, the parameters of the intermolecular potential seem to have been chosen properly, because using them in simulation of the bulk we obtained very good agreement with available results of NMR and Raman measurements Another factor which strongly influences the strength of the interactions is the distance between the molecules. In the experiment the pressure in the Raman cell only was monitored, and itwas assumed that pressure and the corresponding densityin the pores were the same as those outside the sample.It is very probable that this assumption is incorrect, and the density of liquid in the pores is smaller than the value estimated from the density-pressure relation for bulk fluid. This idea is supported by the studies on the partition coefficient, defined as the ratio of the adsorbent

1.o

F;

U

0.4

0

1

2

3

Experimental rotational correlation function of obtained at 296 K and pressure 11’76bar for which density of bulk liquid is 1.9 g/cm3 (solid line), and the simulated correlation functions for densities 1.9 g/cm3 (dashed line) and 1.4 g/cm3 (dottedline).(AdaptedfromRef. 50, Copyright 1992,withpermissionfrom Elsevier Science.)

concentration in the poresto its concentration in the bulk phasein equilibrium with it,forhardspheres in spongematrices [51]. These studies showed that at low pressures the partition coefficient is greater than 1, whereas at high pressures the coefficient is lessthan 1, which was attributed to geometrical restrictions preventing the same degree of packing of molecules inside the pores as outside of the sample. Takingintoaccounttheaboveconsiderations,additional MD simulationsfor densities 1.0, 1.2, and 1.4 g/cm3, lower by 0.5 g/cm3 with respect to those used in experiment, were carried out. In Fig. 12 it is shown that new rotational correlation functions fit the experimental ones much better. In consequence the correlation times comparefavorablywiththeexperimental results, see Fig. 13. Therefore, one may conclude that, for a given pressure, the density of liquid inside the pore is lower by 20-30% than that in the pure liquid phase. The conclusion is also supported by analysis of the experimental results for sulfur hexafluoride inside porous sol-gel glass 1201 and those for the bulk phase [47) at various pressures and temperatures. Direct comparison of the correlation times for the two systems leads to inconsistent conclusions:(1) the timesz20for SF6 in the pores have almost the same valuesin as the pure phase, suggesting the lack of influence of geometrical restrictions and surface interactions on molecular rotations, whereas (2) the times for different thermodynamic states indicate that SF6 molecules in the pores rotate more freely or more hindered. However, when one assumes reduced densities inside the pores, similarly as in Fig. 13, both the times and show that molecular reorientationin the pores is slower than in the bulk phase.

c A

A

A-"

0.1 1.0

Rotational relaxation times, rZO,and angular velocity correlation times r,, for SF6 molecules in the pore of diameter 2.3 nm as a function of density. The full and open circles denote results obtained from the MD simulation for the deep and shallow surface potential, respectively. The full and open triangles indicate experimental data for densities resulting from pressure-density relationship for bulk SF6 and for densities lowered by 0.5 g/cm3, respectively. The lines are to guide the eye. (Adapted from Ref. 50, Copyright 1992, with permission from Elsevier Science.)

Most of the MD simulations were limited to electrically neutral particles or molecules, such as C6HI2and and in this section we try to study the influence of electrostatic forces on acetone molecules confined in the pore.To evaluate the role of electrostatic interactions on the structure and dynamics of the confined liquid, different potentials are considered: theCD approximation in which the charges on silica and polar acetone molecules are taken into account, the ND approximation in which silica is electrically neutral and acetone molecules have dipole moments, and the NN approximation where silica atoms and acetone molecules remain electrically neutral. Although acetone without a dipole moment is not aphysical object, this procedure allows us to compare dynamicsof molecules in the pores for different interaction models which are independent of the molecular shape. The simulations were carried out at constant temperature 298 K, i.e., in the canonical ensemble (NVT), and density 0.786 g/cm3 corresponding to the atmospheric pres-

sure. We considered the cylindrical pores with average diameters d 1.5 nm, 2.0 mm, 2.5 nm, and 3.0 nm, and length4.32 nm, and the numberof acetone molecules in the pores ranges from58 to 240. The calculations were performed also for polar and nonpolar acetone in the bulk phase. The equilibration run has 20,000 time steps and production stage 60,000 time steps (150 PS). Dependence of the single-particle density of acetone across the pore on pore diameter and surfaceinteraction is shown in Fig. 14. Thelayeredstructure of acetone is clearly visible, and similarly as before, tracing positions of molecules we were able to divide them into the surface molecules which reside within the contact layer and the center molecules which move inside the innerpart of the pore. When the electrostatic interactions between the molecules and silica are turned on (the CD model), the surface molecules are strongly attracted by the silica, and hence theyget closer to the poresurface and distances between the center molecules are larger than thoseforthe N D and N N interactions. Small differences also appear between thedensityfunctionsforthe ND and N N interactions, In the first case the dipole-dipole interaction between acetone molecules causes the molecules to become tightly packed and the contact layer is further away from thewall than in the case of the N N model, i.e., when the molecules interact only through the potential. Additional insight into the structure of acetone in the pore is given by orientations of the molecules with respect to the pore surface and the distribution functions foracetone interaction sites withrespect to hydrogens, andnonbridging oxygens, on the pore surface. In Fig. 15, there is shown an example of distribution functions of angles between the molecular axis (see structure of acetone molecule in Table 2) and the normal to the pore surface for theCD interactions For the center molecules the distribution functions are almost uniform,

r (m)

The single-particle density of acetone as function of the distancefrom the pore axis. The solid, dashed,and dottedlines are for CD, ND, andNN interactions, respectively. (From Ref. 33, Copyright 1996 American Institute of Physics.)

ynamics 6

Distribution functions of angles between the x,y , and axes of the acetone molecule and the normal to the pore surface for the pore of diameter 2.5 nrn and the CD interactions. The full and open circles are for the surface and center molecules, respectively. (From Ref. 33, Copyright 1996 American Institute of Physics.)

which indicates that those particles do not demonstrate any preferential orientation. For the surfacemolecules the most probable orientationof the y axis, the axis parallel to the dipole moment, is that perpendicular to the pore surface, and consequently the x and axes are set parallel to the surface. When molecule-silica electrostatic interactions are turned off, no orientational ordering for both the center and surface molecules is observed, and the distribution functions are very similar to those observed for the center molecules in the CD model. The radial distribution functions ggD(r), and for oxygen, carbon, and methyl groups of acetone molecules are very similar for all pore diameters, and in Fig. 16 only the functions for 2.5 nm are resented. The sharp first peaks for the acetone oxygen in the functions and indicate that acetone's oxygens are localized close enough. to surface hydroxyl groups to form hydrogen bonding (521. The maxima of these functions for oxygens appear at 0.207 nm and 0.305 nm, respectively, and taking into account that the 0 - H distance is 0.095 nm it is clearly seen that the bond is almost linear. The position of the carbon's maximum in 0.292 nm, and the length of the C= 0 bond in the acetone molecule, 0.123 mm, allow to estimate the angle between the hydrogen bonding H. *Oand the y axis of the acetone molecule, whichis about 120". This most probable orientation of acetone with respect to the hydroxyl group explains the sequence in which the maxima of the functions are observed, first for oxygen followed by close peaks for methyl groupand carbon.When the molecules interact with thesilica only through the potential thefirst peaks of the all acetone sites in the functions and show similar heights, and the broad peaks for carbon.

1.0

r (m) Radial distribution functions of acetone'soxygen(solidlines), carbon (dashed lines), and CH, groups (dotted lines) with respect to the surface hydrogen, ggD(r), and nonbridging oxygen,gED(r), ggD(r), and for the pore of diameter 2.5 nm. (From Ref. 33, Copyright 01996 American Institute of Physics.)

atoms in these functions are the most distant from the nonbridging oxygens. These results indicate that for the N D and N N interactions any orientation of the molecule with respect to the nonbridging oxygen is allowed. Transport properties of acetone molecules in the pores are characterized by diffusion coefficients which were calculated from correlation function of translational velocities parallel to thepore axis separatelyforthesurface and center molecules.Differentintermolecular and surface interactions influence absolute values of the diffusion coefficients, and therefore it is convenient to discuss the diffusion coefficients withreference tothe correspondingvalues in thebulk phase. In Fig. 17, there is shown pore size dependence of the ratio of the diffusion coefficients in the pores to thebulk system values: 3.05 10"' m2/s and 3.30 lo-' m2/s for polar and nonpolar acetone, respectively. When all electrostatic interactions are taken into account (the model) the surface molecules are strongly attracted by the surface and their motion is very restricted. The interactions betweenmolecules and silica are weaker and the surface molecules havemore freedom (the and N N model); however, dipole-dipole interactions between acetone molecules restrict their motion, and when these forces are removed (the N N model) the diffusion coefficient increases slightly. Considering similar interactions for the center molecules, particularly the long-range electrostatic forces, we expected similar dependence, but forthese molecules the diffusioncoefficient values decrease. One maysuggest two possible explanationsof these results: (l) the surface molecules can screen charges placed on the silica atoms and/or the effective pore diameter for the center molecules is larger than in the case when the surface interactions are represented by the potential, which corresponds to the densities

2 1.o

0.4

0.2

The total diffusion. coefficients (diamonds), diffusion coefficients for the center (circles), and for the surface molecules (squares) function of pore size. The full, dotted, and open symbols are for the CD, ND, and NN interaction models, 1996 respectively.Thelines are to guidetheeye. (From Ref. 33, Copyright American Institute of Physics.) depicted in Fig. 14. The totaldiffusion coefficient, which is an average value overall particles, is lower for stronger interactions, which indicates that it is determined mainly by the behavior of the surface molecules. The significance of the screening effect may be evaluated by analysis of molecular rotations, and we calculated the rotational correlation functions of the first rank for each molecular axis. Examples of the functions for y axis, i.e., for the dipole moment, are shown in Fig. 18. The correlation functions for the surface molecules decrease slower than those for the center molecules. Particularly, very long orientational correlationsof the surface molecules are for theCD model, and at long times the correlation functions havestill a large, nonzero value. Therefore, the rotational relaxation times zlowere calculated for the center molecules only, and their changes with the pore diameter areillustrated in Fig. 19. The correlation times change moderately with the poresize; however, the exceptionto this trend is the axis, which changes drastically when electrostatic interactions between the molecules and the pore surface are included. When the pore diameter decreases, the fraction of the center molecules whose motions areaffected by the surface interactions increases and consequence the relaxation timesalso increase. The strong influence of the electrostatic forces originating from silica on rotations of dipole moments of the center molecules indicates weak screening effect of the silica charges by the surface molecules. When the electrostatic molecule-silica interactions are ignored (the ND interaction) the changesof the relaxation times with pore size are alike, and the correlation times for the and z axes show an effect of tight packing of thecentermolecules,whichcausestherotations to become more restricted than for the interaction of type CD. When all electrostatic interactions

1.5 1

0.5

0.0

15

10 0

5 time

Rotational correlation functions of the dipole moment for the surface (solid lines)and center (dashed lines) molecules in the pore of diameter 2.5 nm. (From Ref. 33, Copyright 1996 American Institute of Physics.)

n

P

0.0

0.1

0.2

0.3 0.4 0.5 l/d (m")

0.6

Rotational correlation times for the x (squares), y (circles), and z (diamonds) axes as a function of the pore size. The symbols and lines have the same meaning inFig. 1'7. (From Ref. Copyright 1996 AmericanInstituteof Physics.)

are canceled, molecular rotations change completely. The weaker intermolecular interactions lead to the decreaseof the absolute valuesof the rotational relaxation times, and different relations between the correlation times, zlzc rlZfor the polar molecules and zlZ for nonpolar ones, distinguish theaxis with the dipole moment and indicate a special influence of the electrostatic interactions on reotations of that axis. The above results show that when the electrostatic potentials are included in the molecule-silicasystem interactions the surface molecules formhydrogenbonds with hydroxyl groups on the pore surface, leading to strong restrictions of their translations, and almost frozen rotationsof dipole moments. At the sametime, the surface molecules get closer to the pore surface, and the volume accessible to the central molecules increases. As a consequence, the center molecules diffuse faster than for the system where molecule-silica interactions are defined only by the potential. The results may be helpful in interpretation of the experimental data of molecular diffusion in porous glasses reported by Koone et al. [17,18]. In the experiment, macroscopic transport of different liquids in silica pores of 2.9 nm diameter was studied; the sol-gel samples were initially presoaked and then the diffusion of deuterated molecules was measured. It was found that hindrances of molecular motion for polar molecules, such as acetone or acetonitrile, were smaller than for nonpolar liquids, e.g., cyclohexane. MD simulations show that surface molecules are localized near the pore surface regardless of the interaction model, and incorporation of electrostatic interactions betweenmolecules and silicagives larger values of the coefficient for the center molecules. One may suspect that in the experiment the contact layer was formed during the presoaking, and mostly diffusion of the center molecules was measured.

Experimental studies [l 3-151 showed that confinement of molecules in pores leads to depression of the freezing temperature, and it was found that this effectis inversely proportionaltothepore size. However,forsmallpores of diameter lower than 8 nm it was very hard to detect the liquid-solid transition 14,151, and we used the MD simulation method to examine theliquid-plastic phase transition of cyclohexane in small silica cavities of diameters of 3 and 5 nm. Dynamics of cyclohexane molecules at temperatures ranging from 190 to 333 K, where bulk cyclohexane exists as a plastic crystal and as a liquid (temperature of the liquidplastic phase transition at atmospheric pressure is 279.82 K), was simulated using conventional methods in the microcanonical ensemble. Densities for the plastic phase were obtained from available data at 193, 195, 263, and 268 K [53,54], and the liquid densities were the same as those reportedby Sun et al. Each simulation started from the fcc structure, characteristic of the plastic phase, and theability of the system to melt the initial configuration was studied. The system was equilibrated during 2000 time steps, and then relaxed through thenext 8000 steps (40 PS).

The distributions of cyclohexane molecules around the silica show molecular layers, and similarly as for the cylindrical pores, we were able to divide the molecules into surface and center molecules. Further considerations refer to the center molecules whose behavior determines whether the systemis a fluid or a plastic crystal. Examples of the pair radial distribution functions for the center cyclohexane molecules are presented in Fig. 20. Temperature-induced evolution of the structure of cyclohexane in both cavities show that peak intensities decrease, become broader and shifted slightly towards increaseddistances. For the large cavity (d 5 nm) the distribution functions showfeatures of the liquid phase for temperatures above 235 K, but when the temperature decreases to 217 K, and below, peaks appear at distances characteristic of the fcc structure.For thesmall system(d 3nm) in the whole temperature range, 190-333 K, the distribution functions do not show any positional order of the molecules in the cavity. Additional calculations carried outfortemperatures 1’75 and 150 K gave a crystalline structure for the center molecules at the lowest temperature only; however, noncoincidence of the third co-ordination zone with that for the fcc lattice suggests that surface interactions affect the crystalline configuration. The time and VCFs calculated from the simulation results have typical shapes, and here we only point out the main features of these functions. At low temperatures the MSDs tend to almost constant values and the velocity correlation

K”

r (m)

Pair radial distribution functions of the center molecules in the cavities of diameters 3 and 5 nm. The short vertical lines indicate coordination zones for an ideal fcc lattice. (From Ref. 27, Copyright 01992 American Institute of Physics.)

functions show negative parts attributedto strong cageeffect. Molecules oscillate around mean values and deflection of a molecule from its equilibrium position results in strong forces reversing the translation direction of the molecule. At higher temperatures the slope increases with temperature,which indicates that molecules diffuse. Greater mobility of molecules leads to the dynamic nature of a local cage formed by the nearest neighbors and the cage effect is less pronounced in the velocity correlation functions. These qualitative results are supported by the diffusion coefficients calculated from the slopesof the MSDs, and in Fig. 21 their for the centermolecules in the cavities are compared with those obtained fro simulations for bulk cyclohexane[19]. At higher temperatures, above 280 K, behavior of the centermolecules isalmost the same as in pure cyclohexane,and poresize dependence of the diffusion coefficientis rather small. Below this temperature, unlike the bulk phase where valuesof L) are almost zero, the diffusion coefficients diminish monotonically with decreasing temperature; however, for molecules in the cavity of 5 nm diameter, a sharp drop of the diffusion coefficient appears between 235 and 217 K. The diffusion coefficients for cyclohexane in thelarge cavity (d 5 nm) correspond nicely to the pair radial distribution functions. These results agree well with the melting temperatures obtained from differential heat capacity measurement for cyclohexane in porous Spherosil of pore diameter from 8 to 125 nm [15]. In Fig. 22, experimentaltemperaturesfortheonset and peak of melting endothermsareextrapolatedtoporediameter of 5 nm, and thetemperature range where the solid-liquid transition should appear coincides with that obtained from the MD Simulations. In the case of the small cavity (d 3 nm), a small diffusion coefficient of the center molecules is observed even at 150 K, indicating

200

150

300

T (K) Temperature dependence of the diffusion coefficients for the center cyclohexanemoleculesinthecavities of diameters 5 nm (full circles) and 3 nm (open circles). The solid lines represent the MD simulation results for bulk cyclohexane. (From Ref. 27, Copyright 1992 American Institute of Physics.)

220 200 0.00

0.05

0.10

0.20

l/d (m") Experimental melting temperatures of cyclohexane in porous materials: Spherosil (circles) 1151, the controlled pore glasses (triangles) and Spherisorb (squares) [l31 as a function of the pore size. The open and full circles denote temperatures for the onsetand peak of the melting endotherm, respectively. The vertical line indicates the temperature range in which the MD simulations predict the liquidplastic phase transition for the cavity of diameter 5 nm.

supercooled liquid in the pore, and this result seems to be at variance with the ordered structure of cyclohexane for that temperature. However, the presence of crystalline defects, suggested by the radial distribution function, may provide a mechanismthrough which amoleculecanundergotranslational diffusion. It must be noted that the highly defective nature of cyclohexane crystallites forming in porous silica was discovered by neutron diffraction 131. The lack of conclusive results concerning cyclohexanein the small pore correspondsto the observationsof Jackson and McKenna [l41 who did not detect any melting endotherm for cyclohexane in the control pore glass of mean pore diameter 4 nm. At the end of this section one must note that confinement of cyclohexane in the cavity also introduces restrictions in molecular rotations; however, the rotational correlation time, increases smoothly with diminishing temperature. This observation indicates that at lower temperatures,wherethepairradialdistribution functions and diffusion coefficients forcyclohexane molecules in thecavity of diameter 5 nm show "frozen" translational motion, the molecules rotate, and the transition may be identified as the liquid-plastic phase transition.

In this chapter we have described applications of the MD simulatio~method to study behavior of polyatomic moleculesinside the silica pores of irregular surfaces. It is hoped that the results achieved from the MI) simulations are sufficient to

demonstrate thecapabilities of the method for characterizationof molecular structure and dynamics ofmolecules in restricted geometries. Theinformationand insight concerning microstructural details and local dynamics generated by MD often are not obtainable by other means. On the other hand, one must remember that real porous media are hard to characterize and may have acomplex and tortuous pore geometry. Thus, becauseofidealized shapes of the model pores, thesimulation results maynot be directly applicable to experiments;however, computer simulations are able to elucidate and clarify the interpretation of the experimental data.

1. W. Steele. Chem. Rev. 93:2355 (1993). 2. I. K. Snook and W. van Megen. J. Chern. Phys. 72:2907 (1980). 3. Y. Antonchenko, Ilyin, N. N. Makovsky, N. Pavlov,and P. Sokhan. Mol. Phys. 52:345 (1984). 4. J. J. Magda, M. Tirrell, and H. T. Davis. J. Chem. Phys. 83:1888 (1985). 5. J. P. R. B. Walton and N. Quirke. Chem. Phys. Lett. 129:382 (1986). Bitsanis, S. A.Somers, H. Davis,and M. Tirrell. J. Chem.Phys.93:3427 6. (l 990). 7. M. Lupkowski and F. van Swol. J. Chem. Phys. 93:737 (1990). 8. M. Schoen, D. J. Diestler, and J. H. Cushman. J. Chem. Phys. 87:5464 (1987). 9. M. Schoen, J. H. Cushman, D. J. Diestler, and C. L. Rhykerd Jr. J. Chem. Phys. 88:1394 (1988). 10. S.-H. Sub and J. M. D. MacElroy. Mol. Phys. 58:445 (1986). 11. A. Z. Panagioto~oulos.Mol. Phys. 62:701 (1987). 12. B. K. Petersen, K. E. Gubbins, G. S. Heffelfinger, U.M.B. Marconi, and F. van Swol. J. Chem. Phys. 88:6487 (1988). 13. J. C.Dore, M. Dunn, T. Hasebe,and J. H. Strange. ColloidsSurfaces36:199 (1 989). 14. C. L. Jackson and B. McKenna. J. Chem. Phys. 93:9002 (1990). M. Malhotra. Phys. Rev. B 44:4296 (1991). 15. R. Mu and 16. L. Nikiel and T. W. Zerda. J. Chem. Phys. 93:8464 (1990). 17. N. Koone, Y. Shao, and T. W. Zerda. Opt. Appl. 24:55( 1994). 18. N. Koone, Y. Shao, and W. Zerda. J. Phys. Chem. 995676 (1995). 19. Brodka and W. Zerda. J. Chem. Phys. 97:5669 (1992). 20 Brodka and W. Zerda. Mol. Phys. 76:103 (1992). 21. Brodka and T. W. Zerda. J. Chem. Phys. 104:6313 (1996). 22. B. P. Feuston and S. H. Carofalini. J. Chem. Phys. 89:5818 (1988). 23. A. Brodka and T. W. Zerda. J. Chem. Phys. 95:3710 (1991). 24. M. P. Allenand D. J. Tildesley, C o ~ ~ ~ t e r S i m ~ l aLiquids, t i ~ n Clarendon, Oxford, 1987. 25. C. W. Gear, ~ u ~ e rInitial i c Value P r o b l e ~ sin Ordinary ~ ~ f e r e n t~quations, i~l Prentice-Hall, Englewood Cliffs, 197 1. 26. R. K. Iler, The Chemistry Silica, Wiley, New York, 1976. A. Br6dka and W. Zerda. J. Chem.Phys.97:5676(1992).

28. T. W. Zerda, W. L. Vasconcelos, and L. L. Hench. J. Non-Cryst. Solids 121:143 (1 990). 29. A. Brodka. Mol. Phys. 82:1075 (1994). 30. M. T.Dove and G. S. Pawley. J. Phys. C. 165969 (1983). 31. W.L.Jorgensen, J. M. Briggs, and M. L. Contreras. J. Phys.Chern.94:1683 (1990). 32. R. Nielson and L. Pierce. J. Mol. Spectrosc. 18:344 (1965). 33. A. Brodka and W. Zerda. J. Chem. Phys. 104:6319 (1996). 34. H. Margenau and N. R. Kestner, Theory ~ ~ t e ~ ~ o l e cForces, u l f f r Pergarnon Press, Oxford, 1969. 35. A. G. Bezus, A. V. Kisielev, A. A.Lopatkin, and P. Q. Du. J. Chern. Soc.Faraday Trans. 2 74:367 (1978). 36. A. Kisielev and P. Q. Du. J. Chern. Soc. Faraday Trans. 2 771 (1981). 37. A. Rahrnan and F. H. Stillinger. J. Chern. Phys. 55:3336 (1971). 38. W. L. Jorgensen, J. Chandrasekhar, J. D. Madura, R. W. Irnpey, and M. L. Klein. J, Chem. Phys. 79926 (1983). 39. ~ y ~ e ~ C 3.for ~ e~ ~ i n d o ~Autodesk s, Inc., 1992. 40. L. A. Nafie and W. L. Peticolas. J. Chem. Phys. 57:3145 (1972). 41. J. P.R.B. Walton, D. J.Tildesley, and J. S. Rowlinson. Mol. Phys. 48: 1357 (1983). 42. J. R. Henderson and F. van Swol. Mol. Phys. 51:991 (1984). 43 A. Brodka. Mol. Phys. 83903 (1994). 44. E. Boek, W. J. Briels, J. van Eerden, andD. Feil. J. Chern. Phys. 96:7010 (1992). 45. W. J. Ma, J. R. Banavar, and J. Koplik. J. Chern. Phys. 97:485 (1992). 46. M.P. Allen and A. J. Masters. Molec. Phys. 79:435 (1993). 47 T.W. Zerda, J. Schroeder, and J. Jonas. Chern. Phys. 75:1612 (1981). 48. J. DeZwaan and Jonas. J. Chern. Phys. 63:4606 (1975). W. Oxtoby. J. Chern. Phys. 77:2153 49 V. Guissani, D. Levesque, J. J. Weis, and (1982). 50. A. Brodka and T. W. Zerda. J. Non-Cryst. Solids 139:215 (1992). 51. L. A. Fanti and E. D. Glant. AIChE J. 35:1883 (1989). 52. C . C. Pirnentel and A. L. McClellan, The Hydrogen Bond, Freeman,San Francisco,1960. 53. R. Kahn, R. Fourrne, D. Andre, and M. Renaud. Acta Cryst. B 29131 (1973). 54. E. Schudt and G. Weitz, ~ a n d o ~ t - ~ ~ r ~ ~ t e i i 2 Data ~ u ~ and e ~ i ~unctional cal ~ ~ l a t i o n s hin~ sScience and Technology, NewSeries, Croup 111, Sa, Springer, Berlin, 197 T. F. §un, J. A.Schouten, N.J. Trappeniers, and S. N. Biswas. J. Chern. Therrnodyn. 20: 1089 (1988). I)

Applied Geology and Geochemistry, boratory, Richland, Washington School of Earth and Atmospheric Sciences, Georgia Institute of Technology, Atlanta, Georgia

I. Motivations for Investigating

Silica-Water Interfaces

11. Model of the Silica-Water Interface A. Structure of a silica-water interface B. Electrical double layer of the silica-water interface C. Ionization and surface charge

277 278 279 280 28 1

111. SurfaceStructures and Properties A. Properties of water at the interface B. Thickness of structured water at the interface

284 284 285

IV. Evidence for Water Behavior at Silica Surfaces from Dissolution Studies A. Models based upon changes in near-surface water properties B. Models based upon solid properties C. Response to modifying effects of electrolytes

288 289 289 290

V. ConcludingRemarks References

292 293

The extensive occurrence of Si02 in crystalline and amorphous forms has motivated numerous investigations in the earth and materials sciences. In natural systems,vastdeposits of quartz in sedimentaryenvironmentsmake up important reservoirs of oil, natural and fresh water. In some cases, the secondary formation of silica scales from the minerals hosting these environments curtails oreven prevents the extraction of economically important fluids. Biogenically mineralized silica depositsconstituteanenormous reservoir in theglobalbiogeochemical

cycling of silicon [l]. Silica is also of broad economic interest as a commercially importal~tmaterialfordiverseapplicationsincluding optical waveguides,hight e ~ p e r a t u r ethermal insulators, laboratory products for corrosive and high-temperature environments, illumination devices, and specialty components (e.g., space shuttle windows) [2]. Silica sols, gels, clathrasils, and zeosils are used extensively for industrial processes, taking advantage of these materials’ high cationexchange potential and sorption properties [3]. In the cement industry, amorphous silica, SiOz(am), is commonly used as anadditive to enhance the performanceof concrete [4]. Becauseofextensive exposure ofsilica to workersduringsomeprocessing operations, a number of groups are focused on environmental investigations of risk factors forsusceptibility to lung diseases, such assilicosis and silicotuberculosis [5]. Silica exposures have also been implicated in a variety of other medical maladies, such as renal disease, autoimmune deficiencies, and cancer An understanding of controls on the behaviorof this end-member oxide is also important to broader issues involvingthe large family of silicate minerals and glasses. The fundamental unit, the Si-0 tetrahedron, is central to the complex borosilicate glasses that willbe usedworldwide tocontainnuclearwastes [7]. Similarly, this understanding is applicable to the largest constituent of the Earth’s crust-the broad class of silicate minerals. A central issue in each of these systems is the reactivity of Si02 in contact with aqueous solutions and the response Si- 0 bonds to a variety of chemical stresses such as solution pH (i.e., bulk solution vs. interfacial), te~perature,and solute composition. More specifically, reactivity can be quantified in terms of the kinetics and mechanisms by which Si 0 bonds rupture in these mineral-solution environments. This relatively simple oxide-water interface iswidely recognized to have physical and chemicalproperties which are uniquefromthebulksolution or solid (e.g., see other contributions inthis volume). It these properties that govern what we refer to as “reactivity” inthecontext of dissolution (durability) and sorption processes. Yet, we will show that a full understanding of the microscopic properties of this dynamic environment continues to be elusive. Thepurpose of this reviewis to summarizethecurrentstate of knowledge regarding our understanding of SiOz-water interfaces. At the risk of some overlap with other contributions to this volume, our discussion begins with abrief summary of spectroscopic and modeling studies. From this basis, we present inferential evidence for local-scale structure and properties of solutions at silica interface provided by recent studies documenting the effects of solute chemistry on silica reactivity.

Recentadvances in quantifyingthe reactivity ofsilica and perhaps all mineral phases emphasize the role of surface structure and chemistry.Silica surface chemistry is extensively discussed elsewhere ([8,9 and references therein). For the purposes of this review, a brief discussion introduces ter~inology andconcepts useful to subsequent discussion of solute effects on interfacial water structure and reactivity of Silica-Water interaction processes.

One means of understallding the structure of water at silica surfaces to examine the surface complexesthat form at thequartz-water interface. The stepwise development of an interface from a freshly cleaved fragment to a fully wetted silica surface is discussed by Parks [lo]. When quartzis fractured in vacuum, the freshly exposed initial surface (Fig. 1.1) is composed of highly reactive underco-ordinated silicon and oxygen bonds. evidenced by their ability to rapidly scavenge water under ultrahigh-vacuum conditions, these “sites” and the corresponding surface structures arehighly unstable and quickly hydroxylate. Water molecules are rapidly dissociated at the surface with protons bonding withoxygen and hydroxyl radicals with Si atoms. These surface species can be detected by a variety of spectroscopic techniques, including FTIR, Raman, and NMR.Investigators have reported both silanol S O H ) and siloxane SiO”) groups (where the sign indicates a surface species) at thesurface (e.g., 1,121). Silanol groups canbe subdivided based

Ill.

1 schematicdiagramillustratingthestepwisewater-surfaceinteractions with freshly cleaved or fractured quartz. (From Ref. 10.)

on the number of hydroxyl groups bonded to each surface silicon atom. Evidence has been reported 131 for approximately 93% of the silanol groups comprising species (Si bonded to three oxygen atoms in bulk crystal) with the balance made up of species (Si bonded to twooxygen atoms). Examples of species include isolated and vicinal types, defined by the absence or presence of hydrogen bonding between adjacent hydroxyl groups, respectively, although the above 1131 study did not distinguishbetween the two. Geminal silanol groups, in which oneSi atom is bonded to two OH radicals, was the likely “Q2” species detected in this study. A complicating factor in determining the distributionsof surficial silanol groups ce of surficial physisorbed water, which is a strong absorber of IR moval of surficial water is difficult, and current techniques are fraught with uncertainties. popular procedure is to heat a silica sample above the ternperature at which water is thought to desorb. The sample can then be quickly placed into a spectrometer chamber and surficial silanol groups can then be measured under vacuum. Care mustbe taken, however, to avoid heating the sampleto the point at which siloxane groups form, rendering the surface hydrophobic (e.g., Ref. 14). schematic illustration of hydroxylated silica shows surface complexesthat are (Fig. 1.11) dominated by silanol groups (represented as SiOH). this adsorbed water film increases beyond approximately three monolayers(Fig. l XI), its propertiesbecome more like bulk water [lo]. In the presence of molecular water, the silanol groups ionize, producingmobile protonsthat associate-dissociatewith the surface to impart an electrical conductivity to the surface. these groups dissociate, hydronium ions are producedwhich diffuse from the surface to develop a pH-dependent surface chargeand potential. This surface charge, in turn, attracts a diffuse cloud of counterions to preserve electroneutrality. The resulting interface that exists at virtually all wetted mineral surfaces (Fig, l.1V) is called the electrical double layer (e.g., Refs. 15,16). Discussions of the mechanisms that develop the interfacial field gradients found at silica-water interfaces are found elsewhere (e.g., Refs. 8 , 17-19).

Surface complexation theory has formalized the chemical and physical structureof the electrical double layer into a theoretical model that assumes surface reactions mimic aqueous complexation (e.g., Ref. 20 and references therein). A consequence of the deprotonation of surface silanol groups at the interface is the generation of an electrical charge gradient. The corresponding electrical potential is highest at the solid surface and diminishes in a near-exponential form with distance toward bulkphase water. One such construct that has been used extensively to describe an oxide-water interface is called the triple-layer model [l 5,161. Three electrostatically charged regions are defined as the and d imaginary planes. The positioning of each plane is determined by the distance of cations from the mineral surface structures by direct binding plane) or electrostatic approach of solvated ions plane). shown in Figs. 2a and 2b, idealized planar surfaces correspond to the decreasing

charge distribution and electrical potential that occurs with increasing distance fromthe surface. Hydrogenionsco-ordinatewiththeunsaturated sitesof the interface at the innermost layer (as an “inner-sphere complex”)). Sodium and other weakly bound hydrated cations are positioned at the layer (as an “outersphere complex”)or the layer (near the bulksolution) whilst negative counterions such as Cl- are denizens of theoutermost d region. Low-temperaturesurface complexation modelsdo not permit sodium to specifically interact with thesurface. This becausepotentiometrictitration dataandco-ordinationtheory suggest sodium is prevented from specifically binding to the surface because of shielding by its own solvation sphere. However, sodium may exist in the layer in an “ion pair” co-ordination with the surface that promotes its partial dehydration and water dissociation within the chargedinterface [21,22]. Although hydration spheres are implicitly ignored in this representation, the same relative surface-metal distances are predicted for the surface co-ordination of H’ and Na’ at and planes, respectively. Representations of Si02 surface chemistry by the triple-layer or other models have produced a generally accepted representation for the distribution of complexes on silica surfaces in aqueous solutions (e.g., Ref. 8). For the purposes of a brief example, quartz in contact with a solution composition containing only an alkali salt (Le., quartz-water-sodium chloride system) has surface structures that can be described by three complexes with a population balance:

where fraction of total sites as >%OH species fraction of total sites as >Si- species O>.Sio-Na+ fraction of total sites as >SiO--Na species 6>SiOH

6,sio-

Note that these “species” are model constructs describing surface titrationdata and the surface charge relationships of oxides. Of the three complexes listed, only SiOH has been directly observed using spectroscopic methods[23-261. The other two complexes may or may not have physical meaning as they describe the timeaveraged degree of surface ionization. For the purpose of later relating surface complex distributions to dissolution kinetic data, the SiO- and SiO-Na’ are codependent upon changes in solutionpH and sodium concentration and cannot be evaluated independently. These terms are addedand referred to as SiOt,,.

Surface complexes describingsilica-solution interfaces are not static, but are rather calculated distributionsreflecting an average electronic state resulting from proton, cation,andhydroxylion interactions withtheundersaturated oxygens atthe mineral surface [27]. The resulting population balance is largely controlled by the relative magnitudes of associatiaon constants for the surface reactions listed in ecause of the strong attraction that the underco-ordinated surficial Si atom has for 0 in the hydroxyl, the bond between 0 and H are relatively weak.

Si-

I

d

Acid-base quartz surface complex reactions and corresponding association constants at 25°C

K, SiOH SiOH' SiOH Na' Si0 -Na' H' SiOH2' SiOH H' H' CI- SiOH,CI

Reference

10-6.8

10-7.1 10-2.3

66

66

Even at solution pH values low 4, protons desorb from silanol groups, yielding the SiO" surficial group. This means that for most of the pH scale (pH 4), the surfaceof silica is negativelycharged. The magnitudeof this negative charge gives rise to a number of properties, one of the most important of which is that water moleculesbecome progressivelyreoriented thepH of solution increases. Titration studies showthe pK, forionization of silanol groupsto SiO,, complexes (see Table 1) is about indicating that the surface only weakly acidic. Recent investigations of surfacecomplexation equilibria by [28] show that surface protonation is correlated with dielectric constant and metalOH Paulingbondstrengthfor several oxides, includingquartz.Their findings suggest the importance of bulk crystal structure in governing the bonding and equilibria of surfaceprotonated species. If correct, interfacial waterproperties may be polymorph dependent. The resulting surface equilibria from these titration and theoretical studies lead to the pH and sodium dependence of average SiO,, distributions represented in Fig. 3. In general, net negative chargeincreases with increasing solutionpH and/ or alkali cation concentration until about pH 10 or 11. Above this pH, further pH increases or addition of alkali have smaller effects on net negative charge, Metals other than the alkali cations also interact withsilica surfaces, some quite strongly. Interaction mechanisms appear to range fromsimple ion exchange in the plane to cation-surface specific binding at the innermost layer. For example, surface association constants vary from 10"7.8 for lithium [29] to 10".8 for ferric iron at Later discussion will show that metal-surface interactions significantly enhance or inhibit silica reactivity. The direction and degree of these effects are partially related to interaction strengths although the specifics are con2 A schematic illustration of thetriple-layermodel of anidealized planar silica-water-sodium chloride interface. (a) Each layer has an associated interfacial potential, h(V), and charge density, (C/m')), that determine the inner (C,) and outer layer capacitance(F/m2), by the relationship: AO~(A@~)-'. The nitude of these parameters decreases with increasing distance from the mineral surface into the solution side of the interface and finally to the bulk solution. In this model,thebulkunchargedsolutionisbeyondthediffuse d layer(see test). (b) Schematicrepresentation of thecorrespondingcharge distribution and potential decay away from the surface. (Modified from Ref. 19.)

1

04

2

6

8

1 0 1 2 1 4

Illustration of the distribution of surface sites onquartz as a function of pH. (From Ref. 66.)

siderably more complex.It is significant to note that silica surfaces can “scavenge” stronglysorbedionssuchas ferric iron and aluminum from low-concentration his process of “superequivalent sorption” so extensive as to exceed the amountrequired to satisfy silica surface charge suchthat the aluminum reverses charge from negative to positve [31]. In quantifying the capacity of the interface to attract or bind ions for a specific silica material, it is important to bear in mind that most of what is known about these surfaces comes from studies of amorphous silica and quartz. ~omparisons between polymorph structures should account for density differences between the phases and the resultanteffect on silanol site densities on the surfaces For finegrained or porousforms of silica, particle size or pore radiusof curvature mayalso significantly perturb the sorptive properties per the Kelvin effect (e.g., Refs. 32-34).

Numerous studies have proposed that adsorbed water adjacent to surfaces has properties that are different from bulk water. The early evidence indicated that water is highly oriented with different entropy, mobility, dielectric constant, dissociation constant, viscosity, rates of water exchange, and proton transfer

Until recently, these studies of properties of interfacial water properties have been difficult to confirm by direct measurement. High resolution 'H NMR spectra have detected the structure of (deionized) water at the interface [33]. In the first monolayer of water away from theinterface, they found evidence for a tetrahedral arrangement of water molecules held together by strong hydrogen bonds. Within each tetrahedral unit, water molecules may reorient themselves or exchange with other water molecules from bulk solution, with reorientationand exchange a function of temperature. The exchangeof OH groups at thesurface with water wasnot detected, attestingtothe very strongbonds betweenSi and silanol groups. Hydrogen bonds connect the network of water molecules to the silanol groups at the silica surface, such that the overall structure of water at the interface is so different from bulk water that there is a nearly 40 K freezing-point depression. Evidence for reorientation of water molecules with changes in bulk solution chemistry wasrecently obtained by Duet al. who collected infrared-sum frequency generation spectra of the quartz-water interface over a wide range of solution pH. They found that atlow pH (1.3,interfacial water has a time-averaged tetrahedral structureheld together by hydrogen bonds. Withan increase in pH (3.8 to 8.0), interfacial water beconles increasingly polarized; the change appears to be very sensitive to small changes in pH suggesting that small perturbations of the electric field, caused by the addition of dissolved cations, would disrupt the water structure further. AtpH values that exceed 8, they showedthat the quartzsurface is completelypolarized by SiO- surfacegroups,producingstrong electric field gradients. One result of this polarization is that it causes water to become reoriented such that protons are able to approach the silica surface. Over the range of pH values studied, water molecules appear to reorient an average of 180" with the result that protons can interact with bridging oxygens in the solid. Correlations of these interfacial solvent properties and silica reaactivity suggest that the properties of near-surfacewatermayplay an indirect role in modifyingthe kinetics of mineral-water reactions (see later discussion).

Models of the interaction of water at the interface have all assumed that the first layer of water is immobile, and the recent spectroscopic work cited above appears to confirm this. The mobility of water beyond this first monolayer, however, is uncertain, with someinvestigators arguing for the absenceof additional layers while others contend that a thickshell made up of many layers of water molecules exists [49]. Although there is no agreement on the thickness of the hydration layer, the existence of a layer at least three water moleculesthick is assumed in the triplelayer model discussed above [lo]. The thickness (number of layers) of this structured water region has profound implications for models of sorption, flocculation of colloids, and reactivity of surfaces. The notion of a layer of oriented water molecules comprising more than a single monolayer well ensconced in the literature and is supported by experimental evidence. For example, it has been argued that the hydration layer surrounding colloidal silica is thick 13 nm) and increases slightly 14 nm) with increasing pH of solution [50]. These authors suggest that the dipole-dipole interaction

between silanol groups at the surface and water molecules generates the hydration layer, and that the increasing thicknessof the hydration layer with increasingpH is a reflection of greater charge densityat thesurface withpH. They also showed that the specific volume of the hydration layer is 0.993 cm3/g, which is similar to that of ice 1.09 cm3/g), in good agreement with the spectroscopic evidence cited above. It has been reiterated recently that only the first layer of water beyond the surface is immobilized, so that the properties of watermoleculesfurtheraway from the interface may be identical to those of bulk water [51]. This postulate stems from the measurement of forces between mineral plates (silica and mica) where the separation between the planes are very small 2 nm) [52-541. These studies indicate that astwo mineralsurfaces approach each other (within2 nm), the DLVQ model.The DLVQ forces between them do not conform to the theory proposed that the forces that act on two particles, such as colloids, will switchfrom repulsive (electrical double layer) to attractive (van der Waals) as thetwosurfacescome together. In experimentswithmica plates, experiments found that the forces become oscillatory betyeen repulsion and attraction, with the wavelengthof the oscillations being 2.5 or roughly the diameterof a water molecule (Fig. 4). shown in Fig. 5 for a pair of silica plates, the repulsive forces between the plates at small distances increases to values not predicted by theory [54].

100

10

0.1 3

Diagram of the measured forces between mica plates inmM KC1 solutions as a function of distance. 'up to 2 nm separation, the forces measured between the surfaces conform to DLVO theory, but at smaller separations the forces are oscillatory. Note that the periodicity of the oscillations is 0.25 nm, which is approximately the diameterof a water molecule. Thesedata are indicative of water structure between the plates (see text). (After Refs. 52,

10

0

100

200

300

500

(A 5 Diagram of forces measured between (hydrophilic) silica surfaces in solutions of variable concentration as a function of distance. The data (represenied by dashedlines)conform to DLVO theory(solidlines) at distance 40-50 A, but exhibitforcesgreater than that predicted by theory at smaller separations. This additional force has traditionally been interpreted as evidence for a thick hydration layer, but it is proposed that protruding silica “hairs” account for an apparent additional force (see text). (After Ref. 54.)

Prior investigators have attributed this excess repulsive force to a thick shell of oriented water molecules between the surface and bulk water. In contrast, the nonDLVO repulsions at short distances observedbetween silica plates are argued to be related to steric forces due to the presenceof silica bbhairs” that protrud? normal to the surface [51,54]. The “hairs” act to shift the surface charge 5 A from the surface (the latter of which is coincident with the van der Waals plane), which is theoretically sufficient to correct for the difference in DLVO theory and measured forces. The observation of fibers or rods of silica has been reviewed by Iler [8]. During the formation of silica gels from hydrosols, for example, colloidal particles link together to form chains, which may further condense to force a three-dimensional network gel ([s],p. 225). However, we suspect that these features are not predominant at the silica surface in the force measurementexperimentsdescribed above, for the following reasons. A protruding strand or “hair” of silica would represent a high-energy feature that would not likely survive for long durationsin a solution that is grossly undersaturated in silica, such as the solutions used in the above experiments. our knowledge, other laboratoriesutilizing sensitive surfaceimaging devices have not reported evidence for such features on crystalline surfaces in undersaturated solutions. It has also been argued [51] that the oscillatory forces measured between two mica plates provide evidence for a single (rather than multiple) layer of structured

water. This ideareceived additional support from measurementsof the viscosity of interlayer water and it was found that only the first layer of water closest to the solid has a greaterresistance to flow than bulk water[48]. This effect is attributed to the oscillating forces to multiple layers of structured water between the two plates This interpretation is consistent with ordering caused by strongly sorbed cations and their surrounding shells of solvated waterat themica surfaces, which is not observed (orat least not detected) in deionized water. The extentof ordering of the solution, observed as variations in non-DLVO behavior, is dependent on the identity of thecation dissolved insolution [57]. Thesolutions are so strongly ordered that they resist the closing plates; repulsive forces build up until a layer of solution is forced out of the area between theplates and the forceis reduced [573. These observations wouldseem to suggest that theproperties of the solid in contact with electrolyte solutions governs the orderingand hence the behavior of solutions at the interface. This leads to the conclusion that thick hydration layers are not present at mineral-solution interfaces, and ordering of solutions, when present, are a function of mineral, and not solution, properties [5l]. We conclude from this suite of papers that the thickness of oriented interfacial water, for silica or any other surface, is not known with any certainty. Because the force measurements between plates are different foreachmineral, this may be evidence that the characteristics of the bulk solid controls the behavior of water at the interface. Therefore, the thicknessof the hydration layer, the mobility of the water molecules, and the sorption behavior of cations and anions at the interface may be variable for each mineral.

The concept of mineral “reactivity” is fundamentally tied to reactions that take place between surface species and bulk water across thesolid-solvent interface. The s~sceptibilityof any solid to aqueous corrosion arises from the afore men ti one^ surface charges. The generation of surface charge, whose magnitude depends on the relative difference between the point of zero charge of the mineral and the solution pH, creates changesintheproperties of adjacentwater molecules, as previously discussed. It is the properties of these interfacial water molecules-in terms of dipolar orientation, intermolecular bonding, decreased mobility, etc.thatare implicated in creatingreactionpathways that facilitate rupturing of bonds. Therefore, knowledge of the dissolution reaction on a fundamental, molecular level goes a long way towards unravelling surface reactivity and vice versa. Addition of a variety of cations to the solvent act as excellent probes of the reactivity of surfacegroups. Since each cation in the first twocolumns of the eriodic Table exhibit different effects on the reactivity of silica, these data taken collectively yieldimportant insights into thebehavior of solvent-solute dynamics in the interfacial region. Because the effects of these cations on silica reactivity also share some notablesimilarities, the data also provide some equallyimportant constraints on models linking interfacial properties with reactivity. These themes are discussed below.

Until recently, investigators have concentrated on propertiesof the solvent, both in the bulk and in the interface, to elucidate mechanisms for the dissolution of silica and other solids. In this context, complexation reactions (e.g., Ref. 58) serve analogs for the dissolution reactionsthat are thought to takeplace at the mineralwater interface. An alternative modelhas beensuggested [28,59] inwhich the physicalproperties of the solid exert fundamental control on the dissolution behavior of crystalline and amorphous materials. In particular, these investigators argue, based on the earlier work 181, that the dielectric constant of the solid is key component to understanding surface-protonation equilibrium constants in water: As we discuss more fully below, both of thesemodelshavesomeoverlapping assumptions and may not be as different from each other as it would initially appear.

r The assumption of this model is that the interactions of dissolved silica and other cations with aqueous solutions are similar to those that take place at solid-water interfaces (e.g., Ref. 58). When cation dissolves in aqueousmedia,it is surrounded by shellof watermolecules.Recentspectroscopicworkhasshown that the strength of the forces between the cation and the solvated shell depends on the charge densityof the cation [60]. The strength of the bond between a metal and oxygen across the interface can be assessed by comparing the frequency of water exchange(k,,). In-the case of Si4', the exchange rateof water at the surfaceis 100.3 IO5 which is much slower comparedto other metals, such Mg2+ (k,, 5-l) or Na' (k,, lo9 [60]. The strength of the bond appears to be related to the ionic potential defined the charge divided by the radius The of Si4' is 10, for example, whereas has value of 2.8 (both compared with octahedral co-ordination). Itis not surprising, therefore, that there is correlation between Ip and k,,, but it not clear to what extent this relationship can aid in predictingsurfaceproperties between solids and aqueous solutions. Dissolution rates of orthosilicate minerals correlate extremely well with k,, values 1611, suggesting that these data may be useful under some circumstances. The data,however, canbe potentially misleading if the nature of the complex in solution is different than that at the surface. As an example, the aquocomplex of A13' has k,, that is similar to that Si This would suggest that small amounts of A1 sorbed onto the surface of silica has no effect on the dissolution rate of silica. This is not the case and suggests that the k,, the aquocomplex is different than the k,, of the surface species [8].

A number of researches have arguedthat the data obtained on aqueous complexes may yield incorrect conclusions concerning mineral surface complexation Ref. 28). These investigators propose thatthe dielectric constantandtheaverage Pauling bond strength (ionization potential divided by the cation-OH distance, or: of the solid provide the most useful data to infer reactions that take

place at the interface. They introduced a model that allows investigators to predict surface-protonation equilibrium constants for variety a of oxide and silicate minerals, Note, however, that the charge densityand the polarizability of the surface are similar to the parameters used by investigators proposing aqueous complexes as surface analogs (see above). The model predicts the free energy of sorption onto the surface of a solid from solution for monovalentand divalent cationsand complexes. Surfacec o m p l ~ ~ a t i o n is a function of both the properties of the solid and the properties of the solution, which are related in the following expression:

where AS$ is the Born solvationcoefficient (kJ/mol), is the dielectric constant of the solid, log is the ion-intrinsic contribution, 13 is the gas constant [JjmolK)], and T i s the temperature in kelvins [S]. The models presented above are an attempt to marshal data, typically determined by experiment, into a paradigmthat allows investigators to makepredictions for systems that lack data. Therefore, the models are onlyas good as thedata from which they are fashioned. In the sections below, we review some recent work on elucidating the nature of the interface. What is surprising, perhaps, that despite the large number of studies on the properties of the surface, relatively little is known with a high degree of certainty.

It is now known that dissolved solutes from the alkali and alkaline-earth groups of the periodic table (e.g., Na', K', Ca"',Ba"') enhance the dissolution rate of quartz (e.g., Refs. 62,63) and amorphous silica[64]. seen in Fig. 6, rates of Si02(am)dissolution increaseby factor of 30 over the low-concentrationinterval of 0.0 to 0.05 M NaG1. Dissolution rates are increasedmarkedly by small electrolyte concentrations 0.01 M), butnotethat rates increase little above 0.10 M. This diagram also shows that Ca- and Mg-bearing solutions display a similar maximumrate-enhancing effect, but when compared afunction of ionic strength rather than molality, sodium ions actually increase rates of dissolution more than magnesium, Thesedata support the notion that even small concentrations of dissolved solutes can strongly affect the dissolution rates of Si02(am). The magnitude of the rate enhancing effect is strongly dependent on the identity of the dissolved cation [63], These investigators found that dissolution rates of quartz increase in theorder Mg2' Na' Ca2' Ba2' forsolutions of equal molality. The electrostatically attracted ions are notbelieved to bind stronglyto the surface and indirectly modify rates through changes in the nucleophilic properties of the reactant, interfacial water. Studies of quartz and Si02 glass support this model by showing that experimental activation energies (Ea,xp)obtained for the electrolyte-bearing solutions are statistically thesameasvaluesobtainedfrom experiments conducted in deionized water. The Arrhenius plot in Fig. 7 shows log k , plotted against the inverse of temperature for silica glass dissolution experiments conducted using deionized water (DIW) and solution of 0.01 and 0.05 M

E?

-10.0 -10.5

0

-11.5

Concentrationof

0.08 0.10 (molal)

0.12

Dissolution rates [mol/(m' of amorphous silica as pure silica glass(SiO') into NaCl solutions with a pH of approximately 5.7 for QSI glass at 60°C. Addition ofsmallconcentrationsofcations into solutionresultsinamarkedincreasein dissolution rates. (From Ref. 64.)

Silica

5.7

-5

-9

k,

7 Rate constant [mol/(m2 S)}, of amorphous silicadissolutionplotted against l / T (K) for three solution compositions. The activation energies of dissolution for experiments in deionized water and 0.01 and 0.05 M NaCl solutions are identical (8032 kJ/mol). These data are indicative of increased dissolutionby changing solution properties at the interface or increasing the frequency of Si- 0 rupture upon the introduction of sodium. See text for details. (From Ref. 64.)

NaCl. The slopes of the lines regressed through the data are proportional to the .Eaqxp of dissolution, as seen in the linear form of the Arrhenius equation: logk,

10gA

l;(

(2.~3R)

where A is the preexponential termthat includes a numberof parameters related to reaction frequency, R is the universal gas constant, and T i s in kelvins. The Ea,xp’s for all three experiments are 80 kJ/mol, within experimental uncertainty. From these and other data, one recent model proposes that the frequency of Si -0 bond rupture is controlled by modifying effects of the electrolyte solutes on solvent properties. relationship is discussed whereby dissolution rates correlate with the exchange of water between the solvation shell surrounding electrolytes and interfacial water 1631. The network of water molecules in the first monolayerbeyondthesurfacegroups are displaced by watermolecules of the solvation shell. The resulting loss of waterstructure, by thedisruption of the hydrogenbondednetwork of watermolecules,causes an increase in entropy terms (AS’) and/or the accessibility of water to the silica surface within thereactionfrequencyterm, A (e.g., Ref. 64). Cationsthatexchangesolvated water rapidly may be expected to greatly enhance dissolution rates of silica, and those that have slower exchange rates will enhance rates to a lesser degree. The value of the frequency of water exchange between the cation and bulk solution, correlates well with dissolution rates of quartz 1631.By this model, cations increase the frequency of successful Si- 0 bond hydrolysis reactions without a change in the mechanistic pathway. These data appear to be in accord with the findings that k,, is proportional to dissolution of rates of ions released to solution from orthosilicate minerals 1611.

The investigations cited above indicate theextent to whichtheproperties and dynamics of the silica-water interface are understood. What is surprising, given the intensity of scrutiny that the interface receives,is how little is known with certainty. The increasing sensitivity of analytical devices described above continue to confirm or overturn popular theories of surface reactivity. With major strides occurring in both detection and computer simulation tools, new findings will be uncovered in the near future. It seems likely, therefore, that the elusive qualities of the interface will only slowly be revealed and that this branch science will continue to fascinate and frustrate for the foreseeable future.

Thiswork was supported by fundingfromthePetrology and Geochemistry Division of theNational Science Foundation (EAR-9405362) andthe US Department of Energy,EnvironmentalManagement Science Program(DEFGO7-96ER14699).

1. H. Graetsch, in Silica: Physical Behavior, Geochemistry and ater rials Applications (P. J. Heaney, C. T. Prewitt, and G. V. Gibbs, eds.), MSA Rev. Mineral., Vol. 29, 1994, pp. 209-232. 2. G. H. Beall, in Silica: Physical Behavior, Geochemistry and ater rials Applications (P. J. Heaney, C. T.Prewitt, and G. V. Gibbs, eds.), MSA Rev. Mineral., Vol. 29, 1994, pp. 469-505. r y ate rials Applications 3. J. B. Higgins, inSilica: Physical Behavior, G e o c ~ e ~ i s tand (P. J. Heaney, C. Prewitt, and G. V. Gibbs, eds.), MSA Rev. Mineral., Vol. 29, 1994, pp. 507-543. 4. A. M. Neville, Properties Concrete, 4th ed., John Wiley, New York, 1996. 5. M. Ross, R. P. Nolan, A. M. Langer, and W. C. Cooper, in Health ~ i n e r a ZDusts (G. D. Guthrie and B. T. Mossman, eds.), Rev. Mineral., 1993,pp. 362-407. Silica: Physical Behavior, Geochemistry and ater rials 6. D. F. Goldsmith,in Applications (P. J. Heaney, C. Prewitt, and G. V.Gibbs,eds.), MSARev. Mineral., Vol. 29, 1994, pp. 545-606. 7. E. U. Vernaz and J. L. Dussossoy. Applied Geochern. Suppl. 13 (1982). 8. R. Iler, The C h e ~ i s t r yof Silica, John Wiley and Sons, New York, 1979. 9. G. A. Parks, in ~ i ~ e r a l - ~ a Interface ter Geochemistry (M. F. Hochella Jr., A. F. White, eds.), MSA Rev. Mineral. Vol. 23, Washington, D.C., 1990, pp. 133-175. 10. G. A. Parks. J. Geophys. Res. 89:3997 (1984). 11. A. P. Legrand, H. Hommel, A. Tuel, A. Vidal, H. Balard, E. Papier, P. Levitz, M. Czernichowski, R. Erre, H. vanDamme, J. P. Gallas, J. F. Hemidy, J. C. La Valley, 0. Barres, A. Burneau, Y. Grillet. Adv. Coll. Int. Sei. 33:91 (1990). 12. H. Bergna, in The Colloid Chemistry Silica (H. Bergna, ed.), Am. Chern. Soc., Washington, D.C., 1994, pp. 1-50. 13. C. A. Fyfe, G. C. Gobbi, and G. J. Kennedy. J. Phys. Chem. 89:277 (1985). 14. C. M. Koretsky, D. A. Sverjensky, J. W. Salisbury, and D. M. D’Aria. Geochim. Cosmochim. Acta 61:2193 (1997). W. Healy. J. Chem. Farad. Trans. 711: 1807 15. D. E. Yates, S. Levine, and (1974). 16. J. A. Davis, R. 0. James, and J. 0. Leckie. J. Coll. Interf. Sei. 63:480 (1978). 17. H. C. Li and P. L. de Bruyn. Surf. Sci. 5:203 (1966). 18. R. 0. James and T. W. Healy. J. Coll. Interf. Sei. 40:42 (1972). 19. K. F. Hayes and J. 0. Leckie. J. Coll. Interf. Sei. 115564 (1987). y F. 20. J. A.Davis, D. B. Kent, in ~ineral- at er Interface G e o c h e ~ i s ~ r(M. Hochella, Jr. and A. F. White, eds.), MSA Rev. Mineral., Vol. 23, Washington, D.C., 1990, pp. 177-250. E. Conway, Ionic Hydration C h e ~ i ~ s t rand y Biophysics, ElsevierScientific, 21. Amsterdam, 198l. 118 22. L. G. J. Fokkink, A. De Keizer, and J. Lyklema. J. Coll. Interf. Sei. (1 990). 23. J. H. Anderson and K. A. Wickersheim. Surface Sci. 2:252 (1964). 24. E. Gallei and G. A. Parks. J. Coll. Interf. Sei. 38:650 (1972). 25. B. A. Morrow and I. Cody. J. Phys. Chem. 80:1995 (1976).

26. 27. 28. 29 30. ?l. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43 44. 45. 46. 47. 48. 49 50. 51. 52. 53. 54. 55. 56. 57. 58.

B. A. Morrow and I. Cody. J. Phys. Chem. 80:1998 (1976). M. Prigogine and J. J. Fripiat. Bull. Soc. Royale Sci. Liego 43:449 1974). N. Sahai and D. A. Sverjensky. Geochim. Cosmochirn. Acta 6132827 (1997). D. L. Dugger, J. H. Stanton, B. N. Irby, B. L. McConnell, W. W. Cummings, and R. W. Maatman. J. Phys. Chem. 68:757 (1964). P. W. Schindler, B. Furst, R. Dick, and P. U. Wolf. J. Coll. Interf. Sci.55:469 (1976). G. R. Wiese, R. 0. James, D. E. Yates, and T. W. Healy. in Internationul Science, Physical C h e ~ i s t r y6, Butterworth, 1976, pp. 53-85. P. W. Atkins, Physical Chemistry, Freeman, New York, 1994. M. B. Kenny and IS. S. W. Sing, in The Colloid Chemistry Silica, Advances in Chemistry Series 234, Am. Chem. Soc., Washington, D.C., 1994, pp. 505-516. IS, K. Unger, in The Colloid C h ~ ~ i s t r ySilica, Advances in Chemistry Series 234, Chem. Soc., Washington, D.C., 1994, pp. 147-164. J. J. Fripiat, A. Jelli, G. Poncelet, and J. Andre. J. Phys.Chern.69:2185 (1 965). J. H. Anderson and G. A. Parks. J. Phys. Chem. 72:3662 (1968). G. Peschel and G. K. H. Aldfinger. J. Coll. Interf. Sci. 34:505 (1970). M.J. Tait and F. Franks. Nature 230:91 (1971). N. V. Churaev, V. D. Sobolev, and B. V.Zheleznyi.Russian J. Phys.Chem. 46: 1320 (1972). W. Drost-Hansen. J. Coll. Interf. Sci. 58:251 (1977). J. TT. Davies and E. K. Rideal, I~terfacial Pheno~eiza, Academic Press, NewYork, 1963. N. K. Roberts and G. Zundel. Nature 278:726 (1979). P. A. Sermon. J. Chem. Soc. Faraday 765385(1980). G. Sposito, The Chemistry Soils, Oxford University Press, New York, 1989. S. B. Zhu and G. W. Robinson. J. Phys. Chem. 94: 1403 (1991). N. Israelachvili, I ~ t e r f ~ c i ~ l Swfuce Forces, AcademicPress,New York, 1992. Q. Du, E. Freysz, and Y. R. Shen. Phys. Review Lett. 72:238 (1994). R. G. Horn, D. Smith, and W.Haller. Chem. Phys. Lett. 162:404 (1989). W. Drost Hansen, in ~ i o ~ h y s i c s (F. Franks and S. Nlathias, eds.), Wiley, Chichester,1982, pp. 163-169. S. Sasaki and H. Maeda. J. Coll. Interf. Sci. 167:146 (1994). J. N. Israelachvili and H. Wennerstrom. Nature 379:219 (1996). J. N. Israelachvili and R. M. Pashley. Nature 306249 (1983). R. M. Pashley and J. N. Israelachvili. J. Coll. Interf. Sci. IO1:51l (1984). G. Vigil, Xu, S. Steinberg, and J. N. Israelachvili. J. Coll. Interf. Sci. 165:367 994). B. Derjaguin and L. Landau. Acta Physicochirn. URSS 14:633 (1941). E. J. W. Verwey and J. Th. G. Overbeek, in Theory the Stubility yop phobic Colloids, Elsevier, Amsterdam, 1948. R. M. Pashley. J. Coll. Interf. Sci. 83:531 (1981). W. H. Casey, in ine era/ S~rfuces,(D. Vaughan and R. D. Pattrick, eds.), University Press, Cambridge, 1995, pp. 185-217.

59. D. A. Sverjensky. Nature 364:776 (1993). 60. J. Burgess, Ions in Solution: Basic Principles Chemical Interffctions, Ellis Horwood, Chichester, 1988. 61. W. H. Casey and H. R. Westrich. Nature 355:157 (1992). 62. P. M. Dove. Am. Sei. 294:665 (1994). 63. P.M. Dove and C. J. Nix. Geochim. Cosmochim. Acta 61:3329 (1997). P. Icenhower and P.M. Dove. Geochim. Cosmochim. Acta (submitted). 64. 65. P. M. Dove and D. A. Crerar. Geochim. Cosmochim. Acta 54:955 (1990). 66. D. B. Kent, S. Tripathi, N. B. Ball, J. 0. Leckie, and M. D. Siegel. U.S.Nuclear Regulatory Commission Report NUREG/CR-4807, p. 113 (1988). 67. P. Schindler and H. R. Kamber. Helv. Chim. Acta 51:1781 (1968). 68. P. W. Schindler and W. Stumm, in Aquatic Surface C ~ e m i ~(W. t r ~Stumm, ed.), John Wiley, New York, 1987, pp. 83-110.

This Page Intentionally Left Blank

Universiti: de Franche Comt6-UFR Techniques, Besangon, France

Sciences et

1. Introduction 11. Surface Structure of Amorphous Silica A.Crystallographicsurfacemodels B. Silica surface structure and adsorption of water molecules C. The dissolution state of surface groups on silica surfaces D. Theoretical determination of proton association constants E. Experimental protonation constants ofsilicas

298 299 299 30 1 30 1 302 304

111. Mechanisms of Hydroxyl Ion Adsorption A.Hydroxideionadsorption B. Competition for adsorption sitesby hydrated ions C. Relation between silica structure and hydration strength of ions D. Cation adsorption onto porous silica surfaces E. Alkaline-earthcationadsorption

305 305 306 308 309 309

IV. Description of the Electrical Interface A. Surface ionization and complexation models B. Determination of the point of zero charge (PZC) C. Analysis of surface titration data D. Deter~ination of intrinsic ionization and complexation constants E. Specific interaction of counterions

313 313 316 318 320 32

Electrokinetic Properties ofSilica-Based Systems A. Electrophoresis B. Electrophoreticchargemeasurements C. Specific interaction of counterions VI.The

Effective ChargeConcept

323 324 325 328 331

ersello

A. Equilibrium properties of colloidal silica systems B. Osmoticpressuremeasurements C. Scatteringexperiments

332 334 336

M I . Conclusion

339

References

339

CTI Silica, one of the most abundant chemical compounds on earth, can be either in crystalline or amorphous forms. The principal crystalline forms are quartz, tridymite, and cristobalite. The amorphous forms ofsilicas are known as pyrogenic silica, precipitated silica, silica gel, and colloidal silica sols. Pyrogenic silica differs from the other types of amorphous silica, in that it is obtained from the hightemperature flame hydrolysis ofSiC14, whereas precipitated silica is prepared by neutralization of a water-soluble sodiumsilicate with an acid. Precipitated silicas do not differ from silica gels except that the latter can have a great diversity of rnicroand rnesoporosity. This includes again the “silica sols,” which are nanometric silica particles dispersed in an aqueous medium. Thecommonpoint between all these silicasis that theelementarybuilding blocks are the tetrahedral (Si04)4-. The polymorphismof crystalline silicas results from the different linkages of the tetrahedral units. A random packing of tetrahedral units, with a nonperiodic structure, results in the various types of a ~ o r p h o u s silica (Iler, 1979). Silica surface chemistrybecomes clearly understandable when we considerthe result of the different arrangements in tetrahedral bonding at the surface. The properties of these silicas may be expected to be directly related to their surface chemistry and a comprehensive understanding of the reactivity of silica surfaces, the dispersion in aqueous solutions, the interaction with polymers, is important for a large set of applications, including chromatography, catalysis, polymer reinforcement, adsorption of liquids, or pigment industry. Several general reviews have been published during the past two decades on the characterization of the surface chemistry of silica (see, e.g., Iler, 1979, or Legrand, 1998). In this chapter we will describe someinvestigations of the interfacial structure of the silica-water interface. These studies are essentially performed with high-purity silica sols of radius in the range 10-14 nm, as the model system. First,the silica surface is describedon an atomic scale. Similarity withthe structure of crystalline silica is used in order to express the structuralcharacteristic of the silica in terms of density of OH groups, bond length, and proton affinities. Surface acidity may be calculated from Coulombic and non-Coulombic interactions between protonsand oxygen atomspresent on defined crystal planes of cristobalite. The second part reports some data conceringinterfacial ionization of silica sols and ion adsorption. Potentiometric measurements of the surface charge densityare used in order toinvestigate the interfacial chemistry. Our purpose is to highlight the complicated chemical processes involving adsorption aswell as the dissociation of ionogenic groups and complexation of metallic cations, which appears to govern

l ~ t e ~ a Str~ct~re ce Silicas

the silica surface charge. Investigationsof the role of hydration phenomena maybe accomplished by studies the adsorption of series of ions which differ in terms of charge and electrical charge. In the third part of this chapter, our understandingof silica surface chemistryis based largely on knowledge of the dynamic and phase behavior of colloidal silica particles. Electrokinetic processes maybe used to characterize the electrical properties of silica particles in aqueous solutions. Electrophoretic mobility derives from interactions between macroscopic motion and diffuse electric charge and is used to describe the interface. We focus our interest on the effects of interaction energetic of ionswiththesurface species, on thechargedistribution.Othertechniques, including small-angle neutron scattering (SANS) and osmotic stress, are used to investigate the structure of the silica-water interface and its significance in the effective charge concept in aqueous solutions. Knowledge of precise description of the distribution of ions in the interface is of considerable importance to better understanding of the nature of surface silica reactions and stability of dispersions of silica.

I!. Since Si4' is verymuch smallerthan 02-,the structureof silicamay be described close-packed cubic arrangementof oxygen ions, with the silicone ions distributed intetrahedral interstices. Silica structure becomes clearly understandable when silicon is joined to oxygen in tetrahedral bonding. The tetrahedral building blocks (Si04)4- can be joined together with two or four oxygen corners in common. The sharing of two corners leads to chain structures; with three common corners we can proceed to construct bands and planes; and with four joined corners, varieties of space structures result. Since each complete tetrahedron is an anionic group with four negative charges, the sharing of two corners reduces the net charge to 2- per tetrahedron, for three common corners to I-, while, with all corners joined, the ratio of silicon to oxygen reaches the neutral compositionSi02.The surface chemistry of silica maydepend on the way the silica structure cleaves and similarity with the structureof crystalline silica is usedin order to understand the surface structure of amorphous silica.

The bulk structure of amorphous silica resembles that of the various crystalline forms. Analysis of the surface structure of hydrated amorphous silicon oxide by scanning force microscopy (Tillmann, et al., 1992) shows that the surface exhibits no long-range periodicity and has glasslike character with next neighboring repeat distance of O S nm. The cristobalite structure is good starting model for understanding the surface structure of amorphous silica. In the cristobalite crystal (Galasso, 1970), the silicon atoms are arrangedat the center of four neighbors, and the oxygen atoms midway between the silicon atoms: the Sidistance is 0.159 nm and the Si 0 Si bond angle 150". The electropositive component, silicon, is tetrahedrally co-ordinated by oxygen atoms and the correspondingSi04 tetrahedra

are the primary building units of their complex structures. Both silicon ions and oxygen ions can be present at the surface of silica. Thus, it is expected that oxygen ions, more polarizable and less charged than silicon ions, will preferentially protrude from the surface plane. The neutralization charges of oxygen ions at the surface plane is realized by the adsorption of protons, leading to the formation of silanol surface groups. Asrevealedby several different techniques (Iler, 1979), there aretwomain types of silanol groups on the silica surface: those with a single hydroxyl group attached to silicon (single silanol) and those with two hydroxyl groups attachedto a single silicon (geminal silanols). If we use the surfaces of ,&cristobalite as model for segments of a silica surface, then one kind of single silanol hydroxyl group can be modeled by the 11) face of B-cristobalite. In this model, shown in Fig. la, the shortest distance between twohydroxyl oxygensis 0.5063 nm (Wells, 1986). Theotherkind ofsingle silanol hydroxyl group is the so-called vicinal silanols, which result from the dehydration of two adjacent geminal silanols on the (100) face of ,&-cristobalite(Fig. 1.lb). The distance between adjacent hydroxyl oxygens that are both attached to the same geminal silanol group of the cristobalite (100) face is0.27 nm,whereasthe shortest distance between two hydroxyl oxygens of adjacent geminal silanols is 0.23 nm. The distance between hydroxyl oxygens of vicinal silanols is about 0.33 nm. At their faces, singly coordinated surface groups are present in isolated pairs. The Si-OH groups protrude from the surfxe. The schematic representations of the (100) and (I 11) faces of cristobalite are given in Fig. 1. In this representation, the oxygens are shown as large circles and silicones are indicated as small circles, The oxygens indicated with solid circles protrude from thesurface and the open circles indicate a lowered position in the lattice. The oxygen ions in the interior of the structure are surrounded by two Si ions (doubly co-ordinated), each neutralizing one half of the negative divalent charge. At the interface the situation is different. Here not all the oxygen ionsaredoublyco-ordinated and therefore also singly coordinatedionsare present. These singly co-ordinatedionsareneutralized by theadsorption of a proton, forming an -Si -OH surfacegroup, called

The schematic representation-of the (100) and (1 11) faces of cristobalite.

silanol. Thedoublyco-ordinated oxygen ionspresent at thesurfacearethe relatively unreactive siloxanes. The number density of surface groups at the surface may be estimated by the number of singly co-ordinated ions locatedin a square of nm2 (dashed square in Fig. l>. As pointed out, at the 11) face the average number density of silanol groups is closeto 5-6 per nm2. This result was in good agreement with conventional surface analysis, which found that the hydroxyl groups have an average density of 5 per nm2, which corresponds to anaverage spacingof 0.5 nm. The surface ionization process on silica involves the adsorption of sodium cations onto the surface sites. With the diameterof the hydrated sodium ionset to 0.702 nm. Geometrical limitations proscribe the presenceof more adsorbed ionsin the immediate vicinity of an ion than its co-ordination number. This number depends on the size of the ion and the distances betweenbetween two neighbor sites. Each surface oxygen has six oxygen neighbors, all situated at 0.506 nm from each other. All the watermolecules are tetrahedrally surrounded.

The interaction of the molecule can be of two types: Si OH group and a water one in which H 2 0 acts as a proton donor in H bond to theoxygen of the -Sithe otherin which H 2 0 acts as a proton acceptor inHabond to thehydrogen of -Si -OH. The energy of interaction between HzQ and the surface of silica outgassed at high temperature hasbeen determined by Basset et al. tobe 25 kJ/mol by heat of immersion technique. Molecular dynamics computer simulations are used to study adsorptionof water onto the silica surface (Garofalini, 1990). The simulations provide an atomic-level view of surface reactions, such as the specific mechanisms of water adsorption, the dissociation of adsorbed water molecules, the siloxane bond rupture, and thesilano1 formation. Water-induced siloxane bond rupture was supposed to occur via an electrophilic-nucleophilic attack of the Si 0 bond (Budd, 1961). These simulationsindicate that water-induced siloxane bond rupture and silanol formation are much more complex. They indicate that the most reactive sites are located within the top 0.20.3 nm of the silica surface. This means that the hydroxyl groups have an average density of 5 per nm2. Given separation of 0.5 nm between surface silicon atoms, one obtains cross-sectional area of 0.20 nm2 per bondedsilanol. Isolated silanols, with only siloxane nearest neighbors, then mightbe at least nm apart, so that the maximum density must be less than one isolated silanol per 0.80 nm2.

Irrespective of the form of silica employed, the mechanism of the reaction may be regarded a depolymerization of the 3D covalent network of silica as a result of the breaking of Si-0-Si bonds:

erselI

-Si-O-Si-"+

-Si-O"Na++ -Si-0"Na'

a function of the strength and the amount of the base employed, the depolymerizationmayreducetheconnectivityprogressivelythroughsheet and chain structures to the monomer itself.

K1,l and we shall use the In order to evaluate the proton association constants modeldeveloped recently by Hiemstraetal. (1989a) for crystallized oxides. Basically we consider that in any form ofsilica (crystalline or amorphous) the silicon atom is tetrahedrally surrounded by four oxygen atoms. The principle of electroneutrality implies that the positive charge of silicon is shared by the charge of the four surrounding oxygen anions. This implies that positive charge is distributed over the anions that co-ordinate the central cation. This concept leads to the definition of a formal bond valence, v, as the charge, of the silicon cation divided by its co-ordination number, CN. In the case of silica, the silicon ions distribute their charge (4 over four surroullding oxygens, neutralizing on the average half a unit charge per -Si bond. It implies a bond valence, v, of

At the surface of silica both silicon and oxygen atoms may be present. Because of their higher polarizability, oxygen atoms are more likely to protrude from the surface plane. The excess charge of oxygen ions atthe surfaceof silica isneutralized by the adsorption of protons forming silanol groups. ons side ring from the bulk structure that the formal bond valence 1 between silicon and oxygen atoms, the local effective charge of a surface group depends on the number of adsorbed protons and also on the number of co-ordinating cations. Siz is extremely low 16.9), meaning that the predicted log K2, for protons do not adsorb on these sites, we may neglect the protonation reactions onto the doubly co-ordinatedspecies. In the caseof singly co-ordinated species, the formal charge may be --Si-O",

-"Si-OH',

-Si-OHl

The intrinsic free' energy of the proton adsorption on the surface group can be calculated by assuming two contributions; a local electrostatic AG,,1 and specific contribution AGchem(Hiemstra et al., 1989a). The local AGcOuIdue to the approach of a proton to the surface must include both the lateral part of the electrostatic interaction due to the interaction of a proton and the oxygen in the case of an oxo site -Si-0" or the hydroxyl in the case of a neutral site -SiOH. The general equation reads

where is the number of silicon ions co-ordinated with a surface group, 1 is the valence of the adsorbed proton, Zo is the valence of the surface oxygen, is the valence of the surface hydroxyl,and is the effective bond valence. are effective microscopic dielectric constants, N is the Avogadro number and L is the distance of charge separationdefined as the sumof the silicon-oxygen and oxygen-proton distances. The proton association constants of various types of surface groups can thenbe calculated using AG!nt

ln

The effective microscopic dielectric constants and the specific contribution AGchem are not directly accessible. Equations (2) and (3) are therefore combinedand rewritten as

For ,j (first protonation of a oxo site) and SiOH) these parameters are respectively

At

ZoNe2

A

2 (protonation of a free silanol

ZHZoHNe2 AG:.,

and

The constantsA l and represent the chemical term for the first and second acidity of a surface site at the surface and B is the coefficient of the electrostatic term. A I and A2 and B are deduced from the ionization properties of the dissolved hydroxide monomers:

Si(OH)30-

H'

Si(OH)4

Taking into account all cations with an electronic con~gurationsimilar to that of argon we obtain linear relationships between log and log and v/L. and K;,? are the first and second ionization constants of the surface hydroxide.We can thus determine numerical values for A I and A2 and B, which are A2 and B (see Foissy, Using L nm for -Si(OH)4 and taking the oxygen-hydrogen distance equal to the distance found in water: nm, we obtain log for thefirst protonation constant forthe monomer,which isin close agreement with that determined potentiometrically (Busey and Messmer, Concerning surface ionization, the hypothesis is made of a similar correlation between all metal oxo-hydroxides, with the same B slope. The log value of the surface silanol groups can be calculated from the solution monomer analog by takingintoaccountthe difference in silicon-protondistancefor silanol in the interface compared with the distance in the solution monomer. The first protona-

tionconstant, -SiOH groups

and second protonationconstant,for singly co-ordinated on the silica surface, can be simply represented by

-SiOH: SiOH H:

t-\ -Si-OH Si

K1,l

OH:

Thelog Kg values of the protonation reactions ofsingly groups rnay be calculated from

logKl,Z

17.3

co-ordinatedsurface

L

The predicted log value for the surface group would be log K1,l 9.70 and log -4.61 taking L 0.155 0.096 0.251 nm. Values for log (9.70) and log (-4.61) are consistent with the observation that SiOH is quite stable and that Si- OH; form only in very acidic media. Because of this difference, we rnay not expect two kinds of ionized sites simultaneously at a given pH. Si OH; are stable in the acidic domain (pH and the densityof Si sites becomes significant at about pH 6. The pH of zero charge (about 3) corresponds also to fully undissociated silanol surface.

The predicted log value for the first protonation of silica is, however, high in comparison with otheracidic surfaces. One of the reasons is that the mean distance between two silanol groups (0.5 nm) on the surface is smaller than the Bjerrum length (0.72 nrn).Table 1 shows the first protonation constantsof the silanol groups of various silicas published by anumber of researchers. The first protonation constant of the silanol groups is in agreement with the potentiometrically determined value of log K1,l.This value is not much different from the first protonation constant for polysilicic acid, which again close to the log of monosilicic acid.

First Protonation Constants of the Silanol Groups of Various Silicas Silica

1% Source &,l

Silica gel 7.2 Crystallized silica 795 Silica gel Spectroscopy 7.1 Colloidal silica 9.3 Colloidal silica 9.8 Silica gel 9.3 Polysilicic acid 9.5-1 0 ~ o n o s i ~ iacid ~ i c Titration. 9.86

Method Spectroscopy Calculation Titration Titration Cation exchange Titration

~ a r s h a l l 1974 , Hiemstra, 1989b Hair, 1970 Allen,1970 Foissy,1998 Peschel,1987 Belyakov, 1974 Ingri, 1959

The spectroscopically estimated values are about two units too low compared with the corresponding protonation constants determined by titration methods. It seems very likelythat theintrinsic acidity log K1,of about 7determined by spectroscopy reflects in fact, the log K1,lof the most acidic silanols. This increased acidity of some surface groups, is probably associated with the effect of impurities like aluminum atoms onto the acidic strength of hydroxyl groups.

According to the current stateof knowledge the silica surface can be considered covert with silanol groups. In the presence of an electrolyte it can be expected that some protons are already exchanged with cations. The important consequence of the cation exchange is that thesilica surface is modified in such away that portion of the exchangeablesites iscovered by hydrated cations. Other possibleinteractions with surface groups are the acid-base reactions. Adsorption of OH" by a silanol group results in a negatively charged surface group, as described by -Si-OH+OH-

~-Si-O-+H20

This negative chargeis compensated by the adsorptionof a counterion.Particularly in the case of sodium hydroxide, the global adsorption process can be written as -Si-OH+OH-+Na'

c-"Si-O-

Na'+H20

There a subtle distinction between OH- adsorption and ion exchange mechanisms. The exchange of a weakly acidic proton for a counterionis stoichiometric in ion exchange mechanisms, but is not necessarily so in OH" adsorption mechanisms. The OH- adsorption is a much more extensive function of p strength than the ion exchangeprocess.

In a typical OH" adsorption experiment, thepH of a silica sol adjusted at the point of zero charge (PZC 2.5) increased by addition of an alkali-metal hydroxide aqueous solution. The adsorption of OH" ions is performed on a high-purity silica sol. The synthesis process is described in Persello (1999). The colloidal silicais dialyzed against water in order to remove electrolytes and impurities. The pH of the silica sol adjusted at pH 2.6, by addition of a proton exchange resin, in order to replace all the sodium ionsby protons. This pH corresponds to the pointof zero charge of silica which is the pH at which all the silanol groups are in a protonated form. As illustrated in Fig. 2, thesequence of increasingadsorptionreadsas Li'
0.6

[OH-leq in M/I

The effect of increasing OH" activity on the adsorption of alkaline cations by col~oidalsilica. Circles denote experimental data and solid lines are computed curves from the model.

understand what are the drivingforces controlling the ionization of silica surfaces in presence of alkali cations, we may calculate the dissociation constants of the surface sites.

In orderto investigate the adsorption of metal ions onto silica surfaces, we used the approach developed by Ninham (Miklavic and Ninham, 1990). We consider here two-di~ensionallattice of surface sites having a density of N sites per unit area; each site may be empty or occupied by one ion (a proton or a monovalent cation). We may use the canonical ensembleand the ~ a x i ~ u m - t e rmethod m for the calculation of the partition function (Hill, 1960). Due to different hydration cation sizes, the combined surface species occupy effective areas on the surface different to that of the underlyinglattice site area. We may expect that the competition for adsorption sites by hydrated cations will be governed by the relative degrees of hydration and the size of the cations. The surface dissociation may be described by the two equilibria

On the surface subject to geometrical constraint, we can write

where @h, and are the ratios of occupied area to lattice site area for the species -Si-O-, Si-OH, and Si0 M,and ns, nh, and nm are the number densities of occupying species for the respective components. Following Miklavic and Ninham (1990), the coupled mass action equations are written as e

and

and Si0

refer to the fractionsites occupied by the surfacespecies M, respectively and are given by and

Si -OH

and

n"/N

[H'& and [M'Is represents the concentrations of protons and metal ions at the surface of silica. The surface concentrationsof the mobile ions are described by the Boltzmann distributions

are the concentration of the metal ions and protons in the bulk state, e is the electronic charge of an electron, k is the Boltzmann constant and is the surface potential of silica. The surface potential may be approximated by the isolated surface potential, determined from the surface boundary condition

That is N(1-62 where

is the inverse

-4 ebye screening length, givenby

Using Eqs. (l 3)-(16), the fractions of surface -Si -OH sites, @h, and M sites, Om, may be evaluated in terms of the dissociation constantsKM a nd metal ion concentration, The density of adsorbed cations as a is determined by numerical calculation, taking KM and ICH as adjustable parameters, until a good fit of the experimental results may be obtained (Fig. 2). Some results obtained at pH 9 are given in Table 2. These values are l or Ah 0.20 obtained assuming fixed parameters: N 5 sites per nm2, nm2, and pKH 9-70. In Table 2, we compare the occupying area for the series of alkali ions obtained with our model and compare them with values calculated from thecrystal ionic radius, the radius of the first hydration sphere and the Stokes radius. We can see il~mediately that the greater the degreeof hydration of the alkali species, as in the case for the sequence K", Na", Li', the larger the pK value. We also show that Na' and Li' the occupying areas are equal (for Na') or greater (for Li') to the first hydration sphere area. ForK' and the occupying

Results for piu,, pKH, and Occupying AreasA M for the Seriesof Alkali Ions, educed from This Analysis are Compared with the Nonhydrated and Hydrated Surface Area of Alkali Cations

A M (nm') work) per (this pKM Ions Li' Na' K' CS'

4.69 4.80 4.90 4.98

(number density of Stokes layer First adsorbedionsIonicareahydratedarea(fullhydrated) (nm') nm') (nm2> area (nm')

0.159 0.166 0.171 0.174

0.60 0.70 0.80 0.88

0.017 0.033 0.060 0.09 l

0.136 0.174 0.246 0.310

1.705 1.547 1.406 1,390

areas are much smaller than the sizes of the hydrated ions. This suggests that the cations Na' and Li' may be adsorbed in their hydrated form, while for K' and dehydration is necessary. The driving potential in the alkali ion adsorption on silica surface sites may be the free energy barrier for solvent exchange in the primary hydration shell of the ion. The free energy barrier can be determined from the potential of average force between the ion and a surrounding water molecule (Koneshan et al., 1998). We report in Table 3 the free energy of water-ion interaction in the primary shell (primary minimum in the ion-water potential of average force as a function of the ion-water distance curve: Figure 16 in Konesham et al., 1988), the position of the primary m i n i ~ u m(corresponding on the mean distance between ion and the first hydration shell) and the barrier for dissociation of the ion-water pair (barrier height of the curve), for the different ions K',Na',Li'. The barrier for dissociation of the ion-water pair varied from 3.3 to 1.9 in units of kir from Li' to Na' and 0.5 to 0.2 from K' to

It is well known that the ionic species presentin an aqueous solutionhavea profound effect on the structureof water in close proximity. This canbe attributed Results from Ion-Water Interaction Potentials and Radius of the Ions Wi

Ions

' i L Na' K' CS'

in

in units -2.667 -2.000 -1.611 .222

4-w

~ydrated

Ionic radius

0.1955 0.2356 0.2798 0.3139

0.097 0.148 0.173 0.204

KT radius (nm) units (nm)

3.333 1.889 0.556 0.222

0.092 0.138 0.173 0.209

to the strong electrical field that exists around ionic species, which leads to both reorientation and compression of the surrounding water molecules. This interaction between ions and wateris known as ion hydration,and hasbeen treated extensively by Marcus (1985). The ion hydration can be characterized in terms of the hydration number, h, which represents the number of moles of water molecule bound by l mol of the electrolyte under consideration at infinite dilution. The inclusion of solvent molecules as hard spheres with a dipole moment has some interesting effects in regard to the adsorption of ions onto the surface of the silica. The water molecules are treated as hard spheres of diameter d, 0.28 nm carrying a point dipole. The electrolytes Na' and K+ are assigned diameters of 0.84d, and 1.08ds,respectively. Since the diameter of the Na+ is smaller than that of the solvent, it is quite sensitive to the local solvent structure about the silica.

The interface for precipitated silica appears to be permeable to ions. The rate exchange is in this case controlled by the rate of ion diffusion within the particle and this is related to thesize distribution of the pores. The size distribution is found to beverysensitive to heat treatment. The porosity and specific surface area is found to reach a maximum at some particular temperature; about400-600°C. The specific surface area decreases with increasing temperature as a result of sintering and silanol condensation in the gel-like layer region within the surface. In view of this change in pore structure on heating, thekinetic of cation exchange, as well as the adsorption capacity, may be influenced by heat treatment. Such effects have been analyzed by Yates and Healy (1976). When heated at less than 500"C, the silica is porous to protons but not to counterions, whereas heated to 800"C, the silica is impermeable to both protons and counterions.

Adsorption divalent cations involvesproton exchange. An important characteristic of this adsorption process is the number of protons released for each cation adsorbed. Thisis referred to as theH+,""+ stoichiometry ratio. Thedirect method of determing this ratio is the back-titrationof the released after cations addition (Kinniburgh and Jackson, 198l). For the adsorption of alkaline-earth cations the H+/M2+Stoichiometry is close to l, which indicates that exchange is closer to being equimolar than equivalent. This may reflect that the adsorption of a divalent ionis related in a general way to its ability to displace protons from the surface. The fact that the H+,"+ ratio is close to 1 for divalent cations, meansthat the adsorption of a cation results from a complexation between the metal ionand two vicinal surface groups: asilanol group and a sodium silanolate group respectively, as

ers Figure 3 depicts results dealing with theeffects size and hydration number on the formation surface complex of a series of alkaline-earth ions, on a lligh-purity colloidal silica at pH 9.0 in absence of added monovalent electrolyte. In the low concentration limit, it is observed that small metal ions such asCa2' and Me2+ are well complexed by silanol than the larger metal ions, Ba2+ and Sr2'. The selectivity sequence is Mg"" Ca2' Sr2' Ba2+. The inverse sequence is, however, found in the presence of a high concentrationof electrolyte (Tadros and Lyklema, 1969). The interpretation here is that in the first case a surface complex with a particular bonding structure is formed, whereas in the second case specific adsorption occurs with non-Coulombic contributions. Themetalion size response of complex stability tothepresence of oxygen donors relates to steric crowding.TheM bondstrengthforthe alkalineearthmetalionsmay be Mg2+ ea2+ Sr2+ Ba". However,the level of steric crowdingmust become larger as the metal ions become smaller and the selectivity against metal ionsvaries as the reverse sequence. In thecase of bidendate complex formation, the net result is that the stronger M bond strengths outweigh thegreater steric dif~c~llties of thesmallermetalions and the order of The reverse situationappears complex stability is Mg2+ Ca2+ Sr2'Ba2'. inthecase ofspecific adsorption.Weanalyze now thesituation of complex formation. The dominant factor in determining how a metal ion will respond in terms of complex stability is largely determined by the steric strain. The interpretation here is that the favorable entropy contributions associated with the formation of sixmembered chelate rings lead to the formation of bonding between one divalent cation and two vicinal silanol groups. The lowest strain energy geometry of a six-

0.25

0.1

0.10 0 L

0.05 E

.oo e ~ u i l i ~ r i uconcentration m of cation

Adsorption of alkaline-earth ionsonto colloidal silicaat pH 9.0 in absence of added electrolyte.

memberedchelate ring, calculated byusing Hancock, 1985), is illustrated by

molecularmechanics(Thom

and

The optimization parameters areSi distance 0.160 nm and -Si 0 -Si angle 120". With small metal ions, such as Mg2' or Ca2', no steric problem arises as the ligand 02-is about the same size as the displaced water molecule,while for larger metal ions, disruption of the solvation sheath may occur. Calculation of the lowest strain energy geometry for the six-membered Ca chelate ring shows that the distance Ca-0 must be equal to 0.240 nm, with theideal -0 bond angle 60". This bond length corresponds to the Ca-0 bond length in the gas phase, as shown in Table 4. For amorphous silica, the typical oxygen-oxygen distance at the surface may be 0.506 nm. In this case, a minimum strain energy geometry is that with Mg2' ions, where the Mg 0 bond length is equal to0.20 and the 0 bond angle is equal to 102".

The Role Numerous studies have shown the importance of pH as a factor controlling the extent of cation adsorption by silica (Kinniburgh and Jackson, 1981). Adsorption experiments are performed with divalent metal ions, on precipitated silica (surface area 175 m2/g) athigh electrolyte concentrations. Results for calcium ions and ionic strength l M NaCI, are shownin Fig. 4. These results demonstrate two quantitative adsorption mechanisms. Thefirst mechanism implies the exchangeof calcium with twosodiumionswithasodium andaproton.The Ca2' exchangeprocess is emphasized from the slope of the plot, which indicates that in this range of Ca2+ concentrations, 100% of Ca2' ions are bound at the surface, corresponding to the "specific adsorption" part.

Valuesfor M -0 Length and IonicRadius(Marcus, Measured Metal Ion Densities at the Plateau (Fig. 3)

T

Number of complexed cations at the plateau per nm2 Mg-0

Sr-0 Ba-0

0.24 0.072 0.19 0.100 0.15 0.250 0.1 1 0.140

M 0 length in the gas phase (nm) 0.206 0.235 0.260 0.290

1988) and

Ionic radius (nm)

rsell

2.5

1.5

.o 0.5

.o 0.5

.o

1.5

2.5

Adsorption density of Ca2+ ion onto a precipitated silica at different presence of 1 M of NaCl.

in

The maximum densityof specificadsorbed calcium for the different pH values is well correlatedwiththenumber of ionized sites (see Fig. 5), asdetermined by titration. In the second region (Fig. 4), the adsorbability of the ions estimated from the slopes of the adsorption isotherms, is significantly reduced. This process is pH independent and suggests that ions may be exchanged with two protons.

1.o

0.5

1.o

1.5

number density of ionized sites per nm2

Linear relation between the adsorbed Ca2+ ions at the plateau of isotherm and the number density of ionized sites for a precipitated silica.

Int

ilic

For transition-metal ions, the slope is greater than 1.2, suggesting adifference in H'/Mn' exchangestoichiometry(Corey, 1981),whichmaybe considered precipitation process.

The origins of interfacial charges on silica may be explained by several mechanisms, divided into four classes: 1. Differences between the electron affinities of two phases (silica in various solvents) 2. Adsorption of ionizable species (polyelectrolytes, surfactants, ions) 3. Ionization of surface amphoteric hydroxyl groups 4. Isomorphous substitution of Si4'byA13' or other multivalent cations. In this section, the interfacial charging processes of silica in aqueous ionic solutions are discussed. In aqueous solutions, the interface between the silica surface and ionic solutions acquires charge which alters the distribution of free ions in solution. The charging processof silica surfaces is closely related to the interactioll between the surface silanol groups and cations presentin the solution. required by the electroneutrality constraint, the surface charge is compensated by adsorption of counterions. The interaction of a counterion with the ionizedsilica surface sites can occur under two different modes; either a true contact ion pair is formed so that the hydrationshells are perturbed, or the counterion is attracted by the electrostatic potential in the vicinity of the chargedsurface. The former interaction typeis called site binding (Yates etal., 1974) or surface complexation, while the latter is referred to as atmospheric condensation. The term surface complex doesnot imply a particular bonding structure. It indicates only that non-Coulombic interactions (chernical, complexation) can lead to ion adsorption. Many equilibrium models designed to predictthe interactions of various solutes with silica surfacehave been described (Morel et al., 1981). The most complete model to be presented here is basically that of Davis (1977), who integrated the modeling approachof Bowden et al. (1977) with the site-binding model of Yates et al. (1974).

The triple-layer model (TLM) is used for describing the surface chemical reactions silica (James, 1981). The characteristics of this model are that H' and OH" ions are allocated to separate inner plane of potential, that electrolyte ions are both specifically adsorbed and allocated to the diffuse layer, and that specific reactions are written for assumed reactions with surface sites. The surface-solution interface is split into three planes. H' and O sidered to react in the surface layer of the solid, while bonded electrolyte ions are located at some distance, from the surface; typically, is of the order of the size of the hydrated sorbed ion. Remaining ions for the charge balance diffuse layer where the charge distribution is described by the Poisso model. A schematic picture of the TLM is given in Fig. 6.

ersello 1

VI3 I

Na+

Cl

Na' I I

Schematic representationof potential and surface charge density as a function of distance from the surface accordingto the triple-layer model.Protons and hydroxide ionsadsorb atthe surfaceand electrolyte are assumed to adsorb atthe B-plane. The three layers are separated by regions of dielectric constants sl and and are modeled as a parallel plane capacitor of capacitances Cland

The charge on the silica surface is attributed to adsorption or desorption of protons on the surface of silica, and subsequent ionization of the silanol groups, schematically -Si-OH+W+ -Si-OH+W"

-Si-OH'2 -Si-O"+H20

The surface charge mustbe balanced by adsorption of counterions. One part of the counterion is attracted by the electrostatic potential in the vicinity of the charged surface; this interaction is referred toas atmospheric on dens at ion. Theother counterions are adsorbed by non-Coulombic interactions and this type of interaction is called site binding or surface complexation. In aqueous NaCl solutions, the site-binding relations are expressed as -Si-OH"+Na+ -Si-OHi+Cl-

-Si-0 -Si-OH~

Na' Cl"

The stoichiometric constraintson theelectrical double layer can be written in terms of the balanceof charge densities. The surface chargedensity, which includes all sites that have reacted with a proton or a hydroxyl ion, is given as

ao=F{[-Si-OHf]+[-Si-OH~~~~Cl-]-[-Si-O-][-Si-O-*.

~ a + ]

(21)

The site densities are given in moles per square meter, is the Faraday constant and is given in Coulombs per square meter. Some of the adsorbed counterions are bound to the surface as surface complexes. The charge density neutralized by complexation, is given as F{[-Si-O-Na+]

[-Si-OHiCl-])

(22)

required by the electroneutrality constraint, the remainderof the surface charge is compensated in the diffuse layer, hence the diffuse layer charge density, o d , is given as ad

P{[--S~--O-I

[--s~--oH:]J

(23)

Electroneutrality of the whole interfacial region requires that 00

(24)

ad

The total site density is given by the sum

all possible forms of surface groups,

NS [----Si-OH]+[-Si-OH~]+[-Si-OH~Cl-]+[-Si-O-]+ Si

o-N~+]

(25)

where the total site density, is expressed in moles per square meter. The adsorption isotherm for protons or hydroxides be may written as a masslaw expression and the thermodynamic equilibrium constants are based on Eqs. (17) and (18). The adsorption of a proton can be written as [-Si-OHi] [-Si-OH](H+) Similarly, the adsorption Si-OH-](H') Si-OH]

K+~w(F) OH- ions can be given as e x p ~ ~ Q)

wherethe (H') is the proton activity in the bulls solution in molespercubic decimeter and the brackets denote the site densities in moles per square meter. and Q- are the apparent affinity constants for H+ and OH- respectively. and are the intrinsic constants of the ionization reactions 7) and (18). Thecomplexationreactionslead to two adsorption isotherm equations. The subsequent adsorption of a hydroxy and a sodium ion is given by [-Si-OH-Na+](H+) [-Si-OH](Na+)

exp

kir

QC

Similarly, the subsequent adsorptionof a proton and a chlorite ion can be given by

is the potential at the surface and is the mean smeared-out potential at the plane of complexation situated at a distance from the surface. The potentials and must be related tothe measurablesurface charge density, The charge-potential relationships are derived from the assumption that the interface is electrically equivalent to a series of capacitors. The capacitance of the capacitor formed by the surface layer and plane is given by (30)

c 1

and that of the capacitor formed by plane -Od/(@p

@dl>

and plane

is defined as 1)

The relation between the diffuse layer charge density, o d , and q d may be obtained from the ~ouy-Chapman theory for planar surfaces, which for symmetrical electrolytes, at 2,5"C, is given by -1 1 . 7 4 4 sinh where c the electrolyte concentration in the bulk (Russel et al., 1989). In s u ~ m a r ya, theoretical calculation (Westall and Hohl, 1980) of the distribution of charges in theinterface, requires a minimumof seven parameters. They are the density of the surface sites (Ns),the ionization and complexation equilibrium constants and the inner- and outer-layer capacitance Cland The four equilibrium constants are taken as adjustable parameters. given an estimated value, usually 0.2, F/m2. can commonly be obtained from crystallographic consideration (Zhuravlev,1987) or fromindependent experimental methods ( M o r r o ~andcFarlan, 1994) and is generally set to 8.3 mol/m2.The other parameters may be derived from potentiometric titration data of the system using a numerical method (Westall and Morel, 1977).

Calculation of the charge density requires knowledge of the point of zero charge (PZC). It is usually determinedby the commonintersection point of titration curves at different ionic strengths or is the isoelectric point from electrophoretic mobility measurel~ents (Tadros and Lyklema 1968). However, in practice, we are limited by the experimental difficulty of measuring a " t r ~ e PZC, ' ~ which would be the PZC if the electrolyte ions are not specifically adsorbing. The PZC may be deduced from theoretical considerations. Explanation of the PZC of silica requires not only theuse of electros ( ~ o o et n al., 1979), but also solvation theory (Jamesand Healy, 1972). ing the Born solvation theory for surfaces and the electrostatic theory of Parks (1965), it has been shown (Sverjensky, 1994) that the PZC can be estimated from the inverse dielectric constant of the silica (1,'~~) and from the bond strength per unit bond length (u/rsioH) of the silicon cation in silica. The equation derived from

ilic

regression of experimental PZC values from the literature Sverjensky, 1977) PZC

21.1158

42.9148

is given as (Sahai and

(33)

( r s ~ ~ 14*4866 )

According to the Pauling theory, the bond strength,U, is equal tothe co-ordination number divided by the valency of the central silicon cation in the oxide. For silica U 1. Values of rsioH are calculated from the mean values of the Si-0 bond length around the cation and the bond length in water (0.096 nm), as rSiOH rsio 0.096. Predicted values of PZC for differnet types of silicas, using Eq. (33), are summarized in Table 5. It is possible to estimatethe PZCfromtheproton affinities constants. Considering that zero surface chargeresults from equal surfacedensities of positive Si- OH:) and negative sites Sicombination of Eqs (26) and (27) readily gives for a surface with a single type amphoteric site: PZC

l -(logK+ 2

log

(34)

Assuming the corresponding ionization constants: log logK1,l -9.70 and log log K1,2 -4.61 (where log K1,l and log are the theoretical proton affinity constants determined in Section 11), it follows that the calculated PZC is equal to 2.55. This result is consistent with that obtained with Eq. (33): Comparison between amorphous silicas and crystalline silica (quartz) in Table 5 shows that PZC decreases for materials with long-range structural order. In other words, acidity of the silanol group increases with crystallinity of the sample and polymerization, as shown by Strazhesko et al. (19'74) and Milonjic (1987). Experimental PZC values for awide variety of silicas are available in the literature. Table 6 containsPZC values fordifferent types of silicas with relevantcriteria in terms of material purity and crystallinity. It should be remarked that experimental PZC values are identical to the ones calculated. Comparison between precipitated silicas and quartz shows that PZC decreases for materials with higher crystallinity, as predicted by the theoretical models. It may be shown again that the PZC increases for silicas with lower surface site density (pyrogenic silica) and for silicas containing aluminum ions as impurities (co~parisonbetween Ludox, Z 175, and S 22). In summary, the PZC is one of the key points of informa-

The Calculated pH of the Point of Zero Charge for Several Silicas. &k, are Taken from Shannon (1993) and Q~OH are from Smith (1954) Silicas Quartz Amorphous silica Vitreous silica

&k

4.5780 2.95 3.8073 3.8073

(nm) 0.245 0.251 0.258

PZC 1.60 3.40

ersell

Experimental pH of Zero Charge for Various Silicas Silica

Source

pH at PZC

Silica sol Ludox TM Bolt, 3.5 Pyrogenic silica Aerosil 1994 Casey, 3.0 3.0 Pyrogenic silica Cab-0-Si1 Precipitated silica 175 3.4 3.0 Precipitated silica Michael 2.0 Crystalline silica Min-U-Si1 2.5 High-purity silica sol S 22

1957 Abendroth, 1970 Zerrouk, 1988 Yates, 1975 and Williams, 1984 work This

tion for determination of surface acidity by titration. We may therefore expect some change of PZC from any chemical or structural modification altering the acid-base properties of the silanol groups.

urface T i t r a t i ~ ~ The experimental surface charge, is determined by potentiometric titration. The method was pioneered by Bolt (1 957)and by Parks and de Bruyn (1962). It consists of potentiometric titration of the silica suspension with H' or OH" ions using a convenient set of pH/reference electrodes and gas-tight vessel to avoid contamination from O2 and in the atmosphere. An extensive review of the influence on the measurement accuracy of several experimental parameters has been given by Sjoberg and Lovgren (1993). An aqueous silica suspension was prepared at a pH and ionic strength low enough so that the surfacecharge is zero, say at pH 3 and 5 IOm3 M. Adding alkali to the suspension will have two main effects: increasing the pH of the suspension and ionizing the surfacesilanol groups, In order to have a situation where the surface consumption of OH- ions reduces to about 20% of the total addition of alkali, the surface area concentration of the silica dispersion must be higher than 2000 m2/kg. The surface uptake of OH" can be deduced from the difference between two titration curves, one using the suspension and one using thesolutionafterthe solid particles have been eliminated by someseparation procedure (dialysis, filtration, centrifugation). The surface charge density of silica is determined by the net uptakeof H' and OH-, and is equal tothe product of the Faraday value with the ionized silanol density, i.e.,

where is the Faraday constant and (AI? is the change in relative surface excess of proton over hydroxide ions dueto each increment of acid or base added. Extensive reviews of surface charge density measurement procedures are given in Foissy (1998) and Huang (198 1). Thesurfacecharge datafrom different sources ofsilicas, suchasLudox, pyrogenic silica, precipitated silicas, quarts,andporous silicas, at 0.1 M NaCl

ce

of Silicas

concentration, are presented in Fig. 7. The data of the various silicas are quite different, which confirms that these cro/pH curves are significantly affected by the composition and structure of silicas. The main difference between silica and other oxides is that there is a flat part between pH 3 (the PZC) and pH 6, where the surface charge density is very low. A strong increase is shown above pH 8. The next step is to analyze surface titratiaon data on a high-purity silica at different ionic strengths. An example of their cro/pH curve is given in Fig. 8 for NaN03 concentrations between 2 and 1 M. The results indicate clearly thatthe ion exchange reaction with the surfacesilanol groups is the sum of the adsorption of a hydroxyl ion and the association of the metal cation with the surface ionized groups. It is instructive to compare quantitatively thenumberdensity of exchanged protons, calculated from the experimental surface charge density, and the total number of surface hydroxyls,which is five hydroxyls per nm2, basedon deuterium exchange experiments (Zhuralev and Kiselev, 1962). This comparison is presented in Table 7 for typical values of pH and ionic strengths. The number of exchanged protons is much lower (one proton exchanged at pH 9.5 and 0.05 M) than the density of hydroxyls on the surface (five sites per nm')), suggesting that the repulsions between electrostatic charges limit their surface density. The reason is that the mean distance between two silanol groups (0.5 nm) on the surface is smaller than the Bjerrum length (0.72 nm). Therefore, it might be expected that no more than one hydroxyl may be ionized per nm2.

1-

m S04 S19

S22

10 Comparison of surface charge density versus pH, by various silicas. and Cab-O-Si1 are three pyrogenic silicas, is a precipitated silica containing pprn Al, and Ludox are two colloidal silicas containing sulfate ions as impurity. Stober silica is a silica obtained from hydrolysis of tetrarnethoxysilane, quartz is a crystallized silica, and S22 is a high-purity colloidal silica with nitrate ions as counterions.)

Surface charge density versus pH for ultrapure silica in the presence various concentrations of sodium nitrate.

of

Surface titration of silica S 22 in N a N Q solutions is now analyzed with a numerical optimization algorithm. The surface charge, may be fitted with the triplelayer model (see Koopal et al., 1987) together with a set of initial guesses of constants KA,Kc, and capacitance Cl. iVs is set to 8.3 low6mol/n? and is given the value of 0.2 IF;/m2.The triple-layer model accounts specifically for the effects of ionic strength and the nature of electrolyte. The numerical opti~ization method used is a nonlinear least-squares procedure for which the root-mean-square difference, between the measured and the calculated surface charge, is minimized. The optimization (Westall and Hohl, 1980)is done by searching for the minimul~in by simultaneously adjusting the valuesof KA,Kc, and Cl.

Experimental Estimate for Numberof Exchangeable Proton Obtained from Surface Titration Data on Ultrapure Silica S 22, asa function of pH andIonic Strengths Number of ionized silanol groups per nm2 PH 9.50 9.00 8.00 6.00 4.00

1M

1.50 1.16 0.77 0.30 0.09

0.1 M

1.17 0.9 1 0.58 0.20 0.05

0.05 M

1.07 0.83 0.52 0.17 0.04

0.01 M

0.84 0.66 0.38 0.10 0.02

0.005 M

0.002 M

0.74 0.59 0.33 0.07 0.00

0.61 0.49 0.25 0.03 0.00

In order to obtain unique set of triple-layer model parameters, the computerfits of surface charge density shouldbe carried out with the condition: ApKis equal to 5, where A + (i.e., PZC 2.5). The calculated curves representing surface charge density a function of pH are shown in Fig. 8. It can be seen that the separate curves at each ionic strength fit the data reasonably well. The optimized set of triple-layer model parameters for ultrapure silica 22) are reported in Table 8. As an example, Table 8 gives again the calculated distribution of surface species, surface charge density, and potentials at pH9andfor an ionic strength5 M. The important characteristics of silica surface chemistry are that 90% of the counterion charge resides inside the shear plane and that more than 80% of the surface sites are in nonionized silanol form. It can be shown again that the difference between the PZC and theintrinsic acid constant, pkl, is large. This means that at pH higher than the PZCthe existence of positively charged sites may be neglected. This agrees with the experimentalresults of Ahmed, where,in the consideredpH range, no positive surface chargewas found (Ahmed, 1966). The surface ofsilica may be described by acid sites alone. Another point is that the density of ionized silanol is negligible near the PZC.That means that near the PZC and overlarge range of pH (2 to 6 atlow ionic strength) the surface charge on silica is close to zero.

The effects of the size and valence of electrolyte ions on the charging processes of silica have been largely discussed in the literature (Kinniburgh and Jackson,1981).

Optimized Triple-Layer Model Parameters for Silica S 22, 9.0, Salt Concentration (NaN03) 5 M Equilibrium constants 9.60 4.60 p& 7.60 PICA 2.60

PIC pK+

Surface species number density 5.000 4.400 -Si-OH 0.074 -Si-0" 0.526 -Si-O-Na+ Surface charge density (C/m2) Surfacecharge -0.0936 Adsorbed charge c@ +0.0840 Diffusecharge od +0.0096

Interfacial parameters 8.3 X 0-6 mol/m2 1.2 F/m2 0.2 F/m2 PZC 2.5

-Si-OH; -Si-OHlCI"

Potentials

8

We now analyze theeffects of sizeand valence of counterions on the surface charge measured by titration methods. We will successively consider some relevant aspects of interactions of monovalent and multivalent metallic ions with silica surfaces.

1. on ova lent Cations Titration measurements were made on a series of silica sols synthesized (Persello, 1999) directly from the corresponding alkali silicates and purified by dialysis. The results are plotted in Fig. 9, which shows the effects of alkali ions on the (pH) curves. Thisexample illustrates that the surfacechargecandependstrongly on the electrolyte type and composition. It can beseen that there is asurfacecharge increase as a functionof the ionic radius of ions and that thesequence of increasing surface charge reads as Li"Na' K' The simple explanation is that small and strongly hydrated cations do not approach charged sites as closely as bigger and poorly solvated ones (see Section 111). The reasoning implies that adsorbed counterions retain their primary hydration sheath. ~ u ~ t i v a ~Cations ent Thissurfacechemistry ofsilicamaybe markedlymodified, when a A1 atom replacesa silicon atomat the surface. Thepresence of adsorbed A13' metal ionscanpolarizetheadjacent silanol groupsandincreasethe acidity of the adjacent silanol groups. Aluminum ions may also create basic -AlOH groups onthe surface, whichgenerate positive sites by protonadsorption.Thismay result to a shift of the PZC. Table 9 illustrates the influence of aluminum ions, introduced as sodium alurninate in the silicate solution during the silica precipitation process, on the PZC. In these experiments the ratio of positive to negative sites was also meausred at pH 7

-0.3

10

Comparison of the effectof a series ofalkaline ions on surface ionization for different silicas obtained by neutralization of the corresponding alkali silicates, in solutions 4 M of thesamemonovalentcounterions.

Effect of Aluminu~Content onthe Basic to Acidic Sites Ratio and on PZC

YOweight of A1 in g A1 per 100 g silica PZC 0.00 0.03 1.oo 3.OO 19.00

2.5 2.7 3.2 4.0 4.5

0.000 0.003 0.042 0.180 0.550

by adsorption of molecular probes(Persello,1989). expected, the densityof basic sites increases with aluminum content, resulting in an increase of the

The concept of surface charge is an ambiguous notion in that the charges must be correlated with the interaction energy between the counterion and the surfacesite. The interfacial charging processesof silica in aqueous solutions has been identified as the ionization of the amphoteric hydroxyl groups on silica surface or the isomorphous substitution of Si4' by A13+. The charge localized on the surface alters the distribution of free ions in solution and an equivalent number of counterions must be adsorbed in order to satisfy the electrical balance in the interface. The presence of counterions leads to the formation of a double layer. This is because of the presence of counterions, which randomize around the macroion to maximize entropy. The thickness of the double layer is of the order of the ~ebye-Huckel screening length given by

where is the density of free electrolyte, zi is the valence of the ionic species, is the permittivity offree space, is the static dielectric constant of the aqueous medium and e is the electronic charge (Russel et al.,1989). The diffuse-layer thickness affects the potential at the inner edge of the layer, so an accurate picture requires some knowledge of the chemistry of the interfacial region. The introduction of finite ion size produces a geometric exclusion of an ion from the surface region because of the hard-core interactions with other ions. It should be expected that a large fraction of the ions are bound at theparticle surface, then the particle charge shouldbe lowered by this fraction. The emphasis here is on phenomena used to characterize the effect of ions condensation on the structure of double layer, In this sectionthecharge of high-purity silicas measured by electrophoretic mobility is compared with that measured by titration. Additional insight into contribution on an ion to the electrical charge is provided from the studies of types of electrolytes on electrophoretic charge.

The most classical method of measurement of the surface chargeis to measure the velocity of a small spherical particle placed in a constant electrical field. The velocity is in this case directly given by the net charge of the particle. The force on a e immersed in a uniform dielectric in the presence of a uniform field, The charge on a charged sphere can also be written as Q where is thenumber of charges onthe particle. Thecharge density, is obtained by the equation

where is the radius of the particle. The potential on the surface of the particle is as important a parameter as the surface charge. Thereis no rigorous theory relating the surface number charge,ZP, to surface potential, In a linearized Debye-Hiickel theory, is related to the surface charge in the simplest case by

where is the inverse Debye-Huckel screening length and is the radius of the silica particles. The surface potential canbe measured by electrophoresis. This is the potential at the surface called the shear plane that separates the colloidal particle, plus those ions that move with it, from the medium. In that sense, effective charge can be measured by electrophoresis, An extensive presentation of theexperimentaltechniquesandthetheory of electrokinetic phenomena given by Hunter (1981). For spherical particles, where the diffuse layer is thick 0), the balance between the electric force and the viscous drag is expressed as (Huckel, 1924)

where EA, is the viscosity of the dispersion media (typically EA, is equal to Pas, in aqueous dispersion), Q is the particle charge (in C), T is the absolute temperature 298 and Em is the uniform electrical field strength (in V/m). U, represents the particle velocity, and denotes the orientation of the incident field. The different physical constants involvedin the fundamental lawsused here are shown in Table 10. In the situation where the diffuse layer is thin ( m compared with the particle radius the potential at the shear plane is given by the well-known von Smoluchoswski formula (19O3), as

This expression also appears to suggest that the mobility is independent of ionic strengthand particle size. Inordertocomparetitrationmeasuredchargeand

ilieas

Physical Parameters Constant value

Numerical

Avogadro's number Boltzmann's constant Charge on an electron Permittivity of vacuum Dielectric constant of water

Symbol NA

k e

6.02552 1.38054 1.60210 10-19C 8.854 C2/(Nm2) 80

electrophoretic data, this expression may be converted to one involving the charge instead of the potential. Using the Debye-HUckel approximation (Russel et al., 1989), the electrophoretic mobility may be given as

To investigate general situations for particles with equilibrium diffuse layer and small potentials, like silica, we use thecomputationalprogramdeveloped by O'Brien and White (1978). The electrophoretic mobility for arbitrary Debyethickness, is expressed in terms of the particle charge, Q (or c-potential). Typical results calculated for 10 nm spheres in NaN03 solutions, are given in Fig. 10. The figure shows how the potential at the shear plane affects the mobility for different thickness of the double layer.

Experimental measurements are made on a modelsilica particle (sol S44). The silica are nonporous spherical particles in relatively monodisperse form, with a radius

spherical particle

0

40 ear plane potential (mV)

~lectrophoreticmobility of a spherical particle in NaN03 as of #a.

function

12 nm. The silica particles are dispersed in aqueous NaN03 solutions and the pH of thedispersion was adjusted by usingionexchange resins (Persello and Cabane, 1999). Figure 11 depicts results dealing with theeffects of the pH onthe electrophoretic mobility of the silica, as a function of the ionic strength. The effect of pH can be interpreted as theresult of the ionization of the surface groups.We shall define two regions. In the first region, typical pH 6, the increase of the mobility is monotonic owing to the effect of pH on the ionization of the surface groups. With increasingpH, the mobility is no longer monotonicand a plateau appears. the density of ionized sites increases, effects from the diffuse charge cloud begin to compensate. maxi mu^ charge can be expected and all charges added to the particle above this maximum valuewill condense on thesurface. The positionof the plateau depends on the ionic strength of the dispersion and the size of silica particles and is also correlated with the parameter These results indicate that there are different surface states and different surface reactions in these two ranges of It is instructive to comparequantitatively the charge determinedby titration and those measured by electrophoresis. We use the computational program developed by O’Brien and White (1978) in order to calculate the surface charge density from themobility results. Thiscomparison is presented in Fig. 12, for 10 nm silica spheres in 2 M NaN03 solution.Thecharge isexpressed in number of charges per nm2. first region, pH 6, can be identified where the charge densities, obtained from titration measurements and calculated from the electrophoretic mobility measurements, are of the same order of ~ a ~ n i t u d In e . the second region, above pH 6, it can be shown that the charge measured by direct titration

Silica sol S24

Electrophoretic mobilities versuspH, of high-purity spherical silicaparticles (radius 12 nm) in various concentrations of NaN03.

R

titration

El

El

Silica sol S24

m

m=2.7

5 PW Comparison of surface charge density measured titration and calculated from electrophoretic mobility.

exceeds thatfromelectrophoreticmobilitymeasurements.Theelectrophoretic charge reaches a plateau and is maintained constant over the range of pH. The charge density at the plateau is equal to B, 0.06e- per nm’. We may compare this value with the charge density measured at pH 9.5 by the titration method: nt 0.60e per nm’. It may be shown that theelectrophoreticcharge density represents only 10% of the potentiometric charge density. If we compare electrophoretic charge density, 0.06e per nm2, with the structural charge density, n, Se- per nm’, this ratio falls to 1.2%. It can be expected that the electrophoretic charge densityof silica is very sensitive to the amountof impurities present at the surface. The results shown in Fig. 12 highlightthemechanism of charging ofsilica particles, ionization counterions. The balance between entropic and electrical contributions means that,in the case of silica, the charge densitycannot exceed 0.06eper nm’. Of course, increase of the ionization degree, which increases as the ionic strength or pHincreases, is unrealistic and the excess Na’ adsorbed counterions are in condensed state. The physical meaning of the charging processis described in the next section. It has been shown that the electrophoretic mobility is depending on the ionic strength but is constant above a critical value of pH (Fig. 1 l), Further insight comes from calculating the charge and charge density from the electrophoretic mobility at different ionic strengths in the pHregion where the mobility is constant. Calculations were performedusingthecomputationalprogramdeveloped by O’Brien and White (1978). The results are shown in Table 11 for a ultrapure silica sol (S24) in various systems. It clearly appears that the charge or charge densityis almost constant, andbecomes independent of ionic strength and pH of the system. This situation can be explained by the fact that large fraction of the ions are bound at the silica surface; the silica surface charge should thus be lowered by this fraction. The charge densityincreases with the bare charge densityup to the point wherethethermalenergy of an ion balancesthe reversible work necessary to

Comparison of Experimental Charges for a Silica Sol (S24 Different Ionic Strengths number ChargeElectrophoretic Concentration of mobility Charge number NaN03 (mol/L) [lo-*m2/(V density per

0.05 0.005

.34 .95 -3.00 -3.80

6.0

0.061 0.059 0.061 0.060

nrn2

per particle 10 nm)

76 74 76 75

SO nm), at

pH ofplateau the of constant charge

4.5 5.5

6.5

remove the ion from the silica surface. From that point on, the electrophoretic charge remains constant and a maximum effective charge can be expected and all charges added to the particle above the maximum value will condense on the surface. We note that the electrophoretic mobility increases with the diffuse-layer thickness and with the size of silica particles (Persello and Cabane, 1999). This can be interpreted in terms of thedimensionlessdouble-layer thickness, dependent charge (where is the radius of silica particles). The critical pH, pH,, above which the surface charge become constant, is only function of the ionic strength. We find two limited critical pH values; at very high ionic strength pH, PZC 2.5 and at verylow ionic strength, pH, 9.6 (where is the proton association constant defined in Section 111).

cti The effect of the type of electrolyt(: ions on the structure of the silica-solution interface has beenwidelydiscussed in theliterature(KinniburghandJackson, 1981). However, the discussion on their effect on nleasurement of electrophoretic charge gives a conflicting picture of the phenomena. In this chapter we analyze the effects of size and valence of counterions on the electrophoretic charge of highpurity silica sol. We will successively consider some relevant aspectsof interactions of monovalent and multivalent metallic ions with silica surfaces.

1. Ions Another matter for discussion concerns the interaction between the counterions present in the interface. The finite sizes of ions influence their arrangement, leading to localized perturbations in the electric field. The electrophoretic charge measurement is used as an experimental test. Table 12 displays the electrophoretic mobility of a series of high-purity silica sol, modified with different alkaline ions and measured for a constantpH 8.5. The different sols are synthesized from the corresponding alkaline silicates (Persello, 1999). The results set out in Table 12 indicate that the different silica sols behave as though the electrophoretic charge is almost constant for a constant (differences in charge are below 5%). This result leads us to believe that discrete charge effects

lectrophoretic Mobilities and Calculated Surface Charge Densities and Charge Number Densities for Different Monovalent Counteriom at pH 8.5 and lca 2.00 (Measurements Performed at Constant Conductivity) Mobility radius Ion

m'/

Ionic 0.071 0.097 0.141 0.173

Li+ Na'

Conductivity (ilcS.cm>

-1.801 0.0530 -3.963 0.0577 -2.064 0.0607 -2.084

4

575 575 577 579

(Nb/nm2)

(C/m2)

0.0613

-0.0085 -0.0092 -0.0097 -0.0098

are short-range and can be ignored for dilute systems, although a strong dependence on thesize ofcounterion is found in the case of potentiometric surface charge (see Section IV).

Ions We analyzethesituation where the silica particle isin contact with a solution containing divalent cations. It well known that divalent ions of the seriesof alkaline-earth ions, are specifically adsorbedon silica (TadrosandLyklerna, 1969). In general the selectivity sequence for the adsorption is as follows: Ba2' Sr2' Ca2' Mg2+ (Malati and Estefan, 1966). A set of experiments was designed to test the effect of the interactionsbetween the divalent ionsand the silica surface sites on the electrophoretic charge. Measure~entswere made on colloidal silica at constant pH(9.5) at different ion concentrations. Figure13 shows the electrophoretic mobility for thedifferent salts, as a functionof the number of ions adsorbed per nrn2 on the silica surface, First we note that the chargereversal phenomenon is not

sol S37 pH=9,50 [SiO,]=l 0

Ba 0

0.05

0.1

0.15

number of adsorbed metal

0.25

0.35

per nm,

Effect of number density and nature of divalent cations on the electroM. phoreticmobility of sphericalsilicaparticlesin NaN03 2

Persello

observed. The silica particles may benegatively charged even at the plateau of divalent ion adsorption conditions. Theresults also provide a baseline forassessing themechanism of adsorption of these ions. In the first part of thecurve,the electrophoreticcharge is not modifiedby adsorption of cations, ascompared with the charge givingby sodium ions. This result leads us to believe that the divalentionsare specifically adsorbed on the surface ofsilica by formation of binary surface complexes (see Section TV). It can be noted that the number of complexation sites is, however, limited. In the case of calcium, we may adsorb in the complexation form, only 0.02 ions per nm2;that corresponds to the complexation of 0.04 surface site per nm2, seen in Fig. 13, the mobility of the silica sol in presence of sodiumions is equal to -1.80 loe8 m2/(V.s) and corresponding charge number density yields 0.044 charge per nm2. Comparison with the number of adsorbed calcium ions suggests that theeffective charge is given by this type of high-reactivity site. Thedominantfactor in determining how ametalion will respond in terms of complex stability is largely determined by the steric strain, The driving force of complexation may be the formation of a six member chelate ring which has lower steric constraints (see Section 111). Thesecond part of thecurve(Figure 13), shows effects due to the valence character of the electrolyte. Divalent cations are adsorbed in theinterface as counterions in place of sodium ions. We note here that the mobilities are lower as the valence of the counterions increases (from -1.80 to -0.32 10"' m2/(V.s). These results agree closely with predictions of O'Brien and White (1978). Charge density calculations show that the charge number density of silica is halved when divalent counterions replace monovalent ions (0.06 charge per nm2 with sodium ions and 0.03 charge per nm2 with calcium ions).

3. ~ y d r o l y z a ~ ~ e ~ u l t i v aIons lent In order to measure thespecificity of an ion for asurface, the minimum concentration of ions for electrophoretic charge reversal at the infinite dilution ofsilica conditions, has been measured (Wakatsuki and Furukawa, 1974). For silica, relative specificities of the metallic ions based on this concept are as follows:

Ca2'Mg2'

Rb'

K'

Na'

Li'

Further insight comes from the comparison of the critical ion concentration of charge reversal with the ionic potential of the ions. The ionic potential was used as anindex of the polarization powerof an ion andis defined as where is the charge of the ions and r is the ionic radius of the ions. Figure 14 illustrates the correlation between the critical ion concentration of charge reversal and the ionic potential for a silica sol. The lackof agreement may stem from the extreme sensitivity of the chargeto the complicatedchemicalprocesseswithinthedouble layer. In fact, the minimum cation concentrations for charge reversal are unrealistic and only some trivalent cations such as A13+ are able to reverse the electrophoretic charge by adsorption. For the other cations, the charge reversal is explained by coating formation on silicaby precipitation processes. Electrokinetic investigation of charge reversal

Exponential relation between the mini mu^ ion concentration for electrophoretic charge reversal and the ionic potentialof the ions, for a precipitated silica.

accompanying sorption of Co(I1I) onto silica (James and Healy, 1972) leads to the proposal that charge reversal is associated with formation of hydroxide coatings on the silica.

In the preceding sections, many types of surface charges have been defined: structural charge determined from crystallographic considerations, potentiol~etric charge (Zp) deducedfromtitration data, Stern layer and diffuse layer charges calculated using double-layer models, and electrophoretic charge determined fromelectrophoreticmobilitymeasurements.Now we investigate the electrical charge involving the knowledge of dispersion processes, such as the stability of sols, the structure of concentrated silica suspensions, and the aggregation phenomena (Chen and Teixeira, 1986). These phenomena are largely determined by the electrostatic repulsion between the charged particles. We must focus on the longdistancebehavior of the electrostatic potential,andtheDebye-Huc~el-like approach can onlybe used if is replaced by the effective charge, The presence of an electrical charge at the surface leads to an adsorption of counterions, which randomize around the silica particle to maximize entropy. This leads to the formation of a double layer, which thickness may be approximate by the Debye-Huckel screening length given by

where the number density of the colloidal particles and is the density of added monovalent electrolyte. This charge is determined by the number of counterions that desorb from the particles at distances that are comparable to the interparticle distances, as only

these ions can transfer momentum from one particle to another. This charge is referred to as the “effective” particle charge, contrasting the “bare” charge, It should be expected that if a large fraction of the ions are bound at the particle surface, then the particle charge should be lowered by this fraction. The effective charge density for spherical colloidal particles can be predicted from the theory of ion condensation developedby Manning for polyelectrolytes (Manning, 195 1).The theory predicts an effective charge density, which increases with the bare charge density up to the point where the thermal energy of an ion balances the reversible work necessary to remove the ion from thecolloidal particles. From that point on, the effective charge remains constant. According to this theory, a ~ a x i m u meffective charge can be expected and all charges added to the particle above this maximum value will condense on the surface. Let us assume, as is generally believed, that condensation does not occur unless the distance between the charges is less than the Bjerrum length, LB 0.713 nm in water at 298 K with 78.5. If the surface is partitioned into squares of areas 0.713 0.713 0.508 nm2. The surface charge density would be then 0.315 C/m2. This is precisely the value of surface charge density that is predicted from calculations (Conzalez-Tovar and Lozada-Cassou, 1989). For many practical systemsthe experimentalstructuralcharge resulting fromcomplexsurfacechemistry is unknown. For monovalent counterions, a correct approximation of the effective charge, Z,, can be given by (Belloni, 1991)

G where

is the particle radius and LB is the Bjerrum length given by (43)

The Bjerrum lengthis equal to 0.713 nmin aqueous solution at roomte~perature. Further~ore,the charge of silica appears to be dependent on the particle size. Using potentiometric tiration, an analytical expression for the particle size dependence of the surface charge of colloidal silica has been proposed by Sonnefeld (1993).

Theosmoticpressure of colloidal dispersions is determined by thenumber of counterions that desorb from the particles at distances that are comparable to the interparticle distances, as only these ions can transfer momentum from one particle to another. It is also measured as a resistance to a decrease in the volume available to theparticles. This resistance originates both from an increase of energy, due to the interparticle forces, and a loss of entropy, due to a restriction of configurations of the particles. The osmotic pressure, II, is related to the pair interaction potential, U, and the equilibrium structure through the Viriel equation:

where is the number densityof particles and is the pair distribution function. In absence of any attractiveforces the pair interaction potential may be approximated by the electrostatic potential of repulsion between the silica spheres, which arises from the overlapof the electrical double layer. The range of interaction is of the order of the Debye-Huckel screening length From electrostatic considerations, the pair potentialof repulsion between two finite-size spherical particles is assumed to be (Derjaguin and Landau, 1941):

where a is the radius of the spherical particles, is the permittivity of free space, is the static dielectric constant of the aqueous medium, and e is the electronic charge. More thermodynamically consistent equations canbe obtained on the solutions of integral equations, such as Percus-Yevick (PY) or hypernetted chain (HNC) equations (Belloni, 1984). The principal advantageof this approach is that therrnodynamic consistency leads to accurate radial-distribution-function solutions. For purely repulsive potentials, like nanometric spherical silica sols, the solutions of the PY and HNC equationshave been found to be in accurate agreement with experimental results. In addition, it is known (Rosenfeld, 1979) that the HNC approximation is appropriate for large interparticle distances and the PY equation is more appropriate for small interparticle distances. The PU and HNC integral equations can be written Yh2)

4x12)

which define the radial distribution function or exP[Y(r>l exp[--BU(r)l (HNC) (48) The h and g functions are related by h g l. The structure factor, S((?), for aqueous dispersions of monodisperse charged spheres canbe obtained by small-angle neutron-scattering technique( ~ a g n eet r al., 1991). The experimental pair correlation function is determined by Fourier transformation of the measured The pair correlation function, is given in terms of the structure factor by 1 z7t2rp

l]Q sin Qrdr

(49)

rsell

An important feature of these equations is that the effective charge of silica particles can be determined from osmotic pressure measurements, by assuming pairwise interaction potential (45).

Osmotic stress methods are based on the movement of exchangeable massbetween two solutions that have established osmotic equilibrium. As a result of the mass transfer the volume fraction of the colloidal dispersion of interest has changed, which is manifested as work being doneon these colloids. The interparticle spacing can be determined by small-angle neutron-scattering measure~ents.The method, introduced by Parsegian et al. 979), is to place the colloidal dispersion in contact with a stressing polymer solution of known osmotic pressure. The contact may be done via a semipermeable membrane (Parsegian,1986). The silica dispersions (with particle radius 10 nm) are placed in dialysis membrane bags (Vislcing, MW 12,000), which are then placedin aqueous solution of polyethylene oxide(PEO, MW35,000) at varying concentrations which apply an osmotic pressure. The osmotic pressure of the dispersions is determined on the basis of the known osmotic pressure of the surrounding polyethylene oxide (PEO)solutions, with which the bags of dispersion are in thermodyllamic equilibrium. Figure15 depicts experimental osmotic pressure results dealing with the effects of volume fractions of silica, types of counterions, for a silica dispersion of particle size a 10 nm. re 15 compares the experimental osmotic pressuredata for different monovalent ions with those for proton (pH 3) and with numerical solutions of Eq. e note that at pH 3, when all sites are associated with protons, the silica dispersion behaves as a hard sphere system,i.e., a noncharged sphere, as expected.

M

Comparison of the effectof a series of alkaline ionson experimental values of osmotic pressure as a function of volume fraction, for a colloidal silica,at pH 9.0 and 3.0. The computed osmotic curve using the HNC model is represented by the solid line.

This confirms that near the PZC the surface of high-purity silica is covert with silanol groups, which is not the case in the presence of impurities, particularly in aluminum ions. At higher pH (at pH 8.5 in Fig. 15) long-range correlations between charged particles are shown, where theorigin represents Coulombic interactions. The effective charge of the silica particles may be calculated by fitting the experimentalosmoticversusvolumefraction curves. Thecalculations are performed using the algorithm developed by Belloni, where the effective charge is an adjustable parameter, using Eqs. (44), (49, and the HNC equation, to take intoaccount explicitly thepairdistributionfunction of the interacting system (Belloni, 1991). The theoretical curve (solid line) for monovalent ions, scaled suggested by Eq. (44), in Fig. 15, followsthe experimentaldata reasonably well. The good agreement between the measured pressures and those calculated by the models indicates that the silica sols are well described dispersions of charged spheres, which carries an effective charge The simulations reveal an effective charge that is equal to 57 charges perparticle and this independent of the typeand concentrationof ions. The effective charge corresponds to charge density of 0.045 charge per nm2, value very close to that measured by electrophoresis 0.05 charge per nm2) at same ionic strength. As expected, goodagreement between themeasured effective charge 57) and thesecalculatedfrom Eq. (43) is shown. Comparison of the effective charge 57) obtained from osmotic measurements with the structural charge 6283) calculated from the density of surface sites (five sites per nm2) confirms that less than 1% of the surface sites accounts implicitly forthe electrostatic interactionpotential. No significant differences are observed function of the nature of counterionsfor alkali ions and tetramethylammonium ions. The discrepancy for the different counterions arises from the assumption of finite size of the ions and the fact that hydration contributions of counterions are not included in the osmotic pressure calculations. The concept of effective charge, determined by osmotic pressure measurement, is powerful in that it is easier to treat the less charged effective silica system, than to study the highly charged structural silica system, governed by complicated chemical processes. An important feature is that only “free” ions, which can transfer m o m e n t u ~ from one particle to another, account implicitly for the long-range electrostatic forces. Condensed ions, which formsurfacecomplexes,forexample,mayhave secondary effects. Extension of the approach to the effect of surface modification of silica isoutlined in the caseof the effect of adsorbed calcium ions on the osmotic pressure. A calcium-modified silica sol obtained by reaction with calcium hydroxide slurry through dialysis membrane (Persello, 1997) was compressed by osmotic stress at pH 10 and ionic strength 2 M. Theexperimentalosmoticpressure versus particle volume fractionis shown in Fig. 15 in comparison with the pressure versus volume fraction curve obtained for unmodified silica sol at same pH and ionic strength. Figure 16 shows the effects of calcium coverage fractions on the interactions between silica particles. We note that for calcium coverage fraction up to 0.6, the osmotic pressure versusparticle volume fraction curves are similar to

r=O

A X r=O.64

silica volume

Effect of calcium surface coverage fraction on a colloidal silica,on osmotic 0.6 Ca" per nm2 represents the phase transipressure curve. The solid curve at tion induced by the aggregation of colloidal particles.

those obtained for the unmodified silica sols. This is consistent with the expectation that the effective charge is not modified, since Ca2' ions are strongly boundto the silica surface sites. Increasing the calcium coverage above 0.6 resulted in a phase separation between an aqueous solutionand anaggregated silica system. Thisresult suggests that above a calcium coverage of 0.6 the forces become of shorter range and attractive.

Small-angle neutron scattering (SANS) is used to examine the structure and correlation functions of the colloidal silica dispersions. experiments were carried out on neutron diffractometers, D l 1 or D17, at the Institut Laue-Langevin (ILL), Crenoble, on silica sol with a radius14 nrn and atvolume fractionof 0.20. The experimental scattering intensity versus scattering vector curves, for the series of alkali counterions, by are shownin Fig. 17. It is remarkable that there are no effects arising from the typeof monovalent counterions even over a large range of pH 3). The effective surface charge is determined by comparing obtained from SANS experiments, to the calculated from liquid-state theory. This procedure is described in (Wagner et al., 1991). The simulation, illustrated in Fig. 17, reveals a good agreement between theoretical (solid line) and experimental (discrete points) results. The fitting procedure yields an effective surface charge of 78e per particle, which corresponds to the effective charge calculated from Eq. (43). ~otentiometric measurements indicate a number of charges ofclose to 2460. Thediscrepancy betweenthe effective charge and thepotentiometricchargeconfirms thatthe small fraction of ions involving in the effective charge is not determined by the chemical structure of Stern layer, but from physical phenomena.

10000 silica sol a=14nm

1000

9

5

M

0

100

10

lo

7 The small-angle neutron scattering curve for a colloidal silica sol in the presence of different alkaline ions. The solid curve represents the computed curve, using an HNC model. This result can be generalized into the interaction between silica particles in the presence of multivalent ions. SANS experiments were carried out on silica sol with a radius 14 nrn and at volume fraction of 0.005, as a function of surface coverage and for the series of alkaline-earth ions. Figure l8 displays the scattering plot of a silica sol (radius equal to 10 nm) with sodium ions as counterions and which is covered with calcium ions, from zero to the plateau of calcium adsorption, 0.25Ca2' per nm2and in presence or not of free calcium ions (above the plateaulimit). The adsorption characteristics of the different silica sols are tabulated in Table 13. li represents the ratio between the total number of calcium ions added in thesol and the number of calcium ions adsorbed at the plateau (0.022) and R" represents the ratio between the number of calcium ions added in the sol and the total number of calcium ions adsorbed.

Silica sol

nm)

0.01

Effect of calcium surface coverage fraction on small-angle neutron-scattering curve.

a colloidal silica, on the

Number Density of Added and Adsorbed Calcium Ions on a 10 nmRadius Silica Sol at pH 9.00. The Plateau of the Calcium Adsorption Isotherm is Found at 0.22 Ions per nn? CT (total added Ca per nm2)

0.0105

(adsorbed Ca per nm2) Ratio R CT/Cp Ratio R* CA/CT

0.0100

0.0477 0.9524

0.0306 0.0300 0.1391 0.9804

0.2300 0.2200 1.0455 0.9565

0.49’70 0.2200 2.2591 0.4427

1.0050 0.2300 4.5682 0.2289

The main difference between the scattering curves at different surface coverage are shown in the small scattering vector limit. So long as the ratio R* is close to 1, i.e., the total number of added calcium is adsorbed, the silica sol behaves as a repulsive system. The effective charge, determined by scattering curve fitting, is close to 56 charges per particle, comparable to that of the unmodi~edsilica sol (R 0). When the ratio R” becomes smaller than 1, i.e., free calcium ions are present in the sol, a strong increase of the scattered intensity is observed at small Q. The origin of this phenomenon (Schmidt,1982) provides evidence ofthe effect of the free divalent ions on the Debye screening length, reducing the repulsion and inducing flocculation. We now study the effect of the type of adsorbed ions (al~aline-earth ion series) in the situation where thesilica particles are covered with 0.03 divalent cations per nm2. The scattering plots are shownin Fig. 19, at volume fraction 0.005, pH 9.00, and ionic strength It can be shown that thereare no effects arising from the typeof divalent ions. observed in osmotic stress measurements, these experiments confirm that the effective charge is not influenced by the nature of ions “complexed” at the surface of silica.

100

80 p!

0

0.01

0.1

1

Q

The small-angle neutron-scattering curve for a colloidal silica sol in the presence of different alkaline-earth ions. The solid curve represents the computed (Q) curve.

The application of the effetive charge concept to describe the thermodynamic properties and structure of silica colloidal dispersions is well established.

The surface chemistry appears to be governed by complicated chemical processes involving adsorption of ions, dissociation of ionogenic groups, complexation of metallic cations, as well as short-range counterion-counterion correlations inside the interface. A comprehensive interpretation of interfacial properties of silica is provided from the complexationand electrochemical approach,which describes the interface in terms of distribution, nature, and density of the surface groups. The charging process of silica surfaces is closely related to the interaction between the surface silanol groups and cations present in the solution. The driving force is the proton-cation exchange reactions. These cations maybe present in the interface as simplecounterions,adsorbed by Coulombic interactions, or may formsurface complexes. These surface associations must effectively be treated as a “neutral” species that did not contribute to the ionic properties of the silica-solution interface. It is found that “association” is more important for the divalent ionsthan for the monovalent electrolyte. Another feature is that adsorption is better correlated with the hydrolysischaracteristics of the adsorbing ion rather than with the charge properties of silica. Electrophoretic experiments on colloidal silica lead to the conclusion that strong modificationof the interface has a limitedeffect on the effective charge. It has been shown that the electrophoretic charge density of silicarepresents only a few percent of the structural. The osmotic and techniques provide useful physical insight and quantitative trends for the interaction potential for colloidal silica. It is clear that the long-range interparticle interaction is explained by a screened Coulombic repulsive potential, where the charge density is given by the effective charge density. It can be expected that the molecular structure of the silica solution interface has a significant effect on the short-range interactions.

Abendroth, R. P. (1970). Colloid Interface Sci. 34591. Ahrned, M. (1966). Can. Chern. 44:1663. Allen, L. H. and Matijevic, E. (1970). Colloid Interface Sci. 33: 420. asset, D. R., Boucher, E. A., and Zettlernoyer, A. C. (1970) Colloid Interface Sei. 34:3. Belloni, L. (1991) in X-Ray Light ~ c a ~ ~ eP.~Lindner i ~ g , and Th. Zemb eds), Elsevier Science Publishers, p. 135. Belloni, L., Drifford, and Turq, P. (1984). Chern. Phys. 83:147 Belyakov, N., Soltirskii, N. Strzhenko,D. N., and Strello, (1974). Ukrain. Khirn. Zh. 40:236. Bolt, C. (1957). Phys. Chern. 1166. Bowden, Posner, A. and Quirk, P. (1977). Aust. Soil Res. 121. Budd, M.(1961). Phys. Chern. Classes 2111. Busey, R.H. and Messmer, R. E. (1977). Inorg. Chern. 16:2444.

(1994). J. Colloid Interface Sci. 163:407. Chen, S. H. and Teixeira, JA. (1986) Phys. Rev. Lett. 57:2583. Corey, B. R. (1981),in A ~ s o r p t i ~ofn Inorganics S o l i ~ - L i ~ u iInterfaces d (M. A. Anderson and A. J. Rubin. eds.), Ann Arbor, MI, Chap. 4. Davis, J. A. (1977), Adsorption of Trace Metalsand Complexing Ligandsat the Oxide/ Water Interface, P11.D. Thesis, Stanford University. Derjaguin, B. and Landau, L. (1941). Acta Physicochim. URSS 14:633. Foissy, A. and Persello, J. (1998), in The S u r ~ f f cProperties e of Silicas (A. P. Legrand, eds.), John Wiley Sons Ltd, Chap. 4B. Galasso, F. S. 970). Structure and Properties of Inorganic ~ o l i ~Pergamon s, Press, p. 89. Garofalini, S. H. (1990). J. Non-Cryst. Solids 120:1. Gonzalez-Tovar, E. and Lozada-Cassou, M. (1989). J. Phys. Chern. 93:3761. Hair, M, L. and Hertl, W. (1970). J. Phys. Chern. 74:91. Helferrich, F. (1962). Ion Exchange, McGraw-Hill Book Co. Hienlstra, De Vit, J, C., and van Riemsdijk, W.H. (1989a). J. Colloid Interface Sci. 133:105. Hiemstra, T., van Riemsdijk, W. H. and Bolt, G.H. (198913). J. Colloid Interface Sci. 133:9 1. Hill, T. L. (1960). Statistical T h e ~ * m ~ d ~ n fDover f ~ i c s Publications, , New York. Huang, C. P. (1981), in Adsorptio~of I~orgffnics Solid-~iquid I~terfaces,(M. A. Anderson and A. J, Rubin, eds.), Ann Arbor, MI, Chap. 5. Huckel, E. (1924). Phys. Z. 25:204. Hunter, R. (1981) Zeta Potential in Colloid Science, Academic Press. Iler, R. K. (1979). The Chemistry Silica, John Wiley, New York. Ingri, N. (1959). Acta Chem. Scand. 13:758. James, R. and Healy, T. W. (1972a). J. Colloid Interface Sei. 4053. James, R. 0. and Healy, T. W. (1972b). J. Colloid Interface Sci. 40:65. James, R. (1981), in ~ ~ s o r p t of’ i o Inorgffn~cs ~~ Solid-Liquid I?zterfaces (M. A. Anderson and A. J. Rubin, eds.), Ann Arbor, MI, Chap. 6. Kinniburgh, D. G.and Jackson, M. L. (l98 l),in A~sorption o~InorganicsSolid-Li~uid Z n t e ~ f f c e(M. s A. Anderson and A. J. Rubin, eds.), Ann Arbor, MI, Chap. 3. Koneshan, S., Rasaiah, J. C., Lynden-Bell, R. M., and Lee. S. H. (1 998). J. Phys. Chem. B, 102:4193. Koopal, L. K., van Riemsdijk, W. H., and Roffey, M. G. (1987). J. Colloid Interface Sci. 118:l17. Legrand, A. P. (1998), The SurfffceProperties o ~ S i l i c ~John s , Wiley Sons Ltd. Malati, M. A. and Estefan, S. F. (1966). J. Colloid Interface Sei. 24:306. Manning, G. S. (1951). J. Chem. Phys. 51:924. Marcus, U. (1989, Ion Solvation. Wiley Interscience. Marcus, Y. Chem. Rev. 88:1475. Marshall, K., Ridgewell, G. L., Rochester, C. H., and Simpson, J. (1974). J. Chem. Industry (Lond.) 19:775. Michael, H. L. and Williams, D. J. A. (1984). J. Electroanal. Chem. 179:131. Miklavic, S. J. and Ninham, B. W. (1990). J. Colloid Interface Sci. 134:305. Milonjic, S. K. (1987). Colloid Surf. 23:301.

Morel, F. M. Westall, J. C., and Yeasted, J. G.(1981), in~ ~ s o r ~ tof~norgffnics iun at S o l i ~ - ~Interfaces i ~ ~ i ~(M. A. Anderson and A. J. Rubin. e&.), Ann Arbor, MI, Chap, 7. Morrow, B. A. and McFarlan, A. J. (1994), in The Colloid C ~ e ~ i s t rSilica ~ (H. E. Bergna, ed.), American Chemi Society, Washington, D.C., Chap. 9. O’Brien, R. W. and White, L. R. (1978). J. Chem. Soc., Faraday Trans. I1 74:1607. Parks, C. A. (1965). Chem. Rev. 65: 177. Parks, G. A. and de Bruyn, P. L. (1962). J. Phys. Chem. 66:967. Parsegian, A.,Fuller, N., and Rand, P. P. (1979).Proc. Natl. Acad.Sci. U.S.A. 76:2750. Parsegian, A., Rand, R. P., Fuller, N. L., and Rau, D. C. (1986). Methods Enzymol. 127:400. Persello, J. (1989), European Patent 317,378, RhBne Poulenc Chirnie. Persello, J. (1999), French Patent Fr 99 06 Persello, J., Magnin,A.,Chang, J., Piau, J. M., and Cabane, B.(1994). J. Rheol. 38: 1845. Persello, J., Pantaleo,A., and Cabane, B.(1997),inSecondRilem Workshop on Hydration and Setting, proceedings, June 11-13. Peschel, G. and Ludwig, P. (1987). Ber. Bunsenges. Phys. Chem. 91536. Rosenfeld, Y. and Ashcroft, N. (1979) Phys. Rev. A. 20:1208. Russel, W. B.,Saville, D. A., and Schowalter, W. R. (1989), ~is~ersiuns, Cambridge University Press. Sahai, N. and Sverjensky, D. A. (1997). Geochim. Cosmochim. Acta 61:2801. Schmidt, P. W, (1982). J. Appl. Crystallogr. 15:567. Shannon, R. D. (1993). J. Appl. Phys. 73:348. Sjoberg, S. and Lovgren, L. (1993). Aquatic Sci. 55:324. Smith, J. (1954). Acta Cryst. 7:479. Sonnefeld, J. (1993). J. Colloid Interface Sci. 155:191. Strazhesko, D. N., Strelko, B.,Belyakov, N. and Rubanik, S. C. (1974). J. Chromatogr. 102: 191. Sverjensky, D. A. (1994). Geochim. Cosmochim. Acta 58:3123. Tadros, Th. F. and Lyklema, J. (1968). J. Electroanal. Chem. 17:267. Tadros, Th. F. and Lyklema, J. (1969). J. Electroanal. Chem. 22:l. Thorn, V. J. and Hancock, R. D. (1985). J. Chem. Soc., Dalton Trans. 1877. Tillmann, R. W., Radrnacher, M., and Gaub, H. E. (1992). Appl. Phys. Lett. 60:3111. von Smoluchowski, M. (1903). Bull. Intern. Acad. Sci. Cracovie 8: 182. Wagner, N. J., Krause, R., Rennie, R., D’Aguanno, B., and Goodwin, J. (1991). J. Chem. Phys. 95:494. Wakatsuki, T. and Furukawa, H. (1974). Soil. Sci. Plant Nutr. 20:353. Wells, F.(1986), StructurfflInorganic Chenzistry, 5th ed., Clarendon Press, Oxford,

U.K. Westall, J. and Hohl, H. (1980). Adv. Colloid Interface Sci. 12:265. Westall, J. and Morel, F. (1977), FITEQL: A general algorithm for the interpretation of experimental data, Technical Note No. 19, Ralph M. Parsons Laboratory, MIT, Cambridge, MA.

White, L. R., ~ a n g e ~ s d o rC. f , S., and Chan, D. Y. C. (1989), ~ o b i l i t yUser : Manual, Programs and Files, Department of Mathematics, University of Melbourne. Yates, D. E. (1975), The Structure of the Oxide-Aqueous Electrolyte Interface, Ph.D. Thesis, Melbourne University. Yates, D. E. and Healy, T. (1976). J. Colloid Interface Sei. Yates, D. E., Levine, S., andHealy, T. W. (1974). J. Chern.Soc. Faraday Trans. I 70: 1807. Yoon, R. H., Salman, and Donnay, C. (1979). J. Colloid Interface Sei, 70:483. Zerrouk, R. (1988), Surface Chemistry of Silica, PhD. Thesis, BesanCon University. Zhurav~ev,L. (1987). Langmuir 3:316. Zhuravlev, L. T. and Kiselev, A. (1962). Kolloid. Zh.

Department of Electrochemistry, Technical Universityof Lublin, Lublin, Poland, and Department of Physical Chemistry, Abo Akademi, Abo, Finland

1. Introduction 11. Almost-PureOrganicSolvents Admixtures 111. Small

of Water

50

IV. MixturesRich in Water A.Surfacecharging B. Dissociation ofweak acids C. Zeta potential V.

SolventWater A. Temperature B. Water at extremely high ionic strengths C. Heavywater References

Thesurface-chargingbehavior ofsilica in waterhas beenextensively studied. Compared with metal oxides, silica shows very low pristine point of zero charge (PPZC): most sources report PPZCof silica at pH between and Therefore, the surface charge density and the potential of silica is negative over the usually studied pH range, and the positive branch of surface charging and electrokinetic curves is rarely visited. In other words,silica is more acidic than most metal oxides. This property is attributed to the presence of surface hydroxyl groupswhich behave like weak acid: This work is dedicated to the memory of Professor Andrzej Waksmundzki.

whose p& is about 7. The values of the pKa of reaction (1) obtained by different authors are summarized in Ref. These pK, are con~iderablylower than that of monosilicic acid. The acidity of silicic acid in solution increases with the degree of polymerization /2], and the pKa of the highest polymers is close to that found for colloidal silica. This behavior can be extrapolated for solid silica, for which the acidity increases with the degree of crystallinity of the sample, i.e., the piu, of amorphous silica is considerably higher than that of quartz. Not always do the differences in pK, values reported fordifferent silica samples reflect the real difference in their acidity, because: pKa is calculated from the fractionof ionized surface hydroxyl groups, thus the total concentration of surface hydroxyl groups N , must be known. Although most authors agree that there are a few such groups per nm2 of fully hydroxylated silica surface, can vary from one sample of silica to another, and most authors do not determine the N , themselves but rather they accept N , values from literature. In some studies reaction (1) is considered as a single reaction responsible for negative surface charge of silica. In many other studies, an additional surface species is considered: SiQH

Na"

SiQNa+H+

(2)

where Na' can be replaced by any other alkali-metal cation. Certainly, the of reaction (1) calculated for a model with reaction(1) alone is higher than the Ka of reaction (l) calculated for a model with reactions (1) and (2) from the same experimental data. The advantage of models with reactions (l) and (2) is that they are able to properly predict the ionic strength effect on the u0. In such models, the increase of the from Li to CS can be explained in terms of the increase of the equilibrium constantof reaction (2) with the atomicweight of the alkali-metal cation. On the other hand, thereis no direct spectroscopic proof of existence of ESiQNa surface complexes. The surface charge densityis calculated from the difference between the volume of titrant used in titration ofsilica suspension on the one hand, and in blank titration on the other, leading to the same pH value. In the vicinity of the PZC of silica, this is a difference between two large and almost equal numbers, so the value and even the sign of the difference is very uncertain. On the other hand, the potential of silica in very acidic media is clearly positive. The scatter of literature values of PZC and isoelectric point (IEP) ofsilica can also be a result of the discussed above difference between amorphous silica, on the one hand, and crystalline silica, on the other. The following crystalline forms are known: and p-quartz, and p-cristobalite, and p-tridymite, coesite, keatite, stishovite, and each of these forms can have somewhat different PPZC. The literature on surface charging ofsilica in media different from water is rather scarce. In many studies, the magnitude of the surface charge ofsilicais controlled by means of ionic surfactants. In nonpolar media small inorganic ions adsorb onsolid surfaces, leaving someexcess of large organic ions in solution

These large ions can form a diffuse layer responsible for electrokinetic phenomena. This trend is opposite to that observed in polar solvents, in which small inorganic ions remain in solution while large organic ions adsorb. For example, AOT (anionic surfactant) charges mineral oxidespositively in dioxane-water mixtures rich in dioxane and negatively in mixtures rich in water [9]. The above mechanism of surface charging in nonpolar media should lead topositive surface charge in solutions of anionic surfactants and negative surface charge in solutions of cationic surfactants when the concentration of a surfactant is sufficiently high. Simple inorganicionsare not solvated by nonpolarsolvents so theCoulombicattraction between the bare anion and the cation is very strong, and the role of inorganic salts in the surface charging process is rather insignificant. However, most commercial chemicals containtraces of moisture, and silica itselfis also hydrated unless it is subjected to extreme drying conditions. Therefore,in real systems the charging mechanism according to reaction cannot be excluded if the solvent contains traces of a substance (not necessarily water) capable of stabilizing protons in solution. Traces of water and other polar solvents solvate simple inorganic ions and thus stabilize them. Since in microelectrophoresis the solid-to-liquid ratio is low, verylow concentrations of such substances in solution are sufficient to support surface chargedensities corresponding to potentials in excess of (plus or minus) 50 mV. Thus, while the charging mechanism of silica in solutions of surfactants in nonpolar solvents is rather clear, the formation of electrokinetic charge in “pure” nonpolar solvents remains unknown: Is this only a question of purity of silica and the solvent? On the other hand, in polar solvents, the charging mechanism of silica is probably similar as in water, i.e., the dissociated protons [reaction are solvated and the negative surface charge dueto the ESiO- surface groups is neutralized by the ions of the inert electrolyte. When the ionic strength is sufficiently high, the charging mechanism according to reaction prevails and impurities strongly interacting with silica (multivalent ions, surfactants) canbe neglected. This is manifested in stable and reproducible values of surface charge density.

Labib and Williams[lo],Spange et al. and Kasseh and Keh [l21 determined potentials of silicain organic solvents without addition of any chemicals. Labiband Williams and Spange et al. dispersed dry silica in organic media. Glass sphereswere usedin the first study and Aerosil in thesecond.Kasseh andKehgradually replaced waterby an organic solventin a stable dispersion (SnowtexXL) in a series of centrifugation-redispersion cycles and finally removed traces of water from the dispersion by molecular sieves. They succeeded in obtaining a stable dispersion in nearly anhydrous acetonitrile but in not in methanol, as the particles clustered at water content in methanol below and the electrophoretic mobility of aggregates was equal to zero. Kasseh and Keh transferred their silica particles from acetonitrile back to water and the potential was the same as in the original aqueous dispersion. In view of this result it is rather unlikely that the surface charge in acetonitrile was supported by strong adsorbing impuritiesof the organic solvent.

The designof electrophoretic cell applied in Ref. 10is suitable to measure electrophoretic mobilities in waterand other polarsolvents, but the values obtained in apolar solvents may be erroneous because of lack of uniformity of the electric field in such a cell. The results from Refs. 10-12 are compared in Table 1. The authors tried to keep the concentration of water as low as possible by dehydration of commercialchemicals ontheonehand,and ofsilica on the other. These results are denoted as “dry” in Table 1. The results termed “wet” were obtained using a dehydrated solvent and asolid equilibrated with air at 50% relative humidity [lo] or by addition of 1 of water [1l]. The data collected in the third column in Table 1 were compiled in Ref. 11 in one table with the “wet” data for acetonitrile, dichloroethane, andnitromethane, so they were apparently obtainedunderthesame (wet) conditions.However,the solubility of water in some of these solvents is less than 1% and these data are listed in Table 1 in a separate column to avoid confusion, Table l shows a common trend, namely traces of moisture cause a decrease of negative potentials. However, the difference between the effect of 1% water and of the amount of water adsorbed by silica from humid air (probably a fraction of with respect to the amount of solvent) is rather insignificant. Dispersions of silica in methanol show an exceptional behavior [12], namely that the potential is

Potentials of Silicain Nonaqueous Solvents (mV) Snowtex Glass GlassAerosil Aerosil wet dry Aerosil wet Acetone Acetonitrile Benzene Chloroform 1,2-Dichloroethane DMF Dioxane Acetic anhydride Ethanol Ethyl acetate Hexane Nitromethane Propylene carbonate Pyridine Tetrachloromethane THF Water Dichloro~ethane Ethylene diamine

dry

XL

”59 -43

120

-57 -85

-90 -100

-87

-29 -51 -45 -42 -37

-63 -88

-92

-60 -70 -60

-44

-69 -73

-53 -39 -15 -34 -36 -28

-85

-55 -45

62

more negative at higher water contents. A decrease in negative potentials on addition of water was reported for other oxides, up to asign reversal from negative to positive as a result of the presence of traces of water Apparently traces of water stabilize anions in solution and thus facilitate specific adsorption of cations making the surface of silica less negative, but little is known about the nature of these anions and cations. Spange et al. [l l] also report sorne potentials obtained from electro-osmotic measurements and these values were considerably lower than those from electrophoresis. This inconsistency canbe partially caused by difficulties in interpretation of experimental data. However, the reporteddifference by an orderof magnitude is probably due to the difference in solid-to-liquid ratio between these two methods, which is much higher in the electro-osmotic method. A signficant difference in potentials obtained by means of thetwomethodscan a result of strongly adsorbing impurities. Let us assmne that an organic apolar solvent contains traces of a cationic surfactant. discussed above,smallanions adsorb on the solid surface creating negative surface charge. In electrophoresis, the surface coverage is significant, but in electro-osmosis, with much higher solid-to-liquid ratio, the surface coverage is negligibly small and other mechanisms of creation of surface charge prevail. Therefore the potential measured by electro-osmosis is less negative. Independence of the experimental value of potential of the solid-to-liquid ratio would indicate that the potential is really under control. The discrepancy reported by Spange et al. suggeststhat some factors remain beyond control. Similar problems are faced by attempts to measure potentials in pure water: it can be hardly controlled when the ionic strength is not sufficiently high. The valuesof potential from three sources compiled in Table 1 are surprisingly consistent, considering that electrophoretic mobilities show relatively significant scatter even whenthey are measured in aqueoussystems. Ontheotherhand, these values are hardly correlated with any property of the solvent. The authors tried to find a correlation of potentials with the following solvent scales: (donor number), heatof reaction of a solvent with a reference acid, namely, SbClS Heat of reaction of a solvent with another reference acid, namely BF3 The chemical shift of 19Fin BF3 The wavenumber of the absorption maxima of dicyano-bis(1,lO phenanthroline) iron I1 adsorbed on silica. The correlations foundwere rather limited. The magnitudeof potential of silicain a solvent dependson a numberof factors, whose interplay can be very complicated. The acidic or basic character of a solvent (ability to solvate anions or cations) decides the sign of the electrokinetic charge and potential. The dielectric constant of the solvent defines directly and indirectly (via solubility and dissociation constant of ionic species) the profile of potential decay from the surface to bulksolution. Finally, the slipping plane distance (the thickness of the layer of very high viscosity near the surface) can also differ from one solvent to another. In view of the complexity of phenomena affecting the value of the potential it would be interesting to try to find correlations between the above-mentioned sol-

vent scales and the magnitude of electrokinetic charge of silica in organic solvents. At very low ionic strengths, the potential is close to the surface potential and where

is the particle radius and

(4) is the reciprocal Debye length. At sufficiently low ionic strengths, is small so the surface charge density canbe estimated from the potential [Eq. even when the ionic strength is not known. The particle radius in Ref. 10 was In Ref. 11 particle size information is not available, and 1 was arbitrarily assumed to get values of consistent with Ref.10. The surface charge plotted against the DN gives a random cloud of points. On the other hand, the plotof as a function of the solvent polarity parameter ET(30) shows a good correlation between these two quantities (Fig. I). The correlation is almost linear when a log scale is used on the ordinate (Fig. 2). There are two experimental points out of the general trend. One of these points represents water. The potential of -28 mV in "pure water" (Figs. and 2) is rather strange compared with potentials of silica at pH 7 at low ionic Considering that the absolute strengths, which are typically in excess of -50 value of potential increases when the ionic stren th decreases, the extrapolated potential at neutral pH at an ionic strength of 10- mol/dm3 (correspondingto the ionic product of water) should be much more negativethan that used in Figs. and

5

"0.000025

"0.000030 30

60

Surface charge density of silica as a function of ET(30) of the solvent.

70

13

0-5

U

A o

l

=U

13

0

13

0

U

wet

A

1

U U

35

40

50

T Surface charge density of silica as a function of ET(30) of the solvent.

2. The value of -85 mV reported in Ref. 12 for the potential of silica in pure water is more consistent with the results obtained at controlled pH and ionic strength. However, even with a clearly ul~derestimatedvaluefromRef.1 l representing water,the correlations in Figs. 1 and 2 are still better than the correlations of the potential with DN and other solvent scales considered in Refs. 10 and 11. Addition of 1-1 electrolytes NaBr, LiBr) to dispersions of silica in anhydrous ethanol causes reversal a of the sign of the electrokinetic charge of silica from negative to positive 131. The critical electrolyte concentration (isoelectric point) ranges from 0.0003 mol/dm3 for ISBr to 0.01 mol/dm3 for LiBr. Direct measurements show that the adsorptionof cations of a 1-1 electrolyte from ethanolexceeds the adsorption of anions, and that the adsorption of potassium is higher than the adsorption of sodium. This lyotropic sequenceis in line with the affinity of silica to alkali-metal cations in aqueous systems. However, the difference in adsorption between cesium and lithium on silica from diluted aqueous systems is not reflected in any significant effect of the nature of alkali-metal cations on the potential (or electrokinetic charge) [14), only the negative surface charge density increases in series from Li to Another notable difference between aqueous and ethanolic systems is that the adsorption of monovalent anions on silica from a neutral aqueous solution is negligibly small, and the adsorption of Br- from ethanol is about half of the adsorption of sodium. These results suggest that the role of simple salts in surface charging in water, on the one hand, and in nonaqueous ethanol, on the other, is completely different. In dilute aqueous solutions the alkali-metal cations

play the role of counterions, i.e., their adsorption is due to electrostatic attraction and they balance the negative surface charge formed as a result of dissociation of protons. In contrast, the adsorptionof alkali-metal cations from ethanolic solution is specific and it occurs inspite of unfavorable electrostatic conditions. Thisspecific adsorption and ion specificity under anhydrous conditions is related to a much weaker solvation of alkali-metal cations in anhydrous alcohol than in water. The nature of the species responsible for surface charging of silica in anhydrous ethanol is not known. In view of significant association of electrolytes in solvents of low dielectric constants, adsorption of complex ions,e.g. [K2Br]+ ratherthan of simple K’ ions is very likely. A similar reversal of sign of the potential of silica to that discussed above for ethanol was found by the same authors in anhydrous l-butanol. The effect of simple 1-1 electrolytes on the potentials of silica in anhydrous DMSO (up to 0.01 mol/dm3 NaBr and LiBr) and acetonitrile (NaCl up to the solubility limit) [l21 was also studied, but for these solvents no sign reversal was observed.

Dealing with solvent mixtures, one encounters problems which do not occur Gith pure one-component solvents, and which are chiefly related to preferential solvation of molecules and surfaces by one of the solvent components. However, “pure” organic solvents are in fact organic solvents containingsmallbutunknown amounts of water. These small amounts of water are not necessarily negligible in view of (probably) strong preferential solvation of traces of ions by water rnolecules. Therefore in order to control the electrokinetic properties of dispersed systems it may be expedient to add a small but controlled amount of water to the system. This approach was chosen in the studies reported in the present section. When properties of different organic solvents are compared, the concentrationsof water should be equal, and a question emerges as to whether we should compare solvents containing the same weight percentage of water or rather the same rnole fraction. Fortunately, the difference betwen potentials in organic solvents containing, e.g., 1 and 2 % of water rather insignificant, so both kinds of comparison lead to similar results. The electrokinetic behavior of silica in solutions of HCl and KOH in waterdioxane mixtures is similar to that in aqueous solutions. The potential negative when no acid or base is added. Addition of a sufficient amount of HCl leads to reversal of the sign of the potential and thus of surface charge to positive (151. Addition of KOH enhances the negative potential up to some rnaximum value and then the absolute value of the potential decreases. This kind of behavior is typical for other organic solvents and for other strong acids and bases. The results can be presented in form of potential f(electro1yte concentration) graphs, and the HCl concentration at which the sign of the electrokinetic potential is reversed can be referred to asthe isoelectric point. The total HCl concentration in the system is not necessarily equal to theequilibrium concentrationin the solution. These two quantities mustbe clearly distinguished. In microelectrophoresis, when the solid-toliquid ratio is low, thedifference between these twoquantities caused by adsorption

is often negligible. The analytical concentration of HCl in mixed solvent can be used a qualitative measure of activity of protons. In view of significant association in nonaqueous solvents it is not obvious if different acids (HC1, HN03,HC104, etc.) give acharge reversal at thesameconcentration,becausetheassociation constant depends on the nature of the anion. A glass electrode gives a qualitative estimate of changes in acidity in nonaqueous solvents, i.e., its potential is more positive in solutions of acids and more negative in solutions of bases, while the effect of addition of simple 1-1 salts is negligible. However, the display of the pH meter is unstable despite long times of equilibration. Therefore, even if the problem of standardization can be properly solved, the electrometric method to determine the pH in nonaqueous systemsisof limited significancewhen quantitative and reproducible results are required. An addition of neutral salts does not significantly affect the pH-meter reading and does not change the colorof pH indicators, but it has significant effect on the electrokinetic potential of silica. In the most polar solvents 25), the salt effect on the potential is similar to that in water, i.e., addition of a salt causes decrease in the absolute valueof the potential, but its sign is not reversed. This is the result of a decrease in the double-layer thickness. Such behavior will be termed “asymptotic,” since the c(csalt)curve asymptotically tends to zero. Concentrations of simple salts sufficient to reduce the potentials in organic solvents by certain factor are usually much lower than in aqueous systems. in less polar solvents 25), addition of a sufficient amount of 1-1 salt leads to reversal of the sign of the electrokinetic potential of silica to positive [16j. Examples of asymptotic behavior and sign reversal (l-butanol, THF) are illustrated in Fig. 3 . The results obtained in anhydrous organicsolvents discussed in theprevious section suggest that the critical value of 25belowwhich the reversal of sign ofelectrokinetic potential in solutions of simple 1-1salts is observed is also valid for anhydrous systems, and the presence of significant amount of water was not crucial. The concentration of salts leading to such a sign reversal is rather insensitive to the nature of the anion, but itis sensitive to the nature of the cation. The reversing concentration for cesium salts is lower than that for lithium salts by an order of magnitude. On the other hand,Fig. 4 shows that the critical concentrations of KC1 and CsCl in 95% dioxane obtained in two different laboratories with twodifferent samples of silicaand two different zetameters aresimilar. The reversal of sign ofthe potential was observed chiefly fororganicsolventscontainingadmixtures of water, but some data reported for nearly anhydrous systems confirm this trend. In mixtures containing the same organic componentand various concentrationsof water, the asymptotic behavioris observed in systems rich in water (high and at a low thesign of electrokinetic potential is reversed. Thisphenomenon was studied in some detail in water-dioxane mixtures, and the reversal was observed at 19 (70% dioxane) [16]. Thisvalue is significantly lower thanthe abovementioned 25 at which the sign reversal is observed in organic solvents containing a 1% admixture of water. However, the border between asymptotic and sign reversal behavior is not sharp. Even with top-quality commercial zetameters an error of 1 or 2 mV is not unusual, and this enough to get apparently positive

-butanol

0

E-5

Chargereversal l-butanol, 95% THE;) and asymptotic (99% effect of CsCl on the electrokinetic potential of silica.

0-5

0

JJI -40

-80 0-5

The effect of KC1 and CsCl on the potential

of silica in 95% dioxane.

potentials for particles which are in fact negatively charged (and vice versa). The uncertainty of experimental results is also the reason why it is not sure how the highest for which the sign reversal is observed depends on the nature of the inorganic salt. Systematic studies in this direction have not been carried out, but the results presented in Ref. 15 suggest a relatively insigni~canteffect of the anion, and a strong effect the cation with alarge gap between Na (weak sign reversing effect) and K (strong effect). Correlation between the CsCl concentrationat which the sign of the potential was reversed and selected solvent scales listed below (and combinations thereof) was tested: Hydrogen bond donation ability, Hydrogen bond acceptance ability, ~olarity/polarizability parameter,E* Reichardt’s ET(30)parameter Gutman’s donor number, Gutman’s acceptor number, AN Acity Basity K.osower’s and polarity parameters. For more information about empirical solvent scales see Refs. 17 and 18. one-parameter approach was not successful, but relatively good correlations were obtained for a number of two-parameter equations, e.g., logcrev 4.43(Acity)

10.47(Ba~ity) 12.54

r 0.86 for six sign-reversing solvents, for which the data are available. Acity and Basity express the ability of the solvent to solvate anions and cations, respectively. They havebeen obtained by statistical evaluation of many solvent dependentquantities and normalized to give 0 in a completely inert solvent and in water. For most solvents they range from 0 to 1 with very few exceptions. Equation (5) gives high c,,,’for asymptotic solvents. Thus, a highability of a solvent to solvate clliefly cations (but also anions) leads to an asymptotic behavior, while in inert solvents a sign reversal can be anticipated. In a similar study the effect of CsCl on the potential of anatase, all the organic solvents showed a sign-reversing behavior.For other oxides data are notavailable, so the above results cannot be generalized. Interestingly enough, the same critical value of dielectric constant 25 as discussed above for CsCl solutions in different organic solvents was obtained in astudy of the signof potential ofsilica in mol/dm3 Ca12 solutionsin acetone-water mixtures [19]. The potential in mixtures richin acetone was reversed to positive and in mixtures rich in water the sign was negative. In water (without acetone) at pH 8, even 0.1 mol/dm3 Ca12 did not reverse the sign of the electrokinetic charge of silica. However, the critical value of would be probably higher if higher concentratio~sofCaI2were allowed. This result suggests that admixture of organic solvents induces sign reversal in the presence of alkalineearth metal cations. The same authors studied the electrokinetic properties of silica

in the presenceof NaI but no sign reversal was observed even in almost anhydrous (99.7%) acetone at salt concentrations as highas mo1/dm3 Using the terminology from Table 2, both acetone and water are asymptotic solvents when NaI is used as the electrolyte, while in Ca12 solution water is asymptotic and acetone is a sign-reversing solvent. It would be interesting to examine if a similar correlation between asymptotic and sign-reversing behavior and the dielectric constant of the solvent as discussed above for CsClis observed for salts of alkaline-earth metal cations. The negative value of the electrokinetic potential of silica in purely aqueous systems is accompanied by positive adsorption of cations of supporting electrolyte,

Classification of Solvents According to the Effect of CsCl on the Potentials of Silica Solvent“ CsCl is practically insoluble nitro me than^ Nitrobenzene Acetone Butanone

Asymptotic behavior Formamide DMSOd DMAe DMF Acetronitriled Methanol Benzonitrile Reversal of sign Ethanolf 1-Propanol 2-Propanol 1-Butanolf 1-Pentanol 1-Hexanol 1-Heptanol 95% Dioxane 70%

35.9 34.8 20.7 18.5 7.2 109 46.6 37.8 36.7 36.0 32.6 25.2 24.3 20.1 18.3 17.1 13.9 12.5 11.1 10 18.7

organic solvent 1% water except for THF and dioxane. ’Data for pure organic cosolvents except for THF and dioxane. ‘Di~ethoxyethane. dAsymptotic behavior reported also in anhydrous system. ‘Dimethylacetamide. fSign reversal reported in anhydrous system.

while the adsorption of the anions is close to zero. Studies of adsorption of ions of supporting electrolyte on silica from nonaqueous solvents are rare. Adsorption of cations exceeds adsorption of anions from NaBr and KBr solutions in DMSO on negatively charged silica 1211. This result suggests that the surface chargeis formed by dissociation of protons from a hydroxylatedsilica surface, and the Na' and K' cations build the countercharge. In contrast, when LiBr is used as the supporting electrolyte, thepotential is also negative, butadsorption of bromideanions is higher than thatof lithium. The only explanationof this behavior is that adsorption of bromide occurs in the closest vicinity of the surface, behind the slipping plane. We must clarify now, why this particular affinity of silica surface to bromide is only r but not for NaBr and KBr. One possible explanation is formation of complex ions, e.g., [LiBr2]-, whose sorption properties are completelydifferent from those of simple ions. Adsorption of such complex ionsis enhanced by a strongsolvation of Li' ions(muchstronger than that ofNa' or K' ions) in solution.

Comparison of surface charging in water and in mixed solvents must involve a co~nmon pH scale. The studies of solvent effects on the magnitude of electrokinetic charge of silica discussed above were carried out in presumably neutral conditions, which, however, have not been controlled. The techniques used to measure pH in water are not applicable for nonpolar, nonaqueous, and almost nonaqueous solvents. On the other hand, the pH canbe measured in mixed solvents rich in water, and different approaches to this problem have been proposed. A pH meter with electrodesstandardized in aqueousbuffersdisplayssome pH value in a mixed solvent. This apparent pH value must be correctedforthe difference in liquid junction potential between concentrated KC1 solution (salt bridge of the calomel electrode) and aqueous buffer used as pH standard, on the one hand, and the solution in which the measurement is carried, out on the other. It is generally accepted that the difference in liquid junction potentials AEj is primarily defined the solvent composition and rather insensitive to the pH when the salt concentration in the standard buffer and in the unknown solutionis much lower than that in the salt bridge, and the cation and the anion in the salt bridge are equitransferent. Values of AE' for several organic solvents and solvent mixtures canbe found in the literature For dilute (up to about 30%) aqueous methanol and ethanol, AEj is negligibly small, but this is rather a fortuitous coincidence. On the other hand, for almost nonaqueous methanol and ethanol AEj is large (the correction exceeds one pH unit) and very sensitive to the solvent composition. The problem with the AEj can be avoided when the solvent composition in the buffer solution and in the unknown solution is the same. Buffer solutions in mixed solvents are usually assigned pH values in solvent scale. This means that the standard stateis defined as (hypothetical) 1 molar solutionof strong acid, in which ionion interactions are negligible, i.e., whose properties resemble an infinitely diluted solution in the solvent of interest. In contrast, for the aqueous pH scale the standardstate isdefined as molarsolution of strong acid, whoseproperties are

identical with infinitely diluted aqueous solution. The difference between solvent scale, on the one hand, and aqueousscale, on the other,is represented by a transfer activity coefficient. pH(aqueous scale)

pH(so1vent scale)

log

(6)

where AtGg is the standard molar Gibbs energy of proton transfer from water into a given solvent. The AtGk can be experimentally determined for salts but not for single ions. The widely accepted convention to split such a Cibbs energy of a salt into contributions of single ions, is that two large ions of opposite sign but similar structure,havethesame A,Gyon. Thenumericalvalues of foraqueous alcohols have been recently summarized [23].

Potentiometric titration of chromatographic silica (large porous grains of specific surface area of 388 and of CPC E251 (controlled pore glass) in many different organic-water mixtures gives surface charge densities in a mixed solvents lower (less negative) than in pure water. Thiseffect is not observed for polyalcohols (glycerol, mannitol), and the in etythritol-water mixtures is even higher than in water. Two typical examples of effects of organic solvents on charging curves of silica are illustrated in Figs. 5 and 6. The increaseof the concentrationof (Fig. 5) leads to anapproximately linear decrease in the of silica. The relative decrease in the is rather insensitive to the pH. similar behavior was observed for methanol. On the other hand, for %propanol (Fig. 6) and many other organic solvents, the decrease in the of silicais not proportional to the concentration of the organiccomponent,namely,concentrationsontheorder of a few percent of organic solvents cause asignificant decrease in the but the surface charge density isrelatively insensitive to thefurtherincrease of theconcentration of organic solvent. This figure shows that the coincidence of the osolvent/oWater curves obtained values is not always observed. Comparison ofeffectsof various solvents leads to the following conclusion. Atgiven pH, ionic strength and solvent concentration: The increase of the chain length in series of homologs leads to a more pronounced solvent effect (less negative The increase of the number of polar groups leads to aless pronounced solvent effect (more negative The effects of isomers are similar. The solvent effects are rather insensitive to the ionic strength, nature of the 1-1 electrolyte, and pH of the solution. The sequenceof the solvents in order of their effects on the surface charge of silicacoincides with the sequence of solvent polarity expressed by the ET(30)solvent polarity scale with a few exceptions. good coin-

.o pH 7 pH 7.5 pH 8

L

t:

0.8 I=

g

v

0.7

0.6 I= 0.5

DMSO Relative decrease ofthe surface charge density of silica inaqueous

1.o

g

0.7

E c

&propanol Relative decrease of the surface charge density of silica in aqueous panol.

cidence of the surface charge density in mixed solvents with the Kosower's and polarity scales is also obtained, which are linearly correlated with the ET(30) solvent scale. Application of the solvent scales listed in Table in mixed solvents rich in water involves certain difficulties. Solvatochromic shifts in solvent mixtures can be directly measured, using the same procedures as for pure organic solvents, but the problem if such results are representative. When the probe molecule is preferentially solvated by one of the components of the solvent mixture, thesolvatochromic shift corresponds to a solventcompositionneartheprobe molecule rather than in bulk solution. As a matter of fact, theexperimentallyobserved nonlinearity of solvatochromic shifts as the function of solvent composition in binary mixtures has been interpreted in terms of preferential solvation [17]. On the other hand, the solvent composition near the surface of silica can also differ from the bulk composition, and not necessarily the same component is preferentially adsorbed by silica. In the present study the values of empirical solvent parameters obtained for pure organic solvents are used to characterize solvent mixtures. Using this approach,onlysolventmixtures of thesameconcentration of the organic component can be compared. For example, the three empirical solvent polarity scales discussed above, and are nearly linearly correlated withthesurfacechargedensity ofsilica in mixed solvents of thesame weight concentration of different organic solvents. The correlation is even better when linear Combinations of solvent scales listed in Section areconsidered.The following equations:

give r 0.9 for 13 solvents for which thedata are available. The OO,~olvent/OO,w~~er is the ratio of surface charge densityin mixed solvent (the sameweight concentration for different organic components) and in water at given pH and ionic strength, ET(30), and are the values of solvent parameters for pure organic solvents, and (OO,solvenc/OO,water)O, and b are the best-fit parameters dependent of the organic solvent concentration, pH, andionic strength. The parameter is positive, and b is negative, which means that the negative surface charge of silica is high for a high ET(30) and and for a low The absolute values of the parameters a and increase with the solvent concentration, but this increase is not linear (cf. Fig. 6). Figure 7is an illustration of a correlationbetween solvent scales and surhcecharge density of silica. Qualitatively similar correlations are obtained when solutions of equal mole fraction rather thanof equal weight concentration of the organic coniponent are compared. Adsorption of alkali-metal cations on silica is slightly lower than the surface charge density (both canbe expressed in mol/m'). The difference between these two quantities expresses the surface chargewhich is not neutralized by the counterions adsorbed in the direct vicinity of the surface and it can be used to estimate the potential 11261. Addition of organic components causes a decrease of the adsorption of alkali-metal cations in such a manner that the difference between these two quantities is approximatelyconstant [27]. However,adsorption of alkali-metal cationsfrom 50% methanol is approximatelyequal to theadsorptionfrom

.o

cc)

0.9

d 0.8

0.7

0.9

1

- 7 Correlation between the surface charge densityof silica in mixed solvent(0.1 mol/dm3 NaCl, pH 8,20% of the organic component) and linear combination of (30) and solvent scales.

water, while the surface charge densityin this mixed solvent is lower than in water [28]. Extensive data on surface charging and adsorption of alkali-metal cations for the same sample of silica are available only for aqueous methanol and ethanol. Adsorption of lithium on silica in the presence and absence of urea, methylurea, and N-methyl acetamide was measured [29], but potentiometric titrations were not carried out. In view of the above-mentioned correlation between the adsorption of counterions and the magnitude of surface charge, it seems that also these three organic components cause a decreaseof the surface charge density of silica. Unlike silica, the surface chargeof metal oxidesis rather insensitive to admixture of organic solvents (except for DMSO) thus the correlations with solvent scales discussed above [Eqs. (8) and are not valid for other oxides.

Dissociation constants of weak acids in organic-water mixtures are usually lower than in water (301. This result is in line with the following equation [31): 4000nNAr,? 3

where is the association constant (reciprocal dissociation constant) of a 1-1 electrolyte, ri is theion size parameterandthesolvent is characterized by its dielectric constant. Even in heavy water whose dielectric constant is only slightly lower than thatof water, thedissociation constants of acids are lower by a factor of ten than in water, while in methanol and ethanol, thedissociation constant is lower by several orders of magnitude. The logarithm of dissociation constant in mixed solvents is approximately a linear function of the mole fraction of the organic cosolvent (Fig. 8). Silica is treated as a weak acid in reaction In this respect, the decrease of negative surface charge of silica might be due to a low dissociation constant of surface silanol groups in mixed solvents. Different methods have been used to calculate log Graphical methods (double extrapolation) have a disadvantage that the extrapolated segments are often notlinear and show some scatter of data, thus log can be only roughly estimated. An alternative to graphical methods is to calculate the parameters of double layer using a fitting routine, e.g., FITEQL [32]. The problem in such curve fitting is that models with few adjustable parameters fail to properly reflect the pH andionic strength effects on the surface charge and diffuse-layer potential. In moresophisticatedmodelsinvolvingmore parameters, good agreement between theory and experiment is obtained with many different sets of parameters.Modelswithmanyadjustableparameters arenot useful in studies of solvent effects on the because each of these parameters candependonthemedium.Therefore,the diffuse-layer model was selected,

7.5 7.0 6.5 6.0

5.5 5.0 4.5 0

20

40

60

Dissociation constants of acetic acid in aqueous acetone ethanol.

80

and in aqueous

although this model has alimited ability to properly predict ionic strength effect on the surface charge of silica. In this model

and there are only two adjustable parameters,namely the total number of surface sites and Let us assume six sites per nm2 irrespective of the solvent composition. This assumption is notobvious:adsorption of organicmoleculescan block some surface hydroxyl groups and this may be the main reason why the surface charge density of silica in organic-water mixtures is lower than in water. The other assumptionis the supportingelectrolyte is fully dissociated in solventsof interest, i.e., in Eq. (l 1) is solvent independent at constant analytical concentration. With a constant and Ns, Eq. (1 1) predicts that at given 00,mixed solvent

(%nixed solvent)

CTO.water

Ewater

but in fact the relative decrease in is much more significant that it follows from Eq. (12).Values log in some mixed solventscalculatedfromthesurface charging data in 0.1 KC1 using the diffuse-layer model with the discussed above assumptions are summarized in Table 3 [33,34]. Table 3 shows that the decrease of ofsilica in mixed solvents is usually slightly less significant compared with the dissociation constants of weak acids in the same solvent mixtures. For example, the difference in piy between water and 10% methanol is 0.14 for surface silanol groups and for acetic acid and 0.16 for benzoic acid. In this respect, the comparison of dissociation of weak acids with the

The Solvent Effect on the Logarithmsof Dissociation Constants of Surface Silanol Groups (Diffuse-Layer Model) and Organic Acids Solvent

Surface silanol groups Acetic acid Benzoic acid Reference

Water

-7.93

-4.76

-4.20

5% Methanol 10% Methanol 20% Methanol

-8.00 -8.07 -8.37

33 1 -4.8 -4.90 -5.07

-4.29 -4.36 -4.58

5% Ethanol 10% Ethanol 20% Ethanol

-8.07 -8.19 -8.26

-4.93 -5.11

-4.50 -4.76

5% Acetone 10% Acetone 20% Acetone

-8.15 -8.21 -8.36

-4.86 31 -4.99 -5.27

-4.54 -4.88

34

surface charging of silica supports the reaction (l) as the charging mechanism of silica.

Electrophoretic mobilities in systems rich in water can beeasily measured, but calculation of potentials can be a problem. The Smoluchowski equation: mobility Electrophoretic

(1

where is the viscosity, is only valid for 100, i.e., for high ionic strengths and large particles. Such systems show a limited stability and they usually settle down before the measurement canbe carried out. Onepossibility to stabilize the systemis to reduce the ionic strength, another is to reduce the particle size, but the system becomes stable beyond the rangeof applicability of Smoluchowski equation. When c 100, the electrophoretic mobility depends not only on the potential, viscosity, and dielectric constant, but it is also a function of the shape and size of solid particles. In other words, in polydispersedsystems,the particles have different mobilities even if the potential is constant. Fortunately, silica can be obtained in form of monodispersed, submicron, spherical particles, sufficiently stable for electrokinetic measurements The calculations of the potential can be carried out using a commercial computer program [36]. At sufficiently low ionic strengths and concentrationsof the organic component, the potential of silica is independent of the solvent composition, although the electrokinetic mobility is much lower in alcohol-water mixtures than in water due to high viscosity and low dielectric constant 1371. The coincidenceof potentials of silica in mixed solvents and inwater is unlikely to be fortuitous, as itis observed in mixtures of waterwith different cosolvents. Moreover,a similar coincidence is observed for other oxides[8]. These results confirm the relevance of the bulk values of dielectric constant and viscosity in the relationship between the potential and electrophoreticmobility. In other words, there is no need to introduce “local” values of ionic strength, dielectric constant, and viscosity which might be different from the bulk values, e.g., because ofspecific adsorption of one of the solvent components. Direct measurementsindicate positive adsorption of alcohols onsilica from diluted aqueous solutions[26], thus thelocal viscosity would be higher and the local dielectric constant lower than the bulk values. For high ionic strengths 0.01 mol/dm’), the electrokinetic behavior of silicain mixed solvents is different from that in water. The positive branch of the electrokinetic curves is insensitive to the solvent composition, but the negativepotentials in 20% aqueous ethanol or methanol are significantly lower than in water at pH 7. The effect becomes more pronounced for higher concentrations of organic solvents and for higherionic strengths. While for 0.01 mol/dm3 KCI, the IEP in mixed solvents is unaffected by the admixture of organic solvents, a clear shift of the IEP to higher pH values is observed in 0.1 mol/dm3 KCl. Systematic studies on the dependence of the magnitude of this shift on the natureof organic solvent have not been carried out. The solvent effect on the electrokinetic potential of silica is consistent with the surface charging and counterion adsorption behavior discussed in

theprevious section. ~ualitativelysimilar and even morepronouncedsolvent effects on the electrokinetic behavior of metal oxides have been reported [S].

The physical properties of water relevant to formation of surface charge, e.g., the dielectric constant and ionic product are temperature dependent, thus the studies carried out in water at temperatures different from room temperature arerelated to studies in nonaqueous solvents. The negative surface charge of silica and adsorption of alkali-metal cations at given pH and ionic strengthincrease when the temperature increases [38]. One can match potentiometric curves(as well as curves of alkali-metal cation adsorption) obtained at various temperatures by a shift of about 0.012 pH unit per l K. This shift corresponds approximately to th ture effect on the ionic product of water ICw. When is plotted aga pKw/2, the curves obtained at various temperatures are identical. Figure 9 shows charging curves of silica at three different temperatures. The course of the adsorption of alkali-metal cations as a functionof is temperature independent. In view of the difficulties discussed earlier, the temperature effect on the PPZC of silica is not known. For other oxides the PPZC decreases with the temperature (endothermic enthalpy of proton adsorption), butthis enthalpy becomesless endother~icfor

-0.05

-0.10

-0.15

silica

L

m

a, -0.20 L

U

15 25 35 deg

-0.25

Temperature effect on the surface charge density aqueous NaC104.

of silica in 0.1 mol/dm3

oxides whose PPZC falls at low pH values. Extrapolation of the available data to PPZC at pH3 (silica) yields the enthalpy of proton adsorption close to zero or even exothermic (increase of PPZC with temperature). Knowledge about the temperature effect on the electrokinetic behaviorof silica is rather limited [39].

Very high ionic strengths induce a change in physical properties of water, namely the dielectric constant of aqueous solutions of 1-1 salts in the molar range is lower by a few percent than that of pure water or dilute aqueous solutions. As was discussed above, it is difficult to determine the exact PZC of silica from potentiometric titration. Also a pH measurement at very high ionic strengthsis not a standard procedure, that is, special pH buffers are required or a correction for liquid junction potential must be introduced. Anyway, there is no indication that the PZG at extremely high ionic strengths would be different from the PPZC: constant values of PZC are reported for KC1 [40] up to mol/dm3 and for NaCl [41] up to 4 mol/dm3. In contrast,the IEP of silica is clearly shifted to higher pH values at ionic strengths on the orderof 0.1 mol/dm3. Increaseof the ionic strength up to 0.7 mol/dm3 does not cause further shift in iep theof silica [42]. This behavior of silica is different from other oxides [43,44] whose 5 potential is positive over the entire pH range at sufficiently high concentrations of certain salts. The normal lyotropic series observed for silica CS Rb K Na Li is probably reversed at concentrations 3 mol/dm3. These resultsmay suggest that a partial dehydration of small cations enables their adsorption behind the slipping plane, and this makes the 5 potential more positive. A similar mechanism canbe proposed to explain the shift of the ZEP in mixed solvents. In that case, the partial dehydration occurs at lower ionic strengths than in water. Another possible explanation of the enhanced adsorption of lithium at high ionic strengths and of the shift of the IEP is that the thickness of the immobile layer of water at silica surface increases with the ionic strength. Most likely, reaction (l) explains the charging mechanism only at relativelylow ionic strengths and pH values, while at high ionic strengths, surface complexation according to reaction (2) and even formation of positively charged surface complexes =SiOH

Na'

SSiOHNa'

is possible.

An admixture of heavy water (up to 50% by weight) does not affect charge density of silica. In contrast, heavy water signi~cantlyreduces dissociation of weak acids (Table4) [45-501. In mixtures of heavy water with water, the log is approximately a linear function of the atomic fractionof deuterium (Fig. Thus, the comparison of the effect of heavy water on surface charging of silica with its effect on dissociation of weak acids does not support reaction (1) as the charging mechanism.

The Dissociation Constants

Acetic [45] Boric [46] 2,4-Dinitrophenol [47] 2-Nitrophenol [45] Phosphoric [48] Selenic [49] Water A ~ ~ o n[SO] i u ~ Cyclohexylam~onium[48] Hydrazinium [48] Imidazolium [48] ~ o r p h o l i n i u[48] ~ Tris H' [48]

Acids at 25°C

in H20

in D 2 0

4.76 9.21 4.2 7.23 2.15 2.57 14.00 9.25 10.57 7.98 6.98 8.49 8.07

5.31 9.73 4.66 7.81 2.42 2.92 14.96 9.81 11.25 8.60 7.47 9.1 1 8.67

5.30 5.25

acetic acid in

mixtures

5.20 5.1 5 5.1

I

100

Dissociationconstants of aceticacidin from Ref. 46.)

H20-D20 mixtures.(Adapted

In contrastwith organic solvents, in mixtures of heavy water and water either of the solvent components is unlikely to specifically solvate the surface. Thus, the significant effect of organic solvents on the of silica and absence of such an effect for heavy water is probablydue topreferential sorption of organicmolecules of silica. This preferential adsorption is also responsible for the difference between silica and other oxides whose is insensitive to admixtures of organic solvents [8].

1. S. K.Milonjic. Colloids Surf. 23:301 (1987). 2. D.N. Strazhesko,V.B.Strelko,V. N. Belyakov, and S. C. Rubanik. Chromatogr. 102: 19 (l 974). 3. J. Lyklerna. Adv. Colloid Int. Sci. 2:65 (1968). 4. D. Parfitt and J. Peacock. Surf. Colloid Sci. 10:163 (1978). S. A. Kitahara, in ~~eno~ena (A. Kitahara andA. Watanabe, eds., Dekker, New York, 1984, p. 119. 6. rison. Colloids Surf. A 71: l (1993). 7. a, of Academic Press, New York, 1995. 8. M. Kosmulski. Colloids Surf. A 95:81 (1995). andori, K. Kon-no, and A. Kitahara. Bull. Chem. 9. Jpn. 57:3419 (1984). 10. E. Labib and R. Williams. J. Colloid Interface Sci. 97:356 (1984). 11. Spange, Simon, G. Heublein, Jacobasch, andM. Boerner.Colloid Polyrn. Sci. 269: 173 (1991). 12. A. Kasseh and E. Keh. J. Colloid Int. Sci. 197:360 (1998). Zhukov and I. Varzhel. Kolloid. Zh. 52:781 (1990). 13. 14. 15. 16.

17. 18. 19. H. A. Ketelson, R. Pelton, and M. A. Brook. J. Colloid Int. Sci. 179:600 (1996). 20. H. A. Ketelson, R. Pelton, and M. A. Brook. Langrnuir 12: 1134 (1996). 21. I. Varzhel, A. N. Zhukov, L. G. Levashova, and S. V. Uspenskaya. Vestnik LCU 499 (1990). 22. Y. Marcus. Solvation, Wiley, New York, 1985, p. 249. 23. Y. Marcus. Pure Appl. Chem. 62399 (1990). 24. M. Kosrnulski. ColloidInterface Sei.179:128(1996). A. Rodrigues, P. J. Sposito. J. Colloid Interf. Sci. 211:408 (1999). 25. dewicz, and M. Kosrnulski. J. ColloidInterface Sci. 157 (1990). 26. M. Kosrnulski. Colloid Interface Sci. 156:305 (1993). 27. M. Kosmulski. Polish Chem. 67: (1993). 28. M.Kosmulski. Bull. Pol. Acad. Sei. Chem. 41:325 (1993). 29. G. Peschel and P. Ludwig. Ber. Bunsenges. Phys. Chem. 91536 (1987).

30. J. J. Christensen, L. D. Hansen, and R. M. Izatt, Proton Ionization Heats ~ ~ e r ~ o ~ y~uuntities, n u ~ i c Wiley, New York, 1976. 31. R. M. Fuoss. J. Am. Chem. Soc. 80:5059 (1958). 32. A. Herbelin and J. Westall, FITEQL ver. 3.1, Oregon State University, Corvallis, Oregon,1994. 33. L. Avedikian. Bull. Soc. Chim. (France) 2832 (1971). 34. L. Avedikian and N. Dollet. Bull. Soc. Chim. (France) 12:4551 (1967). 35. W. Stober, A. Fink, and E. Bohn. J. Colloid. Interface Sci. 2662 (1968). 2.10, Universityof 36. L.White, D. Chan, and Ch. Mangelsdorf,Winmobilver. Melbourne. 37. M. Kosmulski and E. Matijevic. Langmuir 8:1060 (1992). 38. M. Kosmulski. Colloids Surf. A. 83:237 (1994). 39. A. Revil, P. A. Pezard, and P. W. J. Glover. J. Geoph. Res. B 104:20021 (1999). F. Tadros and J. Lyklema. Electroanal. Chem.Int. Electrochem. 17:267 (1968). 40. 41. G. H. Bolt. Phys. Chem. 61:1166 (1957). 42. M. Kosmulski. Colloid Interface Sci. 208:543 (1998). 43. M. Kosmulski and J. B. Rosenholm. J. Phys. Chem. 100: l681 (1996). 44. W. N.Rowlands, R. W. O'Brien, R. J. Hunter, and V. Patrick. J. Colloid Interface Sci. 188:325 (1997). 45 R. Gary, R. G. Bates, and R. A. Robinson, Phys. Chem. 69:2750 (1965). 46. V. Gold and B. M. Lowe. J. Chem. Soc. A 1923 (1968). 47. L. Pentz and E. R. Thornton. J. Am. Chern. Soc. 89:6931 (1967). 48. M. Paabo and R. G. Bates. Phys. Chem. 74706 (1970). 49. A. Vesala and V. Koskinen. Fin. Chem. Lett. 2:145 (1975). 50. R. F. Prini, C. R. de Pattin, K. Tanaka, and R. G. Bates. Electroanal. Chem. 144415 (1983).

This Page Intentionally Left Blank

ST Microelectronics, Agrate, Italy EniChem, Centro Ricerche di Novara, Novara, Italy stry, University of Perugia, Perugia, Italy

I. The Paradigm

of Heterogeneous Catalysis

11. From Dispersed Silica to Reactive Silica A. Pyrolysis B. Fracture C. Strain D. ~ a d i a t i o ndamage 111. RadiationDamage in Dispersed Silica A. Theimparteddamage B. Silicon-link vacancy C. Oxygen-bridgevacancy D. Environmental stability of radiation defects E. Reactivity of native defects in a controlled atmospherethe cases studied experimentally IV.CO2Addition toIon-Bombarded Si02 A. Expectations B. Experimental results C. Theoretical studies C2H4 ReactionswithIon-BombardedSi02 Expectations B. Experimental results C.Theoretical studies References

l

Surfaces represent a way to impart bulk electronic and vibrational properties to condensed-phase atoms exposed to a gas or liquid atmosphere. As far as nonlocalized electronic states are involved, unexplainable reactivities in atomic or molecularterms are manifested. Wesimply mentiona few examples: For metals in condensed phase the electron affinity and ionization energy coincide and equal theworkfunction of theconsidered face; theworkfunction on tungsten lowindex faces is higher than the fluorine electron affinity; and atomic cesium ionizes spontaneously when it comesin contact with tungstensurfaces (this phenomenon is practically exploited in spacecraft ion engines). For semiconductors, doping canbe used to populate preferentially donor or acceptor states, thus conferring to the surface reducing or oxidizing properties, while bandgap engineering can be used to narrow orwiden the gap. The combination bandgap engineering, doping,and micromachiningshouldeventuallyallowtheintegration of controlledchemical reactivity with suitably designed microreactors. In spite of theextreme potentials of thosetechnologies,themainstream of catalysis has remained (possibly because of cultural gaps) almost uniquely concentrated on the attempt to mimic biological catalytic systems. Enzymes are the prototypes of those systems. All enzymes have a proteic part; in most of them the proteinhostsametallorganiccomplex (referred to as prostheticgroup) which behavesasthe active catalytic site of theconsidered reaction. Theprosthetic group is more or less tightly bound to theprotein; if it can shuttle from one protein to another after reaction with the substrate, it is called a coenzyme. Grafting the active site to the protein inhibits its dissolution in the cellular medium and its eventual loss throughthemembrane.Theenzyme, therefore, conjugatesshape selectivity {resulting from the stereospecificity the enzyme with respect to the reactant) with chemical selectivity (imparted by the prosthetic group): Enzyme

Shape-selective protein

Chemical-selective group

While the protein selects the reactant from among the many with which it is in contact,theprostheticgroupaddressesthereactionalongthe useful pathway among the numerous possible ones. Mimicking this scheme is considered the best way to carry out industrial reactions, too. Hence follows the paradigm of heterogeneous catalysis: Catalytic efficiency

Shape selectivity

Chemical selectivity

Spectacular examplesof industrial catalytic systems which mimic biological ones are the Ziegler-Natta and Kaminsky polymerization catalysts. In both cases the active site is a transition metal, admitting either tetrahedral or octahedral coordination, grafted to a support (crystalline MgC12 for the Ziegler-Natta catalyst or suitably substituted cyclopentadienyl for the Kaminsky catalyst) which imparts a structure to the growing polymer. Reproducing the efficiency and elegance of biological systems is, however, difficult, and the research has been concentrated either on chemical selectivity or on shape selectivity.

C ~ e ~ i c selectivity al has been preferred by peopleworkingonhomogeneous catalysis. Functional groups which have been either predicted or identified to catalyze homogeneously acertain reaction have been grafted to otherwise inert,highly dispersed, supports: grafting to a lleterogeneous support is due to theneed to have, at the end of reaction, the catalyst in a different phase from products; high dispersion guarantees that most of the catalyst is in contact with reactants; and chemical inertness is dictated by the need not to interfere with catalyst activity. Due to its chemical inertness and to the possibility of getting highly dispersed materials, silica is an almost ideal support for heterogenizing homogeneous catalysts. Also, silica polymorphism makes this substance an almost ideal candidate for the preparation of materials with highsize selectivity. Size selectivity is achieved by dispersing suitable templates in silica sols. Removing the templates after gelling (usually done by burning or by supercritical extraction) leaves a structure formed by a silica skeleton containing a number of cavities. The sizeof the cavities is determin~dby the size of the template: small molecular templates produce small cavities, while large cavities can be obtained using micellar templates. The template determines (together with the gelling condition) the amorphous crystalline nature of the resulting structure: molecular templatesrnay produce amorphous aswell as crystalline silicas (the latter referred to as zeolites), whilemicellar templates usuallyproduceonlyamorphous silicas. Though silica polymorphismallows many degrees of freedom in the silica structure, certain sizes and shapes can only be achieved by adding to the gelling silica other species, like cations of the Groups I, 11, and 111, toonto be terminated by nonbridging oxygen, like silanol groups O), Si Thesubstitution of mentsfor silicon in atetrahedral Si02 network rnay impart to the network interesting structural and even catalytic properties (in which case one usually speaks silica ~ o ~ i ~For g )instance, . (1) the substitution of titanium for silicon in silicalite (a pure Si02 zeolite) produces a material silicalite) with redox catalytic activity; (2) the substitution of an element A Group ,for silicon produces a negatively charged stoichiometric defect,

O)$i whose net charge can be counterbalanced by inserting into the zeolites a suitable amount of cations of Main Groups I or 11; or (3) the insertion of A in the network in the vicinity of a silanol may result in extremely strong Br~nstedacids, provided that the oxygen-acceptor distance is such to allow the formation of a bond:

(-O),Si-O

+(-0)3Si-00+-A-(O-)3

(1)

Silanol groups may act as "Trojan horses" to impart wanted properties to the silica by g ~ a ~ t suitable i ~ g groups. For instance, the surface of silica gels (which has a hydrophilic character because of the silanols it hosts) may become hydrophobic after reaction with alcohols ROH (R CH,, C2H5, etc.):

-O),Si

OR

-O)3SiOR

H20

This is just an example; a lot of other possibilities are discussed in

Though numerous, thepossibilities offered either by doping the gel with wanted species or by grafting suitable groups to the silanols remaining after gelling and drying are not endless.A new way to impart chemical reactivity to dispersed silica, with the advantageof separating the stageof preparation of the silica network from the stage of its doping, has been proposed by the present authors in recent yearsOur previous studies have considered, 12-51 or C2H4 both experimentally and theoretically, the bombardment in atmosphere [6,7], providing a certain agreement between experimental data and theoretical expectations.

Highly dispersed materials have become in last years the focus of large attention. The “soft chemistry” required for their preparation[S] does indeed allow new materials to be prepared. For instance, the sol-gel technology allows enzymes to be embedded in dispersed oxides [g]. In this way the fragile catalyticsites are protected against environmental factorslike pH or temperature, still remaining directlyaccessible to reactants. In ordinary matter the atomic densityof any condensed phaseis of the order of cm-3. The preparationof a substancein a dispersed phase with densitylower than the density characteristic of a bulk phase is not trivial and requires the use of “tricks”. For instance, burning of templates is used for zeolites, while supercritical extraction is used for aerogels. Highly dispersed SiQ2 is usually prepared via the sol-gel technology, exploiting either the sodium silicate route [lo] or the alkoxide route [l l]. The Si02 gelis dispersed through the whole original solution in which was prepared, so that its atomic density is controlled by the initial dilution of reactants and by reaction products. The inner part of the gel becomes accessible to the atmosphere only when the solvent and reaction products have been eliminated. This is usually possible either by slow drying (that produces relatively compact structures referred to as xerogels) or by supercritical extraction (which produces structures, referred to as aerogels, where can be as low as 10-3po). Special pore structures (or even shortrange crystalline order) can be obtained by adding to the sol suitable templates which are eventually removed by burning them in an oxidizing atmosphere. Amorphous xerogels and aerogels can be seen as reticulate polymers formed by randomly oriented and coiled chains, each. chain being formed by Si(Q-)4 tetrahedra in which OH terminations are occasionally present (for areview of structure and growth of silica condensation polymers see Ref. 12). Thoughthesiloxanicbridge is quiterobust(the Si-0 bond energy being around 5.5 eV [13]), it is cleaved by water even at room temperature:

Since this process is more efficient in the vicinity of superficial E S i Q H terminations (because they are water adsorption sites) and produces new =SiOH sites, it eventually leads to the autocatalytic destruction of the SiQ2 network.

The silanol groups canbe exploited for the functionalization(i.e., for conferring wanted properties) of silica. This is often obtained through the reactions ESiOH

HC1

ESiOH

ROH

ESiCl ESiOR

H20

H20

(3)

where R denotes an alkyl group. Reaction (2) is effective at temperature around 1000°Cand can be used to eliminate OH terminations, while reaction (3) is effective even at a temperature of 100-200°C in an autoclave and can be used to produce hydrophobic silica. The ways for silica functionalizationhave recently been reviewed in Ref. 1. Together with silanol terminations, silica may contain a lot of other defects. Most of them have been characterized by electron paramagnetic resonance (EPR) and are discussed in Ref. 14. A list is given in Table 1, while their energetics is discussed in Ref. 13. For ourpurposes, it is sufficient to mention that the formation of an oxygen-bridge vacancy (OBV) and atomic oxygen, =S~-O-S~E

ESi""SiE + O

requires an energy of 11 eV, while the formation of a silicon-link vacancy (SLV) and atomic silicon requires eV-neither the OBV nor the SLV are therefore expected to be formed by thermal activation at room temperature. Except for its sensitivity to water, silica is quite unreactive and is often used as catalytic support, in view also of the possibility to control its hydrophilic-~ydrophobic character, that allows the catalyst to be adapted to the reaction medium. Several routes have been tried to impart chemical reactivity to otherwise almost inert silica; they are discussed in the following.

The first way to impart chemical reactivity to porous silica was perhaps discovered at the end of the 1960s. "Reactive silica" was obtained by methoxylation in an autoclave of silanol groups and pyrolysis in a vacuum of the resulting adducts [16]. The structure of the active sites of reactive silica has not been completely clarified, though reactivity has been hypothesized to be imparted to silica by the creation of two close silicon radicals stabilized against recombination by a peroxidic bridge

Observed or Hypothesized Radical Defects ofSilica Defect center Siloxyl center Silico-peroxyl center (unrelaxed oxygen-bridge vacancy) Tetrasiloxylcenter(unrelaxedsilicon-linkvacancy)

Formula

Si" ~Si-0"

Si =Si" *Si= 0")4

Similar results are obtained starting from other terminations (ethoxy, propoxy, etc.); though the pyrolysis produces a reactive silica, whatever the alkoxy group, the nature of the reactive site depends, however, on the precursor [17]. Vacuum pyrolysis of alkoxy groups is not the unique way to prepare reactive silica. For instance, it is possible to prepare a reactive silica by adding SiC14-nH, (PZ 3) molecules to a gelling SiOz solution; this produces the addition to the silica skeleton of SiH, moieties [18,19],

(4

SiOH

SiC14-ttHn

(S SiO),-,Si

which after pyrolysis eventually revert to

H,

(4

n)HC1

(E SiO)4-plSin".

Starting from an original observation of Hochstrasser and Antonini, that freshly preparedSi02surfacescontainEPR-active E' centers [20], Radtsig and his coworkers developed in the late 1970s a technique for the preparation of reactive silica surfaces by crushing in a helium atmosphere[21-251. The fractureis expected to occur through the homolytic cleavage of Si-0 bonds and the formation of strained siloxanic centers (SSCs) ESi 0 Si E. Via a combination of adsorption, calorimetric, and EPR techniques, Radtsig et al. showed that, though only a minor portion of the formed defects survive the spontaneous annealing processes, the survivingdefects preserve their original nature-the remaining reactive sites are just the and siloxyl centers, and the SSCs. The Si 0 bond energy in SSCs was found to be around 3.2 eV. The typical amounts of centers observed after crushing is in the range 0.1-1 of the surface sites, the SSC being the most abundant one ~31. The stability of the center in a COz atmosphere andof the siloxyl center in a CO atmosphere are summari~ed in the following observations: The siloxyl center SiO" reacts with CO, formingan EPR-active center Si[OCO]", which ismore stable than the reactants by 1-52 eV. Thesame=Si [OCO]" center, for which Radtsig et al. proposed the structure =Si-O-C*=O is also formed by reaction of the center with COz, this reaction being an equilibrium with an enthalpy difference of 0.78 eV: Si"

Si[OCO]"

The above energies are consistent with the initio density functional calculations by Re et al. Though the reactionof CO with the E' center is allowed, at 300 K and pressure below 10 Torr the adsorption of CO on "Si" is negligible. No information on the reaction of CO2 with the siloxyl center has been reported. This state of affairs allows the ~ S i 0 centers " to be stoichio~etricallyreverted to "Si" by CO adsorption at300 K at pressure and decomposition at 470 K of the formed adduct =Si[OCO]":

SiO"

CO

Si[OCO]

ESi"

CO2

Even theSSC has highreactivity. For instance, Radtsig and Khalif[23] showed that the SSC reacts with N 2 0 forming a peroxidic bridge,

and with GO2 forming a kind of organic carbonate,

Silica aerogels are characterized by a distribution ofSSCs. In fact, the aerogel network has a structure very close to the silica skeleton in the gel before drying. This structure, in turn, was a quasi-equilibrium configuration for the system containingthesolventand is manifestlyfarfromequilibrium when thesolvent is suddenly removed. A way to restore, at least partially, equilibrium is via straining of the cycles forming the Si02 network. Silica aerogel is expected to manifest a reactivity similar to that identified by Radtsig et al. for the SSC. In astudy of CO conversion toCO2 in an O2 atmosphere,undoped silica aerogels were shown to be able to catalyze this reaction [26]. A catalytic cycle which would allow this process (for which each step ther~odynamicallyallowed for Si 0 bond energy in the interval 2.5-3.5 eV, versus an Si 0 bond energy of 3.2 eV [23]) was described in Ref. 27. Another possible cycle requires an activation stage, I

O),Si.. O

Si(O

CO

O),Si[OCO]"

O),Si[OCO]" "Si(0 O),Si" "Si(0

C02

after which the following catalytic cycle is activated: ()),Si" "Si(O -)3 O2 o),Si.. . O . . Si(O O),Si O),Si

-Si(O CO "[oco]S~(O "Si(O -)3 CO

O),Si[OCO]" "Si(0

O),Si.. (IO. . O . Si(0 O),Si "0 Si(0 O),Si "fOCO]Si(O O),Si "Si(0 6 O),Si[OCO]" "Si(0 O),Si"

2co

0 2

C02

2c02

The above cycle is possible only thanks to the presenceof a ofsilicon radicals. The availability of two radicals in close proximity, but unable to recombine, offers extremely exciting catalytic perspectives.

The possibility of exploiting the radiation damage to impart a catalytic activity to silica was first addressed by Yates and Lucchesi, who demonstrated that y-irradiated SiOz is able to catalyze ethylene polymerization[28]. Neither the active sites nor the poly~~erization mechanism were however identified, and the activities in that direction were stopped. One major difficulty in using massless (like y-rays) or light (like electrons) projectiles for imparting damage tosilica is the fact that the nature and amountof the imparted damage can hardly be predicted. This dif~cultyholds for all projectiles whose stopping mechanismis essentially the electronic energy loss, which results in electronic excitation of the target. The damage imparted by heavy ions, in an energy regime where stopping is dominated by the nuclear energy loss, is much easier to predict, because it results from two-body elastic collisions of the projectile with the atoms it invests. This suggested to the present authors the idea of using ion-beam bombardment for the functionalization of dispersed targets. The reasonswhich motivated us to focus our attention onto the radiation damage, rather than on the implanted species, are discussed in Section 111.

The discovery that irradiated solids have special reactivities is as old as modern physics. Becquerel did indeed discover natural radioactivity by observing the formation of latentimages (Le.,of zoneswitha special nonvolatile reactivity) in photographic plates in the vicinity of uranium ores [29]. In the early period of cosmic-ray physics, photographic plates were systematicallyused asnuclear detectors, and many of thefundamental particles of the physicists’ zoo were discovered by means of nuclear emulsions. The persistence of the radiation damage in insulators is still exploited in dosimetry and for a class of nuclear detectors-nuclear-track detectors [30]. In spite of its central role in 20th century physics, the chemistry of irradiated solids has remained practically ignored by chemists (with the remarkable exception of polymers, whose increase of reticulation after irradiation is exploited practically). In authors’ opinion, this lack of interest was due to the following reasons: (1) most of the radiation damage is imparted via electronic excitation and the resulting excited configu~ations aretruly difficult to predict, and (2) most of the radiation damage is imparted to the bulk of the solid and hence is inaccessible toreactantsotherthanetchants. In recent years, however,thesefactorshave lost their constraintcharacter. Machines such as ion implanters have become of widespread use and can be operated in a regime(of stopping dominated by the nuclear energyloss) where the radiation damage is easily predicted Low-density targets, in which most of atoms are exposed to the surroundings, can now be prepared via sol-gel technologyfora large variety of substances (41 elements, at least, have been reported to have been involved in this technology [32]).

ilic

Though both these technologies have become of widespread use by more than two decades, ion-beam techniques haverarely been used in chemistry, presumably because of cultural and cost considerations-among the very few studies in this field, we mention the increase of catalytic activity resulting from copper implantation into Raney catalysts [33] and the catalytic activity toward HCOO position of the Ar+-bombarded Ti02 (1 10) surface [34]. The idea of studying, and possibly exploiting, the radiation damage imparted to highly dispersed solids bad not been seriously considered until the present group showed that it is possible to merge ion bombardment and sol-gel technologyforthe synthesis of otherwise impossible structures [2]. In general, the dispersed phase is a metastable configuration not obtainable by quenchinga stable configuration at high temperature, so that the route for its production requires nonequilibrium, however soft, processing. a rule, heating at high temperature lir a dispersed phase will produce its collapse to the equilibrium configuration, with the destruction of the structure. This implies that highly energetic (and presumably highly reactive) configurations, like the ones which may be obtained by thermal excitation at high T,cannot be obtained in dispersed structures. This conclusion, however, holdstrue in a situation in which degrees of freedom of the dispersed phase are excited by thermal activation. It is, however, possible to concentrate the energy on a degrees of freedom by high-energy (say, E 0.1 I keV) collisions resulting from the impact of an ion beam onto a dispersed target. The result is an unreactive dispersed network embedded with highly reactive sites easily accessible to the atmosphere-a new, otherwise unattainable, structure,

When an energetic ion beam impacts a target, one has to consider two effects: the chemical effect duetothe implanted species itself, and theradiationdamage imparted to the target. The nature of the radiation damage is almost independent of the implanted species, and depends mainly on the nature of the target. discussed later, the radiation damage imparted to Si02 by heavy projectiles, in the stopping regime dominated by the nuclear energy loss, should be constituted by =Si" "Si= diradicals and by silicon-link vacancies in addition to oxygen and silicon knocked-on atoms. The implantation of a foreign species into an Si02target followed by a moderate-temperature heat treatment in an inert atmosphere may produce a set of new materials. An understanding of the possibility of creating new materials by ion implantatiollintoporous silica requires, however,apreliminaryknowledge of the chemical reactivity of the radiation damage imparted to the target. This is an almost virgin field of research, the only published works being-to the best of our knowledge-a preliminary characterizationby this group of the reactivity of argonimplanted porous Si02 in CO2 or C2H4 atmospheres and their theoretical analyses via initio density functional calculations[5,7] or molecular dynamics[3].

To first approximation, the projected range R, of an ion into dispersed material is obtained from the one in a condensed phase by the scaling rule

R,

(4)

Since at E 100 keV the implanted ions are distributed in bulk SiOz typically at depth 0.1-1 p m [31], implantation in silica with 10-3-10"po may extend to depthintherange 1-103 pm, which implies that the i ~ p l a n t ~ t i o n involve region only. To give more quanrelatively thick samples rather than a titative idea of the in-depth extensionof the region crossedby the projectiles, in an SiOz sample with density of 0.1 &m3 the range of Ar' at 50 keV is 1.2 pm, while the range of at 200 keV is 40 pm. In the implantation into condensed materials, most of the knocked-on atoms remain in the proximity of the site where they were generated, therefore allowing fast recovery of the radiation damage. In dispersed system the knock-onstoo are displaced by larger amount than in condensed phase, as described by Eq. (4). The lower the density, the higher the knock-on range; the higher this range, the more difficult the recovery of the radiation damage. Combining the large separation between the recoil and the center resulting from the collision with the high reticulation of dispersed silica, one mayinfer that the i ~ p a r t e dradiation damage has a high roba ability of surviving the spontaneous recovery ~ e c h a n i s ~ Rather, s, since the sites produced by the bombardment are exposed to the atmosphere, the evolution of'the d a ~ a g e dlayer will be related to the residual a t ~ o s p h e r eirt which the target is i ~ m e r s e ~ . The nature of the damage imparted to silica by heavy projectiles is easily predicted: the diradical center left by an. oxygen recoil, and four vicinal EEE Si- 0" centers left by silicon recoil. These defects are produced via two-bodycollisions in the time scale 10"5--10-'3

The collision sketched in Fig. l is expected to produce an SLV. This vacancy has, in the absence of relaxation, the tetrasiloxyl structure (2) ("native form"). The m i n i ~ u menergy (threshold energy Et)necessary to form in collision the SLV is the one required to cleave simultaneously four Si-0 bonds, Et(SLV) 22 eV. Due to therelative independence of the Si 0 bond energyof the anglebetween ond and the one characteristic of the ideal tetrahedral configuration, is predicted to rearrange to form one or two peroxidic bridges (3). Silicon peroxides are known fromsynthetic chemistry /35]. Appreciable differences are noted amongdifferent calculations of the 0 0 bond energy, ranging from0.9 eV [7j to 2.2 eV 131 in the absence of straining. Straining effects are, however, expected to reduce this value appreciably.

shown in Fig. a collision displacing an oxygen atom is expected to form ilicon radicals usually referred to center. Evidence for the was reported in Ref. for single-crystalline quartz irradiated with

ilica Si

Si

l

I Si=

=Si

0

l

Si

Si

-0

Sketches of the centers produced by elastic collisions of heavy projectiles in Si02 ("native defects").

prays. The minimum energynecessary to formin a collision the center one required to cleave simultaneously two Si bonds, Et( E") l l eV. The center is expected to be stable withrespect to networkrelaxations producing an increase of the silicon-silicon distance. When the relaxation produces decrease of the silicon-silicon distance, the formationof an Si Si bond becomes possible, though steric effects will interfere with this process. The steric effect was studied theoretically by Re et al. Together with the cyclic siloxans (Si(OH)2Q)~ (n 2, 3), they considered the molecule (HQ)3SiQSi(QH)3; then they took away one siloxanic oxygen per molecule and allowed the structure to relax with the formation of an Si-Si bond. The bond energy was found to increase from 1.2 eV for n 2 to 2.9 eV for n 3, the value in the absence of any stress being 3.5 eV (see Table 2). Since larger cycles admit more relaxed configurations, the above-mentioned calculations suggest that an Si Si cannot be formed not only for large silicon-silicon nuclear separation but also in the presenge of strong steric constraints. Inthis case, the center is expected to be stable. The OBV excitation to the ionized state,

requires an energy of about 5 eV (the difference between the ionization energyand electron affinity of 0)3Sio).This energy is much too high to allow the ionized state to be formed by thermal excitation. Inthepresence of base like H2Q,

Dependence of the Si- Si Bond Energy on Strain in Molecular Complexes Simulating OBV with Different Topologies

3.5

(HO)3Si Si(OH)3

0 2.9

(HO)2Si"----

Si(OH)2

0

I \

l

(HO)2Si-Si(OH)2 from

(=Si)20 or =z:SiOH, the oxygen can, however, form a silicon cation, *s~(o-)~

:o(

(-o)~s~+

Lewis adduct with the

+:o(

thus reducing appreciably the formation energy. The existenceof threefold coordinated oxygen at was first considered by Rudra and Fowler and Allan and Teter Boero et al. confirmed the presence of such a center in a quartz via ab initio molecular dynamics simulations accurate calculations in the frameof density functional theorywere reported for simple model molecules by Cerofolini and Re [do]. These considerations suggest that in the a ionic con~gurutionis able to form a dativebond to the siliconcation, the u p p r o ~ i ~ a t us e l ~probable as the radical colz~g~ration.

The following reactions are thermodynamically allowed: silanol and silane terminations after exposure to H20: =Si" "Si= + H 2 0 HO-Si=

formation of vicinal

(2) formation of a peroxidic bridge between silicon atoms after exposu =Si" "Si=

=Si-O-O-Si=

(3) NZ formation of a nitrene bridge after exposure to N2: -Si" "si= N2 =Si--N=N---Si=

Sta~ilityof the Silicon-Link Vacancy While the unrelaxedSLV is expectedto react with almost everything (an important exception is, however, given by because of the poortendency of oxygen toward catenation), its peroxidated form is expected to be stable in dry air. It should, however, revert to a fully hydroxylated silicon vacancy after reaction with H 2 0 , througha set of consecutivereactions *&st theformation of hydroperoxyl terminations and eventually of four silan inationsand H202. In turn, a fully hydroxylated silicon vacancy can regenerate siloxanic bridges after heating at high temperature. The centers so formed are expected to be highly strained and hence to have the same reactivity as the SSC identified by Radtsig et al.; the strain is associated with the fact that the reticulation degree of the Si02 polymer has decreased. After the overall process, therefore, its

3. Exploitingthe n v i r o n ~ e ~ t a l ~ n s t ~of~ the i l i t yNative Defects to Control the Si02 Network Topology A sequenceof cycles formed by argon implantation in anO2 atmosphere, exposure to H 2 0 at suitable temperature, and high-temperatureheating, could therefore be used to vary the topological structureof the initial reticulate polymer. For instance, the original network (characterizedby the existence of closed interconnected loops) can be first modified by reducing the number of loops to which each tetrahedron belongs and increasing the average number of members of each ring, then transformed into a tree (connected polymer withno closed loops), and eventually into a set of chains (unconnected polymer).

The experiments describedin the followingwere performed to study theinsertion of CO2 andC2H4 into argon-bombarded poroussilica at anenergy of 300 keV. Argon was chosen to avoid interference effects due tothe implantationof a foreignspecies. CO2 was chosen because itwas expected not to react with the peroxidic centerand therefore to be a marker of the OBV. G2H2was chosen because itwas expected to react with both the OBV and SLV and therefore to provide information of the totality of the radiation damage. The combination of the pieces of evidencefor CO2 and C2H4insertions was expected to provide substantial information on thestructure and chemical reactivity of radiation-damaged porous SO2.

Though (l) during argon implantationat 300 keV an appreciable amount of energy is imparted to thetarget as electronic (rather than nuclear) energyloss, (2) as many silicon recoils as SLVs areproduced,and (3) at highconcentrationthenative defects interact with each other forming more complicated defects, for the interpretation of the observed data it was sufficient to assume that the native radiation damage is constituted uniquely by OBVs and SLVs.

erofolini et

For the prediction of the nature of the adducts formed by CO2 after reaction with the OBV in the radical state, we considered only the electrophile nature at carbon of CO2 attack on metals of the main groups [41].

CO2-OBV In view of the stability of silicon in a CO2 atmosphere, the relaxed OBV too was presumed to be stable in that atmosphere. For silicon-silicon distance in the center in the range 2.0-5.0 the reaction pathway was then expected to involve the consecutive reactions of each of its constituting radicals. This leads to the eventual formation of an ester-like group: -Si"

"Si=

CO2

~ S i - C ( O ) O " "Si= =Si-C(O)O-SiE

For either shorter or longer silicon-silicon distance, a covalent insertion of CO2 is no longer possible. An ionic pathway like =Si" "Si= CO2 ~ S i - C ( O ) O " "Si= ~ s i - c o o - +si= involving electrophilic attack on the carbon via electron transfer from the other radical tothe -C(O)O" groupandtheformation of acarboxylate group -COO-, is forbidden because endothermicby approximately S eV [5]. The formation of the carboxylate, however, is thermodynamically allowed to OBVs in an ionized configuration:

The confirmation of the expectations (i.e., the formation of the Si-C(0)OSi and =Si- COO- +Si= adducts) was not trivial, because no model molecules containing Si- C(0)O- Si or Si- COO-+Simoietieshad everbeen synthesized.

COS-SLV Since silicon peroxides do not react with CO2, the relaxed SLV was expected to be stable in a CO2 atmosphere. Organic peroxocarbonate or inorganic diperoxocarbosilicate (like or (7) in Fig. 2) have not been considered because: (1) to the best of our knowledge, such moieties have never been reported, and (2) isolated siloxyls do not react with (an extended study of the reactivity of isolated siloxyls towards several gases is given in Refs. 21-25). Though these facts are still insufficient to exclude conclusively the synthesis of SLV-CO2 adducts, ignoring them seems to us a reasonable working hypothesis.

~ ~ o r ~ ton ion

F1 Structural formulas of CO2-peroxoadducts whichmightbeformedafter the reaction of CO2 with an SLY:organic peroxocarbonate (6), and diperoxocarbonate

The experimental analysis was essentially based on infrared (IR) spectroscopy. The IR spectra were determined: (1) in as-implanted conditions,(2) after heating at 500" C, (3) after exposure to air, (4) after reheating at 5OO0C,and after reexposure to air.In all cases themeasurements were carried out in avacuum at Torr. Figure 3(a) compares theIR spectrum of the as-prepared sample(solid line) with the one after high-temperature annealing (dashed line); each of these spectra is actually obtained by subtraction of a reference spectrum determined on an unimplanted sample. The comparison shows that the formed species are remarkably stable with respect to temperature. This stability seems to exclude that the species are inserted as Si 0 -C (this structure is indeed known to undergo pyrolysis at around 400°C 16]), thus supporting our contention that SLVs do not react with Figure 3(b) shows that after exposure to air at room temperature the complex spectrum of Fig. 3(a) disappears and there is the riseof a peak which can be reasonably attributed to a carboxylic acid group -COOH. Thecombination of these facts withreportedattributions (for molecules or adducts-like CF3COOH, CF3COO-, etc.-which might resemble the sites produced during the abovelisted process) [42] led the authors of Ref. 2 to ascribe: (1) the sharp peak at 1515 cm-' to the antisymmetric C 0 stretching mode of the carboxylate group -COO-, (2) the broad peak at 1595 cm-' to carboxylate ions co-ordinated to cation or to hydrogen(via hydrogen bonding), and (3) the broad peak at 1710 cm"I to a carbonyl C 0 local mode. In the absence of molecules containing the -O)3SiCOOand O)3SiC(0)OSi(0 moieties, these attributions should, however, be considered putative. initio quantum mechanical calculations were then performed to provide additional supportive evidence. Density functional calculations actually showed that all the above adducts have certain stability range [S]. For the free carboxylategroup(-O)3Si-COO-thecalculatedvibrationfrequency of the antisymmetric 0- 0 stretching modecoincidedalmostexactly (to within 10 cm"> withtheobservedone,thussupportingthe original interpretation;the

bsotba nce

(b)

(a) IR spectra observed immediately after the bombardment in a COz atmosphere (solid line) and after annealing at 500°C (dashed line; in the spectral region below 1530 cm-' the solid and dashed lines coincide within the precisionof the line width); (b) IR spectra after different exposures to air.

whole band from 1590 cm" to 1750 cm" could, however,be explained uniquelyin terms of C=O stretching in =Si Si" ester-like groups, provided 0 -Si bonds are sufficiently stressed [S]. In this context, the that the Si -Cslight intensity decrease of the carboxylate peak and an almost equal intensity increase of the C=O peak(bothfollowingfromhigh-temperatureannealing) suggest that the thermal treatment is responsible for the transformation of ionic pairs "Si -COO- +Si= into covalent bridges Si-C(0)O Si". In the absence of water, however, the efficiency of this process is modest. Exposure of the sample to air afterheating produced theincrease, with approximately first-order kinetics, of a signal at 1705cm" whose intensity increased almost linearly with time. Another heat treatment at 500°C destroyed this signal and was responsible for the reappearance of a band at 1735 cm". The reexposure of the sample to air resulted in the increaseof a band at 1705 cm-' whose intensity increased with kinetics resembling that of autocatalytic reactions and therefore different from that observed during the first exposure to air. While the isolated carboxylate group is stable after annealing at 500"C, its hydrated form "SiCOOH H O S i z does allow much easier reconstruction because of the numerous possibilities of hydrogen bonding. Because of this a subsequent annealing leads to the formation of ester-like bridges. Via calibration with a carboxylicacid, it 'was also possible to determine the total OBV amount-approximately lo2 centers per impinging ion. Though this value is

still appreciably lower (byoneorder of magnitude) than the totalamount of produced defects, it is, however, appreciably higher (by two orders of magnitude) than the total amount of defects remaining in the condensed target after spontaneous relaxation.

Since COz was presumed not to interact with the on the center.

attention was concentrated

Density- function^^ Description of the Possible Adducts of CO2 to the Center The adducts tothe and centers were studied by Re et al. in the context of density-functional theory [5]. Two CO,-E' adducts can be hypothesized: Si C(0)O" and ESi OC'O. Both of them have a certainstability with respect to reactants: 0.37 eV the former and 0.71 eV the latter The lower stability of Si C(0)O" than that of Si OC"0 does not contradict the electrophilicity at carbon of COz toward metalsof the main groups-the reaction may in fact initiate with the formation of the ESi C(0)O" radical and terminate (after a characteristic time re) withtheformation of the more stable adduct Si-OC"0, as observed in Refs. 21-25. If in the vicinity of Si-C(0)O" there is a second=Si", an ester bridge is formedaccording to reaction (6). If within the second silicon radical is not at the right distance to form the ester bridge, Si-C(O)O"isomerizes to =Si-- 0C"O. Even in this case, however,the final compound is the ester-like bridge:

The energy gained in this process depends markedlyon the distancebetween silicon atoms: the insertion in rings containing two or three silicon atoms is stabilized by an energy 2.55 eV or 3.92 eV, respectively,the stabilizing energy in the absenceof any strain being 4.27 eV. These data show the importance of the environment in determining the final state of the adduct. normal-mode analysis demonstrated that the vibration frequency of the local C = 0 mode in the ester bridge depends significantly on the stress, in such a way that the whole band in the region 15901750cm" can be explainedasdue to a distribution of ester-like groups with different strains. It is interesting to note that all the considered relaxed centers react spontaneously with COz formingester-like bridges, the overall reaction being slightly exothermic. The very fact that 90% of the produced defects did not react with suggests, however, that the reaction of the relaxedOBV with CO2 is kinetically limited by the cleavage of the Si-Si bond. According to Ref. 5, the infinitely separated ions Si -COO- and ESi' are unstable by 5.0 eV with respect to the dissociation:

similar energy can,however, be gained by forminga Lewis adduct between +Si= and siloxanic oxygen [40j, so that thecarboxylategroup is expected to form at ionized OBVs.

To get more information on the initial stages of COz addition, initio molecular dynamics simulationswere performed. The C02-E" interaction was studied via the Car-Parrinellomethodconsideringa system formed by a moleculeintbe vicinity of two (HO)3Si* radicals, originally at a silicon-silicon distance of 3.5 roughly corresponding to the silicon-silicon distance before the collision which produced the El' center. When CO2 was put in a configuration for which the van der Waals radii of its constituent oxygen atoms did not overlap with those of silicon, the system was found to relax via CO2 escapeand the formation of a bond between silicon atoms, with no CO2 insertion in between silicon atoms. Figure 4 shows the total potential energy V and the silicon-silicon distance as a function of time t. When one oxygen atom in COz was posed in at a distance from one silicon smaller than the sumof van der Waalsradii, the system was found toevolve in such a way that the total potential energy showed two sharp minima (Fig. 5). These minima were attributed to the formation of carboxylate or ester-like confi~urations. These attributions came from a comparisonof the C bond lengths, which for the ester-like adduct are different from one another and very close to the C and C=O bond lengths, respectively, while for the carboxylate configuration are almost equal to each other and between the C- 0 and C =O bond lengths.

R 0

(fs)

Time evolutionof the total potential energy V (continuous line) and siliconsilicon distancer (dashed line) fora pathleading to the formation of an Si Si bond.

iiica

9

Time evolution of the total potential energy (continuous line) and siliconsilicon distance (dashed line) for a path showing the insertion of GO2 via the formation first of carboxylate group and eventually of an ester-like group.

Consistently with the fact that the principal mode of reaction in maill-group chemistry is as an electrophile at the carbon center the addition to the center took first place with the formation of an Si bond and only later of an Si 0 bond. Another interesting case was observed when was put with both oxygens close to the silicon atoms in the center. The molecule first bent and then formed an adduct where carbon had a tetrahedral con~guration andwas bonded to two silicon atoms (with the correct bond length) and symmetrically to oxygen atoms, with bond lengthcharacteristic of single C - 0 bond (Fig. 6). This center, described with the formula (HO)~Si-(O*)C(O.)-Si(OH)~, is a unique one in which carbon is found in tetrahedral configuration.

2

Thebombardment in an ethyleneatmospherewasperformedbecauseofthe expected reactivity of C with both OBVs and SLVs.

l The OBV in the relaxed form was expected to react with C2H4.The unrelaxedOBV was expected to react with CzH4 forming one or the other of the adducts shown in Fig. 7 . The first of these occursat ionized OBVs and does not require any bond cleavage and presumably has avery low activation energy. All other adducts, being necessarily preceded by the formation of the adduct EE Si- CH2 C*H2,

Time evolution of the total potential energy (continuous line) and siliconsilicon distance (dashed line) for a path showing the insertion of CO2 via the formation of two Sibonds. require the cleavage of the n" bond and therefore presumably require an activation energy. Becauseof this state of affairs, the expected insertion mode atroom temperature involves the formation of the Lewis adduct which in turn reverts [mainly to (9) or afterheating at hightemperature.Preliminary densityfunctional calculations have confirmed that all the above adducts have a certain stability

C2H4--SLV Figure 8 shows a reasonableradical pathway adding C2H4 to anSLV or a partially peroxidated SLY. The reaction would involve first the cleavage of the n" bond in

Adducts which are expected to be formed after reaction of the unrelaxed with C2H4.

ilie

/si\o/

/si\o H

/H

H

H H

H'

'H

7\ Suggested pathway addingC2H4to an SLV or apartially peroxidated SLY.

by siloxyl in and the formation of a carbon radical in hydrogen abstraction by another siloxyl (attack by free radical at a atom of a substratemolecule is indeed knownto occur in radical reactions and the formation of two close carbon radicals in (1 which eventually recombine via the formation of a bond. This pathway leads to the formation of center (1 which the organic part is grafted to silica through an oxygen bridge, resembling the one characteristic of ethers.

A complex evolution of the IR spectra was observed according to the treatments undergone by the sample. The spectra gave evidence for sharp peaks attributed to CH3, CH2,and C 0 moieties, in addition to the one due to co-ordinatedC2H4. The peaks observed after different heat treatments are shown inFig. 9. Though these peaks evolved in complex way according to the heat treatmentsundergone by the sample, the following general features were observed: (1) the C = C signal, observed in as-implanted samples,was found to disappear completely after heating 00°C; (2) the decreaseof the C C signal was associated withan increase of the signal; and (3) the CH3 and C = 0 varied in nonmonotonicway with the heat eatment, but they were strongly correlated irrespective of their intensities. The pieces of evidence (1) and (2) can easily be interpreted by assuming that most of theadsorbedethylene transformsafterannealing in -CH2 in between the silicon atoms originally formingthe center. The final result of this process is the formation of composite in which the void space in Si02 is progressively filledwith CH2 CH2 moieties covalentlyinserted in between different trees forming the silica skeleton. This composite had a thermal stability up to 500°C at least.

91

1

Relevantinfraredspectrameasured on: a sampleagedfor 100 h in dry nitrogen (top, dashedline) and henceheated at 500°Cfor 30minin a vacuum (top, contilluous line); another sample a few hours after the preparation (bottom, dashed line) and hence exposed for 72 h to air (bottom, continuous line). In theOH spectral region (3000-4000 cm-') the spectra are raw data; in the other regions the data are difference spectra with respect to a sample bombarded in the residual ionimplanter atmosphere.

ationbetweenthe and signal is explained if one is formed or lost rding to thethermaltreatment.The pathway leading from the center to is, however, not trivial to identify and is not conclusively determined by our data. At this stage, and without pretendi~lgto be conclus we favor a mechanism, sketched in Fig. 10, involving the formation of a center (in which the Si bond is weakened by a hydrogen bonding) followed by bond rearrangeme ucing simultaneously a strong d vinyl alcohol CH2 which is known to isomerize above interpretations, it was possible to determine the total and SLV amounts. In the hypothesis thatall SLVs and unrelaxed OBVs react with it was foundthat (1) the amount of OBVs corresponds to approximatelyIO2 per impinging ion; and (2) the concentration of SLVs is smaller than concentration of centers by a factor of 3. Therefore,the amount of

Suggested mechanism accounting for the observed variations of the CH3 and C=O IR peaks.

determined by experiments involving the addition of C2H4 is fully consistent with that determined by experiments of addition, and the SLV-to-OB'V ratio is consistent with their different threshold energies.

The possible adductswhich may be formed by reaction of C2H4 with the and SLV defects have been studied by Re et al. in the context of density-functional theory The major conclusions of that study are summarized in the following. Complexes Formed by C2H4 ~ d d i t i o nto the C2H4 Addition to theRelaxed OBV. For (HO)3Si- Si(OH)3thebond dissociation energy in is 3.5 eV. Though strain effects are responsible for a substantial reduction of this energy in cyclic species (2.9 e'V or 1.2 eV in the three- and twomembered cycles, respectively), C2H4 insertion between the silicon atoms is practically prevented except, possibly, for the most strained species. b. CzH4 A d d i t i o ~to the ~ n r e l a x e dOBV in State. When the two silicon atoms are kept apart by the rigid silica structure, the resulting center can be thought of as constituted by two essentially noninteracting silicon radicals. The reaction of this center with C2H4 is therefore expected to involve the consecutive reactions of each of its constituting radicals, the first step being the addition of CzH4 to one of the two radical centers to give a radical adduct. The silicon center was modeled by a (HO)3Si" molecular species and several pathways were considered. Table 3 reports the reaction energies computed for the addition of C2H4 to (H0)3Si" and the possibleevolutions of the initial radical adduct.Theunique

Reaction Energies for C2H4Reactions at One Site of the Reaction (HO), CH2 CH2 (HO),Si CH2 CH: (€€O),Si CH2-CH; (HO),Si-CH2-CH; (H0)3Si-CH"-CH3 (H0)3Si-CH"-CH3

Center

AE (eV) (HO)3Si CH2 CH; (HO),Si CH" CH, (HO),Si CH CH2 H" (HO),Si-CH=CH"+H, (HO),Si-CH =CH2 H" (HO),Si-C" =CH2 H2

-1.3 -0.2

Unstable

thermodynamically allowed pathway is a shift from the original radical species, with the unpaired electron on the carbon, to an a-carbon radical (H0)3SiCH" CH3. Table 3 shows a high reaction energy for the addition of ethylene to the (OH),Si" radical (l eV), in agreement with the experimentalevidence that the addition of a silyl radical to a C C double bondis a remarkablyfacile process [44] (the addition of silyl radicals to a C= C double bond is the key step in the hydrosilylation of alkenes [45,46]). The hydrogen shift leading to an a-carbonradical is a slightly exothermic process as expectedon the basis of the stability scale of organic radicals [47]. Both radicals can recombine with the adjacent second silicon radical of the centerleading to a -CH2 -CH2 or -CH(CH3)bridged species, (HO)3SiCH2CH2Si(OH)3 and (HO)3SiCH(CH3)Si(OH)3, respectively. The relative abundance of thetwo speciesis determined by a subtle compromise between energetic and geometric factors. Indeed, although the a-carbon radical is 0.2 eV more stable than the one, its formation involves a 1,2 hydrogen shift and is expected to have a substantial activation energy. Therefore, whenever allowed by the geometrical arrangement of the two silicon radicals of the center, the acarbon radical, which is formed firstly, combines immediately with the adjacent silicon radical leaing to the -CH2 -CH2 bridged species; only for siliconsilicon distances too short to allow such a recombination is the lifetime of the acarbon radical high enough to allow the hydrogen shift and finally the formation CH(CH3)- bridged species. of the Hydrogen abstraction, leading to vinyl-silicon a species (OH),Si CH CH2, or a double hydrogen abstraction as leaing to a vinyl-silicon radical (OH)$iCH CH" or (OH)3Si C" CH2, respectively, are thermodynamically unfavored by about 1.5-2.0 eV, and are therefore quite unlikely. The final results of the CZH4 interaction with the center can be obtained by considering the subsequent reactionof each intermediate of the first ethylene addition to one (HO)3Si"radical, considered above, with the secondsilicon radical. A11 the considered reactions are reported in Table 4 together with the computed reaction energies. The reactions leading to the -CH2 --CH2 or -CH(CH,)bridged species are both strongly exothermic (by 4.9 eV and 5.6 eV, respectively), thus confirming the original attributions of Cerofolini et al. [6]. Although most of the other possible reactions are also exothermic, they involve a single or double

T

Reaction Energies for

C2H4 Addition at the

If

Center

Reaction

A E (eV)

2(H0)3Si" 2(HO), Si" 2(HO), Si" CH2 2(HO),Si" 2(HO),Si"

C2H4 C2H4 CH2 C2H4 C2H4

(HO),Si (HO),Si (HO), Si (HO)$i (HO)3Si

CH2 CH2 Si(OH)3 CH(CH3) Si(OH), CH CH2 H Si(OH), CH( CH2) Si(OH), H2 CH =CH Si(OH), H2

-3.8

hydrogen abstraction and a sizeable energy barrier is therefore expected. The corresponding products are thus quite unlikely, in agreement with the absence of any experimental evidence for them c. C2H4 ~ d d i t i o nto the ~ n r e l ~ x eOBV d in State. The addition of C2H4 to the ionized OBV is expected to involve an acid-base pathway; in particular, the nucleophilic nature of the ethylene molecule suggests that the reaction with the ionic pair initiated with the co-ordination to thesilicon cation and simultaneous displacement of the co-ordinated siloxanic oxygen:

(-O),Si++

/Si= :O

\si

C,H,

(-O),Si++"

,,Si= :O \Si=

long as the geometry of the OBVdefect in this ionic configuration remains unperturbed, this ionic adduct is expected to be stable. Upon annealing, the center may relax to a different geometry allowing the formation of a -CH2 -CH2 species probably via a silicon carbocation intermediate O),Si CH2 C Re et al. modeled the ionic configuration with the (H0)3Si+ and(OH)$3i- ions and computed energies for the co-ordination of ethylene to the silicocation and its subsequent evolution to a carbocation and a bridgedspecies [7]. The results of the calculationsare listed in Table 5. Thistableshowsthattheco-ordination of ethylene to (HO),Si+ is an exothermicprocess by 4.0 eV. Ethylene,however, can actually displace thesiloxane or silanol oxygen stabilizing the ionic pair only if its co-ordinationenergy is higher thanthat of these Lewis adducts. Modelingtheco-ordination of asiloxane or silanol group to a silicocation by that a H 2 0 molecule to a (OH)3Si+ gives a co-ordination energyof 5.0 eV This energy is an overestimate of the actual co-ordination energyin silica because steric effects reduce, even largely, the possibility of formingthe Lewis adduct. Actually, the calculations performed on the Si- Si bond formation energy show that this strain effect can easily give contributions up to 2.0 eV, so that ethylene could effectively bind to the silicon cation displacing dative oxygen in strained configurations.

Reaction Energies for C2H4Reactions at One Silicocation of the Ionic Configuration Reaction

A E (eV)

CH2

CH2 (HO)3Si'

CH2

(HO),Si't-

"4.0

CH2 CH2 (HO),Si'

CH2

(HO),Si

+0.05

CH; CH2 (HO), Si CH,

CH2

CH;

(HO),Si

CH'

+0.8

On the basis of the analogy with the carbocation chemistry [47], the z-co-ordinated species isexpected first to evolve to a silicon carbocation and then to undergo a hydrogen shift: CH2 (HO),Sif 11 (HO),Si CH2 CH:, (HO),Si CH' CH3.

CH2 Table 5 shows that the @-carbon carbocation has a slightly higher energy than the z-co-ordinated species (less than 0.05 eV) and is therefore a viable intermediate in the formation of a "-CH2 CH2 bridge. The positive charge transfer from silicon to carbon is also expected to be facilitated by the increase of electrostatic pairing energy CH2' Since the original SiSi separation in the ionic configuration is ca. 4 A, it allows the formation of a second Si-C bond between the @-carbon of silicon carbocation an( the silicon anion (in (HO),SiCH2-CH2 Si(OH), the Si Si distance is 4.2 A).The instability of the carbocation (0.8 eV higher than the carbocation) and the unsuitable geometry of the ionic configuration make it difficult the formation of a -CH(CH,)bridge. These results are in good agreement with the experimental IR evidence showing the simultaneous disappearanceof the adsorbed-ethylene signal and the increaseof the CH2 signal upon annealing.

2. Complexes Formed by Addition to the SLV The relatively low bond energy (calculated in the interval 0.9-2.2 eV for the completely relaxedSLV) suggests that a low activation energy is required to cleave the peroxidic bridge. In turn, this suggests that the reactivity of the SLV can be understood by considering this center in the fully radical configuration. Though the SLV center is constituted by four ESi 0"radicals, it is expected to react with C2H4, involving only two out of the four available siloxyls. The SLV center was consequently modeled by two (OH),Si -0" molecular species. The possiblereactionpathways were studied by assuming that the first step of the C2H4insertion involves one of the two radical centers to give a radical adduct: (HO)3SiO" CH2

CH2

(HO)$iO

CH2-CH;

This reaction, forming a radical with the unpaired electron on the carbon, is exothermic by 1.0 eV. For this radical several evolution pathways were considered: (1) hydrogen shift leading to an a-carbon radical (HO)3SiO- CH" CH3; (2) hydrogen abstraction leading to a vinyl-silicon species (H0)3Si0 CH CH,; or (3) doublehydrogenabstraction (as H2 molecule)leading to a vinyl-silicon radical (H0)3SiO-CH=CH".Table 6 shows that onlythe first process is exothermic (by 0.5 eV), while the remaining two are highly endothermic. Because of these considerations, the radical with the unpaired electron on the carbon is expected to evolve spontaneously to the a-carbon isomer. In principle, even the a isomer may undergo H" or H, abstraction or the formation of CH3CH0, but all these reactions are therl~odynanlicallyunfavored by about 1.5-2.0 eV, thus being quite unlikely. The considered reactions, however, are first steps of more complex pathways involving the other siloxyl too. Table 7 lists the most plausible ones together with the computed reaction energies. Both and a radicals can recombine with the CH2 and CH(CH3) bridges, respecadjacent siloxyls leading to -CH2 tively. Though these reactions are strongly exothermic, there is, however, a more favorable alternative pathway-hydrogen abstraction by the secondsiloxyl, leading to a silanol unit and a silyl vinyl ether: (HO)3SiO (HO),SiO

CH2 -CH; "OSi(OH)3 (HO),SiO CH2 CH2 HOSi(OH), CH" CH3 "OSi(OH)3 (HO)3Si0 CH2=CH:, HOSi(OH)3

Though these processes havebeen calculated to be marginally more exothermic(by or -CH(CH3)less than 0.1eV) thantheformation of -CH2 -CH2 bridges, thelatteronesinvolveageometricalarrangement of thetwo siloxyls thus requiring acertain strain energy; thevinyl silyl ether may instead form without any of these requirel~entsso that it is expected to be the main product. The vinyl silyl ether is, however, expectedto be unstable in the presenceof an adjacent silanol group and to evolve to acetic aldehyde and a siloxanic bond: (HO)3Si

CH

CHz H0

Si(OH)3

(HO),Si

Si(OH),

CH3CHO

probablythroughanattack by the silanol oxygen onthe electrophilic silicon atom of the vinylsilyl ether. Table 7 shows that this reaction is exothermic by

Reaction Energies for C2H4 Reactionsat One Site of the SLV Center Reaction (HO),SiO" CH2 CH2 (HO)3Si0 CH2 -CH; (HO)3S i 0 CH2 CH; (HO)3SiO-CH~-CH; (HO),SiO CH" CH, (H0)3S i 0 CH" CH3 (HO)$iO CH" CH3

A E (eV) (HO),SiO CH2-CH; (HO)$iO CH" -CH3 (HO), Si0 CH CH, H" (HO),Si0-CH=CH"+H2 (HO), S i 0 CH CH2 H" (HO)3SiO C" CH2 H2 (HO)3Si" CH3CH0

"1.0 -0.5 +1.6 +2. +2.1 .4

Reaction Energies for Ethylene Reactions at the SLV Center Reaction 2(HO)3SiO" 2(HO),SiOe 2(HO),SiO* 2(HO),SiO* 2(HO)3SiO* 2(HO)3SiO' 2(H0)3SiO" 2(HO),SiO"

(eV) C2H4 C2H4 C2H4 C2H4 C2H4 C2H4 C2H4

C2H4

(HO),SiO CH2 CH2 OSi(OH), (HO),SiO CH(CH3) OSi(OOH), (H0)3 Si0 CH CH2 HOSi(OH), (HO)$i 0 Si(OH)3 CH3CH0 (HO),SiO CH( CH2) OSi(OH), H, (HO),SiO CH CH OSi(OH), H2 (H0)3Si- Si(OH), CH3COOH 2(HO),SiOH CH CH

-4.8 -4.8 -4.9 -5.6 -3.6 -3.4 -2.6 -3. l

0.7 eV and leads to the most stable products among the ones which have been hypothesized for the ethylene addition to the SLV defect(by 5.6 eV), in agreement with the experimental evidence, interpreted in terms of formation of acetic aldehyde after annealing. All the other possible reactions reported in Table 7 are much less exothermic and involve single or double hydrogen abstraction with an expectedlyhighenergy barrier;thecorrespondingproducts,therefore,are not expected to form, in agreement with the absence of any experimental evidence for them

1. V. A. Tertykh and L. A. Belyakova, in ~dsorptionon New and ~ o d Inorgunic ~ ~ d ~ o r ~ e n(A. t s Dabrowski and V. A. Tertykh, eds.), Elsevier, Amsterdam, 1995,p. 147 2. G. F. Cerofolini, L. Meda, C. Spaggiari, G. Conti, and G. M, Garbasso. J. Chern. Soc. Faraday Trans. 92:2453 (1996). 3. G. F. Cerofolini, L. Meda, A. Anselmino, and G. Ranghino, in ~ i c r o p o r o ~and s ~ u c r o p o r o ~ s ~ u t e r (R. i u l sF. Lobo, J. S. Beck, S. L. Suib, D. R. Corbin, M. E. Davis, L,. E. Iton, and S. I. Zones, eds.), MRS Syrnp. Proc. Ser. Vol. 431, Mater. Res. Soc., Pittsburgh, PA, 1996, p. 303. 4. G. Ranghino, A. Anselmino, L. Meda, C. Tonini, and G. F. Cerofolini. J. Phys. Chern. B 101:7723 (1997). 5. N. Re, A. Sgamellotti, and G. F. Cerofolini. J. Phys. Chem. B 101:9695 (1997). 6. G. F. Cerofolini, G. Conti, C. M. Garbasso, C. Spaggiari, and L. Meda. Nucl. Instrurn. Meth. B. 129:49 (1997). 7. N. Re, A. Sgarnellotti and G. F. Cerofolini. Intl. J. Mass Spectra. Ion Processes, 179/180:27 (1998). 8. J. Gopalakrishnan. Chern. Mater. 71265 (1995). 9. D. Avnir. Acc. Chem. Res. 28:328 (1995). 10. A. Freundlich, Colloid and Capillary C ~ e ~ i s t rDutton, y, New York, 1923.

11. S. J. Teichner, in (J. Fricke, ed.), Springer-Verlag, Berlin, 1986, p. 22. 12. K. D. Keefer, in Polymer M. Zeigler and F. W. Gordon Fearon, eds.), Adv. Chem. Ser. No. 224, Am. Chern. Soc., Washington, DC, 1990, p. 227. 56:7304 (1997). 13. G. Pacchioni and G. Ieran6. Phys. Rev. 14. R. A. Weeks. J. Non-Cryst. Solids 179:l (1994). 15. E. Borello, A. Zecchina, and C. Morterra. Phys. Chem. 71:2938 (1967). 16. C. Morterra and M. J. D. Low. Phys. Chem. 73:327 (1969). E. Rhodes, and P. D. Orphanos. J. Catal. 40:236 (1975). 17. M. D. Low, Pauthe, J. Phalippou, R. Corriu, D. Leclerque, and Vioux. J. Non-Cryst. 18. Solids223:21(1990). Belot, R. Corriu, D. Leclerque, P. H. Mutin, and A. Vioux. Chern.Mater. 3:127 19. (1991). 20. G. Hochstrasser and J. Antonini. Surf. Sci. 32:644 (1972). A. Radtsig and A. Bystrikov. Kinet. Katal. 19:713 (1978). 21. 22. V. A. Radtsig. Kinet. Katal. 20:448 (1979). A. Radtsig and A. Khalif. Kinet. Katal. 20:705 (1979). 23. 24. A. Radtsig. Kinet. Katal. 20:1203 (1979). 25. V. A. Radtsig. Kinet. Katal. 20:1206 (1979). 26. Y. Matsumura, B. Moffat, and K. Hashimoto. J. Chem. Soc. Faraday Trans. 90: 1177 (1 994). 27. G. F. Cerofolini,in (A. Dabrowski and V. A. Tertykh, eds.), Elsevier, Amsterdam, 1995, p. 435. 28. D. J. C. Yates and P. J. Lucchesi. J. Am. Chem. Soc. 864258 (1964). 29. H. Bequerel. Compt. Rend. 122:420, 501 (1896). 30. Fischer and R. Spohr. Rev. Mod. Phys. 55:907 (1983). 31. F. Ziegler, J. P. Biersack, and U.Littmark, The Solids, Pergamon Press, New York, 1985. 32. H. Schmidt and H. Scholze, in (J. Fricke, ed.), Springer-Verlag, Berlin, 1986, p. 49. 33. Durbach, J. Mellor, N. J. Coville, and T. E. Derry. Nucl. Instrum. Meth. 80/81:294 (1993). 34. U.Yamaguchi, H. Onishi, and Y. Iwasawa. J. Chem. Soc. Faraday Trans. 91:1663 (1995). 35. R. Dannley and G. Jalics. J. Org. Chem. 30:2417 (1965). 36. R. A. Weeks and M. Abraham. Bull. Am. Phys. Soc. 10:374 (1965). 37. J. K. Rudra and W. B. Fowler. Phys. Rev. 35:8223 (1987). 38. D. C. Allan and M. P. Teter. J. Am. Ceram. 73:3247 (1990). 39. M. Boero, A. Pasquarello,J. Sarnthein, andR. Car. Phys.Rev. Lett. 78:887 (1997). 40. G. F.Cerofolini and N. Re, in Si(E. Carfunkel, E. Gusev, and A. Vul', eds.), Kluwer, Dordrecht, 1998, p. 117. 41. R. Eisenberg and D. E. Hendriksen. Adv. Catalysis 28:79 (1979). 42. L. J. Bellamy, Vol. 1, Chapman and Hall, London, 1975.

43. E. Huyser, in Physical D. Henderson, and W, eds.),AcademicPress,New

(H. Eyring, York, 1975, Vol.VIZ,

p. 299.

44. Chatgilialogl~l.Chem.Rev.95:1229(1995). 45. R.A. Jackson. Adv. Free Radical Chern. 3931 (1969). 46. H. Sakurai, in (J. K. Kochi, ed.), Wiley, New York, 1973, p. 741. 47. March. C ~ ~ ~ i McGraw-Hill, s t ~ y , New York, 1991.

11,

K Department of Electrochemistry, Technical University of Lublin, Lublin, Poland, and Department for Nuclear Waste Disposal Technology (INE), Forschungs~entrul~ Karlsruhe, Karlsruhe, Germany

I.

399

11. ExperimentalMethods Macroscopic B. Spectroscopic

nstant 111. at Isotherms Sorption IV. Factors AffectingtheAdsorption pH B. Ionicstrength C. Temperature

omplexation V. Surface Correlation of surface binding constants and hydrolysis constants B. Complexation or electrostatic attraction? Comparison VI.

Silica ofOxides Other with

References

404 404 412 417 422 422 424 426 428 431 43 3 436 437

Theoretical and practical aspects of adsorption of heavy metal cations on silica have been studied since the 1950s. Early studies are collected in the monograph on silica by Iler [l]. Applications of sorbents based onsilica gel in radiochemistry have been reviewed by Laskorin et al. [2]. The surface of silica is negatively charged except in very acidic media. Thus, electrostatic conditions are favorable for cation adsorption at pH values down to about 3. Only a few experimental studies of the surface charging behaviorof silica This work

dedicated to the memory of Professor Jerzy Szczypa.

at pH 3 are available. It seems that at pH 3 silica is positively charged like other oxides below their point of zero charge (PZC). Tadros and Lyklema [3] and Bolt [4] report positive values of the of silica below pH 3; also the electrokinetic potential of silica at very low pH values is positive Recently Koopal [6] has suggested a model inwhich the primary chargeof silica can be only negativeand at low pH values the surface charge density and the surface potential asymptotically tend to zero. In this model silica-type surfaces are distinguished from gibbsite-like surfaces (most oxides belong to the latter group) which have a PZC and their surface potential is pseudo-Nernstian in the vicinity of the PZC (at least at low ionic strengths). Apparently, silica should be more efficient as adsorbent of cations than metal oxides, e.g., hematite and alumina whose surfaces are positively charged in the neutral pH range, but actually silica is a weak adsorbent of heavy metal cations compared with metal oxides, and significant adsorption on silica is observed at pH values only slightly lower than the pH of precipitation of the corresponding metal hydroxide in absence of silica. In natural systems, uptake of heavy metal cationsby iron or manganese oxides is often more significant than adsorption on silica even when the latter is muchmore abundant. The fact that metalcations prefer to adsorb on positively charged surfaces of alumina or hematite rather than on a negatively charged surface of silica is intriguing and not fully understood. Compared with the above-mentioned metal oxides, silica shows a significant solubility, exceeding lmmol/dm3in acidic and neutralaqueoussolutions [l]. Thus, the concentrationof dissolved silicate species in solution can be comparable with or higher than the concentration of active surface sites at a sufficiently low solid-to-liquid ratio. The solubility of silica is enhanced at basic pH values but it decreases in thepresence of some ~ultivalelltmetal cations. For instance, the presence of Al(Il1) can reduce thesolubility of silica by many orders of magnitude The followingeffects (which are ratherinsignificant with insoluble metaloxides, e.g., hematite) can effect adsorption on silica due to its solubility. The silicate species in solution can compete with the surface sites for heavy metal cations. The silicate species can interact with adsorbed metal cations to build phasesof sparingly soluble metal silicates. The competiton of silicate ions in solution with the surface of silica for heavy metal cations is a consequenceof similar chemical nature of ligands in solution and onthesurface andthus similar stability constants ofrespective complexes. Compilations of stability constants of metal ion complexes[7,8] provide very scarce information about soluble silicate complexes. Linear correlation between stability constants of complexes of different metal cations with silicic acid and their first hydrolysis constants was found Jensen and Choppin [lo] studied complexation of europium by aqueous orthosilicic acid by means of TRLFS (time resolved laser induced fluorescence spectroscopy). Addition of orthosilicic acid does not affect fluorescence lifetime and excitation spectra of europium (HI) solution, but two water-soluble species were detected by means of ligand competition technique, namely ~u(OSi(OH)~)2' and~ ~ ( o S i ( 0 H ) ~of) stability ~' constants log 7.26

and log .7. This means that in the presenceof orthosilicic acid in millimolar range, conzplexes of silica are the prevailing species in dilute solutions of Eu(II1) (Fig. Formation of a new phase of metal silicate upon adsorption on silica was recognized long ago but analogous processes forheavy metal cationswere not considered until Charlet and Manceau [l21 recently suggested formation of clay minerals upon sorption of and Ni on silica. This phenomenon is accompanied by a slow change from a tetrahedral to an octahedral co-ordination. Sorption of ions on silica is similar to ion exchange and in many early studies of cation sorption on silica the term ion exchange has been used. The number of released protons per one adsorbed multivalent cation is often referred to as the proton stoichiometry factor, r. In sorption of ions on oxides r is variable and it is usually lower than the valence of the cation. The proton stoichiometry Factor can be directly determined from pH measurements for a sufficiently high concentration of heavy metal cations. Even when the proton Stoichiometry factor equals to the valence of the cation, the electrostatic contribution to the adsorption energy is still significant when the heavy metalcationsarenotadsorbed at thesame distance from the surface as protons, thus the apparent equilibrium constant of the reaction:

l

Calculated speciation of 1 nM Eu(lI1) in the presence of 0.2'7 mM*Si(OH)~ in equilibrium with atmospheric CO2 as a function of pcH. (From Ref. 10, R. Oldenbourg Verlag, Munich.)

Men' (in solution) nH' (bound with silica) Me"' (bound with silica) nH' (in solution) depends on the pH. In order to quantitatively account for the effect of the electrostatic field on the adsorption of heavy metal cations twoelectrostatic potentials must be determined: the surface potential to calculate the electrostatic work connected with the desorption of protons and the potential at some distance from the surface, where the heavy metal cations are adsorbed.If the cations adsorb in the surface plane, these two potentials are equal. Let us assume that adsorption is due to the following reaction: =SiOH

H20

Me2'

=SiOMeOH

2H'

where (except in the last term) denotes concentrationsof surface species and denotes the activities of ions in the bulk solution, and is the potential in the surfaceplane (subscript 0) and in theplanewhereheavymetalcations are adsorbed. For (heavy metal cations adsorb in the surface plane) the exponential factor in Eq. (3) disappears. At low concentrations ofheavy metal cations, =SiOH] is practically constant, and Eq. (3) reduces to

eMe

SiO~eOH](H'}2 (Me2')-' where K' is constant for a given solid-to-liquid ratio. Many authors postulate the formationof bidentate surface complexes: 2=SiOH

Me2'

(=Si0)2Me

2H'

The equilibrium constant of reaction (5) is often defined as E;(=Si0)2Me SiO)2Me](H+}2[ SiOH]-2(Me2'}" exp [-F(~@o2@~e)/(RIP31

The second powerof SiOH] in (6) represents the probabilityof finding two surface sites together. This kind of reasoning is valid for independent moleculesin solution but not for fixed surface species. Nevertheless, Eq. (6) has been used to define the equilibrium constant of reaction (5) even in very recent publications. Morel 131 suggests that the probability of finding surface sites together should be represented by SiOH], and the equilibrium constant of reaction (5) should be redefined as

4.

where the factor 2 from SiOH] was incorporated into the equilibrium constant. At @Me and for low concentrations of heavy metal cations both the expressions (6) and (7) reduce to Eq. (4).

Theexperimentallyobserved proton stoichiometryfactorfordivalent heavy metal cations is typically lower than 2. Therefore reaction (2) or (5) (both lead to identical uptake curves for low initial concentrations) do not properly describe the mechanism of adsorption. The proton stoichiometryfactor between S and 2 canbe interpreted as a result of two surface reactions: reaction (2) or (5) and

electric-field-depenWith combinationof reactions (9) and (7) one cannot avoid the dent factor in the relationship between the pH and magnitude of adsorption, e.g., when the exponential factor in Eq. (3) [or Eq. (7)] reduces to unity then the exponential factor in Eq. (9) certainly does not. The surface potential cannotbe directly measured. In absence of specific adsorption or when the concentration of specifically adsorbed species is sufficielltly low, the surface potential of oxides can be roughly estimated from the Nernst equation -(pH

p H o ) S ~ l n ( l O ) / ~-(pH

pHo)

59 m\r at 25°C

0)

for silica is about 3. This equation overestimates the absolute valueof the surface potential, especially for high ionic strengths. Electro~ineticphenomena make it possibleto determine the potential of silica. In the presence of specific adsorption there is no direct relationship between the potential and q0.When ions are strongly adsorbed behind the slipping plane even the signs of these two potentials can be opposite. On the other hand, when the concentration of heavy metal cations and thus their effect on the surface potentialis small, the diffuse-layer model leads to the following relationship between the surface potential and the potential at certain distance d from the surface: -In e

S +goe-Kd S

and

These equations can be used to estimate the surface potential from the potential when some reasonable value of slipping plane distance is assumed (unfortunately this distance cannot be directly measured). A problem in using Eqs.(1 S)-( S 3) is that thebulk dielectric constantandthebulkDebyeparameterarenot necessarily relevant for calculationsof the electric potential in theclosest vicinity of a charged surface.

The diffuse-layer model has a limited ability to properly reflect the dependence of thesurfacechargedensity andpotential ofsilica on the pH and ionic strength. Therefore, more complicated models havebeen introduced. The distribution of electric potential near a charged silica surface can be estimated by fitting parameters of a model of electric double layer using the experimentaldata [usually curves at different ionic strengths] and then to calculate the potential of interest using these model parameters. There is no generally accepted model of electric double layer: onehastomakea difficult choice between modelswith fewer parameters, which donot exactly reflect thechargingbehavior ofsilica (e.g., theabove-mentioned diffuse-layer model) and moresophisticatedmodels for which quite different combinations of model parameters give practically the same (pH) curves. The model parameters must be somewhat different for different 1-1 electrolytes, because the surface charge density and the adsorption of alkali-metal cations on silica increases from Li to This sequence is probably reversed at extremelyhigh ionic strengths Interestingly, the electrokinetic potential of silica at a given pH andionic strength rather insensitive to the nature of alkali-metal cation when the ionic strength is not too high. This property is in line with the triple-layer model (TLM),in which thediffuse charge is due tosurface ioni~ation (not directly influenced by the nature of the counterions) while the surface charge is chiefly governed by formation of surface complexes involving the counterions. The electrokinetic behavior of silica, even in absence of specific adsorption, is not fully understood. For metal oxides the isoelectric point (IEP) coincides with the minimum of colloid stability. In contrast silica shows a maximumof stability in the vicinity of

The most obvious experiment to study sorption of heavy metal ions on silica is to put a certain amount ofsilica into solution and compare the initial and final concentration of heavy metal ions in solution. The choice of analytical methods (radiotracer, atomic absorption spectrometry, and VIS absorption spectrophotometry)dependsonthenatureandconcentration ofheavy metalions of interest. Apparently the experiment is simple, but certain safety measures must be respected in order to obtainreliable results. Dzombak and Morel 161 have listed their criteria to select reliable and representative adsorption data from literature, and reject those which are of little value: Use of standard oxide preparation method Experimental temperature between 20 and 30°C Absence of competing ions Absence of in systems that form carbonate precipitates Use of appropriate equilibration time Use of reasonable method of solid-liquid separation in sorption experiments.

Dzombak and Morel have reviewed sorption of heavy metal cations on hydrous ferric oxide whose surface properties are different from silica. For example, the hydrous ferric oxidestudied by these authors showeda very highsurfacearea (600 m2/g), and the number of surface sites used in calculations was one-fifth of the total number of iron atoms. This is only possible for certain methods of preparation, otherwise only a small fraction of iron atoms is exposed. Different types of natural and synthetic silica show somewhat different adsorption properties toward heavy metalcations. For example thePZC or IEP (and thus electric potential at given pH) differs from one sample of silica to another. The solubility depends on the typeof silica and on theparticle size. Therefore, comparison of the parameters characterizing the adsorption of cations on various types of silica requires caution. Nevertheless the quantities characterizing sorption of given heavy metal cation shouldbe similar for different types of silica when their properties (surface area, pHo, number of surface sites) are properly accounted for. The temperature control is important in view of temperature effects on adsorption of cations, Theseeffects are significant for silica, as well, and will be discussed later in this chapter. Adsorption of heavy metal cations onsilica is more sensitive to concentrationof 1-1 electrolytes than adsorption on hydrousferric oxide, especially when the anions of supporting electrolyte form stable water-soluble complexes with theheavy metal cation of interest. This effectis partially due to competition between anions in solution and surface hydroxyl groups metal cations andpartially it an indirect influencevia thedistribution of electric potentials in the interfacial region. Systematic studies of ionic strength effects on sorption of cations of silica are rare and they are discussed in a separate subsection. The presence of 1-1 electrolytes in sorption experiments canbe hardly avoided. Manystudies have been carried out at high but constantionic strength (the contributions of acid or base addedto establish the pH to the total ionic strength are neglig~ble) rather than low but variable one. The role of alkaline-earth metal cationsin the adsorptionof heavy metal cations on silicais somewhat different than in case of adsorption on metal oxides. In principle there is no competition between heavy metal cations and alkaline-earth metal cations for surface sites of metal oxides (hematite, alumina),and adsorption isotherms obtained in presence and absence of alkaline-earth metal salts are identical. Hayes and Leckie [l71 suggest that heavy metal cations form inner sphere complexes with the surface while alkaline-earth metal cations form outer sphere complexes. With silica, alkaline-earth metal cations can form a layer of sparingly soluble silicate, whose adsorption properties aredifferent from those of silica. The presence of alkaline-earth metal cations depresses the adsorption of nickel on silica. Fig. 2 shows that the pHSO (the pH value at which the distribution of the heavy metal between the solid and liquid shifted by more than one pH unit in the presence of 0.01 mol/dm3 of Ba, but theeffect is insignificant for Ba concentrations 0.001 mol/dm3. The initial concentration of Ni was mol/dm3 and the estimated concentrationof surface hydroxyl groupswas on the orderof 0.01 m01/dm3. For this initial concentration of nickel presence of sodium (upto 0.3 1n01/dm3)does not affect the adsorption of nickel.

2o

8

P Effect of Ba on sorption of Ni on silica.

Underthesameexperimentalconditions,the effect of alkaline-earth metal cations on sorption of Gd was rather insigni~cant(Fig. 3). The presenceof adsorption competition between heavy metal and alkaline-earth metal cations in the former and absenceof adsorption competition in the latter case be caneasilyexplained interms of adsorptionbehavior of al~aline-earthcations. Thesecations similar adsorption edges to those observed for heavy metal cations. For Ni and a (Fig. 2) the pH ranges of the adsorption edges coincide and this is reflected in adsorption competition, but the pHsO forGd falls several pH units below the pHSO for alkaline-earth metal cations. These two examples show that sorption of some heavy metal cations on silica is significantly depressed when the concentration of a l ~ ~ l i n e ~ e ametal r t h cationsis comparable with the concentration of surface hydroxyl groups but sorption of some other heavy metal cations remains unaffected. Many heavy metal cations form insoluble carbonates and stable water-soluble carbonate complexes. Therefore the presenceof carbon dioxide should be avoided or the amount of carbonates in the system should be controlled and accounted for. The results obtained in systems free of are of limited value instudies aimed for technological applicationsor studies of migration in the environment.For example, Pu and Np are stabilized in solution in form of strong carbonate complexes, but silica successfullycompetes forPu(1V) (distribution ratio on the orderof 1000) [18], The timeof equilibration is a somewhat controversial question. A long time may be necessary to obtain an equilibriumvalue. Therefore, in many studies the equilibration timeis many days and weeks. On the other hand, fast titrations reflect

noMg

Mg

Effect of

on sorption of Gd on silica.

chiefly the adsorption process, while longer equilibration can lead to formation of new phases like sparingly soluble hydroxide or silicate. A study of sorption kinetics is certainly more valuable than a single measurement after aselected time; one can also perform adsorption and desorptionexperiments, e.g., by back-and-forth pH titration. Coincidence of results of adsorption and desorption experiments wouldindicate that equilibrium values were obtained. Figure 4 shows the result of an adsorption experiment immediately followedby a desorption experiment. The uptakeof nickel observed at a given pH value in backand-forth titration is very consistent. However, agingof silica with nickel adsorbed on it athigh pH values causes someirreversible process, probably crystalli~ation of nickel hydroxide or formation of nickel silicate. Figure 5 shows similar curves asin Fig. 4, but the system was kept at pH 9.8 for l h before a titration with acid was performed. The sorption is not reversible here, and after desorption of 60% of preadsorbed nickel further acidification of the system does not cause releaseof nickel. Separation of the solid and liquid phase is apossiblesource oferrors. Correctness of experimental procedures must be examined for particular system of interest. Fine colloids can passfilters and remain in solution after centrifugation. On the other hand,filters can adsorb small amounts of metal ions. Figure 6 shows results of three consecutive filtrations of Pu solution in diluted HCl. At highinitial concentrations the radioactivity of the solution is unaffected because the sorption capacity of the filter is lower than the amountof Pu in solution. However, withvery

Back-and-forth titration of silica in presence of nickel.

M M

L

Back-and- fort^ titration of silica in presence of nickel. Back titration after h delay.

Nalgene filter

A Whatman filter open symbols 261-10 Pu filled symbols 2e-7 M Pu

A 2

of f i l t r ~ t i o ~ ~ Effect of consecutive filtrations on the concentration of Pu in solution.

low initial concentration each consecutivefiltration leads to decrease of the concentration of metal ions. Alsoadsorptiononglassware used in theseparation procedure and on pH electrodes can significantly influence the results when the concentration of heavy metal cations is very low. In high-speed centrifugation the solid particles are pressed one against another and this may disturb the adsorption equilibria etablished in suspension. Usually silica contains chemisorbed water. With low solid-to-liquid ratio the amount of water bound with silica surface is negligible in comparison with the amount of water in solution. However, with very high solid-to-liquid ratios the amount of water introduced into the system with silica can be significant and it must be taken into the calculation. This difficulty is not encountered in typical sorption experiments with metal oxides. Most often the magnitude of adsorption is determined from depletion of solution, but direct measurements of concentrations of adsorbed ions arealso possible [19]. Such a measurement requires procedure to remove the liquid without distortion of the equilibrium. Direct measurements of the amount of the cations of interest bound with the solid do not necessarily lead to the same results as adsorption calculated from depletion of the solution [20]. The limited range of radiation makes it possible to determine the magnitude of adsorption of radionuclides without separationof the solid and liquid phase when the solid forms thin layer on the bottom of the vessel when the radiation detector is placed under the bottom [21]. The results obtained by passing solutioncontainingheavymetalcations through a column filled with silica gel are more difficult to interpret than batch experiments, especially when gradient of and ionic strength builds in the

column. On the other hand, results obtained under dynamic conditions are more relevant to specific applications, e.g., migration of radionuclides and other pollutants. A dynamic method has been used to study sorption of U(VI) and Zr(1V) on silica 1221. It has been already discussed in the Introduction that a properly selected adsorption model must reflect the experimentally observed proton stoichiometry factor. The number of protons released by adsorption of one heavy metal cation can be directly determined by potentiometrictitration. Significant results areobtained when the concentration of heavy metal ions is sufficiently high. Some results for silica are summarized in Table 1. Vydra and Galba [23] determined the adsorption of heavy metal cationsin the presence of concentrated acetate buffer. Heavy metal cations form strongacetate complexes (Ref, 7, Vol. 3) and this certainly affects the adsorption equilibria. Spark et al. [24] calculated their proton stoichiometry coefficients fromKurbatovplots (this is an indirect method).Thevaluesabove 2 obtained for Cu and at high ionic strengths are difficult to understand. The heats of adsorption of heavy metal ions canbe determined by microcalorimetry. The interpretation of the measured heats is difficult because the adsorption is accompanied by surface charge formation. In view of limited accuracy of calorimeters, this method is applicable for relatively high concentrations of heavy metal ions. A few experimental studies have been published for metal oxides but not for silica. The presenceof less than 1 mmol/dm3 of heavy metal cations can induce reversal of sign of the potential of silica from negative to positive. In contrast, alkalineearth metal cations do not cause such a sign reversal even at much higher concentrations. For example, 0.1 mol/dm3 of Ca12did not reverse the sign of potentia1 of silica at pH 8 1261. In this respect silica is different from most metal oxides, e.g., reversal of sign of potential of rutile in a millimolar solutionof calcium nitrate has been reported Thus, the effect of different multivalent cations on the electrokinetic behavior of metal oxides is qualitatively similar, but with silica we must clearly distinguish between alkaline-earth metalcations(no sign reversal) and heavy metal cations (sign reversal).

Proton StoichiometryCoefficients Ref. 23

24,

Ref. Ref. 25,

Cation CO Ni

Mn Zn cu Cd

pH

r

0.001 M KN03 0.1 M KN03

6.3-7.3 1.18-1.34 1.16-1.4I .6 2.4 6.3-6.8 1.36 6.4-7.7 1.23 6.01.23 1.191.39 l .33 I .8 4.1 1.6 1.23

M NaCl

3.5

2.9 l .84

r

James and Healy [28] studied the effect of Co(II), La(III), and Th(IV) on the electrokinetic properties of silica. Similar results were obtained by electrophoresis and streaming potential. At pH 4, when the metal cationsare not adsorbed, they do not exert any significant effect on the electrokinetic potential of silica. At pH 4 the presence of multivalent cations shifts the potential of silica to less negative values, and when the concentration of heavy metal cations is sufficiently high, the sign is reversed to positive. The minimum pH value at which the sign is reversed depends on the nature of the cation, the concentration of multivalent cations and solid-to-liquid ratio: this effectis more significant for cations carrying a higher charge. At very high pH values the sign of the electrokinetic potential is reversed to negative again at the pHcorresponding to the IEP of the oxide or hydroxide of the heavy metal of interest. The same type of behavior was found for Pb(I1) (291 (Figs. 7 , 8). Electrokinetic results can be used to test models of cation adsorption.So far this possibility has been exploited for iron and chromium hydrous oxides [30] but not for silica. Comparison of experimental and theoretically calculated electrokinetic potentials is not straightforward and it needs certain assumptions. A study of clustering of 7 nm Ludox particles at pH9.5 in the presenceof yttria provides information about the mechanism of yttrium adsorption onsilica [3l]. The amount of adsorbed yttrium gradually increases within two weeks of aging. The cluster size increases up to 700 nm within the first 10 days and then it suddenly decreases and stabilizes at 130 nm. On the other hand, silica does not show any

e-

z

potentials of silicain the presence and absence of Pb. (From Ref. 29, Elsevier Science.)

potentials of silica in the presence and absence of Co. (From Ref. 28, Academic Press.)

significant clustering in absence of yttria. Thisresult was interpreted as follows. For short aging times (and surface concentrations), the yttrium ions are evenly distributed over the silica surface and this gives rise to clustering (pH 9.5 is close to the IEP of yttria).Whensome critical surfaceconcentration of yttrium is exceeded, yttrium hydroxide particles are formed and the clusters break down.

Therearefundamentalquestions which can be hardlyanswered by studying adsorption only by means of macroscopic methods. For example, model calculations for adsorption, on the one hand, and surface precipitation, on the other, lead to similar dependence of percentage of uptake on the pH. any spectroscopic methods usedin studies of adsorption require complete dehydration of the sample, ultrahigh vacuum, low temperatures, and other operations which are likely to affect the structure of the surface layer formed in contact with aqueous solution. Fortunately, there are a few methods to study a surface in contact with aqueous solution. Such studies are termed in situ. Anderson [32] studied NIR and VIS spectra of silica impregnated with solutions of Cu(I1) and Cr(II1) salts. The concentrations of heavy metal cations or solid-to-liquid ratioarenotreported.Freshlyimpregnatedsamplesdisplaythe samespectraastheimpregnating solutions. Thismeans that themetalcations are present as aquocomplexes. As the impregnated samples are dried, the co-ordi-

nation environment of the metal cations changes, and this is reflectedin thespectra. Anderson interprets the observed changes as replacementof water in the co-ordination sphere by anions on the solutionside and by hydroxyl groups on the surface side. These changesare reversible, i.e., original spectra are restored upon readsorption of water, but calcination leads to irreversible changes. Roose et al. [33,34] used the magneticfield dependence of the protonspin-lattice NMR relaxation time to determine the number of paramagnetic ions: free in solution, on the one hand, or adsorbed, on the other. This method, called nuclear magnetic relaxation dispersion (NMRD) is applicable only for Paramagnetic ions with a long electron spin relaxation time such asMn2' but for Ni2' whose relaxation is faster, the free and adsorbed ions are not distinguishable by means of the proton NMRD.An important feature of this method that aninsoluble hydroxide either precipitated in solution or on thesurface give no proton relaxation enhancement (Fig. 9). In other words, theNMRD makes it possibleto distinguish between adsorption and surfaceprecipitation. ~nfortunatelythis method is only applicable to a limited number of ions. A maximum coverageof 0.43 Mn2' ions per nm2 of Stober silica (monodispersed spherical particles obtained by hydrolysis of tetraethoxysilane) was found at 4°C. This figureislowerby an order of magnitude than the number of surface sites used in modeling of surface charging of silica. Sorption of paramagnetic cations [Cu(II), Mn(II), Fe(III)] can be studied by means of electron spin resonance (ESR), also called electron paramagnetic resonance (EPR). Spectraof adsorbed ions arequalitatively compared with the spectra of model compounds (e.g., metal hydroxide) to determine the oxidation state, coordination numbers and degreeof hydrolysis of the adsorbed species. In their ESR

pH pH pH pH

Water proton 1/TIPNMRD curves for increasing Mn(I1) concentrations. (From Ref. 33, 0Academic Press.)

study, Hronsky et al. [35] found that Mn(II) adsorbs in form of monodentate surfacecomplex at low initial concentrationbut clusters areformed at higher concentrations. Martini et al. C361 interpret their ESR spectra as an evidence for monomeric species of Mn(T1) and Cu(1I) adsorbed on silica. ESR spectroscopy canbe used to distinguish between complexes in solution and immobilized (adsorbed) species. In solution the anisotropic contributions are averaged and in immobilized state they fully contribute to the spectra. The Cu” ion yieldsresolved ESR spectra at room temperature, and in such case the ESR method can be used to study adsorption from solution. The proposed structure of surface ternary complexeswas based on Cu2’ spectra obtained in the presenceof chelating agents [37]. Solid-state spectroscopy has been used to study adsorption of soluble aluminum ionic species on silica[38]. Analysis of 27Al and 29Si NMR spectra shows that aluminum is incorporated into the silica framework. This results in an increase of surfacechargedensity ofsilica and decrease of its solubility. Octahedral aluminum species are associated with the silica surface as inner-sphere complexes. Absorption of x-rays is primarily a functionof the atomicweight of an element. An analysis of the spectrum makes it possible to determine the nearest co-ordination environment. The advantageof this method is its specificity, namely the shape of the absorption edge in given range of energy is due to absorptionby one element. Measurements in situ are possible, but the ions of interest in the bulk solution and adsorbed ions are not distinguished and their spectra overlap. Deconvolution of such spectra wouldbe difficult, especially when the spectrumof ions in solution and ions on the surfaceare similar. At least the first co-ordination sphereconsists of the same number of oxygen atoms for both adsorbed species and a cation in solution. In practice the x-ray absorption studies are carried out in systems where most of ions of interest are adsorbed, otherwise theliquid must be removed by centrifugation or filtration to leave a paste containing very little of the mother liquid. The other limitation of this method is that the total concentrationof ions in the system must be sufficiently high. It was discussed above that both reactions 2 and 5 lead to the same course of adsorption edges. In extended x-ray absorption fine structure (EXAFS) spectroscopy, the difference between one and two silicon atoms in the second nearest co-ordination environmentof central metal ion can be distinguished [monodentate and bidentate surface complexes, reactions(2) and (5)]. An EXAFS study of Cr(II1) sorption of amorphous silica colloid 1391 indicated formation of monodentate Complexes with Si-Cr distance of 0.339 nm at surface coverage 20%. The spectra of adsorbed species are shown in Fig. 10 with a spectrum of chromium hydroxide for comparison. In contrast, the EXAFS spectra of Co(1I) sorbed on quartz (401 are similar to that of crystalline hydroxide and this suggests sorption of large polynuclear complexes of or formation of cobalt hydroxide. Also Cu(1I) seems to form hydroxide clusters on fumed silica at coverages as low as 0.8% as detected by EXAFS spectra [41] (Fig. 11). EXAFS was used to study adsorption of uranyl ion on silica [42]. A dispersion containing 10 g Si02 per dm3 (nearly monodispersed spherical particles, 7 nm in

6

Experimental EXAFS spectraof Cr(1II) adsorbed on silica. (From Ref. American Chemical Society.)

diameter)and 5 mol/dm3 U(VI) was concentrated by ultrafiltration and the EXAFS spectra of wet filters were taken. At pH 3 and 5 the spectra of adsorbed species resemble uranyl hydrous oxide and this result suggests formation of polynuclear complexes. Clause et al. used a similar silica in their EXAFS study of adsorption of Ni(I1) but they dried their samples before the EXAFS measurements. They came to the conclusion that adsorption of nickel in absence of ammonia directly leads to nickel silicate formation. Osthols et al. collected EXAFS spectra of Th adsorbedon Aerosil at surface coverages of 3.8 to 230%. Theyproposedthemostprobablestructure of the adsorbed complex, namely Th4' ion co-ordinated by seven molecules of water in solution and two surface Si04 tetrahedra sharing a corner.No trace of polynuclear

U

weighted Fourier back-transformed EXAFS spectra of Cu(I1) adsorbed on silica at two different surface coverages. (From Ref. 41, Academic Press.) complexes was detected, but the solubility product of T h o 2 was exceeded at high coverages. Probably the structureof the surface precipitate is so disordered (as it is in case of amorphous Tho2) that noTh-Th interactions can be detected. Time-resolved laser-induced fluorescence spectroscopy (TRLFS)has been successfully used to study speciation of lanthanides and actinides in solution in range of low concentrations. The detection limit for Cm3' by this method is 5 mol/dm3 The Cm3' ion yields a narrow peak and the peaks of hydrolyzed species are shifted toward longer wavelengths. The shift is more pronounced for Cm(0H)i species than for CmOH2+ species. isp placement of watermolecules from the inner co-ordination sphere leads to increase in the fluorescence lifetime. Similar shifts are observed for other water-soluble complexes. In spectra of real solutions there is some overlap of peaks of particular species. Peak deconvolution gives unequivocalspeciationsinsystemswithnomore thanfour species. In dispersions ofsilica containing Cm(1II) the Cm3' peakgraduallydisappears when the pH increases and twopeaks of hypotheticalsurface species can be distinguished [46].

ilic

The results of adsorptionexperiments are oftenplotted in form of adsorption isotherms at constant pH and ionic strength. In view of a strong effect of p the magnitude of adsorption over the range of adsorption edge, the course of an adsorption isotherm canbe completely different when the pHis changed by a small fraction of 1 pH unit. Equations of adsorption isotherms derived for adsorption of gases are also used in studies of adsorption of heavy metal ionson oxides. This approachis justified by analogies between dilute solution and gaseous state. For example, Malati et al. used Langmuir isotherms to interpret their results. Most of these adsorption isotherms have arectilinear segment in the rangeof lowconcentrations, corresponding to Henry adsorption isotherms. In the Henry range the percentages of uptake (pH) curves are independent of initial concentration of heavy metal cations and the adsorptionisotherms inlog-log co-ordinates at constant pH are linear witha slope of 1. For higher concentrations of heavy metal ions (but still much lower than a monolayer), the slopeof log-log adsorption isothermsis often lower than Dzombak and Morel [l61 interpret this nonlinearity of adsorption isotherms in terms of surface heterogeneity. They proposed a bimodal distribution: the surface of hydrous ferric oxide has strongand weak sites. The concentration of the former is lower by two orders of magnitude than the concentration of weak sites but the binding constant of heavy metal ions to the strong sites exceeds that of weak sites by three ormoreorders of magnitude.Figure l:! showsadsorptionisotherms calculated for different distributions of surface sites. The distributionsof surface sites were selectedin such a mannerthat the uptake for co~centrationsof heavymetalionsinsolutionequal to mo1/dm3, onthe one hand, and mol/dm3, on the other, are equal for all three distributions. The solid line corresponds to a linear log-log adsorption isotherm of a slope of 0.8 and it serves to guide the eye. For initial concentrations of heavy metal cations below the concentration of the strongest sites, the contribution of the weaker sites to the adsorption is negligible, and the slopeof log-log adsorption isotherm equal to l . Once the strongest sites are saturated, the slope decreases, and even a plateau can be modeled when the difference in binding constantsbetween strong and weak sites is sufficiently high. At high concentrations the slope. oflog-log adsorption isotherms increases again. Thus, model adsorption isotherms for a bimodal distributionhavea characteristic hump.Withmorethantwokinds of surface sites adsorption isotherms with multiple humps are obtained and with a continuous distribution of surface sites, linear log-log adsorption isotherms of slope below 1 can be modeled. Figure l:! shows that with only four kinds of surface sites an almost rectilinear adsorption isotherm can be obtained. Irrespective of the kind of distribution (bimodal, continuous),in order to obtain nonlinear adsorption isotherms at coverages below 1% of monolayer, the difference in binding constants of weakest and strongest sites must be by three or more orders of magnitude, and this corresponds to alarge difference in standard adsorptionGibbs energies. Adsorptionfromsolutioncan be consideredasexchange between solventmolecules andadsorbate.Theadsorptionenergiesat different adsorbates on heterogeneous surfaces are correlated. Therefore the surface hetero-

I

1

0-5

Calculated log-log adsorption isotherms for three discrete distributionsof surface sites (Table 2). geneity following from the surface geometry has aless significant effect on adsorption from solution than on adsorption of gases. On the other hand, the apparent surface heterogeneity may be due to formation ofvery strongternarysurfacecomplexesinvolvingmetalions of interest and ligands presentin the system as impurities. It has been experimentally demonstrated

Distributions of Surface Sites Used to Calculate Adsorption Isotherms in Fig 12 Weakest

Strongest

Sites Concentration

pkl

Two 1.44 Three 1.44 Four 1.44

10 9.9 9.8

Concentration

1.44

pkl

Concentration

7.5

1.44 1.44

1.44

pkl

7.2 1.444.7

Concentration

pkl

1.44

S 4.5 2.5

lo-*

Concentrations are in mol/dm3, pH 8, 10 g silica (144 m2/g) per dm3, and the concentration of inert electrolyte is mol/dm". pklis the negative logarithm of formation constant of SiOMe' inner sphere surface complex where Me is a divalent heavy metal. The electrostatic contribution was calculated using a triple-layer model with stability constants of =SiO- and SiONa surface groups equalto -5.71 and -7.24, respectively and inner- and outer-layer capacitance 2.2 and 0.2 F/ m2, respectively. For more details TLM see Chapter 11.

that addition of certain ligands forming strong complexes with metal cations in solution considerably enhances the adsorption. Enhancement of adsorption of heavy metal cations on silica in the presence of ligands forming strong complexes in solution has been demonstrated by Park et al. [48j. These authors studied the effects 1,10-phenanthroline and 2,2'-bipyridine on theadsorption of Fe(II), Cu(II), Co(II), andNi(II)frommillimolar solutions of nitrates (Fe(I1) sulfate) in 0.1 M NaNOs on chromatographic silica fromMerck at ligand-to-metalratios 0: 1, I: 2: and 3:1. Taking a typical surface area of chromatographic silicas and the solid-to-liquid ratio reported in the experimental part of Ref. 48, the concentration of heavy metal ions in solution is on the order of 1 ion per 10 nm2 of silica surface. Considering a diameter of ions in solution and density of surface hydroxyl groups on silica, this corresponds to a moderate surface coverage. The most characteristic effect of complexing agents was observed at low pH values: the percentage of uptake was in excess of 25% at pH as low as 2 for ligand-to-metal ratio 3:1 for all eight metal-ligand combinations, while in theabsence of theseligandstherewas no detectable adsorptionatpH below 4. Moreover,thepercentage of uptake isrelatively insensitive to the pH at low pH values in the presence of ligands, but when the uptakeapproaches loo%, theuptakecurveassumes its usualshape, i.e., the adsorptionsharply increases withthe pH.Manyotherexamples of enhanced adsorption ofheavy metal cations on silicaby preadsorbed complexing agents can be found in the literature [49,50]. Taking into account the above experimental results the following explanation for enhanced sorptionof heavy metal cations atvery low initial concentrations can be offered. Let us assume that in addition to the components introduced on purpose the experimental system contains unknown anions An (not necessarily one species)whose concentration is on the order of mol/dm3. Thesource of these anions can be different, e.g., silica, water, and salt added as supportingelectrolyte, and the impurities may not necessarily come from one source. At very low total concentrations of the heavy metal cationsof interest, there in an excess of An and the metal cationsare chiefly adsorbed in form of ternary complexes. Thestability of these complexes is very high and the adsorption occurs already at low pH values. Whenthetotalconcentration ofheavy metal cations considerably exceeds the concentration of An, the contributionof the ternary complexes to the total adsorption becomes negligibleand this results in a shift of the uptake curves toward higher pH values. Manyheavymetalcationsformstrongwater-soluble sulfate and phosphate complexes and the hypothesis that these inorganic anions are responsible for enhancedsorption of heavymetalcations at low initial concentrations can be examined by a simple experiment. If such a hypothesis was true, addition of a sufficient amount of sulfate should enhance the sorption also at high initial concentrations of heavymetalcations. For instance, thecurvesshown in Fig. 13 indicate that sulfates do not enhance the adsorption gadolinium on silica. A model involving two surface species, namely SiOHGd3' and SiOGdAn, where An is bivalent anion (whosenature remains unknown), and analogous complexes for yttrium, hasbeen successful in modeling of the effect of initial concentration of

LQ

40

8

A0

A

0

20

4

5

Effect of sulfates on sorption of Gd on silica.

trivalent cations on the uptake (pH) curves [51]. The total concentration of An assumed in themodelcalculationswas4 mol/dm3. Concentrations of impurities in this range are quite probable anddifficult to detect. This model reflects an experimentally observed trend, namely, for initial concentrations of heavy metal cations slightly exceeding the total concentrationof An, the percentageof uptake at very low pH values is significant and almost insensitive to the pH. A similar model can beused to interpret the nonlinearity of the adsorption isotherms of nickel on silica in the range of low nickel concentrations. Nickel was assumed to form two kinds of surface species, namely SiOHNiAn and SiOHNiClO~,whereAn is amonovalentanion (impurity). Figure 14 shows experimental and calculateduptakecurvesforvarious initial concentrations of nickel. Liu et al. [52] studied the adsorption of zinc and cadmium on silica powder at initial concentrations 10-3-10-5 mol/dm3 at pH 6-6.5. They do not specify the exact pH value, probably the pH was variable and the ionic strength was mol/dm3. Their adsorption isotherm in log-log co-ordinates is linear with a slope of about 0.6 up to acritical surface coverage equalto 0.656 Cd2+ and0.81 Zn2+ ion per 1 nm2. Above this critical coverage, further increase of concentration in solution did not affect the surface coverage. The adsorption isotherms were interpreted in terms of so-called Darken plots, namely

initial Ni concentration

40 lines: calculated [An]=le-7 M

5.5

7.5

9.0

9.5

Adsorption of Ni on silica at different initial concentrations.

with 2.75 for Cd and -2.77 for Zn, where.XR is the surface coverageand yR is activity coefficient of adsorbed cationsreferred to Raoultian standard state,defined maximum surface coverage. This activity coefficient was calculated

where CbYb is the activity in solution and the superscript refers to the activity in solution at which XR is reached. The activity coefficients in solution are calculated from Debye-Huckel theory. The same authors (521 calculated the surface concentration of adsorbed species as r'/d where l? is adsorption density andd is the diameterof adsorbed species. This diameter assumed to be 0.5 nm for Cd and0.6 nm for Zn, i.e., on the order of the sizeof hydrated ion. At 1 the surface concentration exceeds the bulk concentration by factor of 3. TheHenrianstandardstateforsurface species was defined hypothetical adsorbed species behaving infinitely dilute at 1. The activity coefficient related to the Henrian standard state of adsorbed species yH should tend to unity for low concentration limit. Extrapolation of the experimental data from Ref. 52 does not lead to Henry's adsorption isotherm. In viewof the above-mentioned slope of about 0.6 of log-log adsorption isotherms, it can be easily shown that both yH and ;yR decrease by factor of per one order of magnitude of The ratio yH/yR calculated from (14) (at 0 In yR is rather insensitive to the nature of the adsorbing cation and toits concentration in solution. The extrapolation to XR 0 used to obtain this result in Ref. 52 is disputable.

For most cations, so-called adsorption edges are observed (e.g., Fig. 3). The sorption is practically equal tozero below a certain pH value. Then a sigmoidal increase of sorption is observed over a range of about 2 pH units and the uptake reaches practically 100%. The position of such a curve can be characterized by the pH50 value, i.e., the pH atwhich 50% of the heavy metal cations are adsorbed. The pHs0 increases when the solid-to-liquid ratio decreases. A low pH50 valuefor given cations means that this cation is strongly adsorbed, provided that the comparison between different cations is made for the same type of silica, and at a constant so1id-to”liquid ratio. In order to obtain a quantity characterizing the affinity of metal cations tosilica independent of the solid-to-liquid ratio, a distributioncoefficient is sometimes used:

R (amount sorbed/mass of sorbent)/(amount in solution/volume of solution)

(16)

The above definition is more appropriate when the sorbate can penetrate thesolid, and in equilibrium it is uniformly distributed inthevolume of the solid. For adsorption, the R defined by Eq. (16) is approximately proportional to the specific surFace area of the solid, and itmay be expedientto relate the amountsorbed to the surface areaof the solid rather than to its mass. Certainly R is a function of the pH, and the plots of log R(pH) are often linear unless hydrolyzed species in solution play a significant role. Figures 15-17 show some important properties of R(pH) plots. All the curve; (R in dm3/kg) were calculated for a 144 m2/g silica with six surface sites per nm-. The stability constants of surface complexes (+M, were chosen to obtain R 100 (50% uptake for 10 g silica per dm3) at pH 7.5. This result was found experimentally for nickel at low initial concentrations. The ionic strength is alwaysmol/dm3,the stability constants of hydroxycomplexesin solution were taken from Ref.7, and the TLM parameters are given in the footnote to Table 2. When the exponential factor expressing the contribution of electrostatic energy to the Gibbs energy of surface reaction is neglected (Fig. 1 9 , the slope of the log curves is equal to the numberof protons released per adsorbed heavy metal (the experimental slope forNi is about 1). This is the principle of calculation of r fromKurbatov plots. At pH 10, R levels out (plateau) for r 2 NiOH surface species for which the Boltzmann factor disappears in TLM), SiONi”surface species) R reachesamaximum andthen andfor r 1 decreases. This pH value is characteristic for nickel: modelcurves hydrolyzing at lower pH values show a plateau or maximum at lower For a model involving TLM (Fig. 16), the slope of log R(pH) curves at pH 10 is about 2 for r and r as well. The difference between the curves obtained for r 1 and r at low pH values reflects the deviation of the surface potential calculated from the TLM model and from the Nernst equation. Figure 16 shows that Kurbatov plots do not give a proper value of r when the electrostatic contribution to the adsorption Gibbs energy is significant.

c

Calculated distribution coefficients for an adsorption model ignoring the electrical double layer.

l

Calculateddistributioncoefficients for an adsorption modelinvolving TLM (see Table 2 for values of TLM parameters).

SiONi+ SiONiOH

4

5

6

7

solution

8

9

Calculated distribution coefficients for an adsorption model ignoring the electrical double layer but considering Ni silicate complexes in solution.

Sharp maxima on R(pH) curves (model curves forr 1 in Figs. 15 and 16) are not observed experimentally for metalcations, but a maximum on the R(pH) curve has been reported for uranium sorption onsilica, probably due to formation of carbonate complex at high pH [53]. On the other hand, the plateauin the model curve forr (Fig. 15) reflectsthe experimentally observed trend that the sorption of cations on silica often does not reach 100% even at very high pH values. Sorption of Pm on colloidal silica at an initial concentration of metalions on theorder of mol/dm3 is an example of such a behavior. A maximum adsorption of 90% is reached at pH 8 and further increase in the pH causes even a decrease of adsorption 1541. Model curves presented in Fig. 17 were obtained by considering competition of soluble silicate complexes in addition to that of hydroxy complexes. The stability constants of silicate complexes in solution were chosen to obtain a plateau at about 90% of uptake and they are not supported by any literature data. This approach gives a flat maximum for a model with r 1 and a 1:1 silica complex in solution and a plateau for r 2 when the silicate-to-metal ratio in the soluble complex equals 2.

For low ionic strengths the activity of ions in solutions can be estimated using Davies’ equation [SS]:

This means that increase of the concentration of inert electrolyte from 0.01 to 0.1 mol/dm3 de resses the activity of divalent ions in solution by a factor of 2 (more exactly 10 and for trivalent ionsthecorrespondingfactor is 4 (10°.56).The effects of inert electrolytes on the activities of ions can be measured by means of ion-selective electrodes; they are also reflected in enhanced solubilities of sparingly soluble salts. Let us recall a typical reaction used in modeling of heavy metal ions adsorption [Eqs. (8) and (9)]: =SiOH

Me' =SiOMe' H' SiOMe'](H')[= SiOH]" (Me"}"

exp [-F(q0

In order to keep the K=SiOMe+ constant, the decreasein the activity of the Me2' ions by a factorof 2 (induced by increase of the ionic strength from0.01 to 0. 1 mol/dm3) should cause a decrease of SiOMe'] and thus of the adsorption by a factor of 2 at a constant pH.Similar effects are expected for other modelreactions. Therefore, if the stability constants of surface complexes are defined in terms of activities of ions in solutionandconcentrations of surface species, theadsorptionshould decrease when the ionic strength increases. O'Day [S61 obtained uptake curvesof Co(1I) at initial concentration of 3 mol/dm3 on quartz independentof the ionic strength over the range0.01 to 1 mol/ dm3. These uptake curves werevery sharp withachangefromalmostzero to almost 100% sorption within 0.5 pH unit. Spark et al. [24] observed a shift in pHso by 0.1 to 0.2 pH unit to higher pH values on increase of the K N 0 3 concentrationfrom 0.001 to 0.1 mol/dm3for Cu(II), Zn(II), Co(II), and Cd(Il). When 0.1 mol/dm3 KN03 was replaced by 0.1 NaCl the shift was similar for Zn and Co. For Cd the shift was more significant (competition between complexing anionand surface sites) but the adsorptionof Cu increases in the presence of chlorides. The effect of NaC104 and Mg(N03)Z (up to 0.1 mol/dm3) on the uptake curves of mercury on silica at pH 6 negligible. ~nfortunately,for the pH range 2-6 only 0.1 mol/dm3 NaC104 data are available, Presence of 0.001 mol/dm3 NaCl shifts the pH where the onset of adsorption is observed by 4 pH units toward higher pH values (as compared with 0.1 mol/dm3 NaC104). Increase of the NaCl concentration from0.001 to 0.1 1nol/dm3 causes a further shift by 2 p pH 9 the effect of chloride concentrationis insignificant In order to explain this phenomenon quantitatively, Tiffreau et al. [58] considered two surface complexes, namely ==SiOHgOH and =SiOHgCl, and determined their stability constants assuming the diffuse-layer model. They adapted the acidity constants for silica obtained in TLM. Such an adaptation is not recommended (see Chapter 1 l), namely it leads to highly underesti~atedsurface charge densities of silica in the range of high pH values and high ionic strengths. Kosmulski and Szczypa [59] observed a shift of pHso for lanthanides adsorbed on silica by 0.3 pH unit for anincrease of ionic strength from0.001 to 0. 1 m01/dm3.

The distribution coefficient decreased by a factor of 4, i.e., less significantly than expected from Eq. (17). Adsorption of cadmium on silica is rather insensitive to addition of NaC104 but it is strongly depressed by 1 mol/dm3 NaCl due to formation of stable chloride complexes which do not adsorb on silica [60]. Adsorption of Gd and Y on silicais insensitive tothe ionic strength [51]. Adsorption of Ni is insensitive to the ionic strength at low initial concentrations but athigher initial concentrations the adsorption decreases when the ionic strength increases [61]. This effect is independent of the nature of the 1-1 electrolyte and the shift of the uptake curves approximately corresponds to that calculated from Eq. (17). The above results can be summarized as follows. For most heavy metal ions Eq. (17) overemphasizes the ionic strength effect on their adsorption on silica. This is probably because theionic strength effect on theactivity coefficient of the adsorbed species is similar as on that of ions in solution. Complexation of Hg(1I) and Cd by chloride ions by in solution clearly depresses their adsorption, but adsorption of Cu(1I) is enhanced in the presence of chlorides.

Adsorption of metal ions on oxides increases and the uptake curves shift toward lower pH values when the temperature increases. This pH shift can beused to quantifytheextent of thetemperature effect onadsorption. For sorption of heavy metal cations on alumina, the shift was found to be independent of the ionic strength and initial concentration of heavy metal cations [62]. However, the most usual way to quantify the temperatureeffect on adsorption is to calculate the standard enthalpy of adsorption which is positive (endothermic) when the adsorption increases with the temperature. Malati et al. [47] calculated the following enthalpies of cation adsorption on silica from the temperature dependence of the slopeof Langmuir isothermsin linear form at pH (in 7 kJ/mol: 23 for Cu(II),13 for Cd, 10 for Ni and 2 for Co. The result that the adsorption of heavy metal cations on silica (and other oxides)is endothermic seems paradoxical, because it is well known that adsorption is always exothermic:since the entropy decreaseswhen molecules or ions are adsorbed, the enthalpy must be negative (exothermic) to be thedrivingforce of the process. Anendothermicprocessmust be entropically driven. Theexplanation of this apparentparadox thatadsorption of heavy metal cations is a complex process involving their dehydration (not written explicitly in reaction equations), which is responsible for endothermic enthalpy and positive entropy of the overall process. Adsorption of heavy metal cations occurs in a complex system whose state is defined by the surface charge density, surface potential and concentrations of the heavy metal ions of interest, protons, and anions and cations of the supporting electrolyte in solution and in the interfacial region. Adsorption and surface charging equilibria are sensitive to the temperature and in order to keep some parameters constant, some others must vary. For example, adsorption of protons on oxidesis exothermic,thus cro is variable when pH is constantand vice versa. Standard enthalpy of adsorption at constant pH can be defined as [63]:

where I is the ionic strength, cads is the concentration of adsorbed cations, and and are their equilibrium activities (concentrations wereused in Ref. 63) in solution at temperatures TIand The pH and ionic strength are independent variables and theycan be directly controlled, and is controlled indirectly, riamely it is a function of the initial concentration of heavy metal cations. When adsorption isotherms for and are available, and corresponding to given Cads can be obtained by inter- or extrapolation. The enthalpy definedby Eq. (18)is endothermic, not only for adsorption of heavy metal cations but also for alkali-metal cations. This is because the surface potential at given pH is more negative and the electrostatic conditions are more favorable for cation adsorption. Standard enthalpy of adsorption at constant is defined

Hall [64] postulates the following definition of standard enthalpy of adsorption:

where cl c, are surface concentrations of all species. Such a definition can be easily applied when the components l y1 are molecules whose concentrations in solution can be adjusted independentlyto achieve required surface concentration. In systems involving ions, adjustment of all concentrations to get the required cl c, at two different temperatures violates the condition of electroneutrality. For example, to keep the surface concentration of anions of the s u p ~ o r t i nelectrolyte ~ constant the Concentration of the supporting electrolyte must be adjusted. This means that the concentration of cations of the supporting electrolyte is already fixed and no freedom is left to adjust the surface concentrationof cations independently. In other words, it is impossible to increase the concentration of anions of supporting electrolyte and simultaneously decrease the concentrationof cations of supporting electrolyte without illtroducing a new component. Therefore the present author does not find Eq. (20) practical and prefers Eq. (18) and (19). Adsorption of alkali-metal cations, Ca and Cd on silica at constant is insensitive to the temperature,i.e., the enthalpy calculated from Eq.(19) equals zero. the other hand, for a positive enthalpy of adsorption on silica at constant of 20-23 kJ/mol was found The difference between enthalpies at constant pH and atconstant approximatelyequalsthenumber of protons released perone adsorbed heavy metal cation multiplied by standard enthalpy of proton desorption from silica (about 23 kJ/mol). Adsorption isotherms of Cd and on silica are linear and the standard enthalpyof adsorption is independent of the concentration of these divalent cations. Anotherapproachto is to find an equilibriumconstant of thesurface reaction responsible for the adsorption and calculate the standard enthalpy of this reaction as: e

e

For example, the following surface reaction =SiOH

Cd2'

SiOCd'

H'

(22)

was found tobe successful ininterpretation of Cd adsorption onsilica as a function of pH and ionic strength [60]. TLM wasused to modelthesurfacepotential involved in the equilibrium constant of reaction (22). Different sets of TLM parameters give almost identical model charging curvesand the numerical valueof && in Eq. (21) depends on the choiceof these TLM parameters. The A a d s H o C d kJ/ mol is independent of experimental conditions (ionic strength, Cd concentration) and of the assumed parametersof TLM. Theinner and outer capacitanceand total number of surface sites were assumed to be independent of There are no firm physical grounds to make such an assumption, although there is also no direct method to determine the temperatureeffects on particular parametersof an electric double layer. This is the main difficulty in application of Eqs. (21) and (22). In view of these controversial assumptions, the AadsHoCd calculated from Eq. (21) is rather a comfortable empirical parameterwhich combines the enthalpyof adsorption and the temperature effect on double-layer parameters. The - - ~ ( P H ~ ~ ) / 0.018 ~ T pH unit/K and AadsHoGd kJ/mol [calculated from the uptake (pH) curves according to Eq. (18)] are rather insensitive to the initial Gd concentrationuptomol/dm3.Forhigher initial concentrationsthe values of both parameters gradually increase and the initial Gd concentration of mol/dm3theyare twice thevaluesfor m01/dm3. Qualitatively similar dependence of Aa&f-IONi on the initial concentration was observed. The critical Gd concentration atwhich A a d s H o G d starts to increase coincides with the concentration at which the Gd adsorption isotherm onsilica shows deviations fromlinearity. This means that (see Section The enthalpy of adsorption on strong sites is less endothermic than on. strong sites, or The enthalpy of formation of ternary complexes involving impurities isless endothermic than that of formation of metal-surface complexes

According to the surface complexation model, adsorption of heavy metal cations can be described in terms of mass law. Surface sites (silanol groups) react with metal ions to form complexes. The stability constants of these complexes are independent of the experimental conditions (pH, ionic strength, solid-to-liquid ratio). The de~nitionsof stability constants of surface complexes usually involve an exponential term expressing theelectrostatic work of transfer of adsorbed metal cations (and coadsorbed hydroxyl and other anions) from bulk solution to the surface an of desorbed protons from the surface to the bulk solution. Different formulae of surface complexes can be proposed considering:

Number of surface sites involved in adsorption of one metal cation (mono-and multidentate surface complexes) Number of protons released per one adsorbed cation Number of anions of inert electrolyte coadsorbed with one adsorbed cation Distance of the adsorbed cations and coadsorbed anions from the surface. The list of possible formulae of surface complexes shown below was compiled according to the following rules: The number of surface sites per one adsorbed cation does not exceed two. The number of protons released per one adsorbed cation can range from zero to z,the valence of the cation. Thecoadsorbedanions of inert electrolyte donot overbalancethe positive charge of the surface complex. Electrostatic attraction is necessary to form outer-sphere complexes. The triple-layer model makes it possibleto model inner-sphereand outer-sphere complexes. In the former the metal cations are adsorbed in the surface plane (the ltr0 potential) and in thelattertheadsorptionoccurs in the layer of chemisorbed counterions of supporting electrolyte (the potential). Inthe diffuse-layer model only inner-sphere complexes can be modeled. The following symbols havebeen used: SiOH, surfacesilanol group; Mv', Di2', and Tr3', mono-, di-, and trivalent heavy metal cation; anion of the inert electrolyte; separates the inner- and outer-sphere parts; y , electric potential in the following expression: Stability constant of surfacecomplex eXp(--.FY/Rr)

n(products)

n(reactants)-l

where is a product considering the stoichiometric coefficients (but the exponent of the reactant is always set to l); r, number of protons released per one adsorbed heavy metal cation.

Species Monovalent cations SiOHMv' SiOHMv' XSiOHMvX SiOMv MV+ Divalent cations Monodentate inner-sphere complexes SiOHDi2+ SiOHDi2'..X-

Y

0 0

" N O

Y

0 0 0

0 0

Species

Y

SiOHDi2'..(X"), SiOHDiX' SiOHDiX'. X" SiOHDiX, SiODi' SiODi' XSiODiX SiODiOH Outer-sphere complexes SiQ- Di2' identate inner-sphere complexes (SiOH)2Di2'

r

2@p

1 1

-2@0

"2@O

(SiQH)*DiX' (SiQH)*DiX' X(SiOH)2~iX2 (SiOH)(SiQ)Di' (~iOH)(SiO)Di' X(SiOH)(SiO)DiX (SiO),Di Trivalent cations Monodentate inner-sphere complexes SiOHTr3' SiOHTr3' XSiOHTr3'. (X"), Si~HTr3'. (X-)3 SiOHTrX2' SiOHTrX2'. XSiOHTrX2'. (X"), SiOHTrX: SiOHTrX: XSiOHTrX3 SiOTr2" SiOTr2'. XSiOTs2'. (X"), SiOTrX' SiOTrX'. XSiOTrX, SiOTrOH' SiOTrOH'. XSiOTrOHS SiQTr(QH)2

-@0

1

-@0

1

-3@0 "2@O "2@O

0

-@0

-2@0

"2@O

1 1 1

-@0

0

Species Outer-sphere complexes SiO- Tr3' SiO-. Tr3'XSiO-. Tr3'(X-)? Bidendate inner-sphere complexes (SiOH)2Tr3' (SiOH)2Tr3'.. X" (SiOH)2Tr3'.. (X-)z (SiOH)2Tr3+..(X-), (SiOH)2TrX2' (Si0H),TrX2'.. X(SiOH)2TrX2'.. (X-), (SiOH)2TrXi (SiOH),TrX:.. X" (SiOH)2TrX3 (SiOH)SiOTr2' (SiOH)SiOTr2'.. X(SiOH)SiOTr2'.. (X-)2 (Si0H)SiOTrX' (SiOH)SiOTrX+.. X(Si0H)SiOTrXz (SiO)2Tr' (SiO),Tr'. (Si0)zTrX (SiO)zTrOH

Y

llro llro llro

$8

0 0 0 0 0 0 0 0 0 0

-2.11'0

-2llro -2llro $0

0 "2.11'0 "2.11'0

2.11'8

llro

0

0 0

The factor and the number of protons released per one adsorbed heavy metal cation define the shape of the percentage of uptake (pH) curves. The number of anions coadsorbed with one heavy metal cation defines the ionic strength effect on the percentage of adsorption at constant pH. The long list of surface species makes it possible to find a surface complex for which both these factors agree with the experiment. The number of surface sites involved in the surface complex responsible for adsorption of heavy metal cations can be assessed by means of spectroscopic measurements. When a model with one kindof surface complex is not able to properly reflect the experimentally observedtrends, combination of two or more different surface complexes canbe considered. Interpretationof spectroscopic data obtained for a combination of two or more surface complexes may be very difficult.

Stability constants of metal complexes withsimilar ligands often show some degree of linear correlation. Surfaces of oxidesarecoveredwithhydroxylgroups.

Therefore, a linear correlation between hydrolysis constantsof metal cations and calculated stability constants of surface complexes canbe expected. The most reliable method to examine such correlationsis to take sorption data formany metal cations on one type of silica and at identical experimental conditions. Apparent stability constants are influenced by many factors, e.g.: surface area of silica, density of surface hydroxyl groups, and solubility, which are difficult to control or account for when results obtained with different samples of silica or at various experimental conditions are compared.Dtrgger et al. [67] were probably the first to demonstrate the correlation between adsorption on silica and the first hydrolysis constant. Schindler 1681 reports an excellent linear correlation between hydrolysis constants of four metal ions: Fe3', Cu2', Cd2', and Pb2' (these hydrolysis constants are well known in the literature [7,8]) and stability constants of surface complexes of these ions with amorphous silica (Fe with silica gel) obtained from author's own experimental results [69] (at ionic strengths ranging from 1 to 3 m01/dm3 NaC104). The results can be summarized in the form of the following two equations:

where the left-hand sides represent the stability constants of mono- and bidendate surface complexes and right-hand sides represent the first and second hydrolysis constant, i.e., all eight points (p*Q;, pK1) and (p*Q$,ppz) are located on the same straight line passing through the origin of the ordinates. This linear plot has been reprinted in several subsequent publications. The quotients *Q; and *Q$ in Eqs. (23) and (24) are related to the stability constants of corresponding surface reactions by the following equation:

*Ks

exp[(z

(25)

where is the valence of the metal cation and r is the number of protons released per one adsorbed cation. Schindler explains his approach (using quotients rather than equilibrium constants)by approximate equalityof and As a matterof fact the experimental data on Cd adsorption on silica reported in Ref. 69 can be reproduced using the following reactions: Cd2' Cd2' Cd"

(26) -=CdOH3.7 pp2 (27) 7.7 20H" Cd(OH)2 SiOH SiOCd' H' p*Qy 6.09 (28) 2 SiOH SiO);?Cd 2H' p*Qi 14.20 (29)

corresponds to the solid-to-liquid ratio used in the adsorption equilibriumconstants of reactions (28) and (29)were taken the hydrolysis constants are well known [7,8]. According to the convention used by Schindler the equilibrium constant of reaction (29) has square of concentration of surface groups in the denominator (see Introduction). Figure18 shows that one experimental uptake curve is not sufficient to determine *Q; and simultaneously. While with models using one surfacespecies the stability constant for this species can be unequivocallydeterminedfromexperimental data

Adsorption edges for Cd calculated for two different stabilityconstants of bidentate surface complex. The stability constant of the rnonodentate surface cornplex was fitted to match the experimental data from Ref. 69.

provided that the density of surface sites is known, models with two or more surface species lead to ill-posed problems, i.e., many different sets of model parameters lead to almost identical model curves. The model curvesin Fig. 18 are very similar (considering a scatter of results in this kind of measurements) in spite of difference in by two orders of magnitude. Calculations of stability constants of surface complexes always involve some assumptions which may be disputable. Therefore, demonstration of correlation of pHso obtained for different heavy metal cations of the same valence and their hydrolysis constants would be more convincing. Fig. l9 shows the Schindler’s data and results from more recent publications 124,291. The data from Ref. 24 show some degree of linear correlation between pHso and the first hydrolysis constant, but the data from Ref. 49 are not correlated. Ref. 68 shows a somewhat idealized picture, while in fact the correlationbetween the hydrolysis in solution and sorption on silica is moderate.

According to the Boltzmann equation, the concentration of cations near a negatively charged silica surface is higher than the concentrationin the bulk solutionby a factor of exp where is the local electric potential. For divalent cations this gives an increase of concentration by a factor of 100 for -59 A purely electrostatic adsorption shouldbe sensitiveto the valence of the cation but

1st hydrolysis constant The pH50 from Refs. 24, 49, and 69

a function of the first hydrolysis

constant.

not to its nature. Interestingly, the adsorption behavior of all divalentcations (heavymetal and alkaline-earth metal)reported in Ref. 67 was very similar. Reference 52 reports identical adsorption behavior of Zn and Cd and Ref. 23 reports identical adsorption behavior of Mn, and Ni. A limited pH range was covered in these studies. Their results show no ion specificity and they apparently confirm the electrostatic mechanism of sorption. The difference between the highest and the lowest pH50 reportedin Ref. 49 for various divalent cationswas 1 pH unit, and the corresponding difference in Refs. 24 and 69 exceeds pH units, so these results clearly show ionspecificity. In thepurely electrostatic adsorption some degreeof ion specificity may be induced by difference in ionic radii, but the differences demonstrated in Refs. 24, 49, and 69 are too significant and they rather support the surface complexation model. According to thewidely accepted imageof the electric double layer, the absolute value of electric potential at certain distance from the silica surface decreases when the ionic strength increases. It is known from experimentsthat the adsorption of heavy metal ions is rather insensitive to the ionic strength, thus, the heavy metal ionsareadsorbed at a very shortdistancefromthesurfacewherethe electric potential is not influenced by the ionic strength. The distance between adsorbed heavy metal ions and the surface silicon atom canbe derived from EXAFSspectra: 0.38 nm for Th-Si[44], and 0.34 nm for Cr-Si[39]. Probably the distances are similar for other heavy metal cations. Such short distances are not accessible for

cations of the supporting electrolyte, so the electric potential +Me [Eq. is close to the surface potentialwhen the concentrationof heavy metal ions is not too high. Let us try to estimate this potential assuming a purely electrostatic mechanism of adsorption. For simplicity let us assume that the thicknes of the layer of adsorbed cations is always 0.3 nm, i.e., somewhat less than typical diameters of hydrated cations. When thesolid-to-liquid ratio and the specific surface area of the solid are known, the concentrationof metal cations in the surface layer and thus the apparent electric potential (for a purely electrostatic model) can be estimated. Some results (calculated for pH50 are shownin Table 3. These values are lower than surface potentials estimated from theTLM model but higherthan potentials measured at low concentrations of heavy metal cations. The above electrostatic model ignores therelease of protons from the surface when the heavy metal cations are adsorbed. Thereis no direct method to find the number of protons released per one adsorbed heavy metal cation atvery low initial concentrations. Let us roughly assume that one proton is released per one adsorbed heavy metal cation. The surface potentials estimated from hdsorption data of heavy metal cations using an electrostatic model with qMe and r 1 are shown in the last column of Table 3. Figure 20 compares these surface potentials with the potentials calculated for model (for parameters see Table 2). Considering the simplicity of the model, and the fact that the same parameterswere used for different types of silica and for differnet heavy metal cations, the correlation between the lines (TLM) and the symbols (Table 3) in Fig. 20is quite reasonable with an exception of two data points from Ref. 24. This discrepancy can be explained by surface precipitation of Cu(Z1) hydroxide, which is likely in view of relatively high initial concentration of

Electric Potential $Me and Estimated from Purely Electrostatic Model of Adsorption Using the Boltzmann Equation. An Adsorption Layer Thickness of 0.3 nrn Assumed

6.8 8.9 73 8.6 5.6 7.7 5.8 5.2 5.4 7.3 7.4 6 7.2 5.8

cu Zn Cd CO cu Gd Y Cd Ni cu Cd Pb

24 24 24 24 39 40 50 51 51 60 61 69 69 69

144 144 144 144 94 126 61 66 66 87 99 80 80 80

288 288 288 288 141 25 1 122 99 99 173 199 161 161 161

50 5

The surface potential of silica calculated from adsorption of heavy metal cations (symbols) and from model for 0.01 (solid line) and 0.001 mol/dm3 NaCI04 (dashed line).

Cu and the tendency of this cation to cluster at silica surfaces even at low surface be refined coverages as detected by EXAFS 1411. The above electrostatic model can by using different thicknesses of the surfacelayer and different numbersof protons released per one adsorbedheavy metal cation for different cations to obtain a better correlation than is presented in Fig. 20. Ion specificityis also governed by the contribution of dehydration of adsorbing cations to their Gibbs energy of adsorption.

Different aspects of sorption of heavy metal cations on oxides have been discussed inanumber ofreviews and monographs 1701. Manystudies were devotedto adsorption on iron 161 and aluminum oxides and hydroxides, fewer to adsorption on manganese(1V) and titanium(1V) oxides, and studiesof sorption of metal cations on other oxides and hydroxides are rare. Most of the features discussed in the present chapter are not unique to silica, but they are common for different metal oxides. Certainly, electrostatic attraction is not the driving force for adsorption of cations positively chargedsurfaces of alumina.The affinity series of metal cations is similar for different oxides and it corresponds approximately to their first hydrolysis constant. However, the correlation between adsorption on different oxides is far from linear. For example, the uptake curves for alumina and silica

obtained at the same solid-to-liquid ratio (in m2/dm3)were almost identical for Zn and Cd, while adsorption of CO and Cu on alumina was much higher than on silica [24]. Generally the interactionof heavy metal cations with the surface of alumina is stronger, e.g., complexation of Cd by chlorides exerts more significant effect on sorption of Cd(I1) on silica than on sorption on alumina[60]. Adsorption of Cu on A100H is strong enough to prevent clustering 1711 while the same metal cations forms precipitates on silica even at low coverages [dl].In spite of a few differences, the following properties are common for adsorption of heavy metal cations on silica, on the one hand, and metal oxides, on the other. The courseof uptake curvesis rather insensitive to the ionic strength (adsorption of some cations on aluminaslightly increases when the ionic strength increases). The course of uptake curves is insensitive to the initial concentration of heavy metal cations when this initial concentration is sufficiently low.. The uptake at constant pH increases with the temperature. The point of zero charge of the oxide shifts to lower pH values and the isoelectric point shifts to higher pH values (compared with thepristine value) when the concentration of heavy metal cations is sufficiently high. Formation of inner-sphere complexes canbe detected by EXAFS spectroscopy, but some metal cations form surface precipitates already at low surface coverages. The pH andionic strength effect on adsorption canbe modeled using the surface complexation model.

A grant from the Alexander von Humboldt Foundation edged.

is gratefully acknowl-

1. R. K. Iler, The Chemistry Silica, Wiley, New York 1979, p. 673 ff. 2. B. N. Laskorin, V. V. Strelko, D. N. Strazhesko, and V. Denisov, sorbent.^ on the Basis of Silica in ~ u d i o c h e ~ i s t rAtomizdat, y, Moscow, 1977. 3. T. F. Tadros and J.Lyklema. Electroanal. Chem.Int. Electrochem. 17:267 (1968). 4. G. H. Bolt. J. Phys. Chem. 61:1166 (1957). 5. P. Ney, Zeta Potentiule und Flotierbarkeit von Miner~len,Springer, Vienna, 1973, p. 153. 6. L. K. Koopal. Electrochimica Acta. 41:2293 (1996). 7. R. M. Smith and A. E. Martell, Critical Stability Constants, Vols.4-6, ~ l e n u m Press, New York, 1976. 8. E. Fogfeldt, Stability Constunts Metal Ion Complexes, Pergamon Press, Oxford, 1982. 9. G. Pokrovski, PhD. Thesis, Toulouse, 1996. 10. M. P. Jensen and G. R. Choppin Radiochim. Acta. 72143 (1996). 11. D. B. Kent and M. Kastner. Geochim. Cosmochim. Acta 49:1123 (1985).

12.

Charlet and A. Manceau. Geochim. Cosmochirn. Acta 58:2577 (1994). F. M,Morel, ~ r i n c ~ l e s o f A q ~Chemistry, atic Wiley, New York, 1983, pp. 387-388. 14. K. Milonjic.ColloidsSurf. (1992). €5. J. Depasse, J. Colloid. Interf. Sci. 194260 (1997). 16. D. A.Dzombakand F. M.Morel, Swface Complexation ~ o ~ e l i n g : Ferric Oxide, Wiley, New York, 1990. 17. K. F. Hayes and J. 0. Leckie. J. Colloid Interf, Sci.115:564(1987). 18. I, C. Pius, M. M. Charyulu, B. Venkataramani, C. K, Sivaramakrishnah, and S. K. Patil. J. Radioanal. Nucl. Chern. 299:1 (1995). 19.W.Srnitand C. L. Holten. J. Colloid Interf. Sci.78:1(1980). 20. Janusz. J. Radioanal. Nucl. Chem. (1989). 21. A. Kolics,E.Maleczki, K. Varga, and G. Horanyi. J. Radioanal. Nucl.Chem. 158:121(1992). 22. IS. Milonjic, D. M. Cokesa, and R. V. Stevanovic. J. Radioanal. Nucl. Chem. 158:79 (1992). 23. F. Vydra and J. Galba. Coli. Czechosl. Chern. Comrn. (1969). 24. M. Spark, B.B. Johnson, and J. D. Wells. J. Soil Sci. 46521 (1995). 25. R. Alexander Clauss and K. Weiss, J. Colloid Interf. Sci. 62:577 (1977). 26. H. A. Ketelson, R. Pelton, and M. A. Brook. Langrnuir 12:1134 (1996). 27. H. M. Jang and D. W. Fuerstenau. ColloidsSurf.21:235(1986). 28. R. 0 . James and T. W. Healy. J. Colloid Interf. Sci. (1972). 29. R. Herrera-~rbinaand D. Fuerstenau. Colloids Surf. A. 98:25(1995). J. R. Crawford, I.H. Harding, and D. E. Mainwaring. J. ColloidInterf.Sci. (1996). 31. M. Yasrebi, M. Ziomek-Moroz, W. Kernp, and D. H. Sturgins. J. Am.Chem. Soc. 79:1223 (1996). 32. J. H. Anderson. Catal. 28:76(1973). 33. P. Roose, J. van Craen, C.Pathmamanoharan,and H. Eisendrath. J. Colloid Interf. Sci. 188:115 (1997). 34. P. Roose, J. van Craen, G. Andriessens, and H. Eisendrath. J. Magn.Reson. A 220:206 (1996). 35. V. Hronsky, M. Rakos, J. Belak, and D. Kazar. Czech Phys. B28:1277 (1978). 36. G. Martini. J. Colloid Interf. Sci.80:39(1981). L. Burlamacchi and G. Martini, in ~ a g n e t i c in Colloid and ~ n t e r f ~ ~ e Science (J. P. Fraissard and H. A. Resing, eds.), Reidel, 1980, p. 621. 37. A. vonZelewsky and J. M. Bemtgen. Inorg. Chem.21:1771(1982). 38. W. E. Stone, G. E1 Shafei, J. Sanz, and S. A. Selim. J. Phys. Chem. 9710127 993). 39. S. F. Fendorf, G, M. Larnble, M. G. Stapleton, J. Kelley, and Sparks. Environ. Sci. Technol. 28:284 (1994). 40. P. O'Day, C. J. Chisholm-Brause, S. N. Towle, G. A. Parks, and G. E. Brown. Geochim. Cosmochirn. Acta 6122.5 15 (1996). 41. Xia, A. Mehadi, R. W. Taylor, and W. F. Bleam. J. Colloid Interf. Sci. 185:252 (1997). 42. A. W. Dent, J. D. Ramsay, and S. W.Swanton, J. Colloid Interf. Sci. (l 992).

43. Q. Clause, M. Kermarec, L. Bonneviot,F. Villain, and M. Che. J. Am. Chern. SOC. 114:4709 (1992). 44. E. Osthols, A. Manceau, F. Farges, and L. Charlet. J. Colloid Interf. Sci. 194:lO (1997). 45 T. Fanghanel, J. I. Kim, P. Paviet, R. Klenze, and W. Hauser. Radiochim. Acta 66/67:81(1994). 46. IS. H. Chung, R. Klenze, K.K. Park, P. Paviet-Hartmann, and J. I. Kim. Radiochim. Acta 82:215 (1998). 47. G. C. Bye, M. McEvoy, and M. A. Malati. J. Chern.Soc. Faraday Trans. 1. 79:231l (1983), Chem, Tech. Biotech, 32:781 (1982). M.A. Malati, McEvoy, and C, R. Harvey. Surf. Techn. 17:165 (1982). J. Park, K. H.Jung, K. K.Park, and T. Y. J. Colloid Interf. Sci. 160:324 48. (1993). 49. A. R. Sarkar, P. K. Datta, and M. Sarkar. Talanta 43:1857 (1996). 50. Y. J. Park, K. H. Jung, and K. K. Park. J. Colloid Interf. Sci.171:205(1995), 172447 (1995). N. N. Vlasova and N. K. Davidenko.ColloidsSurf.A119:23 (1996). N. N. Vlasova, N. K. Davidenko, V. I. Bogornaz, and A.A.Chuiko. ColloidsSurf.A104:53(1995). R. Kocjan and R. Swieboda.Microchim.Acta 122:175 (1996). 51. M. Kosmulski. J. Colloid Interf. Sci. 195:395 (1997). 52. J. Liu, S. M. Howard, and K. N. Han. Langrnuir 9:3635 (1993). 53. K. H. Lieser, S. Quand-Klenk, and B. Thybusch. Radiochim. Acta 57:45 (1992). 54. B. Satmark and U-. Albinsson. Radiochim. Acta 58/59: 155 (1992). ~n, London 1962. 55. C. W. Davies, Ion A s ~ ~ o c i a t iButterworths, 56. P. A. Q’Day. PhD. Thesis, Stanford University, 1992. 57. M. G. MacNaughton and R. Q. James. J. Colloid. Interf. Sci. 47:431 (1974). 58. C. Tiffreau, J. Lutzenkirchen, and P. Behra. J. Colloid Interf. Sci. 172532 (1995). 59. M. Kosrnulski and J. Szczypa. J. Radioanal. Nucl. Chern. 144:73 (1990). 60. M. Kosmulski. Colloid Surf. A 117:201 (1996). 61. M. Kosrnulski. J. Colloid Interf. Sei. 190:212 (1997). 62. M. Kosrnulski. J. Colloid Interf. Sci. 192:215 (1997). 63. M. L. Machesky, in ~odeZingin Aqueous Systems 11(D. C.Melchior and R. L. Bassett, eds.), ACS Symposium Series No. 416, p. 282 (1990). 64. D. G. Hall. Langmuir 13:91 (1997). 65. M. Kosmulski.ColloidsSurf.A83:237(1994),Ber.Bunsenges.Phys.Chem. 98: 1062 (1 994). 66. M. Kosmulski, J. Colloid Interf. Sci. 211:410 (1999). 67. D. L. Dugger, J. H. Stanton, B. N. Irby, B. L. McConnel, W. W. Curnmings, and R. W. Maatman. J. Phys. Chern. 68:757 (1964). 68. W. Schindler. Qster. Chern. 86:141 (1985). 69. P. W. Schindler, B. Furst, R. Dick, and P. U. Wolf. J. Colloid Interf. Sci. 55:469 (1976). 70. P. Schindler,in A~sorl?~ion Inorg~nic,~ at ~olid-LiquidInterfaces M. A. Anderson and A. J. Rubin, eds.), Ann Arbor, 1981, Chapter 1. D. G. Kinniburgh and M. L. Jackson,in Adsorption Inorganics at SoZid-Li~~id Interfaces (M. A.Anderson. and A. J. Rubin.eds.).AnnArbor.1981. D. 91.

J. Lyklerna. Croat. Chem. Acta, 60 (1987) 371. P. W. Schindler and Stumm, in Aquatic Surface Chemistry Sturnrn. ed.), Wiley, 1987, p. 83. K. F. Hayes and L. E.Katz, in Physics C ~ e ~ i s t r y~ i ~ e rSurfizces al (P. Brady, ed.), CRC, Raton, 1996, p. 147. 71. F.J. Weesner and W. F. Blearn. J. Colloid. Interf. Sci.196:79(1997).

Henry Krumb School of Mines, Columbia University, New York, New York

I. Introduction 11. Silica-Water Chemistry 111. Mechanisms of Surfactant and Polymer Adsorption on Silica IV. Adsorption of Surfactants V. Adsorptio~of Polymers VI. Adsorption of Surfactant-~olymer Mixtures VIE. Summary References

Surfaceproperties ofsilica controlmanyindustrialprocessesthat utilize this mineral. These properties can be markedly altered not only through its state of hydrolysis and pretreatment, but also by the adsorption of surfactants, polymers, and surfactant-polymer mixtures. In this chapter, mechanisms of adsorption of different types of surfactants and polymers on silica are reviewed. Effects of surface modification of silica by adsorption of surfactants, polymers, and surfactant-polymer mixtures are examined with relevance to industrial processes. Synergistic and antagonistic effects of surfactant polymer interactions at silica-water interface are also discussed.

Silica is the most abundant mineral of the earth's crust and it is encountered in various industrial processes. To control the performance of these processes, it is often necessary to modifyitssurfacepropertiessuchas zeta potential, surface energy, hydrophobicity, and adsorption capacity. Surface properties of silica are a function of its state of hydrolysis as well as pretreatment. Adsorption of surfactants and polymers canalso lead to marked changes in its interfacial properties. In this chapter, adsorption of surfactants and polymers on silica and its influence on the interfacial properties and interfacial processes ofsilica suchas flotation response, hydrophobicity, and grindability characteristics are discussed.

Silica particles in aqueous solution possess a surface charge due to preferential dissolution and interfacial ion exchange at the solid-liquid interface, The magnitudeand signof thesurfacecharge is dependent on theconcentration of the potential determining ions and is a function of pH and ionic strength of the solution. The dissolution process of silica in aqueous solution is illustrated in Fig. 1 [l]. When fresh silica is contacted with water, hydrolysis of the surface species takes place giving rise to silanol groups atthe surface (Fig. la). When silica is exposed to water for longerintervals, its surface tendsto hydrolyze further to form =Si(OH)2 (Fig. l b). Formation of more silanol groups is possible, particularly under low bulk concent~ationof alkali metals (Fig. IC). Further hydrolysis of the above surfacecan result in the dissolution of the surface species to form bulk silicic acid and expose new silicon atoms to the bulk water (Fig. Id). The above surface hydrolysis followed by dissolution will take place until soiid-solution equilibrium is reached. Figure la-d represents a continuous increase of surface hydroxyl groups on silica to form a silicic acid surface. When silica is in contact with fresh aqueous solutions, the surface will form successively Si(OH), Si(OH)2, Si(OH)3 at the interface and Si(OH)4 in solution until equilibrium is reached. The difference between a fresh silica surface in waterand the subsequently formedsilicic surface is that the number of ionizable sites per silicon atom that generates the negative surface charge is higher for the latter than for the former. Themaincauseforthesurfacechargegeneration is the dissociation of the silanol groups at the interface: -Si-OH

H' +-Si'+H20 or

-Si-OH

OH-

-Si-OHzS

-Si-O"+H20

Due to the acidic nature of the silanol groups, silica possesses a negative chargeat neutral pH. The pointof zero chargeof silica is about pH The number and type of silanol groups on the surface and the degreeof hydration of the surface are the major factors that determine the surface propertiesof silica particles. The rate of silica hydrolysis is a function of the solution pH, ionic strength, and temperature. The surface properties are affected also by prolonged contact with

HOH

The hydrolysis of the surface species

of silica. (From Ref. 1.)

solutions. It is to be noted that pretreatment ofsilicaby reagents such as HI;, HNOs, and NaOH, often employed in research, can produce drastic effects on its surface properties. These effects have been demonstrated by electrokinetic measurements. It has been shown thattreatment of quartz withhydrofluoricacid solution followed by warmsodiumhydroxidesolutionhasamarked effect on the zeta potential of the quartz (Fig. 2). The isoelectric point can be shifted by as many as four pH units. At pH5.8, the zeta potentialshifted gradually from 3 mV to -37 mV in about 22 h Thus the isoelectric point of the treated silica can change depending on the aging time.

/TEMP.

ST

l

H

H

c"'

c-

z

c-

6

Zeta potential of quartz treated with hydrofluoric acid and then with hot sodium hydroxide solution asa function ofpH without aging,and atpH 6 after 4, 6, and 22 h of aging; dotted curve is for equilibrium values obtained for HNO, treated quartz for comparison. (From Ref. 2.)

Surface properties can be modified by the adsorption of polymers and surfactants. Adsorption itself can essentially be considered as preferential partitioning of the adsorbate into the interfacial region. It is the result of one or more contributillg forces arising from electrostatic attraction, chemical reaction, hydrogen bonding, ~ ~ d r o p h o binteractions, ic Coulombic interaction, and solvation effects. Surfactantscan beclassified as cationic, anionic, nonionic, and zwitterionic depending on the nature of their hydrophilic group. This classification can also be applied to polymers according to the functional groups on the polymer. Cationic surfactants and polymers adsorb readily on silica due to electrostatic interactions, since the pH of most practical systems is above 2, and there can be such interactions between these cationic reagents with silica, which is negatively charged under these conditions. Anionic surfactants or polymers do not adsorb onsilica at neutral pH due to the presence of similar charge on the solid and the adsorbate. owever, anionic sur-

factants can adsorb onsilica in the presence of multivalent metal cationsin the pH range in which metal ions hydrolyzeto their first hydroxyl complexes. This has been attributed to the chemisorptionof the first hydroxyl species of metal ions on silica and modification of its surface (31, and the surface precipitation of hydroxides [4]. Nonionic ethoxylated surfactants and polyethylene oxide polymers have been found to adsorb on silica. Because of the absence of ionic components in these surfactants and polymers, the driving force for the adsorptionof these reagents is considered to be hydrogenbondingrather than electrostatic interaction.The hydrogenbondingoccurs between theether oxygenof thesurfactant/polymer and the surface hydroxyl groups silica: of -§iOH O(CH2CH2)2=. Interestingly, another group of nonionic surfactants, the alkyl polyglucosides, is found to exhibit minimal adsorption on silica

Adsorption of surfactants and polymers on solids is of major technical importance, since it can lead to changes in a variety of interfacial phenomena such as colloid stability (dispersion, flocculation) and wetting behavior (flotation, detergency). §urfactant molecule contains both hydrophilic and hydrophobic moieties. They can modify theinterfacial properties significantly at very low concentrations. Also, they can form aggregates in solutions and at the solid-liquid interface by hydrophobic interactions at certain concentrations. The conformation of surfactants at solid-liquid interfaces governs the interfacial properties. Effect of adsorption on interfacial parameters (contact angle, zeta potential) and consequently flotation recovery are bestillustrated in Fig. 3 for the cationic surfactant, dodecylammonium acetate (DAA), on quartz For DAA and quartz at neutral pH, increase in adsorption due to association of surfactants adsorbed at the solid-liquid interface into two dimensional aggregates called solloids (surface c ~ Z Z ~ or ~ ~hemimicelles s) occurs at about 10-4M DAA. This marked increasein adsorption densityis accompanied by concomitant sharp changesin contact angle, zeta potential, and flotation recovery, suggesting that these phenomena depend primarily on the adsorption of surfactant at the solid-liquid interface. The surface phenomena that reflect conditions at the solid-liquid interface (adsorption density and zeta potential) can in many cases be correlated directly with the phenomena that similarly reflect conditions at the solid-liquid-gas three-phase triple contact (contact angleand flotation). Increase in the length of the hydrophobic chain of a surfactant is expected to cause an increase in its adsorption owing to increased lateralinteractions between the hydrophobic chains. An effectively lower dielectric constant of the interface would also promote adsorption of a long-chain material. Flotation of quartz at neutral pH clearly shows a systematic increasein adsorption with increase in chain length. [6] (Fig. 4). Adsorption of surfactants can alteralso the strength propertiesof the solid, for example, the minimum stress required for fracture, by changing its surface energy. In grinding mills, the surfactant adsorption can influence the friction characteristics of the particles and the grinding media and affect the flow characteristics of the mill. Addition of dodecylammonium chloride to grinding mills has been found to

Correlation of adsorption density, contact angle, flotation response, and zeta potential forquartz as a function of dodecylam~oniumacetate concentrationat pH 6 to 7, 20 to 25°C. (From Ref. 6.)

The effect of alkyl chain length on the relative flotation response of quartz in the presence of alkyl ammonium acetate solution. (From Ref. 6.)

produce both beneficial and detrimental effects in the case of grinding of quartz depending on the pulp pH and theamine dosage (Fig. 5). The increase in grinding efficiency in the alkaline and neutral pH ranges was attributed to the enhanced fluidity of the mineral pulp in the mill, while the decrease in the low pH range is due to the increase in the floc strength induced by surfactant bridging between particles [7]. Adsorption of nonionic ethoxylated surfactant on silica has been studied to determine the effect of ethoxylation and hydrocarbon chain length. The dependence of adsorption on ethoxylation of the surfactant is illustrated in Figs. 6 and 7 using surfactants withdifferent EO numbers but the samealkyl chain. Theresults show the slope of the ascending part of the isotherms to increase with decrease in ethoxylation, indicating increased interactions between the hydrocarbon chains of the adsorbed surfactant, possibly due to reduced steric hindrance to hydrocarbon chain association by the ethoxyl group. For adsorption in the low-concentration region where chain-chain interaction is not dominant, the adsorption is higher for the longer ethoxyl chains due to the cumulative hydrogen bonding between the surfactant and the solid. Thus the adsorption density in the ascending part at low concentration is greater for the surfactants with the higher ethoxylation than those with the lower ethoxylation, and this trend is reversed at higher concentrations

Effect of amine on amount of -200 mesh product by wet ball milling of quartz. (From Ref. 7.)

om

kmoi/m3

Adsorption tertiaryoctylpheno~yethoxylatedalcohols (EO10 and 8.) on silica.

U

A m=

km~l/m

Adsorption nonylphenoxy~thoxylatedalcohols (EOlo,EOzo,and EO40) on silica. 8.)

(near theplateau). The plateau adsorptionof the surfactant decreasesprogressively as surfactant ethoxylation increases [8]. The packing area of the surfactant estimated from the plateau adsorption is plotted function of ethoxylation number in Fjg. 8. linear relationship is obtained and from the slope parking area of 9.2 per -(OCH2CH2) segment is obtained, implying adsorption of the whole ethoxyl chain on silica. The adsorption isotherms of surfactants of different chain length but the same degree of ethoxylation are given in Figs. 9 and l0 for EOlo and E040, respectively.

40Q

Effectofdegreeofethoxylationonmolecularparkingareascalculated from plateau adsorption densities for alkylphenoxyethoxylated alcohols on silica. (From Ref. 8.)

krnol/rn3

Adsorption of EO10 alcohols (C8,

and Clz) on silica. (From

pH

kmol/mS

Adsorption of EO4@alcohols

on silica. (From

8.)

8.)

The slopes of the ascending part of the isotherms are comparable, and the maximumplateauadsorptiondensity is relatively independent of thehydrocarbon chain length. The onset of the plateau adsorption is found to shift to lower concentrationsin line withtheobserveddecreasein critical micelle concentration (CMC)with increase in hydrocarbon chain length. The effect of surfactant adsorption on the hydrophobicityof silica is illustrated in Figs. 11 and 12. The hydrophobicity of silica was determined by mixing toluene with surfactant-silica system in a separatory funnel. The silica-surfactant-toluene dispersion was shaken and then allowed to settle. The two phases (aqueous and toluene) were allowed to flow out of the funnel separately and weight of the silica in these phases measured. The hydrophobicity was calculated as the percentage of solid that is in the toluene phase. In these figures,while in the absence of the surfactantthe silica surface exhibits complete hydrophilicity, withincreasein adsorption of both C8@EOloand C8@E040on silica, the surface becomes hydrophobic.Whereas silica particles in C8@EOlo solutionremainhydrophobic as adsorption is increased, the C8@E040surfactant makes the particles hydrophilic above a certain adsorption level. At saturation adsorption of Cg@E040the hydrophilicity of silica particles is totally restored. The change in hydrophobicity of particles is proposed to be the result of a corresponding change in the orientationof the hydrocarbon chains at the interface. In the low concentration region, the hydrocarbon chainslie flat on thesilica surface and render the silica hydrophobic. As adsorption increases, the flat hydrocarbon

0 D

40

g

kmol/ms

Changesinthehydrophobicity C8@EOI0.(From Ref. 8.)

ofsilica

particleswith

adsorptionof

an

QO40.

Changesinthehydrophobicity (From Ref. 8.)

ofsilica

particleswith

adsorption of

chain will tend to be squeezed out. In the caseof C8Q>E040 adsorption, theeffective surface area covered by the hydrocarbon chains is reduced, causing a decrease in hydrophobicity. At the maximum plateau adsorption, all the hydrocarbon chains are considered to be perpendicular to the solid surface, resulting in reduced coverage and low hydrophobicity, In the case of C8Q>EOlo,hydrophobicity increases with adsorption density and remains constant in the plateau region. This is due to the high density of hydrocarbon chains at the interface. At the plateau adsorption, the number of hydrocarbon chains with Elo four times that of Anionic surfactants do adsorb on silica but only in the presence of the multivalent metal cations. This phenomenon has been utilized in the flotation of quartz wheremetalcations areaddedas activators. Significant adsorption of anionic surfactants on silica can also be induced by the other surfactants that can adsorb on silica. Thus the adsorption of the ethoxylated surfactant, octaethylene glycol mono-n-dodecyl ether (CI2EO8),on silica can induce adsorption of the anionic surfactant sodium ~-octylbenzenesulfonate The anionic surfactant adsorbs on silica either through interaction with the adsorbed ethoxyl chains of the nonionic surfactant and/or through hydrocarbon chain-chain interaction with neighboring CI2EO8.Theadsorptionbehavior of Ci2E08 inthe low concentration region is only slightly altered by the presence of but shows a decrease in the plateau adsorption density (Fig. suggesting that the coadsorption of does not alter the environment in which C12E08adsorbs [9]. The change in zeta potential of the silica due to C12E08adsorption is shown in Fig. 14. The curve closely resembles the adsorption isotherm. This is attributed to the masking effect of the surfactant: the adsorbed surfactant shifts the position of

Adsorption density of octaethylene glycol mono-n-dodecyl ether (C12E08) and ~-octylben~ensulfonate from 1:1 C12E08/C8fpS mixture on silica. (From Ref. 9.)

-10

-40

-50 10-2

Zeta potential of silica particles in octaethylene ether (C12E08)solution. (From Ref. 9.)

glycol mono-n-dodecyl

the shear plane towards the bulk and increases the thickness of the stern layer, and consequently the valueof zeta potential. In contrast, coadsorptionof anionic with CI2EO8is observed to maintain the zeta potential between -40 and -50 mV throughout the entire test region (Fig. 15). The masking effectof the nonionic CI2EO8is offset by the adsorption of anionic

Adsorption of polymers is quite different from that of surfactants because a polymercan adsorb in variousconfigurationsdepending onthesolution and solid properties. Multiple attachment of eachpolymerchainfurthercomplicatesthe conformation of the adsorbed polymers. Also, due to its macromolecular nature, the energy changes associated with adsorption are much larger than those of surfactants. The adsorption isothermof polyethylene oxide (PEO) on silica is shown in Fig. 16 along with the effect on flocculation. The adsorption rises linearly in the beginning and then reaches a plateau. Interestingly a few ppm dosage of the polymer is sufficient to cause complete flocculation of the material [IO]. The effect of solid concentration and suspensionp on adsorption of polyethylene at silica-water interface is illustrated in Fig. 1’7.All points are from the adsorption plateau region. The pH is found to have strongeffect on PE0 adsorption even though is anonionicpolymer. PE0 adsorbson silica throughhydrogen

0

Zeta potential of silica particles in1 molar mixtureof octaethylene glycol mono-n-dodecyl ether (CI2EO8)and ~-octylbenzensulfonate (From Ref. 9.)

too

W/Silica Gel

70

Woter

g

16 Flocculation of silicagel (From Ref. 10.)

as a function of polymerconcentration.

Solidsloading,vol% Effectofsolidconcentration adsorption on silica. (From Ref. l l

and suspension pH on polyethyleneoxide

bonding between the ether oxygen and the surface hydroxyl group of silica. Since pH has a strongeffect on the hydrolysisof silica and decides the hydroxyl groupat the interface, the adsorption density of PE0 is pH dependent [l l]. Solid concentration in the suspension is also found to affect the adsorption of PEO. At high solid concentrations the adsorption decreases. To understand the cause for the decrease, a mixture of pyrene-labeled polymers and unlabeled polymers was used in the adsorption studies to enable in-situ probing of the polymer conformation at the interfaces by fluorescence spectroscopy. The conformation of the polymer is given as the coiling index, defined as the ratio of intensity of the excimer peak (480 nm) to that of themonomerpeak (373 nm) in thepyrene fluorescence spectrum. The results in Fig. 18 show that the coiling index of the adsorbed polymer drops as the solid concentration increases. This suggeststhat the polymer becomes stretched as the solid concentration is increased. Such stretching correlates well with the observed decrease in adsorption density, as the stretched polymer will occupy more area at the interface. The observation has a significant bearing on polymer performanceas most industrial processes are conducted at high solid concentration.

Polymers and surfactants are used together sometimes to obtain desirable effects. Polymer-surfactant interactions in solution and at the interfaces can change the interfacial properties of the solid directly or indirectly. It is shown that depending

0.5

L

E! U

0.2

10

20

30

40

Solids loading, v01 Correlation between the polymer conformation and the polymer adsorption density. (From Ref. 1 1

on the nature of the polymer and the surfactant, polymers can affect flotation of quartz by affecting the adsorption of the surfactant on it Flotation of quartz by cationic surfactant dodecylamine hydrochlorideis shown in Fig. 19. The depressing of flotation at high amine concentrationis attributed to the reversed adsorption of surfactant and possibly the armoring of bubbles by particles and a decrease in the net levitation force. The effect of a nonionic polymer, polyacrylamide (PAM), on quartz flotation by amine at pH 6.4 is illustrated in Fig. At 3.6 kmol/m3amine,PAM, which does not adsorb on quartz and does not cause flotation by itself, has no measurable effect onthe flotation either. At 6 kmol/m3amine, flotation is slightly increased by the polymer addition. Thiseffect is attributed to the uptakeof water molecules by the polymer for hydration rather than the adsorption of polymer on quartz. The hydration of polymer causes an increase in the effective concentration of amine. Anionic sulfonated polyacrylamide (PAMS), a copolymerof acrylamide and acrylamido-2-methylpropanesulfonic acid, is found to increase amine flotation of quartz (Fig. 21). PAMS does not adsorb on thenegatively charged quartz. There is no direct activation of amine adsorption by polymer. However, the polymer-surfactant electrostatic interaction can lead to the formation of complexes. The polymer-surfactant complex can reduce the armoring of bubbles and lead to flotation. Theanionicpolymercan also bridgetheadsorbedamine to the amine on the bubble surface and enhance flotation under saturated adsorption conditions. The

Flotation of quartz as a function of dodecylamine~ydrochlo~ide concentration. (From Ref. 12.)

kmol/m3

A

6.0 x 2

The effect of nonionic polymer PAM (From Ref. 12.)

The effect of anionic polymer PAMS dodecylaminehydrochloride. (From Ref. 12.)

on the flotation of quartz using

on the flotation of quartz using

hydration effect of the polymer may also be responsible for the enhancedflotation in this case. It is to be noted that the extent of salting-out by is significantly higher than that by due to its higher hydration tendency. Flotation obtained in solution containing the amine and a cationic polymer, PAMD,is illustrated in Fig. 22. P A M D is a copolymer of acrylamid and methacrylamido propyltrimethyl ammoniumchloride. P A M D can completely depress the cationic flotation of quartz, possibly due to masking of the surface by the large polymer molecule. Competitive interactions between the polymer and surfactant at the adsorption sites on the quartz surface mayalso contribute to the decrease. The polymer itself does not float the quartz, as can be seen from the figure, since the adsorbed polymer will expose a hydrophilic surface. In the anionic flotation of quartz, activators such as multivalent metal ions are required to provide adsorption sites for surfactant. Cationic polymers can also adsorb on silica and may offer adsorption sites for surfactants. The flotation of quartz by dodecylsulfonate in the presenceof cationic P A M D is shown in Fig. 23. Dodecylsulfonate does not adsorb on quartz and renders no flotation due to the similar charge of the surfactant and the solid. PAMD, on the other hand, adsorbs on quartz but does not cause flotation by itself. The adsorbed P A M D provides adsorption sites for the sulfonate and activates the flotation. In polyl~er-surfactant-solid systems, the sequenceof addition of surfactant and polymer is also found to affect the adsorption of the surfactant and the polymer. In a study of the interaction of an anionic surfxtant, sodium dodecylsulfate (SDS), and a nonionic polymer, polyethylene oxide(PEO), at the silica-water interface, it

(kmoi/m31

The effect of cationic polymer PAMD on the flotation of quartz using dodecylaminehydrochloride. (From Ref. 12.)

kg

0 10 H

0.6

(kmol

The effect of nonionic polymer sodium dodecylsulfonate. (From Ref. 12.)

on the flotation of quartz using

was found that adsorptiondensity of SDS on P E 0 preadsorbed silica is consistently higher than the adsorptionof SDS and P E 0 simultaneously at the same surfactant and polymer concentrationE131 (Fig. 24). The anionic SDS does not adsorb by itself on silica. Interestingly, P E 0 activates SDS adsorption. The mechanism of interactions between SDS and P E 0 is proposed to be association between the ether oxygenof thepolymerandtheheadgroup of thesurfactant. In the case ofsilica preadsorbed with PEO, the SDS adsorption can be considered to be due to the binding of SDS on PEQ at the solid-liquid interface. This PEQ-SDS interaction is stronger at the solid-liquid interface than that in bulk solution. Thus the adsorption density is higher in this case than that from PEO-SDS mixture in solution. The above studies show that dependingon the type of surfactant and polymer, the adsorption of them on silica showseither synergism or antagonism. These results are si~nificant for the ~odificationof surface propertiesof silica by polymer and surfactant, and have important implications to practical systems containing polymers surfactants.

Silica undergoes hydrolysis inaqueoussolution.Therate of the hydrolysis depends on pH, ionicstrength, andprevioustreatmenthistory of thesample. Dissociation of thesilanolgroupsresults in pH-dependentinterfacialcharge. The resultant zeta potential is function of pH and ionic strength. The isoelectric point of silica is about pH 2.

Adsorption of sodium dodecyl sulfate on silica from mixtureof polyethylene oxide and sodium dodecyl sulfate. (From Ref. 13.)

Surfactants and polymers can adsorb on silica through different mechanisms. While cationic surfactant and polymer can electrostatically interact withsilica, the nonionic surfactants and polymer can adsorb on silica through hydrogen bonding. Anionic surfactant or polymer, though not adsorbing on silica by itself due to electrostatic forces, can adsorb on silica surface in the presence of multivalent metalcations,cationic polymers andsurfactantsand even nonionicsurfactants or polymers. The adsorption of surfactants and polymers can drastically alter the interfacial physicochemical properties of the silica, such zetapotential and hydrophobicity. These phenomena can be utilized in industrial processes such as adhesion, flotation, and ~occulation/dispersion. The interactions of different types of surfactants and polymers at the silicawater interface can promote or depress each other depending on their properties such as charge, molecular weight, and even mode of addition. These interactions can be utilized to modify the silica surface to desirable state, or to force adsorption of surfactants onto surface where it normally does not adsorb. Practical implications of these effects are to be noted.

The authors acknowledge financial support of the National Science Foundation (CTS-9622781, CTS-9632479, and ~ ~ C - 9 8 0 4 6 1 8 ) .

an

1. R. D. Kulkarni and P. Somasundaran, in Oxide-~1ectroZyte~ n ~ e r ~ a c(R. es Alwitt, ed.), Am. Electrochem. Soc., Princeton, NJ,1973, p. 31. R. Kulkarni and P. Sornasundaran. Int. Min. Proc. 489 (1977). 3. M. C. Fuerstenau and B. R.Palmer, in FZotation-A. M.~ a ~ d~i en~ o r i uVol. l, I (M. C. Fuerstenau, ed.), AIME, New York, 1976, pp. 148-196. 4. K. P. A~anthapadmanabhanand P. Somasundaran. Colloids Surf, 13:151 (1985). Zhang, P. Somasundaran, and C. Maltesh. J. Colloid Interf. Sci. 191:202 (1997). 6. Fuerstenau, T. W. Healy, and P. Somasundaran. Trans. AIME, 229:321 (I 964). 7. H. El-Shall and P. Somasundaran, in Advances inFine Particks ~ ~ o c ~ s ~ i n g Hanna and Y. A. Attia, eds.), Elsevier, New York, 1990, pp. 41-57. P. Somasundaran, E. D. Snell, and Qun Xu. J. Colloid Interf. Sci. 165 (1991). 9. P.Somasundaran, E. D. Snell, E. Fu, and Qun Xu. Colloids Surf. 63:49 (1992). Somasundaran, andC. Maltesh.Powder 10. E. Koksal, R. Ramachandran,P. Technoi. 62:251 (1990). 11. T.Chen and P. Sornasundaran, submitted. P. Somasundaran and L,. T. Lee. Separat. Sci. Technol. 16:1475 (1981). 13. C. Maltesh and P. Somasundaran. J. Colloid Interf. Sci. 1.53998 (1992).

Institut fur Technische Chemie, Technische Universitat Munchen, Garching bei Munchen, Germany

I. Introduction, General Features Polyelectrolyte Layers

of Adsorbed Polymer and

11. Theoretical Basics Polymer A. and polyelectrolyte adsorption of charged colloidal particles DLVO Theory/Stability C. Steric stability and flocculation of polymer-covered colloids D. Flocculation kinetics

463 464 472 474 475

111. Experimental A. Silica particles of different morphology and surfacechemistry B. Adsorption measurements of polymers and polyelectrolytes C. Stability-flocculation measurements

475 475 476 480

TV. Results and Discussion A. ~dsorption-characterizing parameters B. Suspension-stability-characterizingparameters

483 483 504

References

513

The regulation of the stability of dispersed particles by adsorbed polymer and polyelectrolyte layers is of high scientific and technological significance in all fields of application where surfacesand interfaces are critical for material properties ,2]. Stabilization and flocculation are mainly influenced by thestructure of the adsorbed layer and the chemicaland electrostatic interactions. The segment density profile and especially the extension of the adsorbed layer into the solution play a

decisive role for steric stabilization. The confornlatiol~of neutral macromolecules in the adsorbed stateis determined by the chemical structure, the concentration, the molar mass and the molar mass distributionof the polymer, the surface chemistry, and the properties of the dispersion medium. The conformationof adsorbed polyelectrolytes is additionally dependent on the charge density the polymer, the surface charge, the pH, and the ionic strength of the solution. The investigationof the stability of colloidal systems based011 the adsorption of polymer and polyelectrolyte layers refers to problems of phase separation, coagulation, flocculation, retention, and drainage, as well as to stability of dispersions and foams, corrosion, and adhesion. Destabilization of colloidal suspensions by coagulation and flocculation is, e.g., applied in water purification, refinement of wine, beer, and other beverages, filtration and drainage of sewage and pulp, avoidance and decomposition of aerosols, and concentration of radioactive waste. ~tabi~ization of colloids in dispersion is accomplished in the production of pigments, colors, inks, detergents, pharmaceutical and cosmetic preparations, paper, food agrochemicals(pesticides, insecticides), in intermediates in the manufactureof ceramics, clay, cement, gypsum, magnetic tapes, photographic dispersions, catalysts, gas adsorbers, chromatographic adsorbents and membranes in mineral flotation, drill sludge, in the improvement of soils, and in galenical drug design. The knowledge of the adsorption behavior of neutral and charged ~acromolecules on silica particles and the influence on stability is interest because of the applic~tion aspects of silica as well as of the use of silica particles of different morphology as model colloids to study the adsorption and stability behavior.

In the last few decades many theoretical treatments havebeen made to describe the conformation and the concentrationprofile of adsorbed polymers at the solid surface. E n r i c h ~ ~ e noft the polymer in the adsorbed layer by higher affinity of the polymer segments at the surface, in comparison to the solvent molecules, is named adsorption. In thereverse case, negative adsorption, ordepletion, occurs by impoverishment of the polymer segments in the boundarylayer. If the adsorption energy of the segment high enough, a flat conformation is obtained. In real systems a large number of interacting polymer chains compete for the interface places, and expanded layers with loops and tails are present. Exact theoretical solutions of the distribution of the polymer segments at the surface and in the solution can be made The adsorption of a macromolecule was treated as the Conformation changeof a statistical coil reflecting at the adsorbing surface [3-91. Monte Carlo simulations [10,1l], mean field lattice models 112201, or analogous c o n t i n u u ~considerations [21-251, train-loo~tail models [26291, diffusion equation approaches [21,22,30], and ground state and two eigenfunction approxi~ations[12,14] belong to first group of theories based on conformational statistics. A second group based on the density profile expands the Flory-Huggins theory of homogeneous polymer solutionsto inhomogeneous systems with a concentration

gradient [24,31-331. Information on chain conformationsis lost. The scaling theory of de Gennes [34-371 derives power laws in definite concentration regions. Lattice models [13,24,31-331 and continuum considerations [21-291 are used in both theoretical groups. Thetheories based on conformationalstatistics seem to be more efficient. The Monte Carlo simulations consider eigenvolumeinteractions, and but arerestricted to a small number of chains becauseof insufficient computer power. 111 comparison, the methods basedon the self-consistent mean field consider the total ensemble of chains and give the thermodynamic equilibrium between surface and bulk phase. In the mean field lattice model the walks of the polymers are established in a potential field, produced by the present polymer chains. The field is a function of the concentration profile and also determines the segment concentration. Therefore the potential field is called a self-consistent field (SCF). In most cases numerical iterations are necessary to solve the system of equations. All possible conformations are considered and the structure of the adsorbed layer is characterized by the numbers of trains, loops, and tails and the numbers of segments in trains, loops, and tails. The self-consistent mean field model of Scheutjens and Fleer [18,38,39] can be applied to different systems of polydisperse samples, polyelectrolytes [39-42], and block and statistical copolymers [12,43]. self-consistent field approach exists also for grafted polymer brushes [44] as well as for grafted polyelectrolyte brushes [45] regarding the influences of surface charge, grafting density, and indifferent electrolyte on segmentvolumefractiondistance profile, brush height, and free energy. Also the probability of bridging by adsorbed polymers on two surfaces can be calculated 121. The principal adsorption characteristics of neutralmacromoleculesdeduced from theoretical treatments are demonstratedin Figs. 1-3. The typical high-af~nity adsorption isotherms with increase of affinity by molar mass from theta solvents (x 0.5) and good solvents(x 0) with the adsorption energy parameter 1.0 [x, (U; U?, U$ are the fcee adsorption energiesof the solvent andof the segments respectively] are shown in Fig. 1. At high solution concentration a pseudo plateau is obtained which depends strongly on molar mass.

0-4 +b

Isotherms of polymers of different chain lengthr with x hexagonal lattice. (From Ref. 12.)

xs

0.5(-)

and x

0

Adsorbed amounts versuschainlength r for for different ies x and different adsorption energies xs. Pointsrepresentexperimental system ~S/cy~lohexane/35~C/SiO~( results:system PS/CC14/Si02 Ref. 12.) Because of the reduced segment repulsions bigher amounts 8 are adsorbed from theta solvents than from good solvents. The adsorbed amount 8 verus the chain length r is shown in Fig. forthesolutionconcentration for different solvent qualities x and adsorption energies xs.In good solvents (x 0) is small and nearly independent of chain length for long chains(r 100). In theta solvents (x 0.5) larger adsorbed amounts increasing linearly with log r are postulated. Increasing xs to 3.0 results in an increase of the adsorbed amount. The volume fraction (btr and the fractionof segments in the trains is shown in Fig. 3 for different chain lengths r, versus the adsorbed amount and p are independent of r at low 8. With increasing the fraction of segments in trains becomes smaller. Above 1, (btr increases very slowly, and more and longer loops and tails areformed.Largeradsorptionenergiesandpoorersolventquality (x 0.5) increase (btr and p. The fraction of segments in loops and tails and therefore the average thickness of the adsorbed layer is larger in good solvents. The adsorbed amount is mainly determined by the train andloop segments. The hydrodynamic thickness 8 h of the layer, predominantly influencedby the tails of the adsorbedmacromolecules, is shownversus the adsorbedamountforgood (x 0) and theta (x 0.5) solvents and for different chain lengths r in Fig. The bulk concentration (bb is not constant and changes along the full curves. For

amount

fraction and fraction of segmentsp in trains versus the adsorbed (From Ref. 13.) for different chain lengths r and different x and

Hydrodynamic layer thicknessSh versus the adsorbed amount for different chain lengths r and two solvencies x. (From Ref. 12.)

small unique increase of &h with is postulated for all chain lengths. At constant (S)b (dashed curve) clear dependence of 8 h on molar mass is found at large adsorbedamounts.At low adsorbedamountsflat layers exist. At high adsorbed amounts the first surface layer becomes saturated, tails are developed and the hydrodynamic thickness &h is strongly increased. Especially interesting adsorption behavioris demonstrated by block copolymers. Usually one block is preferentially adsorbed an anchor. The other block, with lower adsorption affinity, is excluded from the surface and forms dangling tails into the solution. brushlike structure is formed with solvated shell which provides repulsive stabilizing barrier. The detailed conformation is determined by the solvent, which can be selective or nonselective for the two polymer parts and by the lengths of the blocks [12]. The main energetic and entropic factors determining the adsorption behavior of neutral polymers and polyelectrolytes are the segment adsorption energy, the loss of chain conformation entropy, the entropyof mixing and the mutual interactions of segments, especially in the case of polyelectrolytes [18]. With the useof the ~cheutjens-Fleer model the dependency of the adsorption on adsorption energy, solvent quality, salt concentration, and surface charge has been studied. Using numerical calculations on the basis of the self-consistent field theory the adsorption of polyelectrolytes on oppositely chargedsurfaces has been studied and shows subtle balance of forces influenced by segment charge,surfacecharge density, salt concentration, and specific chemical surface segment interaction. In the caseof pureelectrostaticadsorption at moderate segment charge,rather extended conformations are predicted with decreasing amounts at increasing salt concentration. In this screening-reduced regime criticalsaltconcentrationfor total desorption exists. A transition to screening-enhanced regime occurs if additionalnonelectrostaticinteractionstake place leading to ma~imum in the adsorbed amount, depending on electrolyte concentration. Isolated chains of partially ionized polyelectrolytes on planar charged surfaces were examined by Monte Carlo study without nonelectrostatic adsorption energy Train-loo~tail distributions deviate depending on charge and ionization of the chain, surface charge density, ionic strength, and solution and surface properties. The characteristicsof ion, polyelectrolyte, and protein. adsorption deduced from theoretical models are summarized in Ref. 4’7. The principal characteristics of the adsorption of polyelectrolytes of various charge densities surfaces of differentcharge at low and high ionic strength obtained from theoretical treatments are presented structure models in Fig. 5. The adsorption of neutral macromolecules mainly determined by the segmental interaction parametersof adsorption and of solvency x. For polyelectrolytes, additional electrostatic interaction considered by an additional term dependent on the chargedensities of surface and polyelectrolyte and onthe ionic strength K A1 a ionization degree, K Debye-Huckel parameter). The effective energy parameter becomes

~onfornlationmodels of adsorbed polyelectrolytes: (a, c) macroions of low charge density; (b,d) macroions of high charge density. In Fig. 6 the adsorbed amount6exis plotted as a functionof log c, (c, electrolyte concentration). If the surface is uncharged the adsorbed amount is small (curve 1). The mutual segmental repulsion prevents a strong coverage of the surface at low ionic strength. If the macroion and the surface have the same charge no adsorption takes place (curve2). At low c, strong repulsion results in depletion (dotted curve). At different charge signs charge compensation occurs (curves 3-5). In the case of pure charge compens~tionthe adsorbed amounts are independent of molar mass. The macroion becomes desorbed at high ionic strength becauseof ion exchange

Adsorbed amounts of polyelectrolytesversusthesaltconcentrationfor various combinations of segmental and surface charge and x,. (From Ref. 12.)

the small counterions. At high electrolyte concentration the electrostatic interactions are screened and additional chemicalinteractions and x become dominant. Adsorption does not exist if is smaller than the critical energy parameter Adsorption occursif exceeds the critical value (curves 1,2,4,5), the adsorbed amount approaches a limiting value with increasing electrolyte concentration in a good solvent, whereas phase separation takes placein a poor solvent (curve4”). In the case of specific adsorbing counterions the polymer segments are displaced from the surface and the adsorbed amount shows a m a ~ i ~ uthat m depends on the salt concentration (curves 2’, 4’). A more detailed description is given in Ref. 12. The segment densities versus the distance from the surface of adsorbed polyelectrolytes withsegmentnumber 2000 at bulkconcentration rjhb onan uncharged surface were calculated for different salt concentrations using the Roe theory Figure 7 shows the profiles for the interaction parameters 2, x 0.5 in comparison to the uncharged polymer. At low ionic strength a minimum results, produced by the strong repulsive segment interactions compared to the monotonic decreaseof the segment densityof the uncharged polymer. The macroionsareadsorbed in very flat conformationwithsmalladsorbedamount. Reduction of repulsion at high ionic strength eliminates the minimum. In the case of pure electrosorption without chemical affinity [42] the adsorbed amount decreases with increasing ionic strength because of competition by the coions. A maximum of the adsorbed amount in dependence of charge density is produced which decreases andis directed to higher values with increasingionic 0.6 changes the depenstrength (Fig, 8). Specific chemical adsorption with

Semilogarithmic volume fraction profilesof a strongpolyelectrolyte at various salt c ~ ~ c e n t r a t i or~ s ;2000, l x 0.5, 0.71 nm,hexagonal lattice. (From Ref. 12.)

8

0.5

n 0

trations; 41

0.04

0.08

Adsorbed amount Qex versus the charge density 1 r 100, x 0.5, 0.6 nm, U

at different salt concen0.01 C/m2. (From Ref.

dences on ionic strength and charge densitydistinctly. The increaseof the adsorbed amount with ionic strength up to a charge density 0.02 is attributed to the screening of the segments and a transition from flat to more coiled conformation. Polymers of very lowcharge density show a decrease of the amountwith c, because of the exchange of the macroions by the coions. The adsorbed amount of weak polyelectrolytes (Fig. 9), the length of trains and loops (Fig. lo), as well as the pH because of increasing ionizafraction p in the trains are strongly dependent on tion of the macroion segments with pH [40].

4

5

7

8

PH

Adsorbed amounts of a poorpolyelectrolyte versuspH atvarious surface charge densities; pK 5. (From Ref. 40.)

5

4

7

8

PH

Length of trains loops and tails of a poor polyelectrolyte versus pH at various surface charge densities;pK 5. (From Ref. 40.)

The validity of scaling and mean field models is discussed for adsorbing and nonadsorbing polymers [14]. The influence on colloidal stability by bridging, steric stabilization, and depletion flocculation is treated. The mean field theory is more appropriate than scaling to describe the main experimental results.

The stability of colloidal dispersions determined by the interactions of particles during a collision. Attraction is caused by van der Waals forces. Repulsion is produced by the interaction the charged double layers around the particles. Additional the particle solvent affinity and the increasing solution viscosity near thesurfaceplaysa role in the stability of suspensions.Derjaguin and Landau [49,50] as well as Verwey and Overbeek E511 developed independently from each otheraquantitativetheory(DLVO), which describestheaggregation of dispersed particles as consequence of the energy profile at approach of two particles. The attraction energyby London-van der Waals forces and the repulsion energyby the overlap of the electrical double layers were calculated as a function of the particle distance. Theenergydistance profiles are shownschematically in Fig. 11. The repulsive energy is an exponential function of the distance dependenton the thickness of the electrical double layer. The attraction energy inversely proportional to thedistance. The totalinteraction potential for twospherical particles of radius if given by

Interaction energies between two particles as a function of the distance.

where attraction potential, repulsion potential, Y surface potential, K Debye-Huckel screening parameter, H distance between the particle surfaces, A Hamaker constant (specifically dependent onthedispersedmaterial and the medium). The total interaction energy (Fig. 11) demonstrates primary minimum at very small distances. Under certain conditions potential maximum, Vmax,and also a secondary minimumexist. The height of the energy barrier, Vmax,is responsible for the stability of the suspension. Addition of electrolyte decreases the energy maximum. The critical coagulation concentration (CCC) is defined by the condition Vmax 0. Aggregation can also be induced by external energies as by shearing. In the case of a secondary minimum a reversible, loose metastable aggregation could occur, especially with large particles. Additionally to the DLVO energy contributions the hydrationenergy expressed a double-layer exponential function is considered in an expanded theory [52]. Because silica particles of small size, in deviation from the DLVO theory, do not coagulate when theelectrophoreticmobility becomes zero at acidic p increasing electrolyte concentration. Repulsive hydrationforces are especially suggested an inherent property of silica. This theoryis suited to describe thestability of silica suspensions over wide range of pH and electrolyte concentration. Also the ordering of liquids on the surface of colloidal particles plays an important role in stabilization [53].

Addition of polymer to asuspension changes the physical properties of the medium as well as of the particle surface by adsorption. The thermodynamic properties are influenced by the particle interactions with the medium and by the interparticle forces. They determine the equilibrium structureof the suspension and the osmotic pressure. Together with the hydrodynamic interactions they control the dynamic processes in the system as diffusion, sedimentation, electrophoresis, and aggregation kinetics. Steric and electrosteric stabilization of the suspension is imparted by polymers and polyelectrolytes adsorbed on the colloidal particles. Amphiphatic block and graft copolymers are special steric stabilizers because the insoluble anchor polymer is attached to the surface, whereas the solublepart projects into the solution as the stabilizing moiety [54,55]. Osmotic and mixing effects as well as enthalpic and entropic contri~utionsare involved in the stabilizing mechanisms [55-591. Depletion stabilization can be produced by free polymer in solution [59-631 caused by the unfavorable energy ~ontributionwhich is necessary to deplete the polymer coils from the gap between the colliding particles. On the other hand,flocculation can be produced by addition of polymers to the colloidal suspensions. Possible reasons for the aggregation of the particles are: Interpenetration of the adsorbed layers under worse than 0 conditions, according to the ~ l o r y - ~ u g g i ntheory, s where a correlation between the critical flocculation point of the suspension and the point of the solvent is observed [S91 ridging flocculation by the adsorption of macromolecules on two particles under low coverage conditions 164-701 An excess of free polymer in solution which reduces the potential barrier by changing the attraction and the electrostatic repulsion term [71,72] Depletion flocculation [73] due to theenergetic unfavored depletion of the polymer chains from the space between the particles which results in a reduction of the space volume by a closer approach. In thecase of adsorbed polyelectrolytes, surfacecharge ne~~tralizatiolland reduction of the electrostatic repulsion produces flocculation [74]. At small coverage the adsorbed polyelectrolyte molecules can form patchesof positive and negative charge on the surface. Coulombic attraction of those mosaic structures on different particles can also lead to flocculation during Brownian collision [75]. The Scheutjens-Fleer mean field model of adsorption has been applied to a systemof two parallel flat surfaces in a distance of a number of lattice layers [12]. With this theory the free enthalpies of interacting polymer layers and their additional contribution to the total interaction energy could be calculated. The interaction betweenparticles with adsorbed ~olyelectrolytes bas been examined by Monte Carlo calculations [76,77]. Attractive and repulsive forces by electrostatics and bridging as function of separation, electrolyte, polyelectrolyte structure, and colloid charge are reported.

The theory of Smoluchowski [78] for rapid aggregationis based on thediffusion of the particles. The aggregation rate corresponds to the number of collisions and the number of aggregated particles with the aggregation numberp as a function of time can be formulated by

with t' t / t , l/(kno) half-life, k 4kBT/(3q) second-order rate constant, concentration of spherical particles at thebeginning (particles/m3), and k 6.18 10-1s[m3/(s.particle)] in water of viscosity at 25°C. This theoretical model is restricted to spherical particles without any repulsive interaction. With progressive aggregationlarger particles deviating from the spherical shapeareformedwithincreasingdistribution in size and shape. So the Smoluchowski theory is only valid for the beginning of the aggregation process. In some published results the flocculation rates reach values appreciably lower than the theoretical values. Hydrodynamic and steric interactions, viscous effects [79-821, repulsive potentials [83,84], and the reversibility of aggregation (84-861 are supposed to be the reasons for thelower measured aggregation ratesin comparison to the perikinetic process. The theoretical treatments [79-861 of slow aggregation takeinto account that the repulsing electrostatic interaction is not completely screened. The potential-distance curve has a barrier, which prevents successful aggregation during thecollision process. A factor called stability ratio W depending on the potential barrier is introduced, which considers the fraction of successful collisions:

k where

with

W

corresponds approximately to the energy barrier.

l. Preparation and Pretreat~entof Colloidal Silica, Preparation of Aqueous Suspensions Amorphous, microporous silica particles of relatively low polydispersity can be precipitated by hydrolysis of Si(OC2H5)4in a mixture of water and ethanol by catalysis withammonia [87]. The Si02 particles arewashed, centrifuged, and dried. Redispersion of the dry pulver in water is accomplished by ultrasonication [88-941. Pyrogenic silica produced by hydrolysis of silicon halides in a hydrogenoxygen flame consists of aggregates of primary particles of different size and size distribution [95-971. Weak acidic SiOH groups and siloxane groups are present on the surface of the particles. Their surface concentration can be altered by thermic and acid or basic treatment [98-1011.

Chemical modificationsof the silica surface canbe accomplished by reactions of groups 11021, e.g., by: t acid-base reactions with amines [l021 with alcohols or diazomethanes 103,1041 h (CH3),SiC14-, 1051 in gaseous form or in chlorinated hydrocareaction of alcyl trimethoxysilanes [l061 Chlorination of the SiOW groups by thionylchloride and Friedel-Crafts alkylationor reactionwithlithiumorganic orGrignardcompounds 107,1681. Precursors for radical polymerization can be obtained in this way. Preparationproceduresforaqueoussuspensionarereported [90,93-95].

in manypapers

haracteri~ationof Particles, Particle urface chemist^, Surface Charge, The particle size, the size distribution, and the second moment of the distribution respectively can be measured directly by electron microscopy[88,91,92-94,108,1091 and indirectly determined by photon correlation spectrometry (PCS) and ultracentrifugation (IJC) [88,91,93-95,101,110], and by small-angle x-ray scattering [l 1l]. The specific surfacearea can be determined by N2 adsorptionwiththe Brunauer-Emmet-Teller (BET) method [88,93]. The surface charge density dependence on pH can be obtained by p [l 12-1 141 and by titration with apolyelectrolyte of opposite charge(polyelectrolyte titration) 113,llS, 1161. The zeta potential at the shear plan of the moving particles can be measured by electrophoresis, electro-osmosis, streaming potential and sedimentation potential 117,1181, Laser Doppler anemometry measures the electrophoretic mobility with high accuracy 113,1191. An additional methodis the electrokinetic sonic amplitude potential 1971. calculate the zeta-potential, retardation and relaxation effects of the electric double layer have to be considered for particles of different size and shape.Additionally,surfaceconductivity and polarization of the electrostatic double layer should be taken into account 120,1211. Rheological properties of the suspensions are obtained by viscosimetric measurements in a rotational device [loll.

A survey on the experimental methodsused in polymer adsorption is given in Ref. 122 and in particular, results on silica particles are presented as examples. dsor~tionIsotherms by Direct and Indirect In special cases the adsorbed amount can be measured directly on the surface by attenuated total reflection (ATR) on theflat surface of a transparent solid medium [123], ellipsometry on a macroscopic flat surface [124-1261, surface plasmon resonance [127], surface enhanced Raman scattering [128], and quartz crystal micro-

balance 1291. Removal of the excess polymer in solution can be achieved by centrifugation or filtration and washing with puresolvent. It has to be ensured that no desorption takes place in this procedure. Extreme conditions are required in the purity of the solutions because of the preferential adsorption of contal~inants on the small macroscopic surface area. The usual procedureto determine the adsorption isothermsis the indirect determination of the adsorbed amount on colloidal particles from the concentration difference in the solution before and after the adsorption process. Separation of the dispersed coveredparticles from the solutionis accomplished by sedimentation or filtration. The adsorbed amount A (g/mz) is calculated by A Ac( VIA,), with V (cm3) solution volume, A, (m2/g) surface area, and (g/cm3) change in solution concentration. Sources of error are small Ac and flocculation. The concentration in solution is measured by titration of the total organic content [130], by spectroscopic methods, fluorescence, IR [98,99,129,131-1 331, FTIR, NMR, refractometry 1341, radiometry 1351, and by turbidity [92,136]. Radio-, spin-, or chemical labeling of the polymer to improve the sensitivity of the analysis has been carried out. Changes of the adsorption behavior of the labeled substances have to be taken into account. In the case of polyelectrolytes, polyelectrolyte titration is very sensitive method to determinethesolutionconcentration [47,114,115]. Complexation of PVAwith boric acidand potassium iodide 1371 and of P E 0 with tannic acid1361 are special methods for specific polymers. Desorption measurements by displacers have been undertaken in some cases.

Bound The fraction of bound polymer segments is an important parameter deciding the strength of the binding, the structure of the adsorbed layer, and the desorption of themacromolecule.Thespectroscopicpropertieschange if asegment becomes adsorbed in comparison to the properties of the free segment. The techniques for measuring bound fractions on dispersed particles and on flat surfaces are reported. (IR). Becauseof hydrogenbonding or dipole-dipole interactions between polymer and surface groups the frequency and intensity of absorption bands of the polymer segments well of the surface groups change. By calibration with standard materials the fractionof segments involvedin adsorption can be determined quantitatively if the resolution of the bound andfree bands sufficient. After its introduction13l], the method has been used bymany authors on colloidal suspensionswith different polymers in organic [99,100,110,138143,2741and in aqueous media 3381. Double-beam Fourier transform IR methods in reflection mode 1381 and the FTIR-ATR (attenuated totalreflection) method on flat surfaces [144,145] have been used. Simultaneous determinations of adsorbed amounts and specific segment surface interactionsare made by a special technique in a vacuum cell with transmission FTIR 11467. NMR can beused intwo different ways to decide between segments of an adsorbed macromolecule locatedin bridges at the surface, loops and tails and of soluted macromolecules. The change in the resonance frequency, chemicalshift 147-1493, well as the lifetimes of the nuclear

states, spin-spin and spin-lattice relaxation 150,15l] can be used as characterizing quantities, differentiating between rigid adsorbed and mobile segments in loops and tails and in solution.Thismayleadto an overestimation of boundsegments because also neighboring segments are reduced in mobility. On the other hand, rapid exchangeof segments in trains and segments in loops and tails can apparently increase the measured mobile segments. c.ElectronSpin (ESR). Inthesame way asNMR,the reduced mobility of adsorbed segments leads to a broadening of the ESR lines of spinlabeled polymers 140,149~ 52-1541. In order to get reliable results the labels are not allowed to change the adsorption behavior of the corresponding polymer. d. ~ i c r o c a l o r i ~ ~ t r yBecause . of theenthalpiccontributionaccompanyingthe surface binding of a polymer segment, microcalorimetry is principally suited to measurement of the number of adhered segments. Enthalpies of wetting 11.551 as well as enthalpies of adsorption [99,100,152,156-159,2743 replacing solvent by segments, are used for this purpose. ~ ~ e c t r o c h e ~ i Titration. cal The composition of theStern layer located immediately adjacent to the particle surface is changed by the train segments of the adsorbed macromolecules and therefore the electric properties of this layer and consequently of the electrical double layer are modified. In electrochemical studies the capacitance, defined as the ratio between surface charge density and surface potential and obtained from titration curves, can beused to measure the train density of the adsorbed polymer layer on silver iodide 1601 and on oxide surfaces [161].

3. Thickness of the Adsorbed Layer The segment profile and the adsorbed layer thickness are determining factors for the influence of adsorbed polymer layers on flocculation and stabilization of colloidal suspensions. Static measurementsby ellipsometry and scattering and hydrodynamicmeasurements by diffusion, sedimentation, viscosity, flow through capillaries, and disjoining pressure measurements result in different moments of the segment density profile and corresponding different thickness averages. a. Static eth hods by Scattering, R e ~ e c t o ~ e ~ r y , ~ l l i p s o ~ e t ~ yThicknesses . of adsorbed layers can be obtained by neutronandx-ray scattering. Thishas especiallybeen possiblefor silica particles coveredwithdeuteratedpolystyrene layers in a brominated solvent to optimize contrast matching [12]. Reflectometry and ellipsometrymeasurethe light intensity and polarization changes by the adsorbed layer on a flat surfaceusingfor their evaluationthe Fresnel equations 162,163,164). b. ~ y d r o d y n a ~ i c ~ h i c k nby e s ~s ~ ~ ~ ( Ps CiS )o, Sne d i ~ e n t a t i o n U C j , Viscosity, andElectrokinetic eth hods Potentialj. By adsorbedpolymer layers the flow rate in capillaries is reduced and the friction coefficients of colloidal particles are increased by the increaseof the hydrodynamiceffective diameter The latter results in a decreaseof the diffusion coefficient [88,89,164-1681 and of the sedimentation coefficient of the suspended particles [91,131,1641 and in an increase of the intrinsic viscosity of the suspension 165,169,1701.

o~ye~e~trolyte

By quasielastic light scattering (QELS) orphotoncorrelationspectrometry (PCS) of scattered light using monochromatic laser the intensity fluctuations produced by diffusion are measured by Doppler broadening or by the intensity correlation function. The diffusion coefficient can be obtained from the exponential decay of the correlation functionand related to the effective radius of the polymercovered spherical particles by the Stokes-Einstein equation. The difference from the radius of the bare particle corresponds to the hydrodynamic thickness of the adsorbed polymerlayer. Conditions for accurateresults are stable, not flocculating suspensions, low particle concentrations to avoid electrostatic, steric, and hydrodynamic interactions, and narrow particle distributions [88,91]. In the case of polyelectrolyte layers electrostatic interactions have tobe avoided, and suitable systems with salt addition but no aggregation have tobe selected [l 15,1711. Sedimentation in an analytical ultracentrifuge leads to the same hydrodynamiceffective thickness. In special cases also electrokinetic measurements can be used to get the layer thickness. The adsorbed layer disturbs theelectrostatic double layer and the flow of charge along the surface163,172-1 741. The electrokinetic thickness is the shift of the shear plane needed to give the same current with 6, In double-layer thickness). If the current i is compensated by conductive countercurrent, then 6, In Vs/Vso)(V, streaming potential). The electrokinetic slipping plane dependson the extensionof the doublelayer and, hence, on theionic strength. So the electrokinetic and hydrodynamic thicknesses are not in agreement if the double-layer thicknessis smaller than the hydrodynamic thickness. A strong dependence of 6,on ionic strength exists, which disappears in the hydrodynamic limit if then 6tjfh. This treatment is only possible if no distortion of the double layer is caused by the adsorbed polymer segments. This may be plausible for the low layer concentrations of loops and tails but not for the large segment concentrations in the trains. When counterions are displacedby train segments an increase of the electrokinetic potential is produced 1751. Surface Charge of Covered Particles Potentiomet~i~ pH The surface charge density of the bare and polymer-covered silica particles can be determined by titration with thepotential-determiningions H' or OH" in inert N2 atmospherewithpotentiometricand conductometric measurements. Two titrations of the silica suspensions and of the bare solutions are performed. The titrated amountsof H' or OH- ions at different electrolyte concentrationsaresubtracted at thesame pH [120]. This difference divided by the applied surface areaof the silica particles corresponds to the surface charge density. In the case of the polymer-covered particles reference solution with the same polymer and electrolyte concentrations is measured for comparison [171].

4.

5. Volume Fraction Profile By Small-An~leN e ~ t r oS ~ a t t e r i n ~ Small-angle scattering techniques, using neutrons, are applied to determine the segment densityprofile and the thickness of an adsorbed polymer layer in particular suspensions. In these SANS studies the scattering from the adsorbed layer can be detected by contrasting the

scattering of the particles with a suitable H20/D20ratio in the solvent medium (contrast matching) [176,177] and in organic solvents, e.g., CC14 [148]. Neutron reflectometry has beenused also on flat macroscopic surfaces ([12], chapter 3.5). Small-angle neutron scattering has also been applied to examine the organization of silica spheres (370 nm) in particles flocculated by cationic copolymers of acrylamide and N‘, N”-tril~ethyla~noethylacrylate 1781. Highly charged polymers flocculate by charge compensation in flocs produced by random collisions, whereas lightly charged polymersflocculate without charge compensationby interparticle bridging and flocs of liquidlike order are obtained. By NMR spin-spin measure~ents areprincipally suited to determine the relaxation decay of the segments in the adsorbed polymer layer 1791. The different concentration regimes in the layer can be described as a sum of exponential relaxation decays from very short of the segments directly fixed at the surface to longer T2 of the segments in the tails. Withhigh-resolution solid-state NMR spectroscopy,structural and dynamic studies of adsorbed ”C-labeled polymers are possible [130]. c. ~Zuurii~etry.By fluorescent labeling of thepolymers the segment density profile can be determined by evanescent wave-induced fluorimetry 1801.

tly imaged [114,181]. of the silica suspension of very small particle concentration (0.01 g/L) is placed on a filter (Anodisk, Whatman) of 100 nm pore size. Then the solvent is sucked off by a subpressure of 800 mbar and washed twice with pure water. The small poresize and the narrow poresize distribution of the filter prevent high suckoff streams of the solvent. After drying, thefilters are pasted with silver solder and connected conductively on the object support. Thefilters are then evaporated with gold and imaged in theREM.The precipitated colloid particles are magni~catedby a factor of 10,000. The aggregation can also be imaged by transmission electron microscopy after fixing the dispersion state, e.g., by cryofixation /182]. An additional possibility is the counting of the aggregated particles in an ultramicroscopeafterstoppingtheaggregationprocess 1831. Alsoultramicroscopy under flow conditions measuring thesize of single particles by laser light scattering has been applied 184.1.

iffusion Coefficient (PCS) Photon correlationspectrometry (PCS) or quasielastic light scattering (QELS) measures the average diffusion coefficientof particles between 3 nm and 3 p m diameter and can be used to measure the coagulation and flocculation behavior of silica particles under different conditions [94,185-1911. After rapid mixing of particle suspension, polymer, and electrolyte solution the diffusion coefficients are

measured by PCS as a function of time. Particle concentrations in the rangeof 0.1 to 10 g/L are suitable to get enough scattering intensity and to avoid multiple scattering. The adding sequence can be changed in order toinvestigate the influence of mixing [94,191]. The time-dependent diffusion coefficients can be treated with the second-order kinetics of aggregation of Smolucho~ski[78,94]. special evaluation procedure has been worked out [l861. Particle ~ o u n t i n g Using an electrical pulse counter (Coulter counter) the number of particles in a known volume are counted during the passage of the suspension through a small orifice. With this passage the current flow and the resistance between two electrodes on opposite sides changes and the corresponding voltage is amplified and registered. The llumber of particles is measured by the pulse frequency and the particle volume by the pulse height. By analyzing the pulse height the size distribution can be obtained. Calibration with suspensions of known particle size is required. This method isuseful in the upper colloidal size range 921. In size measurements of flocculated particles [l931 the stability of the flocs under shearing in the orifice and little current flow through the liquid inside the flocs has to be guaranteed. Light ~ c a t t e r i n ~ Turbidity , According to the scattering theories of Rayleigh and Mie the scattered light of colloidal suspensions shows strong dependence on particle size, mass concentration, scattering angle and wavelength of light. At constant concentration andscattering angle the intensity will increase with particle size. This property the basis for the scattering and turbidity measurements used extensively in aggregation studies [194]. In particular, the critical aggregation concentrations are determined by turbidity. Aspecial optical particle analyzer hasbeen developed to measure thefirst steps of flocculation [195]. Thedependence on wavelength isused to measure sensitively the increase of the particle diameters [196,197]. Perikinetic and orthokinetic conditions are appliedby different apparative devices 198,1991. Laser optical devices and pulse height analyzers are applied 1991. Also polymer-induced phase separation is studied by turbidity analysis [200]. Flocculation is also measured by the simple jar test with controlled agitation where the clarification of the supernatant is observed by light scattering [201].

5. ~ e d i m ~ ~ t a t Volume, ion ~edimentationRate The simplest method is the observation of the turbidity and sedimentation of the formatted flocs in a test tube or cuvette by eye. Different mechanisms of flocculation result in varioustypes of sediments.Tightandloosesediments, gels and orderedstructurescan be formed by sedimentation,Fromthe ti~e-dependent change in the concentration height profile, the sedimentation rate, the profile of the boundary layer and from the final sediment volume characteristic i n f o r ~ a t i o n about the flocculation mechanism and the type of the aggregated dispersion canbe obtained [92,93,101,202,203].

The estimationof the particle size by sedimentation rate is a widely used indirect method. Gravitational settling is only applicable forvery dense particles in the high colloidal size range becauseof disturbing diffusion and convection influences in the case of small particles under low density differences compared to the medium. The concentration profile produced by sedimentation can be obtained by direct sampling, by measuring the absorption of light and x-rays (e.g., laser-photosedimentometer) or thehydrostaticpressureat definite heights in acuvette 1202,2031. Solvated particles of high lyophobility and flocculated particles with entrapped fluid show special characteristics of a composite. Subjecting the particles to centrifugation the Sedimentation force and velocity becomes increased by the increased centrifugal acceleration, With the Stokes law the radius of spherical particles can be calculated from sedimentation coefficients S obtained by sedimenation velocity experiments:

dxldt with

S

or

In-

X1

SW

t,)

Refraction of visible light, Schlieren technique using the refractive index gradient and absorption are the usual techniques to follow the Concentration profile of the sedimentingparticles. This Sedimentation procedureand evaluation is restricted to low particle concentration. In concentrated suspensions hydrodynamic interactions with neighboringparticles influence the sedimentation. Extrapolationsto zero concentration or approximation formulas have tobe applied. Sedimentation equilibrium experiments are not used in colloidal particle size determination. With the development of disk centrifuges in recent years an extension to small particle diameters of 10 nm, and high resolutions in the size distributioll could be achieved by measuring the turbidity of light or x-rays at definite positions of a transparent disk in a rotor [204]. An interesting new device is also the field flow fractionation method, which combines the centrifugal field with a flow in perpendicular direction [204].

Filtration Cuvette viscometers with two coaxialcylinders, Ostwald capillary viscometers, and cone-plate viscometers are used for measuring the complicated flow behavior and the viscosityof colloidal suspensionscovered by adsorbedpolymer layers. Influences of particle concentration, size and size distribution, shape, surface properties, and solvation have to be taken into consideration. Results on silica suspensions withneutralpolyethyleneoxide and cationic polyelectrolyte layers are obtained [95,101,205]. Laminar tubeflow isa meansof applying shearto a suspensionin order tostudy orthokinetic flocculation 11991. The filtrability of the filtrate is also used to investigate the structure of flocculated suspensions 12061. The stability of pyrogenic nonporous silica with adsorbed layers of PMMA and ~ S / P Mcopolymers M~ was studied by oscillatory rheological techniques in order to quantify theelastic moduli and hence the strengthof the flocculated network as a function of surface coverage, molar mass, architecture, and solvency 12071.

er

lectrical P r o ~ e ~ i e s Electro-osmosis, streaming potential and electrophoresis are used to measure the electrokinetic behavior at interfaces. Electrophoresis is especially used to measure the motion of colloidal particles in suspension under the influence of an electric field. A special sensitive and accurate method is the laser Doppler electrophoresis [l 16,199,208,2091. The zeta potential of the particles can be evaluated from the measured electrophoretic mobility pE by the formula 1541

The function considers the retardation and the relaxation of the electrical double layer in the electric field [210]. An analytical formula has been derived for zeta potentials smaller than 40 mV [211] and is tabulated for dielectric and conductive particles of different size. Theoretical considerations and approximative forn1ulas for special cases are available 1212-2141. Theories for the zeta potential of particles with adsorbed polymer 172-1741 and polyelectrolyte [21l] layers also exist. ~ o r c e - ~ i s t a n c e~ e a s u r e ~ e n t s With force-distance measurements, interactions between two adsorbed polymer layers on macroscopic silica or mica surfaces formed as crossed cylindrical filaments can be investigated [215-2191. Force-distance profiles have been obtained in a wide range of solvency conditions for several polymers, block copolymers, and polyelectrolytes. Attractive and respulsive situationscan be demonstrated.The influence of polymer coverage, molar mass, rate of approach, and distance, and especially the effectof polymerbridging at low coveragehas been studied. Comparisonsbetwenexperimental results and theoretical structuremodelsare made. This is a possibility to get direct insight into the contributing energies and into the mechanisms of stability and flocculation of dispersed particles covered by polymers.

A ~ s o r ~ t ~Isotherms on of NeutralPolymersandCharged Polymers In earlier investigations the adsorbed amountwas the unique measured quantityto characterize the adsorptionlayers of polymers. From the adsorbed amount and the correspondingadsorptionisotherms,basicinformation ontheadsorption is obtained.Theothermeasuredcharacterizingparametershaveto be related to the area specific amount adsorbed and discussed in relation to it. These relations result in additio~alinformation about the structureof the layer, the conformations of the adsorbed macromoleculesand the segment distributions in trains, loops, and tails.

1

In most studies the adsorption isotherms are obtained indirectly by measuring theconcentration of thesupernatantsolutionafterseparation of the colloidal particles from the surfacelayer by filtration or centrifugation. In rarely used direct methods the amount of polymer in contact with the surface is measured. Theadsorptionisotherms of polyethyleneoxide (PEO) on precipitated and pyrogenic silica in water show high affinity [12,90,220,2741. The plateau amounts increase with the molar massof P E 0 on nonporous precipitated silica (Fig. 12). comparison of the molar mass dependenceof the measured adsorbed amount with the theoretical dependence is shown in Ref. 12 (Fig. 2, Fig. 13). Slightly lower plateau values (g/mz) are obtained on pyrogenic silica under the same conditions [90,163,203]. The adsorbed amounts decrease with molar massin contrast to the normal increase [221] because of the lessaccessible parts in the porous structure. With increasing molar mass up to 5000, low molar mass P E 0 shows also an increase of theadsorption affinity onmacroporous silica demonstrated by the rising initial slope of the isotherms. The adsorbed amountsof polyvinylalcohol (PVA) and of P E 0 decrease withpH [88,89,222,223] because of the increasing dissociation of the SiOH surface groups (Fig. 14). The OH- groups act as displacers for the polymer segments. ~imethylationof the OH end groups of PE0 400 and P E 0 5000 leads to a decrease in the affinity and an increase of the plateau value because of the high affinity of the hydroxyl end groups in comparison to the methylated groups(1571.

Dependence of the adsorbed plateau amounts of polyethylene oxide on molar mass on unmodified and modified precipitated silica and on latex; Csilica 5 g/mL, g/mL.

l'

l000

Dependence of the adsorbed plateau amounts I'of polystyrene, polytetrahydrofurane, and polyethylene oxide on molar mass M ,silica, (a) for solvents, for good solvents. (From Ref. 12.)

On precipitated silica modified by octyl groups, larger adsorbed amounts and a stronger dependence on molar mass is observed [l681 (Fig. 12). This can be attributed to stimulated loop formation on the places where hydrophobic octyl groups exist. The results for latex support this explanation. The adsorption isotherms of statistic and block copolymers of ethylene oxide and propylene oxide from CC& onpyrogenic silica were measured as a fraction of the ethylene oxide fraction in combination with IR ~easurementsof the adhered segments on the same systems [274] (Fig. 15). The results are discussed in Section IV.A.2. ~oly(viny1pyrro~idone) has been adsorbed pyrogenic silica as a function of time, molar mass, and mass distribution from water and dioxane solutions [l55]. The special polydispersity effect could be explained theoretically [20],

I

5

0

Displacement measured by adsorbed amount I', hydrodynamic thickness and nuclear spin relaxation of the solvent rt,.with OH- as displacer; system H2O/SiO2. (From Ref. 220.)

0

40

50

80

100

mole fraction

Dependence of the adsorbed plateau amounts of EO-PO-EO triblock copolymers on the EO mole fraction; pyrogenic silica, CC14.

The kinetics of adsorption of polymers was also investigated in some papers 140,141,2241. The equilibrium is obtained in minutes orin hours depending on the system. Adsorption measurements with poly(viny1pyridine) have also been carried out [225]. olystyrenes of molar mass ranging from 600 to 2,000,000 with narrow mass distribution were adsorbed in theta conditions from cyclohexane [226]. The adsorption isotherms of polystyrene (PS) and are of the high-affinity type, the adsorbed plateau amount increases with molar mass [log, 139,2251. The different adsorption affinity of isotactic and syndiotactic PMMA is used to separate these two tactic components by competitiveadsorptionfromCHCl,solutiononthe nonporous silica Cab-0-Si1 [227]. The adsorption of other carbonyl-group-containing polymers can be measured by IR [100,129,131,224]. Adsorptionisotherms of PS, poly(dimethy1 siloxane) (~DMS), polyethylene oxide (PEO) and block copolymers of and S as well as of EO and S on the silica-trichlorethylene interface have been determined as function of molar mass and copolymer composition [228]. Whereas adsorbed PS and PDMS have substantial fractions of the segments in loops the P E 0 adopts a Rat conformation with short loop lengths. At equivalent molar mass the plateau values of the copolymers of S and DMS are between the adsorbed amounts of the homopolymers, suggesting that both blocks contribute to adsorption. Plateau amounts of PE0-PS block copolymers of high styrene fraction are much larger than the values of P E 0 or PS of the same molar mass. Segregated conformations with P E 0 blocks in trains e silica surface and exclusion of the PS blocks in tails are tentatively explained. ock and random copolymers of styrene and methacrylate were adsorbed on precipitated silica from a mixed solution of carbon tetrachloride and n-heptane. The adsorption of the block copolymers varies steadily with the nonsolvent content, while the adsorptionof the random copolymersbecomes influenced onlywhen approaching theta conditions and phase separation [202]. For comparison, FTIR reflection measurements of PEO-b-polystyrene block copolymers adsorbed on mica

sheets exist [143,215] measuring the adsorbedamount afterrinsing with the solvent toluene and drying the covered mica. Polymersof different block ratios were measured. Surface densities and areaspermoleculecould be calculated [215]. The results were compared with results from surface force measurements. The adsorbed amounts 4.9 mg/m2 for PS and 2.3 mg/m2 for PEO-B-PS are in good agreement with those obtained gravimetrically (5.2 mg/m2) for PS [229] and by surface force (2.0 mg/m2) for PEO-b-PS [230]. The effect of anchor and buoy block sizes of poly[(dimethylamino)ethylmethacrylate]-b-poly(~-butylmethacrylate)~ on the adsorption behavior onsilica has been measured by the amount adsorbed and by the hydrodynamic layer thickness 12311. Adsorption isotherms of poly(isopropy1 acrylate) of different tacticity from trichlorethyleneshow, that stereoregularity hasnow influence on theadsorbed amounts 12327. Random copolymers of vinyl acetate and ethylene and in comparison polyvinylacetate havebeen adsorbed on methyl- and cyclohe~yl-modi~ed silica surfaces from CC14 1481. The adsorption of copolymers of acrylarnide and vinylimidazole, a pH dependent cationic polyelectrolyte, on precipitated silica in H 2 0has been measured in dependence of pH and cationicity by the determination of the total organic carboncontent 12011. For comparisonoptimal flocculation is measured by the clarification of the supernatant by light scattering. Adsorption isotherms of polyelectrolytes have been meaured by polyelectrolyte titration of thePEconcentrationinthesupernatantsolution [47,114,115, 120,17 1,18l]. The adsorbed amounts of the cationic polyelectrolyte poly(dially1dimethyl-ammoniumchloride),PDADMAC, of different molar mass on precipitated silica particles increase with thepH (Fig. 16) because of the increasing surface charge densityand with ionic strength (Fig.17) because of the reduced electrostatic

Dependence of the adsorbed plateau amounts of PDADMAC of various molar mass on pH without electrolyte; precipitated silica.

0.0

l

I-

4

0

Dependence of the adsorbed plateau amounts of PDADMAC of various molar mass on electrolyte concentration at pH 5.8; precipitated silica.

repulsion by screening of the segments 115,1711. At high ionic strength the adsorbed amount and the layer thickness increases with the molar mass. The charge density of the PE was varied by copolymerization of DADMAC with the neutral ~-methyl-~-vinyl-acetamide (NMVA).Theadsorbedamountsincrease with decreasing chain charge density because of charge compensation (Fig. 18). The decrease of the adsorbed amount of the neutral PNMVA with pH can be attributed to the increasing dissociation of the silanol groups and less hydrogen bonding between the carboxy groups of NMVA and the silanol groups or to less hydrophobic interactions of NMVA and the surface siloxane groups because of adsorption of OH- ions. The adsorption on negatively charged polystyrene latex shows the same tendencies. But thepH dependences are nonexistent because the surface charge and thePE charge do not change with pH. Interesting adsorption features are obtained with poly-l-lysine on latex because in this case the PE charge is dependent on pH 11811.

raction of A ~ ~ e Se~ments r e ~ By IR the fraction of occupied SiOH groups of pyrogenic silica at adsorption of statistical and block copolymers of ethylene oxide (EO) and propylene oxide (PO) and the corresponding homopolymers are measured in CC14 and CHCl, from the bands of the free and hydrogen-bonded Si0 groups [99,100,110,274]. A linear correlation between the calorimetric-determined area specific adsorption enthalpies and the IR absorption of the hydrogen bridges hasbeen established demonstratingthattheadsorptionenthalpy is a definite measure of the bounds.

sor

4

7

8

PH

Dependence of the adsorbed plateau amounts of P ~ A ~ ~ A C - c o - N ~ ~ A of various compositions at pH 5.8 without electrolyte; precipitated silica.

Competition and displacement studies demonstrate that the EO segments are preferentially adsorbed compared with the PO segments. The fraction of adhered segments decreases with increasing adsorbed amount, molar mass,EO fraction, and block length The dependences can be interpreted by eonformational models of the block copolymers with central EO PO blocks. The statistical polymers show homopolymer adsorption behavior similar to PPO. Thecoverage of SiOHgroups by theadsorption of polystyrene (Fig. of acrylnitrile-butadiene copolymer and of styrene-methycrylate copolymer has also been measured by IR. The fraction of boundsegments of poly(vinylpyrrolidone), PVP, on silica measured by IRand microcalorimetry in chloroform,dioxaneanddeuterium oxide are compared to literature values from NMR and EPR and the differing pvalues are discussed critically (Fig. 20) [loo, 138, 155, values decreasing withand independent of adsorbed amount are observed. NMR and ESR values are discussed also as a function of coverage 140,l The number of surface SiOH groups and of the adsorbed C = O groups of polycaprolactone, polybutylmet~acrylate,polyvinylacetate, and polyvinylpyrrolidone are notin agreement. A multiple interaction quotient depending on. the molar mass has tobe considered (Fig. The bound fractions PIR of polycarbonate and oligo(ethy1ene glycol adipate) fromsolutions in dichlorethylene and toluene on Aerosilmeasuredfromthe C = O stretching band decrease with increasing adsorbed amount at low concentration. At high concentrations aggregates are formed and adsorbed

.o

l

l

I

10

l000

10,000

Fractions of adsorbed segments of polystyrene measuredby IR as a funche chain length; system: cyclohexane, SiOz, The curve corresponds to the theory [l401with x 0.5 and 0.6.

.o

0

Fractions of adsorbed segments of 120,000 measured by ESR and IR as a function of the coverage CC14, (From Ref. 12.)

Stereoregularity, surface coverage, and agitation effect the pIR-values of poly(isopropyl acrylate) of different tacticity from trichlorethylene [232]. Flocculation by bridging influences the values. With IR and UV the adsorption of polystyrene (PS) and poly(methy1methacrylate) (PMMA) from trichlorethylene and from cyclohexane under theta conditions has been studiedonnonporouspyrogenic silica (Aerosil 130). Theadsorbed amount, the fraction of occupied silanol groups and the fraction of adsorbed polymer segmentsp have been determined [139,226]. The 6-valuesof both polymers are independent of molar mass, whereas the p-values in the plateau decrease with molarmass. Theadsorption of PMMAhas also been measured by EPRand competitive desorption experiments havebeen performed 1521. Competitive and displacement adsorption of PS and poly(ethy1ene oxide), PEO, has been investigated on Aerosil 130 from carbon tetrachloride and dioxane from the shift in the IR frequence of the SiOH groups [236]. 6- and p-values have been determined. The polar EO groups adsorb preferentially well as large molecules. PS becomes displacedcompletely by PEO.Thedegree of displacement ofPS depends on the fraction of styrene units attached to the SiOH groups and on the molar mass of both polymers. Adsorbed amounts and the specific segment surface interactions of polystyrene have been measured with a special transform FTIR technique in vacuum cells. The surfacesilanol structure and the interaction of the styrene groups with the free silanol groups could be quantified accurately [146]. The fraction of adsorbedsegmentsfor ester group containing the polymers polybutyl methacrylate, polycaprolactone, and polyvinylacetate were determined by IR and microcalorimetry [99,100]. Different dependences of poH and on surfdce coverageresulting in a multiple interaction quotientQ were found (Fig. 21). From thecorrelation of thebindingenthalpies and thefrequency shift of IR absorption of the SiOH groupsby adsorption of low molar mass model substances could be concluded that more solvent molecules have tobe displaced by the polymer segment compared to the model substance because of the restriction in the chain. PMMA adsorbed on Aerosil from CC14, butyl chloride and dichlorethane show at a small chain lengthof up to 50 segments, a fractionof adhered segments, pIR 0.85, corresponding to very flat adsorbed layer; in the range pIR dropsto 0.4. Achange of solventproduces hysteresis effects. With solvent competition in mixed solvent systems, regulation of the pvalues is possible (2371. Bound fractions of PMMA and PS from CC14 on silica measured by electron spinresonance(ESR)showadecreasewithcoverage 1521. Acomparison of bound fractions of PMMAfromIRandESR was made [149]. pIR showsa generaltrend that gives an underestimation of theboundfraction.Apossible explanation is thatadsorbedsegments do not affect the IR signal (e.g., if an MMA segment is adsorbed but cannot form a hydrogen bond). Another explanation is that the surface heterogeneity prevents the optimal matchof silanol groups and polymer segments. But on the other hand, the PNMR and PESR may be overestimated because also segments neighboring the bound segments are restricted in their mobility.

n U

U

0-

.o

Fractions of adsorbed segments pOH, and multiple interaction quotient versus the coverage A/A,. (a) Polycaprolactone, CC14. (b) Polybutyl methacrylate, CC14. (c> Polyvinylacetate, (d) Poly~inylpyrrolidone,CHC13.

.o

a1

0.4

0.8 Coverage

~overoge

The bound fractions, p, of poly(vinypyrro1idone) adsorbed on Aerosil under different coverages in comparison to terminally anchored polymers has been measured in D 2 0 by pulsenuclearmagneticresonance and compared to p-values obtained by ESR [151]. Bound fractions of random copolymers of vinylacetate and ethylene fromCC14 on methyl and cyclohexylmodifiedsilica particles areobtained by NMRand compared to values deviated from the segment density profile from SANS 11481. The preferential adsorption of the acetate groups is found. In qualitative comparisons with Monte Carlo simulations the contributions of trains, loops, and tails are discussed. NMR linewidth and time relaxation measurements are made to investigate the segment mobility of grafted P E 0 on porous and nonporous silica 1154,2381. Structural and dynamicstudies of adsorbed "C-labeled polyethylene have been accomplished with high-resolutionsolid-state NMR spectroscopy [l 301. An irnmo-

bile adsorbed part with majortrans conformation and minorgauche conformation (ratio 16:1) and a mobile part away from the surface with fast transition between gauche and trans, but more restricted and of smaller chain lengththan in solution, was observed. The influence of chloroform benzene mixtures on the conformation of grafted poly~ethyleneoxide) chains labeledat the end with nitroxidefree radicals has been studied with ESR[153]. The chains extend moreinto solution if the CHC13 content is increased. The folding of the grafted macromoleculesreflects the composition of the mixture. The integral adsorption enthalpies (displacement) of P E 0 adsorbed from aqueous solution represent the number of adhered segments directly on the surface:

1. The enthalpies decrease with coverageresulting in the same segment fractionin the trains at a given adsorbed amount for all molar mass with M 5000. 2. They have the same value at very low coverage (demonstrating flat conformation). 3. The plateau enthalpies decrease with increase in molar mass (longer tails and loops) [157,16Oj. Adsorption enthalpies of P E 0 and PP0 homopolymers and of triblock copolymers of EO and PP0 (EPE, PEP) on pyrogenic silica from CC14 and H 2 0 solutions have been measured together with the corresponding adsorption isothernls 11591. The fraction of adhered segments and the binding enthalpies were evaluated. The values in CC14 decrease with increasingamount adsorbed, molar mass,EO fraction and blocklength. Preferentially adsorbed hydroxyl end groups increase the adsorption enthalpies, and the EO middle and end block segments are preferentially adsorbed in comparison to thePO blocks. Conformation models are deduced. Adsorption enthalpies of oligomeric and polymeric vinylpyrrolidone on pyrogenicsilica from CHC13 measured by microcalorimetry were related tothe adsorbedamounts [100,132]. Fractions of adheredsegments were derived and compared to IR determined pIR fractions. The incollgruity of p values measured by IR, ESR, NMR, and microcalorimetry is explained. Adsorption enthalpies of polyvinylpyrrolidone on pyrogenic silica from aqueous and dioxane solution were obtained as a functionof molar mass and coverage and the fraction of surface occupancy is calculated from the enthalpiesby a calibration procedure /12]. A direct comparison to IR and NMR values is made. Layer Thicknesses of Adsorbed Homo~olymers, and Polyelectrol~es thicknesses of adsorbed polyethylene oxide, polyvinyl alcohol, and ock copolymer layers have been measured on particles of sufficient monodispersity as precipitated colloidal silica [90-93,163,1681 and on polystyrene lattices [91,239] by diffusion with photon correlation spectrometry andby sedimentation with ultracentrifugation. Low silica concentrations csio, 5 g/cm3 have to be used in order to avoid bridging flocculation, interference, and multiple scattering effects [go].

The hydrodynamic thickness increases with the adsorbed amount of PE0 (Fig. 22), intensively neartheplateauregion of theadsorptionisotherm,andwith increasing molar mass (Fig. 23). The lengths of the tails are responsible for both influences. From the dependence of the layer thickness on coverageasegment density profile can be derivedbased on thecrudeassumption of an additive adsorbed layer. Thedependences of theplateauthicknesses on molar mass of P E 0 are compared [l21 with values from different authors [88,90,92,16’7,240,247] in Fig. 24. A comparison is also made in Fig. 25 with root-mean-square thicknesses obtained from SANS segment density profiles. Hydrodynamic thicknesses obtained from viscosity measurements of Si02 dispersions show also an increase with coverage of P E 0 [241]. In comparison to profiles from SANS thereal profiles will extend further into the solution becauseof the dangling tails of the adsorbed macrornolecules (Fig. 25) [242]. In diffusion measurements by PCS, electrolyte friction or electroviscous effects produced by the polarization of the electrostatic double layer have to be considered at small ionic strength. This has been experimentally and theoretically established [161,239,242,243].Hydrodynamic effective thicknesses are mainlyinfluenced by the outer segments in the tails of the adsorbed macromolecules, which are also dominantly responsible for the stability and flocculation. Influences of chemistry and polarity of the surface, of preferential adsorption of segments of the copolymers

silica

0.0

0.2

0.4

0.6

0.8

1

1.2

1.4

1.6

1 ~ 3 ~ / ( ~ / ~ 2 )

of

Dependence of the hydrodynamic layer thicknesson adsorbed amount A 900,000 ondifferentsurfaces; cSilica 5 g/mL, clatex 5 g/mL.

X

Dependence of the corrected hydrodynamic plateau layer thickness molar mass of on different surfaces; c,i1ica clatex 5 lo-’ g/mL.

on

Double-logarithmic plot of the hydrodynamic plateau thicknesses of in dependence of the molar mas n various substrates from different authors; lower line, on silica: 0 [90], upperline, on polystyrenelatex: 17] [SS], (From Ref. 8.)

0

10,o

40'0

dis~nce Comparison of the segment density profile (SANS) and layer thicknesses of P E 0 of different molar mass from PCSand ellipsometry.

and of the hydrolysis degreeof PVA could be recognized aswell as the influence of particle curvature on the hydrodynamic thickness. The surface particle curvature influence can be eliminated by use of a correction formula [244] leading to corrected thicknesseson flat surfaces. comparison of the dependences of those corrected values on molar mass with theoretical the approach [245] leads to discrepancies in the scaling factors [242]. Hydrodynamic thicknesses of PE0 on octyl-group-modi~edsilica and on bare silica are compared [166,168]. The plateau thicknesses corrected for flat surfaces and their dependence on molar mass are the same on bothsilicas despite of different adsorbed amounts.Because the thicknesses are determined by the length of the tails, it can be concluded that the tail formation is not influenced by the hydrophobic parts of the surface. Also the thicknesses on both silicas are not affected by the sequence of mixing of polymer solution and silica suspension and also not by stepwise polymer addition in comparison to addition in one step 1681. Time effects and influencesof thesurplus of thepolymer in solutionare observed [166,168]. These thickness alterations can be attributed to long-lasting structure changes in the layer forced by competing polymer segments from the solution. Thedependence of the PE0 layer thickness on pH as a parameter for the binding strengthon small particles (Ludox 32 mm) shows constancyof the thickness up to pH10.5 and then a sharp drop tozero, whereas the decreaseof the adsorbed

amount is more gradual. The SiOH groups are dissociated at highpH and can no longer form hydrogen bonds with the EO segments. The dependence of the hydrodynamic layer thicknesses on the total molar mass and on the molar mass of the P E 0 blocks leads to the conclusion that the hydrophilic PE0 blocks are able to adsorb on the SiOH surface groups of silica but not on the hydrophobic latex surface, where only the hydrophobic PP0 segments are adsorbed (Fig. 26). Therefore very different adsorbed conformations on silica and latex exist with fiat layers on silica but extended layers on latex. Analogously the smaller thicknessof PVA on silica and the decreaseof the thickness with increasing degree of hydrolysis is evidence both for favored binding of the alcohol groups to the silanol groups and for the preferred adsorptionof the acetate groups on the apolar latex surface [91]. Layer thicknesses of adsorbed polyelectrolytes are rare in the literature. In the case of adsorbed cationic PDADMAC, of copolymers of DADMAC and NMVA, and of poly-L-lysine polyelectrolytes, layer thicknesses couldbe only obtained with stable systems by careful elimination of electrostatic interactions and electroviscosity effects at moderate ionic strength l 5,171,181]. The dependence on molar mass and adsorbed amount of PDADMAC and also on ionic strength could be measured under selected conditions (Figs. 27 and 28) 171,1751. At plateau adsorption the hydrodynamic thicknesses of all investigated polycations increase with increasing ionic strength because of the screening of the segment charges by the NaCl ions. At high ionic strength the thicknesses increase with the molar mass becausethelongermacromoleculesformmoreextended tails thantheshorter

ob

o5

Dependence of the hydrodynamic plateau layer thicknesses on the molar mass of the PE0 blocks of PEO-PPO-PE0 block copolymers, EI],precipitated silica, 0, polystyrene latex.

L:

K2

10

0 0.0

1.0

Hydrodynamic layer thickness of PDADMAC of various molar masses versus adsorbed amount, pH 5.8.

35

25

1 372000 100000

1.o

l

I mol

Hydrodyllanlic layer thicknesses of PDADMAC of various molar masses versus electrolyte concentration.

ones. Chemical affinity in addition to electrosorption hinders the displacement of thesegments by competingsmallcounterions.Atlowcoveragethethicknesses increase with the adsorbed amount independent of molar mass. There are only short tails andthethicknesses are small. With rising adsorbedamountlonger tails are formed and the increased thicknesses become molar mass dependent. In the case of pure electrosorption the layer thicknesses of polyacrylamides 0.1) [246] increase with coverageand approach a limiting value. The decrease of the thickness aswell as of the adsorbed amount withionic strength is caused by the increased competition the small counterions [12]. Thicknesses measuredby ellipsometry on flat surfaces are much smaller because this measuring method averages over all adsorbed segments including trains and loops. ~ e Density ~ t Profiles Volume fractionprofiles were obtained by SANS forPE0 layers adsorbed onlatex. With the profiles root-mean-square thicknesses between 2 and 3 nm were calculated, whereas the hydrodynamic thicknesses range between 30 and 40 nm. is not able to detect the very small segment concentrations of the tails at large distances from the surface(Fig. 25). On bare silica in water SANS profiles of polymer layers were not measured, because contrast matchingof silica and incoherent scattering from water could not be sufficiently accomplished. ~egmentdensity profiles of random vinylacetate-ethylene copolymerswre measured by SANS in CC14 on methyl-, butyl-, and cyclohexyl-modified Stober silica surfaces and compared with theory. Bound fractions of segments are deviated[148]. e Density of Bare and Covered Silica Particles Surface charge densities of bare pyrogenic and precipitated silica particles as function of pH were obtained by potentiometric pH titration 112-1 14,171,2221 and by titration with apolyelectrolyte of opposite charge(polyelectrolyte titration) [l 13,114,1711, pH titration showed that the increase of the charge density of bare particles with pH becomes steeper with increasing NaCl concentration. This can be explained by the increased screening of the dissociated SiQ- surface groups and consequently more SiQH groups become dissociated. The point of zero charge, PZC, is located in the region 2 pH 4. At pH 5.8 a charge density of-0.007 Cjm2 was measured corresponding to dissociation of 1% of the total SiQH groups and an average charge distance of 5 nm. At H 12 in NaCl solution the group surface density corresponds to 3.5 QH((nm) Also surface charges of particles covered with adsorbed polymer and polyelectrolyte layers were determined. PVA-coveredparticles show small decreaseof the charges of the acidic silanol groups in comparison to bare particles, but the pH dependence is not altered [247]. This canbe attributed to a simultaneous adsorption of PVA segments and of potential-determining ions. Coverage with neutral PNMVA does not influence the charge density [171]. In the charge density determination by polyelectrolyte titration with PDADMAC charge reversal by excess adsorption of has to be considered. It is recommended to determine the adsorbed amount at the zeta potential 0 where charge co~pensationoccurs. In this way the surface chargeof octyl-modified silica

has been measured and asurface charge densityof 80% in comparisonto baresilica has been found, caused by the reduced number of surface SiOH groups. The adsorption of the cationic polyelectrolytes PDADMAC and P(DADMACCO-NMVA)results in an increase of the surface charge at pH 8 and at low ionic strength produced possibly by an additional dissociationof SiOH groups 1711. At pH 8 the charge densities of PE-covered and bare particles are the same because of no additional SiOH group dissociation. At pH 9 the charge densities of bare and covered particles are very high and correspond to the dissociation of all SiOH surface groups with an average charge distance of 0.5 nm. Additional dissociation of SiOH groups in the porous surface region, adsorption of OH-ionsonthe hydrophobic Si-0Si bridges or slight dissolution of silica are supposed to be the reason 120). eta Potential of P o ~ y m e r - ~ o v e Particles re~ From l~easurementsof the electrophoretic mobilities the zeta potentials of bare particles at different pH and ionic strength as well as the isoelectric points are obtained [171]. At pH 3 the zeta potentials of bare silica particles are positive attributed to the adsorptionof H' ions. The isoelectric point with the electrophoretic mobility U 0 is located at PH 4, agreeing well with the PZC. At higher pH the dissociation of the SiOH groups results in a negative zeta potential asymptotically reaching a plateau. With increasing NaC1 concentration the positive zeta potentials at low pH and the negative zeta potentials at high pH are reduced by screening through the corresponding counterions. From the electrophoretic mobility of silica particles with adsorbed electrokinetic effective thicknesses are determined 1631. The discrepancy between the hydrodynal~ic andelectrokinetic layer thicknesses and the contradiction to the theory may be due to the approximate theoretical treatment and the influence of P E 0 adsorption on the surface charge causedby the change in the dielectric constant of the layer and by the dissociation of the OH groups. The surfacepotentials in the presenceof the adsorbed layer are much lower than the potentialof the bare surface and refer to aconsiderable changeof the surface chargeby P E 0 adsorption 1631. The zeta potentials of poly(vinylalcoho1)- and poly(acry1amide)-covered silica particles decrease due to the shift of the slip boundary by the adsorbed layer 1247). Thicknesses of the layers were calculated from the displacementof the slip boundary with the assumption of unchanged surface charge and electrical double layer. Electrophoretic measurements of the effective charge and of the electrophoretic velocity of Stober silica particles with adsorbed polyelectrolyte layers of sodium poly(ethy1ene sulfonate) and poly(styrene sulfonate), poly(4-vinyl-~-ethyl-pyridiniumbromide), poly(4-vinyl-~-benzyl-pyridinium bromide), and copolymers of 4-vinyl-~-ethyl-pyridinium bromide and 4-vinyl-~-benzyl-pyridinium bromide have been carried out 12481. The results are explained by hydrophobic, dipoledipole, and electrostatic interactions. At adsorption of the countercharged PDADMAC and copolymers of DADMAC and NMVA at pH 5.8 the negative zeta potentials at different NaCl concentrations decrease to zero and are overcompensated positive to zeta potentials by excess adsorption of the polycations leadingto positive zeta potential plateaus at

plateau adsorption. The plateau valuesof the zeta potential decrease at highionic strength because of screening (Fig. 29) [171,191]. The dependence of the electrophoretic mobilities of poly(divinylben~ene), poly(vinylbenzyl chloride) and of copolymer-coated particles on pH and the isoelectric pointsmeasured by microelectrophoresisarestrongly influenced by thesample history and the coating conditions [249].

isp placement andExchange of Polymers a. ~ o ~ o m e r Polymerscan be displaced by monomeric displacers from silica surfaces [250]. Above a critical displacer concentration determined by the surface affinity of the displacer the polymer is desorbed from the surface. Displacement isotherms of P E 0 were measured by batch adsorption, NMR, and hydrodynamic layer thickness. values were determinedforthe displacers dimethylformamide and dimethylsulfoxide [220]. PMMA on silica in toluene suepension has been displaced by dioxane [251]. Adsorption energy parameters xs of the polymers PEO, PVAC, PMMA, PBMA, PS, and PVC relative to the solvent CC14 have been obtained from displacement studies [252]. Polymer Becauseof theco-operative interactions the displacer efficiency of polymers is very high, but mostly the process becomes slower than with monomer displacers. Very few polymer-polymer displacement studies exist. Poly(buty1 methacrylate) in CC14 was desorbed by poly(tetrahydr0furane) 12511. In a mixture of hydrogenated and deuterated polystyrene a complete adsorption preference was found for the deuterated polymer on oxidized silicon [253]. Also tacticity differences may lead to adsorption fractionation.

7.

!t

-1 0.01

0.1 Imgl-1

1

Zeta potentials of silica a function of the concentration of PDADMAG 428,000 for various salt concentrationsat pH 5.8. The marks on the abscissa indicate the concentrations which correspond to the adsorption plateaus.

It is possible to displace P E 0 of low molar mass by P E 0 of high molar mass from the silica surface. This is demonstrated by the hydrodynamic thickness. If P E 0 900,000 is added to silica fully covered with P E 0 50,400, the layer thickness increases within 10 min to nearly the same value at direct addition of P E 0 900,000 [168]. A reduction of the thickness with timeis measured with P E 0 layers of high molar The reduction becomes lower with increasing excess of the polymer in solution and upon critical excess concentration time-independent thickness become existent (Fig. 30). Two reasons, polymer degradation. and reorganization of the polymer layer, may be responsible. The existent time-dependent flocculation by bridging is hint for the changeof the layer structure in period of hours. Measurements of the adsorbed amountsof mixture of two polystyrene samples (PS) of different molar masses 40,000 and 280,000, in cyclohexane by tritium tracer technique showthat the small molecules are displaced from silica by the long molecules if their concentration is high enough to cover the surface [254]. Exchange studies of adsorbed PS labeled with3H on silica by the same unlabeled polymerin solution gives very small exchange rates with long lifetimes in the adsorbed state. Thermodynamic equilibrium is guaranteed. The kinetics is dependent on the concentration in solution [165].

Dependence of the hydrodynamic layer thickness MW 900,000, at differentexcessconcentrations silica, 1 0 - ~g / m ~ .

of PEO, on thetime t; modified

Kinetics of competitive adsorption of PS of different molar mass was measured on porous silica showing preferential adsorption of the low-molar-mass chains in the pores at the beginning followed by a long process exchange by the larger polymer leading to equilibrium within 35 h. Measurements with samples of high polydispersity result in a preferential adsorption of the 1~ediunl"molar-mass polymer [255].

coagulate when theelectrophoreticmobility becomes zero at acidic pH or by increasing electrolyte concentration, the coagulation of silica sols is described to be anomalous in terms of the electrolyte-pH control [256]. Repulsive hydration forces are especially suggested as an inherent property of silica. As reasons for the anomalous stability, also the steric repulsion of polymer silicate species in layers at the water interface and hydration effects of the cations in this layer are discussed. Thehydrationforceparametershave been estimatedfrom turbidity data. The results of this treatment show that the hydration force decreases with increasing NaCl concentration up to0.2 mol/L. This theoryis suited to describing thestability of i 'ca suspensions over a wide range of pH and electrolyte concentration [52]. o the ordering of liquids on the surface of silica particles plays an important role in stabilization. Whereas in pure water,silica suspensions are stabilized by the diffuse double layers [53] the role of Coulombic forces becomes negligible for the stabilization in water-2-butoxy-ethanol 22571 and in water-2,6-lutidine mixtures [258]. The van der Waals attraction and stability of Stober silica in mixtures of ethanol~yclohexanewas investigated by rheology, adsorption excess isotherms, smallangleX-ray scattering (SAXS), and by optical dispersionmeasurements. WiththeobtainedHamakerconstantsthe theoretical treatment by multilayer adsorption and a single-sheet, hard-spheremodel was ingoodagreementwith the experimental results [259]. ~oagulationand flocculation pyrogenic and precipitated silica has been measured by sedimentation and rheometry [92,95,260]. Aggregation state and aggregation kinetics can be evaluatedfromthedecreasing diffusion coefficientsof the aggregating particles withtimeunder different polymeradsorptionconditions [91,94,168,190,191]. Critical flocculation concentrationscould be determined. Correlationstothe extensively studiedadsorptionbehaviorcould be achieved. The state of flocculation is discussed by comparing the diffusion coefficient after a certain time aggregation under different conditions, such as concentrations of silica, polymerand salt, pH and molar mass.The kinetics of aggregationwas obtainedfromthetimedependence of the diffusion coefficient by determining the halflife value and the rate constant with the second-order Smoluchowski theory [78] using a special evaluation procedure [94,186]. Precipitated amorphous silica with low polydispersityin particle size distribution iswell suited forprincipal investigations of this type.

The coagulation of bare precipitated silica particles occurs at pH 3.5 without addition of salt andatneutralpH by salt additionat 5 mol/L [92,94]. The second-order kinetics and the approach of the rate constant of coagulation at high salt concentration to the theoretical Smoluchowski value6.2 10"' m3/(s. particle) is measured.But also deviationsfromsecond-order kinetics andrate constants of 40 to 90% of the theoretical value arefound [82,261-2631. Hydrodynamic and steric interactions, viscous effects [79,80,261], repulsive potentials [83,84], orthe reversibility of aggregation (85,861 aresupposed to be the reasonsforthedisagreementandareintroduced in the theoretical derivations. PCS coagulation measurements in the concentration range of to 10lg Si02 m3/(s. particles result in decreasingrateconstantsfrom 3.2 10"' to 1.7 particle) fordecreasing particle radii from 515 to 45 nm (1851. The lower rate constants are explained by hydration forces which increase the stability. The flocculation of precipitated and pyrogenic silica (Aerosil) by polyethylene oxides (PEO) was investigated under continuous addition of the influences of time of addition, stirring time, stirring intensity, polymer coverage, concentration ofsilica, pH and ionic strength were measured by turbidity and sedimentation [93,101,203]. At sufficiently highconcentrationtheaggregated Aerosil particles can be flocculated by P E 0 of molar mass 6000 and 40,000 even in the absenceof electrolyte, leading to reproducible flocculated systems at standard mixing and stirring. With additional rheological measurements theclassical model of polymer bridging was identified as the most likely mechanism of destabilization [loll. A transition from granular flocs to structured flocculates has been observed with growing surface coverage if the particle concentration is sufficiently high. At low particle concentration stabilization occurs at full coverage with P E 0 40,000. In contrast to pyrogenic silica consisting of aggregated particles the monodisperse spherical particles of precipitated silica can only be flocculated in granular flocs by decreasing thepH below 3.5 or by addition of electrolyte, eNaC1 0.01 mol/ L. The flocculation of pyrogenic silica at high pH and without salt is markedly dependent on mixing conditions, and the structure changes from granularflocs to structured flocculation at increasing coverage with PEO. At high salt concentra3.5 the precipitation of both silicas becomes independent of mixing. Modified precipitated silica with surface octyl groups becomes unstable at cNaCl 0.008 mol/L,the diffusion coefficient decreasesfrom 3.68 to 3.21 m2/s and the rate constant of coagulation increases from zero to 5 10"' m3/(s. particle) near the diffusion controlled value of 6.2 10"' m3/(s. particle) at cNaC1 0.03 mol/L 11681. The lower critical coagulation concentrations (CCC) of modified silica in comparison to unmodified silica can be attributed to thelower surface charge. The higher-charged particles need a larger salt concentration to become unstable because of the electrostatic double layer and the electrostatic repulsion. From measurements ofsilica particles withadsorbed P E 0 layers the subsequent characteristics were obtained [94]. The critical flocculation concentration (CFC) of NaCl is characterized by the steep transitions from high to low diffusion coefficients (Fig. 31) at increasing electrolyte concentration after a cer-

Dependence of the diffusion coefficient on log c after t 40 min for precipitatedsilica(138nm)bare and coveredwith at fullcoverage, 160,000 (A), 325,000 900,000 1

tain flocculation time. oflow molarmass, 160,000,producesa distinct decrease of the critical ~aC1-concentration CFC, in comparison to bare silica and the same rate constants are obtained. At low salt concentrations near the CFC with P E 0 of molar 27,000 and 50,000 only small flocs are formed, which donot growfurther.Plateauadsorption of high-molar-massPEO, MW 160,000, leads to stabilization with very smalldecrease of thediffusion coefficients (Fig. 31) and considerable lower rate constants (Fig. 32). Thehigh layer thicknesses caused by the tails of the adsorbed P E 0 chains are responsible for the steric stabilization by elastic and osmotic repulsion. Particles covered with small adsorbed amounts of P E 0 having small layer thickness are not stabilized and flocculate in rate and state like uncovered particles. The correlation between rateconstantandhydrodynamicthickness of the layer demonstratesthe steric influenceof the tails of theadsorbedmacromolecules on stability and flocculation. At low coverage,with flat adsorbingchains of layer thicknesses smaller than 6 nm, aggregation occurs. At high coverage and plateau adsorption, tails are developed resulting in layer thicknesses larger than 6 nm, and the particles become sterically stabilized. The dependence on thickness d can be forrnulated by k exp(-/?d), where 0.28 nm-' for all salt concentrations. The different suspensions of bare silica, octyl-group-modified silica, and latex become fully stabilized at the same hydrodynamic thickness of 10-15 nm of the adsorbed P E 0 layer independent of the substrate and the way of establishing the thickness in time, at different excess concentrations, and with different molar mass (Fig. 33) [168].

Double logarithmic plotof the flocculation rate constant k of concentration cSilica I g/mL, 25°C; 27,250

The followingsteps are of decisive influencein the kinetic model of flocculation, mixing of polymer and dispersion, collision of polymer and particle to first adhering, conformation change of the adsorbed macromolecule, collision of covered particles, followed by bridging of the adsorbed macromolecules. Dependent on the rate-determining step (fast or slow conformation change) equilibrium or nonequilibrium flocculation occurs [64,94,195]. Flocculation of P E 0 silica particles is also interpreted by bridging flocculation under nonequilibrium conditions and by the formation of loose aggregates by small long-range interactions [241]. The influence of the adsorption of cationic polyelectrolytes of different charge density poly(dially1-dimethyl-a~~oniumchloride),PDADMAC,and of copolymers of DADMAC with ~-methyl-~-vinylacetamide on the stability of precipitated silica suspensions has been extensively studied by measuring the diffusion coefficientof the particles with photon correlation spectrometry [171,191]. Raster electron micrographs demonstrate the aggregation state of bare and covered particles under different conditions directly 19l].

1 0

modified silica in molar mass chang~ change in a d s o r ~ amount ~d

D

aging unmodifiedsili latex

Dependence of the diffusion coefficient after t 40 min on the hydrodynamic layer thicknessd for differentways of obtaining the thickness: change in the molar mass and in the adsorbed amount, aging of the suspension. Surfaces: unmodified silica, modified silica, and latex; 5 lo-’ g/mL, 0.1 mol/L, Clatex 5 lo-’ CNaCl 1 mol/L.

A comparison of measured zeta potentials and diffusion coefficients is given [l911 at different adsorbed amounts, molar mass, charge density and ionic strength. An example the dependencesof the zeta potentials is shown in Fig. 29 and of the diffusion coefficients in Fig. 34. The adsorbed amount of the cationic polymers reduces significantly the zeta potential of the covered silica particles. Also the zeta potential becomes screened by increasing electrolyte concentration. The isoelectric point is reached at the same amount of adsorbed charges regardless of the ionic strength and of the charge density of the adsorbed polyelectrolyte. When the zeta potential decreases the repulsionbetween the bare or covered particles becomes too small to avoid aggregation. Aggregation is observed if the amount of the zeta potential is smaller than. 20 Additionally, at low ionic stren~th,mosaic flocculation of the polyelectrolyte covered particles can be observed at about50 to 80% of surface saturation with macroionsof long chain length. At high ionic strength, flocculation canoccur at zeta potentials higher than 20 by bridging. The principle of the mosaic flocculation is described in Refs. 194 and 246 and of the bridging mechanism in Refs, 264 and 70. At larger adsorbed amounts a chargereversal occurs and a plateau valueof the zeta potential is obtained at plateau adsorption. This zeta potential plateau rises with increasingsalt concentration due to the increasing adsorbed plateau amounts. At high ionic strength the zeta potential plateau decreases again due to the screen-

0 0 428000

mg

Diffusioncoefficients of silicaas a function of theconcentration of PDADMAC 428,000 at pH 5.8 for various salt concentrations. The marks on the abscissa indicate the concentrations which correspond to the adsorption plateaus.

ing of the charges of the adsorbed macroion segments. i n the adsorption plateau the suspensions arestabilized electrostatically at low ionic strength, electrosterically at medium ionic strength, and sterically at high ionic strength. The steric stabilization by osmotic repulsionof the tails of the adsorbedpolyelectrolytes needs macroions of long chain length11911. The aggregationby van der Waals attractionis very fast and diffusion controlled at verylow zeta potentials of the polyelectrolytecovered particles (Fig. 29). ridging and mosaic flocculation are much slower because repulsive potentials between particles exist. A dependence of flocculation on the Si02 particle concentration and on mixing conditions was observed [69,275]. The flocculation mechanism is changed by the charge density of polyacrylamides 12651. A sensible dependence of the bridging of Si02 particles on the molar massof the polyacrylamides, solvency,and particle size is found in glycerol-water mixtures. The optimum flocculation conditions of precipitated silica by the cationic copolymer of different composition of acrylamide (AM) and vinylimidazole in a two-step process have been studied by the clarification of the supernatant [201]. Reversible flocculation by particle neutralization and bridging was found by pH adjustment. 0ptimum flocculation occurs under charge reversal conditions measured by the zeta potential. From the molar mass dependence additional bridging is concluded. ~tr~ct~re Systems. The topology of flocculated silica particles and the conditions fortheir formation has been investigated by analyzing the angle dependence of small-angle neutral scattering intensity (SANS) of nonaggregated spheres with radius 19 nm and of flocs [266]. Bridging flocculation has been regulated by the adsorption of polyethylene oxides and of strong cationic polymers of

large chains lengths. Phase diagrams characterizing the conditionsof stabilization, flocculation, and gel formation are deduced. Addition of polyvinylalcohol toaqueous silica suspensions of Ludoxtype results incoacervation at a definite PVA/Si02 ratio of2.5 chainsegmentsper nm2, leading to an oil-like viscous phase [267]. The particles are just covered by a monolayer of PVA, where thehydroxylgroups are oriented and hydrogenbonded to the SiOH groups of the surface and the hydrocarbon chains form a hydrophobic coating. The stability of pyrogenic nonporous silica with adsorbedlayers of P M ~ and A PS/PMMA copolymers in xylene and toluene was characterized by the elastic modulus from oscillatory rheological techniques in order to quantify the strength of the flocculated network as function of surface coverage, molar mass, architecture, and solvency [207]. The flocculated network is weaker at complete surface coverage and when diblock Copolymers are used. Higher-molar-mass polymers lead to weaker flocculation but no change in the elastic modulus. ~ ~ ~ r e ~ aKinetics tion Second-order flocculation kinetics asafunction of the concentration of bare particles havecommonly been found [244,261,263,268] under unstable conditions. Only in one paper is deviation from the second order reported. PCS results on the kinetics of coagulation are accomplished by several authors [94,168,185,187-19l]. Therateconstantsnearlyreachthe theoretical Smoluchowski value of k 4 kT'/3q 6.16 x 10"' m3/(s. particle) for water at 25°C and are independentof the concentrationif the silica concentration is verylow E941. Second-order flocculation kinetics were also measured for silica particles covered withP E 0 and PDADMAClayers of different molar mass and charge density at different pH andelectrolyte concentrations (Figs. 32 and [94]. P E 0 layers of low molar mass result in a decrease of the critical flocculation concentration.Therateconstantsarethesameathigh electrolyte concentration. Polymerlayers of high molar mass 325,000, 900,000) reduce the rate constants at full coverage drastically even at high electrolyte concentrations, and stabilization occurs. At low coverage, adsorption of high-~olar-mass polymer leadsto flocculation with rate constantsof the same valuesas with P E 0 of low molar mass. The correlation between rate constant and hydrodynamic layer thickness demonstrates the steric influence of the tails of the adsorbed macromolecules on stability and flocculation [94]. The rate constants of silica particles coveredwiththepolycations P ~ A D M Aare ~ different inthecases of mosaic, bridging, and van der Waals flocculation, Plateau coverage with PDADMAC of high molar mass leads to stabilization and to a corresponding strong decrease of the rate constantsalso at large ionic strength due to electrostatic or steric stabilization, as discussed already 17 1,19l].

3. ~tabilizationbyPolymer Adsorption Additionally to the cases discussed before, the stability behavior of silica particles with adsorbedlayers of block and random copolymersof styrene and methacrylate

0

Flocculation rate constants of silica as a function of the concentration of PDADMAC 428,000 at pH 5.8 for various salt concentrations. The marks on the abscissa indicate the concentrations which correspond to the adsorption plateaus.

in carbon tetrachloride-n-heptane solutions was investigated [2027. For random copolymersthe stability is decreased by loweredsolventquality and becomes increasingly solvent dependentaround the theta conditions. For block copolymers of more than 20%styrene, decrease of solvent qualityinitially reduces stability, but thereafter the stabiliy improves because of better anchoring of the blocks.

Polymer layers adsorbed on colloidal particles prevent aggregation and flocculation by theformation of steric barrier [55,94,171,191]. On theotherhand, addition of polymers accomplishes destabilization and flocculation of dispersions withtheformation of granular flocs orstructured flocculates [260]. Vander Waals attraction, direct bridging of the particles by the macromolecules, attraction by interaction and entanglement of polymerchainsadsorbedon different particles, and interaction by mosaicstructureshavealready beendiscussed possible mechanisms. Phase separation may be also induced by depletion interaction in case of nonadsorbing polymers, e.g.,poly(dimethylsi1oxane) on hydrophobic silica particles grafted by n-C18 chains in cyclohexane suspensions [200]. The observed trends of this depletion behavior can be explained theoretically in qualitative way. Because of the short duration of particle collision, not only the thermodynamic equilbrium of adsorption between two particles determines the steric stabilization

but essentially there are also kinetic effects connected with the competing processes of desorption and reorganizationof the adsorbed layer and the rheological properties [17]. The rate of adsorption and the approach to the thermodynamic equilibrium in comparison to the collision rate,the influenceof dynamic effects as stirring and shearing have to be carefully examined in order to interpret flocculation tests [198]. An interesting efficient complex flocculation of fibers was obtained for very small Si02 particles (6 nm) in combination with cationic potato starch [269].

5. Interaction by Force-~istance ~easurements sions depends on factors such as the kinetics of Brownian collisions, the hydrodynamic interactions, the state of the surface layers in relation to the thermodynamic equilibrium, and predominantly on the existent surface-surface forces. The forces between atomicallysmooth mica surfaces immersed in organic and aqueous liquid media were determined as a function of distance from 0-300 nm by direct force-distance measurements, Layers of polystyrene in cyclohexane under different solvent qualities, and of polyethylene oxide and poly-L-lysineinwaterwith different electrolyte contents were measured [216,269]. Attractive and repulsive interactions between the adsorbed layers are present depending on the extent of adsorption, the quality and nature of the solvent, and the cationic polyelectrolyte for combined electrostatic and steric forces. Force-distancemeasurements ofpoly(vinylalcoho1) layers oncrossed quartz filaments [215] show that long tails are responsible for stabilization, and the loop and tail concentration in the layer may be rather low. At distances smaller than 20 nm the steric repulsion of the loops becomesefficient. Adsorbed layers on mica surfac in 0.1 mol KN03 watersolution at partial and full coverage are investigated. ecause of the overlap of the opposing segments on the two surfaces a repulsive osmotic interactionis measured under the good solvent conditions at full coverage, whereas at lower surface coverage an initial attractionattributedtobridgingfollowed by ashort-rangerepulsion is found [217]. Theoretical treatments based on mean field and scaling arguments [13,270,271] are in qualitative agreement with the results. Force-distance results for symmetrical and highly asymmetrical poly(styrene4ethylene oxide) layers from toluene on mica are compared with theoretical scaling laws [215]. Forces between poly(2-vinylpyridine)-covered mica surfaces are measured as function of polymer charge and ionic strength [272]. At low pH and low ionic strength adsorption is independent of the chain length and causes surface charge reversal by a flat adsorbed layer. More loopy conformations at high salt concentration, 0.1 mol/L, lead toan additional steric force. At neutralization of the polymer a large expansion of the layers is measured. The results highlight the crucial role of segment surface binding and of intersegmental repulsion on the extension of the layer. Depletion effects caused by electrostatic repulsion or a double"1ayer osmoticpressuregradient are also observed by kinetic adsorption effects.

~fiInd~.strial a Chemistry, 1. H. Burkert and J. Hartmann, in ~ l l m a n n s ~ n c y c l o p eo VCH-Verlag, Weinheim, 1988, Vol. A 11, pp. 251-361. Stscherbina, 2. H. Dautzenberg, W. Jaeger, J. Kotz, B. Philip, Ch. Seidel, and Polyelectro~ytes:Formation,Characte~ization and Application, CarlHanser Verlag, Munich-~ienna-New York, 1994. 3. E. A. DiMarzio. J. Chem. Phys. 42:2101 (1965). 4. E. A. DiMarzio and F. L. McCrackin. J. Chern. Phys. 43539 (1965). 5. R. J. B. Rubin. J. Chem. Phys. 43:2392 (1965). 6. R. J. Roe. J. Chern. Phys. 43:1591 (1965); ibid. 44:4264 (1966). 7. R. Sirnha, H. L. Frisch, and F. R. Eirich. J. Phys. Chern. 57:584 (1953). 8. R. Sirnha, H. L. Frisch, and F. R. Eirich. J. Chem. Phys. 21:365 (1963). 9. R. Simha, J. Polym. Soc. 29:3 (1958). 10. I. Carmesin and K. Mremer, Macromolecules 21:2819 (1988). 11. K. Binder. Colloid Polyrn. Sci. 266:871 (1988). H. M. Scheutjens, Cosgrove, and B. 12. G. J. Fleer, M. A. Cohen Stuart, J. Vincent, Polymers at Interfaces, Chapman and Hall, London, 1993. 13. J. M. H. M. Scheutjens, and G. J. Fleer, J. Phys. Chem. 83:1619 (1979): ibid. 84178 (1980). A. Coben Stuart. Colloids Surfaces 14. J. M. H. M. Scheutjens, G. J. Fleer, and 21:285 (1986); ibid. 32:l (1988). ~id G. J. Fleer and Lyklema, in sorption from Sol~tionat S o l i d l ~ i ~ Surfaces (G. D. Parfitt, C. H. Rochester, eds.), Academic Press, London, 1983. 16. A. Silberberg. J. Colloid Interf. Sci. 38:217 (1972). 17. G. J. Fleer and J. M. H. M. Scheutjens. J. Coli. Interf. Sci. 111:504 (1986). H. Scheutjens and G. J. Fleer. Macromolecules 18:1882 (1985). 18. J. Scheutjens and G. J. Fleer, in The Effect qf’Polymers ~ i . s ~ e r s i o ~ 19. J. M. H. Propertie~s(Th.F. Tadros, ed.), Academic Press London, 1982. 20. M. A. Cohen Stuart, J. M. H. M. Scheutjens, and C . J. Fleer. J. Poly. Sci. Pol. Phys. Ed. 18559 (1980). 21. S. F. Edwards. Proc. Phys. Soc. 85:613 (1965); ibid. 88:265 (1966). 22. P.-G. de Gennes. Rep. Prog. Phys. 32:187 (1969). 23. E. Helfand and Y. Tagami. J. Chem. Phys. 56:3592 (1972). 24. E. Helfand. J. Chem. Phys. 62:999 (1975); ibid. 63:2192 (1974). 25. E. Helfand and A. M. Sapse. J. Chern. Phys. 62:1327 (1975). 26. K. M. Hong and J. Noolandi.Macromolecules14:727(1981);ibid. 14:1229 (1981). 27. Silberberg. J, Chern. Phys. 482835 (1968); ibid. 46:1105 (1976). 28. C. J. Hoeve, E. A. DiMarzio, and P. Peyser. J. Chem. Phys. 42:2558 (1965). 29. C. A. J. Hoeve. J. Polym. Sci. 30:361 (1970); ibid. 341 (1971). 30. P. G. de Gennes, Scaling Concepts in Polymer Physics, Cornel1 University Press, Ithaca, New York, 1979. 31. R. J. Roe. J. Chem. Phys. 60:4192 (1974). 32. E. Helfand. ~acromolecules9:307 (1976). 33. T.A. Weber and E. Helfand. Macromolecules 9:311 (1976). 34. P.-C. de Gennes. Macromolecules 14:1637 (1981).

35. P.-G. de Gennes and P. Pincus. J. Phys. Lett. L 241 (1983). 36. Klein and P. Pincus. Macromolecules 1129 (1982). 37. des Cloizeau. Physique 31:715 (1970). 38. M. A. Cohen Stuart, C. J. Fleer, and J. M. H. M. Scheutjens. J. Colloid Interf. Sci. 97:515 (1984). 39. H. A. van der Schee, and J. Lyklema. Phys. Chemie 88:1637,6661 (1984). 40. M.R. Bohmer, 0. A.Evers, and J. M.H.M. Scheutjens.Macromolecules 23:2288 (1990). 41. H. G.M. van de Steeg, M. A. Cohen Stuart, A. de Keizer, and B. H. Bijsterbosch. Langmuir 8:2538 (1992). 42. Papenhuijzen, H. A.van der Schee, and G. J. Fleer. ColloidInterf.Sci. 104540 (1984). 43. van Lent and J. M. H. M. Scheutjens. Phys. Chem. 94:5033 (1190). 44. S. Milner, T. A. Witten, and M. E. Cates. Macromolecules 21:2610 (1988). 45. S. Misra, S. Varanasi, and P. P. Varanasi. Macromolecules 22:4173 (1989). 46. S. Beltra'n, H. H. Hooper, H. W. Blanche, and J. M. Prausnitz. Macromolecules 243178 (1991). 47 M. A. Cohen Stuart, G. J. Fleer, J. Lyklema,W. Norde, and J. M. H. M. Scheutjens. Adv. Colloid Interface Sci. 34477 (1991). 48. Marra, H. A. van der Schee, G. J. Fleer, and Lyklema, in ~dsorption from S o l ~ t ~ o(R. n H. Ottewill, Ch. H. Rochester, and A. L. Smith, eds.), Academic Press, 1993. 49. B. V. Derjaguin. Trans. Faraday Soc. 36:730 (1940). 50. B.V. Derjaguin and L. Landau. J. Expt. Theor. Physik 15:662 (1945). 51. E. W. Verwey and J. Th. G. Overbeek, Theory of the Stability ~yop~obic Colloi~s,Elsevier, Amsterdam, 1948. 52. H. Yotsumoto and R.-H. Yoon. J. Colloid Interf. Sci. 157:434 (1993). 53. Goodwin, ed.), The Royal Society Th. Overbeek, in Colloid Dispersions of Chemistry, 1982, 1. 54. R. Hunter, ~ o ~ ~ z d ~ t i oColloid n Science, OxfordSciencePublications, Clarendon Press, Oxford, Vol. 1, 1987. 55. D. H. Napper. J. Coll. Interf. Sci. 58:390 (1977). 56. R. Evans and D. H. Napper. Kolloid Polym. 251:329,409 (1973). 57. D. G. Meier. J. Colloid Interf. Sci. 32: 106 (1970). 58. F, Th. Hesselink,A.Vrij, and J. T. G. Overbeek. Phys.Chem.75:2094 (1971). 59. D. H. Napper, Polymeric S t ~ b i l i ~ ~ t of i o nColloid ~ispersions,Academic Press, London, 1983. 60. B. Vincent. Adv. Colloids Interf. Sci. 4:192 (1974); ibid. 58:390 (1977). 61. I. A. Kitchener. Br. Polym. 4217 (1972). 62. Th. F. Tadros. Colloids Surfaces 18:137 (1986). 63. P. D. Pate1 and W. B. Russel, J. Colloid Interf. Sci. 131:192 (1989). 64. R. H. Smellie and V. K. La Mer. J. Colloid Interf. Sci. 23:589 (1985). 65. G. Fleer and A. J. Lyklema. J. Colloid Interf. Sci. 46: 1 (1974). 66. A. T. Clark and M. Lal, in The Effect of P o l y ~ e r s Dispersion Properties, (1-11. F. Tadros, ed.), Academic Press, London, 1982.

67. E. C . M. Pelssers, M. A. Cohen Stuart, and C . J. Fleer. J. Chem. Soc. Faraday Trans. 86:1355 (1990). N. Kiselenko, and A.Moldovanov.ColloidPoly.Sci. 68. Ad.A.Berlin,V. 270:1042 (1992). Chaplain, L. Janex, F. C. Graillat, and R. Audebert. Colloid Polymer Sci. 69. 273:984 (1995). 70. V. Adachi. Adv. Colloid Interf. Sci. 56:l (1995). 71. R. I. Feigin and D. H. Napper. J. Coll. Interf. Sci.75:525(1980);ibid.75:135 (1989); ibid. 75:525(1981). ~ t i o n J. Fleer, J. M. H. M. Scheutjens, and Vincent, in Polymer ~ ~ s ~ ~and 72. D i ~ ~ ~ e r s i o n S ~ a ~ i l i tD. y , Coddard and B. Vincent,eds.),ACSSymposium Series, American Chemical Society, Washington D.C., 1984, 245. 73. S. Asakura and F. Oosawa. Chern. Phys. 22:1255 (1954). 74. R. J. Hunter, R. Matarese, and D. H. Napper. Colloids Surfaces 7:l (1983). 75. J. Lyklema and G. J. Fleer. Colloids Surfaces 25:357 (1987). 76. M. K. Granfeldt,B.Jonsson, and C. E.Woodward. J. Phys.Chem.95:4819 (1991). 77. R. Podgornik, T. &sesson, and B. Jonsson. J. Chem. Phys. 102:9423 (1995). 78. M. von Smoluchowski. Phys. 17:557,585 (1916); ibid. 92:129 (1917). 79 B. V. Derjaguin and V. M. Muller. Dokl. Akad. Nauk. SSR (Engl.) 176:738 (1967). 80. L. A. Spielmann. J. Colloid Interf. Sci. 33:562 (1970). 81. E. P. Honig, G. J. Roebersen, and P. H. Wiersema. J. Colloid Interf. Sci. 36:97 (1971). 82. J. W. Th. Lichtenbelt, C. Patmamanoharan, and P. H. Wiersma. J. Coll. Interf. Sci. 49:281 (1974). 83. N. Fuchs. Z. Phys. 89:736 (1934). 84. B. Derjaguin. Kolloid Z. 69:155 (1934). 85. C. Frens, J. Th. G. Overbeek. J. Colloid Interf. Sci. 36:286 (1971); ibid. 38:376 (1 972). 86. C. Frens. Faraday Discuss. Chem. Soc. 65:146 (1978). 87. W. Stober, A. Fink, and E. Bohn. J. Colloid Interf. Sci. 2652 (1968). 88. M. A.Cohen Stuart, F. H. W. Waajen, T. Cosgrove, T. L.Crowley, and B. Vincent. Macromolecules 17: 1825(l 984). 89. Th. van den Boomgaard, T. A. King, Th.F. Tadros, H. Tang, and B. Vincent. J. Colloid Interf. Sci. 6658 (1978). 90. E.Killmann, H. Maier, P. Kaniut, and N. Giitling.ColloidsSurfaces15:261 (1985). 91. E. Killmann, H. Maier, and J. A. Baker. Colloids Surfaces 31:51 (1988). 92. E. Killmann, Th. Wild, N. Giitling, and H. Maier.ColloidsSurfaces18:241 (1 986). 93. E. Killmann, H. Maier, P. Kaniut, and N. Giitling. Ber. Bunsenges. Phys. Chem. 88:1141(1984). 94. E. Killmann and H. Adolph. Colloid Polym. Sci. 273:1071 (1995). 95. J. Eisenlauer and E. Killman. J. Colloid Interf. Sci. 74108 (1980). und Anwendungen,in 96. HydrophobesAerosil@,Hertstellung,Eigenschaften S c ~ ~ ~ t e n r e i ~ e P Degussa i g ~ e n ~32e (1969); 16-11 (1977); 6 (198l), I

97. 98. 99. 100. 101. 102. 103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130.

H. Maier, J. A. Baker, and J. C. Berg. J. Colloid Interf. Sci. (1987). I. D. Robb. Comprehensive Polymer Sci. 2:733 (1989). M. Korn, and E. Killmann. J. Coll. Interface Sci. 7619 (1980). E. Killmann, M. Korn, and M. Bergmann, in ~ d ' ~ o r ~ tfrom i o n Solution (T. H. Ottewill, C. H. Rochester, A. L.Smith,eds.),AcademicPressLondon,1983, p. 259. J. Eisenlauer, E. Killmann, and M. Korn. J. Coll. Interface Sci. 74:120 (1980). R. Iler. The C ~ e ~ i s t r ySilica, Wiley Interscience, New York. C. C. Ballard, E. C. Broge, R. K. Iler, D. S. St. John, and J. R. McWhorter. J. Phys. Chem. 65:20 (1965). A. K. van Helden and J. W. Hansen. J. Coll. Interf. Sci. 81:354 (1981). E. Schulte, der ~ a p i ~ l Gas-Chromatographie, ar Springer, Berlin, 1983. S. Fritz and J. N. King. Analyt. Chem. 48570 (1976). R. Laibel and K. Hamann. Adv. Colloid Interf. Sci. 6:95 (1974). A. J. Lecloux, J. Bronchart, F. Noville, C. Dodet, P. Marchot, and J. P. Pirard. Colloids Surfaces 19:359 (1986). D. D. Lasic. Colloids Surfaces 20:265 (1986). M. Korn, E. Killmann, and J. Eisenlauer. J. Colloid Interf. Sci. 76:7 (1980). T. Konishi, E. Yamahara, and N. Ise. Langmuir 12:2608 (1996). T. P. Goloub, L. K. Koopal, and B. H. Bijsterbosch. Langrnuir 12:3188 (1996). D. Horn. Progr. Colloid Polym. Sci. 65:251 (1978). R. P. Abendroth. J. Colloid Interf. Sci. 34:592 (1970). D. Bauer,E.Killmann, and W.Jaeger.ProgressColloidPolym.Sci.109:161 (1 998). R. Rehmet and E. Killmann. Colloids Surfaces A 149:323 (1998). R.J. Hunter, Zeta Potential in Colloid Science, Princi~les and ~ ~ p l i & a t i o ~ Academic Press, London, 198l . D. Kerner and W. Leiner. Colloid Polym. Sci. 253:960 (1975). R. H. Muller, Zetapotential und Parti~elladung in ~aborpra~is, Wissenschaftliche Verlagsgesellschaft mbH, Stuttgart, 1996. J. Lyklema, F ~ n d a n ~ e n o t f~ ~rnterfa&e l~~ and Colloid Science, Vols. I and I1 Academic Press, London, 1995. H. Oshima. J. Colloid Interf. Sci. 168:269 (1995). E. Killmann. Makromol. Chem. Macromol. Symp. 1757 (1988). J. C. Dijt, M. A. Cohen Stuart, J. E. Hofman and G. H. Fleer. Colloids Surfaces 51:141(1990). R. R. Stromberg, D. J. Tutas and E. Passaglia. J. Phys. Chern. 69:3955 (1965). E. Killmann and M. Kuzenko. Angewandte Makromol. Chemie 35:39 (1974). A. Takahashi et al. Macromolecules 13:884 (1980). J. F. Tassin, R. L. Siemens, W. T. Tang, G. Hadziiouanno~l~ J. D. Swalen, and B. A. Smith. J. Phys. Chem. 93:2106 (1989). J. A. Creighton, C. G. Blatchford, and M. G. Albrecht. J. Chem. Soc. Faraday Trans. TI 75:790 (1979). T. Fu and C. J. Durning, Polymer Preprints ACS Div. Poly. Chem. 31:519 (1 990). D. Inoue, H. Kurosu, Qun Chen, and I. Ando. Acta Polymer 46:420 (1995).

131. B. Fontana and S. Thomas. J. Phys. Chem. 65:480 (1962). 132. E. Killmann and M. Bergmann. Colloid Polym. Sci. 263:381 (1985). 133. C. van der Linden, R. van Leemput, and G. J. Howard. J. Poly. Sci. 34211 (1971). J. Yeh and H. L. Frisch. J. Polymer Sci. 27:149 (1958). 134. 135. A. K. Mills and J. A. Hockey. J. Chem. Soc. Faraday I 71:2384 (1975). 136. V, A. Attia and J. Rubio. Polymer J. 7:135 (1975). 137. M. M. Zwick. J. Appl. Polymer Sci. 9:2392 (1965). 138. J. C. Day and I. D. Robb. Polymer 21:408 (1980). 139. M. Kawaguchi, Yarnagiwa, Takahashi and T. Kato. J.Chern. Soc. Faraday Trans. 86:1383 (1990). 140. I. D. Robb and R. Smith. Eur. Pol. J. 10:1005, 1008 (1974). 141. G. R. Joppien. Makromol. Chem. 175:1931 (1974). 142. E. Killmann and M. Korn. Progr. Colloid Polymer Sci. 67:41 (1980). P. Tripp and M. L. Hair. Langmuir 8:241 (1992). 143. 144. D. J. Kuzmenka and S. Granick. Colloids Surfaces 31:105 (1988). 145. E, Killrnann and M. Reiner. Tenside Surf. Det. 33:3 (1996). 146. C. P. Tripp and M. L. Hair. Langmuir 9:3523 (1993). 147. F. Blurn. Colloids Surfaces 45:361 (1990). 148. T. Cosgrove, N.Finch, B. Vincent, and J. Webster.ColloidsSurfaces31:33 (1 988). 149. H. Sakai, Fujimori, and Y. Imamura. Bull. Chem. Soc. Jap. 53:3457 (1980). 150. T. Cosgrove and P. Griffiths. Adv. Colloid Interf. Sci 42:175 (1992). 151. K. G. Barnett, Cosgrove, B. Vincent, D. S. Sissons, and M. Cohen Stuart. Macromolecules 141018 (1981). 152. I. D. Robb and R. Smith. Polymer 18:500 (1977). 153. H. Ben Ouda, H. Hommel, P. Legrand. H. Balard, and E. Papirer. J. Colloid Interf. Sci. 122441 (1988); J. Chem. Soc., Faraday Trans. l 843865 (1988). 154. H. Hommel, P. Legrand, P. Tougne, H. Balard, and E. Papirer. Macromolecules 17: 1587 984). 155. M.A. Cohen Stuart, G. J. Fleer, and B. H. Bijsterbosch. J. Colloid Interf. Sci. 90:310, 321 (1982). 156. E. Killmann and K. Winter. Angew. Makromol. Chemie 4353 (1975). 157. P. Trens and R. Denoyel. Langmuir 9:519 (1993). 158. R. Denoyel, G. Durand, F. Lafuma, and R. Audebert. J. Colloid Interf. Sei. 139:281 99 159. E. Killrnann and W. Melchior. Progr. Coll. Interf. Sci. 8334 (1990). 160. L. K. Koopal and J. Lyklema. J. Electroanal. Chern. 100:895 (1979). 161. G. R. Joppien. J. Phys. Chem. 82:2210 (1978). e t ~Polarized y Light, North 162. R. M. A. Azzam and N. M. Bashara, E l l i ~ ~ o ~and Holland, Amsterdam, 1989. 163. N. V. Churaev and E. A. Nikologorskaja. Colloids Surfaces 59:7 (1991). 164. E. Killmann, J. Eisenlauer, and P. Kaniut, Naturforsch. 34a:68 (1979). 165. R. Varoqui and E. Pfefferkorn. Progress Colloid Polymer Sci. 83:96 (1990). 166. E. Killmann, P. Sapuntzjis, and H. Maier. Makromol. Chern. Macrornol. Symp. 61:42 (1992).

167. T. Kato, K. Nakamura, M. Kawaguchi, and T. Takahashi. Polynz. J. 13:1037 (1981). 168. B. Wind and E. Killmann. Colloid Polym. Sci. 276:903 (1988). 169. M.D. Croucher andT.H. Milkie,in The Eflkct Polymers ~is~e~si ~ r o ~ e r t i e(Th. ~ s F. Tadros, ed.), Academic Press, London, 1982, p. 101. 170. F. W.Rowland and F.R. Eirich. J. Polym.Sci. PartA-l 4:2033,2401 (1966). 171. D. Bauer and E. Killmann. Macromol. Symp. 226:173 (1997). 172. R. Varoqui. Nouveau J. Chirnie 6:187 (1982). 173. P. G. de Gennes. C. R. Acad. Sci. Paris 197:883 (1983). 174. M. A. Cohen Stuart, F. H. W. H. Waajen, and S. S. Dukhin. Coll. Polym. J. 262:423 984). 175. A. M’Pandou and B. Siffert. Colloids Surfaces 24: I59 (1987). Interf. Sci. 176. T. Cosgrove, M. A.Cohen Stuart, and B.Vincent.Adv.Colloid 24143 (1986). 177. L. Auvray and J. P. Coton. Macromolecules 20:202 (1987). 178. B. Cabane, Wong, T. K. Wang, F. Lafuma, and R. Duplessix.Colloid Polymer Sci. 266: 101 (1988). 179. T.Cosgrove and K. Ryan. Langmuir 136 (1990). 180. I. Caucheteux, D. Ausserre, H. Hervet, and F. Rondelez. Phys, Rev. Lett.541948 (1985). 181. 0 . Rustemeier and E. Killrnann. J. Colloid Interf. Sci. 190:360 (1997). 182. H. Sonntag, ~ e ~ r der ~ ~ c h VEBDeutscherVerlag der Wissenschaften, Berlin, 1977. 183, A. L. Smith. Spec. Discuss. Faraday Soc. 1:32 (1970). 184. N. Buske, H. Gedan, W. Katz, H. Lichtenfeld, and H. Sonntag. Colloid Polymer Sci. 258:I303 (1980). 185. P. Ludwig and G. Peschel.ProgressColloidPolymerSci.76:42(1988);ibid. 77:146 (1988). 186. H. Versmold and W. Hiirtl. J. Chem. Phys. 79:4006 (1983). 187. C. Cametti, P. Cadoastefano, and P. Tartaglia. J. Colloid Interf. Sci.131:409 989). 188. R. Midmore, Colloids Surfaces 60:291 (1991). 189. R. Arnal, J. R. Coury, J. A. Raper, W. P. Walsh,and T.D. Waite.Colloids Surfaces 46: (1990). 190. D. Axford, T.M. Herrington, and B. Midmore.ColloidsSurfaces69:73 (1992). 191. D. Bauer, E. Killmann, and W. Jaeger. Colloid Polymer Sci. 276598 (1998). 192. B. A. Matthews and C. T. Rhodes. J. Coll. Interf. Sci. 32332 (1970). 193. R. Hunter and J. Frayne. J. Colloid Interf. Sci. 76:107 (1980). 194. Lagaly, 0 . Schulz, and R. Zimehl, ~ i s ~ e r s i und o ~ eE ~~ u l s i o ~Dr. e ~ Dietrich , Steinkopf Verlag, Darmstadt, 1997. 195. E. G, M.Pelssers, M. A. Cohen Stuart, and G. J. Fleer. Colloid Interf. Sci. 137:350, 362 (1990). 196. C. Cowell, R. Li-ln-On, and B.Vincent. J. Chem.Soc. Faraday 74337 (1978).

197. J. Edwards, S. Lenon, A. F. Toussaint, and Vincent, in Polymer Adsorption and Di,~persion Stability(E. D. Goddard and B. Vincent, eds.), ACS Symposiurn Series 240:28 1 (1984). 198. J. Gregory. Colloids Surfaces 31:231 (1988). 199. D. Horn, H. Auweter, W. Ditter, and Eisenlauer, in Organic Coatings, and Technology, Vol. 8 (G. D. Parfitt and A.V. Patsis, eds.), MarcelDekker, New York 1986, p. 251. 200. B. Vincent, J. Edwards, S. Emmett, and R. Croot. ColloidsSurfaces31:267 (1988). 201. S. Vaslin-Reimann, F. Lafuma, and R. Audebert. Colloid Polymer Sci. 268:476 (1990). 202. I. F. Cuthrie and G. Howard, in Polymer A~sorptionand Dispersion Stability, (E. D. Goddard, B. Vincent, eds.), ACS Symposiurn Series 240, 1984, p. 297. 203. E. Killmann, N. Giitling, and Th. Wild, in Polymer Adsorptio~and ~ispersion Stability, (E. D.Goddard, B. Vincent, eds.), ACS Symposium Series 240, 1984, 357. 204. T.Allen, Particle Size ~easurements,Vol. Chapman and Hall, London, 1987. 205. Y. Otsubo and K. Watanabe, J. Coll. Interf. Sci. 122:346 (1988); 127:214 (1989); 133:49 1 (1 989). ion 206. J. Gregory and A. E. L.de Moor, in Polymer Adsorption and D ~ s ~ ~ r sStability, (E. D. Coddard, B. Vincent, eds.), ACS Symposium Series 240, 1984, p. 445. 207. G. S. Grover and S. C. Bike. Langmuir ]]:l807 (1995). 208. E. E.Uzgiris. Adv. Coll. Interf. Sci. 14:75 (1981). durch ase er-~oppler ~ ~ e m o m e t rund ie 209. H. Wachering, Ele~trophorese,Fa. Miitek, 1985. 210. D. C. Henry. Proc. Roy. Soc. A 233:106 (1931). 21 1. H. Ohshima. J. Colloid Interf. Sci 168:269 (1994); ibid. 185:269 (1997). 212. P.-H. Wiersma, A. L. Loeb, and J. Th. G. Overbeek. J. Colloid Sci. 22:78 (1966). V. Deryaguin, ed.),Consult. Bureau, 213. S. Dukhin, Research in Surface Forces 1966. 214. R. W. O'Brien and L. R. White. J. Chem. Soc. Faraday Trans. I1 741607 (1978). 215. H. Sonntag, B.Ehrnke, R. Miller, and L.Knapschinsky,in The Ee'ct of Polymers Dispersion Properties (Th. F. Tadros, ed.),AcademicPress, London 1982, p. 207. 216. J. Klein, Y. Almog, and P. Luckham,in Polymer A~sorptionand Dispersio~ Stability (E. D. Goddard, B. Vincent, eds.), ACS Symposium Series 240, 1984, p. 227. 217. P. F. Luckham and J. Klein. J. Chern. Soc. Faraday Trans. 86:1363 (1990). 218. J. N. Israelachvili, R. K. Tandon and L. R. White. J. Coll. Interf. Sci.78:432 (1 980). 219. H. J. Tauton, C.Toprakcioglu, L. J. Fetters, and J. Klein. Nature 332:712 (1988). 220. G. P.vanderBeek, M. A. Cohen Stuart, and Cosgrove.Langmuir7:327 (1991). 221. E. Kokufuta, S. Fujii, Y. Hirai, and I. Nakamura. Polymer 23:452 (1982). 222. Th. F. Tadros. J. Colloid Interf. Sci. 6436 (1978). 223. J. Rubio and J. A. Kitchener. J. Colloid Interf. Sci., 57: 132 (1976).

5

E Killmann and M. Bergmann. Colloid Polymer Sci. J. M. Herd, A. J. Hopkins, and G. J. Howard. J. Polym. Sci. C M. Kawaguchi, K. Hayakawa, and A. Takahashi. Polymer J. T. Miyamoto, Tomoshige, and H. Inagaki. Polymer J. J. V.Dawkins, M. J. Guest, and G. Taylor,in The ~ o l y ~ e r ~ (Th. F. Tadros, ed.), Academic Press, London, p. H. J. Terashima. J. Coll. Interf. Sci. U.Almog and J. Klein. J. Coll. Interf. Sci. D. T. Wu, Yokoyama, and R. L. Setterquist. Polym. J. R. Botham and C. Thies. J. Coll. Interf. Sci. E. Dietz. Makromol. Chem. G. J. Howard. Appl. Poly. Symp. Vu. Lipatov, L.M. Sergeeva, T.T. Todosiichuk, and T. S. Khramova. Translation from Kolloidnyi Zh. M. Kawaguchi, A. Sakai, and A. Takahashi. Macromolecules K.K. Kalnin’sh, A. N. Krasovskii, B. G. Kalen’kii, and G. A.Andreyeva. Vysokornol. Soyed. H. Hommel, A. P. Legrand, H. Balard, and E. Papirer. Polymer E. Killmann and P. Sapuntzjis. Colloids Surfaces A: Physicochem. Eng. Aspects

M.Kawaguchi, M. M. Mikura, and A. Takahashi.Macromolecules, F. Lafuma, K. Wong, and B. Cabane. J. Colloid Interf. Sci. G. A. Schumacher and Th. G. M. van de Ven. Faraday Discuss.Chem.Soc. H. Oshima, T. W. Healy, L.R. White, and R.W. O’Brien. J. Chem. Soc.Faraday Trans. M. J.Garvey, Th. F. Tadros, and B. Vincent. J. Colloid Interf. Sci. M, A. Cohen Stuart, F. H, H. Waajen, T. Cosgrove, B. Vincent, and T. L. Crowley. Macromolecules G. Durand-Piana, Lafuma, and R. Audebert. J. Colloid Interf. Sci. (1 T. P. Golub, I. K. Dorofeeva, and M. P. Sidorova.Colloid J. T. Okubo. Polymer Bull. R. Sprycha, H. T. Oyama,

Zelenev, and E. Matijevic. Colloid Polymer Sci.

G. P. Van der Beek, M. A. Cohen Stuart, and J. M. H. M.Scheutjens. J. Colloid Interf. Sci. G. P.VanderBeek, M. A.Cohen Stuart, and G. Fleer.Macromolecules G. P.VanderBeek, D.D. G.

M. A.Cohen Stuart, and G. J. Fleer.Langmuir 5:1180

Leonhardt, H. E.Johnson,

and

Granick.Macromolecules

Howard and S. J. Woods. J. Polymer Sci.

er

255. M. Kawaguchi, Y. Sakata, S. Anada, T. Kato, and A. Takahashi.Langmuir 10:538 (1994). 256. Th. W. Healey, in The Colloid Chemistry Silica E. Bergna, ed.), Advances in Chemistry Series, ACS Washington D.C., 1994, p. 147. 257. S. R. Kline and E. W. Kaler. Langmuir 10:412 (1994). 258. V. Gurfein, Beysens, and D. Perrot. Phys Rev. A 40:2543 (1989). 259. Z.Kira’ly,L.Turi, I. Dekany, K. Bean, and B. Vincent.ColloidPolym.Sci. 274779 (1996). 260. E. Killmann and Eisenlauer, in The Ee’ct cfPolymers Dispersion Properties (Th. F. Tadros, ed.), Academic Press, London, 1982, p. 221. 261. A. Lips and E. Willis. J. Chern. Soc. Faraday 69:1226 (1973). 262. Cahill, P. G. Cummins,E. J. Staples, and L. Thompson. ColloidsSurfaces 18:189 (1986). 263. B. Derjaguin. Pure Appl. Chem. 48:387 (1976). 264. R. H. Smellie, and K. La Mer. Colloid Sci. 13:589 (1958). 265. T. K. Wang and R. Audebert. J. Colloid Interf. Sci. 119:459 (1987). 266. B. Cabane, K. Wong, and R. Duplessix, in Polymer Association S t r ~ c t ~ (M. r e El Nokaly, ed.), ACS Sy~posiumSeries 384, 1989. 267. R. K. Iler. J. Colloid Interf. Sci. 51:388 (1975). 268. John H. van Zanten, Menachem Elimelech. J. Colloid Interf. Sci. 154:l (1992). 269. L. Wigberg and T. Lindstrom, Nordic Pulp Paper Res. 249 (1987). 270. M. H. Schentjens and G. Er. Fleer. Macromolecules 18:1882 (1985). 271. Rossi and P.Pincus. Macromolecules 22:276 (1988). 272. Y. Marra, M.L. Hair. Phys. Chem. 92:6044 (1988). 273. C. v. d. Linde, thesis, Universite Libre de Bruxelles, (1976). 274. Killmann, C. Fulka, M. Reiner. J. Chem. Faraday Trans. 86:1389 (1990). 275. U.Spiegler, thesis, Technische Universitat Miinchen, (1999).

This Page Intentionally Left Blank

Institute of SurfaceChemistry, National Academy of Sciences, Kiev, Ukraine

I.

524

11. Methods for Activating Surfaces of Functional Organosilicas Activation A. of a~ine-containingsilica matrices B. Activation of thesurface ofvinylsilicasbymaleic anhydride C. Organosilicas with grafted mercapto groups and activated matrices on their basis D. Activation ofsilica matrices with grafted silicon hydride groups 111. Physicochemical Properties ofEnzymes Immobilized on Functional Organosilica A. Influence of geometrical characteristics of silica matrices on the enzyme kinetics of immobilization B. Influence of pH of amediumonimmobilization of proteins C. Activity and stability of immobilized preparations IV.

Application of Enzymes Immobilized on Silicas for Analysis of Various Application of Immobilized Biocatalysts for Synthesis of Active Compounds A. Role of papain in reactions of formation bonds of peptide Synthesis of Leu-enkephalin using immobilized papain

VI. Adsorption and Chemisorption Equilibrium for Proteins on Surfaces

524 526 533 534 537 537 537 539 540 543 546 547 550

Silica 552

VII. Some Up-to-Date Investigations in the Field of Biologically Active on Ix~mobilized Cornpounds 556 Silicas VIIL.

Conclusions References

557 558

Over the last few yearsinorganic carriers, especially thosebasedondispersed silicas, have been used widely not only for transfer of many organic ligands and metal complex cornpounds [ 1 4 ] and important analytical reagents [5,6] to a heterogeneous state, but also for immobilization of enzymes, coenzymes, hormones, antibodies, antigens, and other biologically active macromolecules [7-101. As distinct from organic matrices, silica carriers offer a number of advantages owing to their mechanical,microbiological,chemical, and thermal stability, feasibility of regeneration, as well as rigidity of their skeleton, which does not change its structure at various values of pH and ionic strength of diverse solvents. Silicas lend themselves readily to geometricalmodification, which allowsone to produce matrices with the optimum structurecharacteristics, namely pore diameter,specific surface area, and particle size. Chemical stability of inorganic carriers makes it possible to effect addition of cross-linking agents in organic solvents, which extends substantially the range of feasible surface activation reactions. As is known, preparations immobilized on silica have typical advantages that are caused by their capability of transferring active compounds to a heterogeneous state, and manifest themselves in ease of separation from reaction mixtures, possibility of repeated use and automation of processes. In addition, such preparations are often characterized by higher stabilities in comparison with native enzymes, which obviously is a matter of interest for practical purposes. It isbelieved thatthe increased stability of enzymes is attributed to multiple-point bonding of protein with surface sites of an activated carrier, which prevents unfolding of polypeptide globules(for instance in a thermal treatment). Of no small importance is also the fact that enzymes immobilizedon silica are suitable to study their kinetic characteristics at high concentrations of an organic solvent, whichis impossible to do in thecase of soluble enzymes. In view of this, when working with immobilized preparations in a number of instances it is possible to gather important data furnishing a better insight into the m e c h a n i s ~of action of biocatalysts. The aim of this chapter is to analyze the experience acquired in developing efficient and versatile enough methods for activation of surfaces of functional silicas; to discuss results of research into adsorption and chemisorption equilibria for enzymes and a number of other natural macromolecules on activated silica matrices; to consider possibilities of application of immobilized-on-silica preparations of biological catalysts in analytical practice and fine organic synthesis.

Although during the earliest work devoted to immobilization ofenzymes some success was attained using unmodified silica matrices for adsorption of a corresponding enzyme, researchers cameto a conclusion that itwas necessary to provide a stronger binding between protein and surface [l 1-13]. In particular, more stable preparations were produced by applying approaches that involved cross-linking. For example, some enzymes whose adsorption was often effected in combination

with other preliminarily adsorbed proteins couldbe cross-linked by glutaraldehyde [14]. In a number of cases a success was scored when employing silica matrices impregnated with solutions of transition metal salts, and attachment of protein macromolecules to these matrices seemed to proceed by the co-ordination mechanism 1151. Use was also made silica matrices that prior to adsorbing a suitable protein were treated with cyanogen bromide or cyanuric chloride 16,171. Although the mechanism of reactions proceeding during such a treatment with participation of silanol surface groups hasnot been studied in detailyet, one can contend thatin most cases the treatment is accompanied by formation of hydrolytically unstable bonds Both in acidic and basic media such bonds are very susceptible to breakdown, which may be the cause of washout of an attached enzyme. A more promising way to produce active and stable preparations containing immobilized enzymes consists of applying modified organosilicas, i.e., silicas with an organic layer that is chemically grafted to the surface through hydrolytically stable bonds. Application of such ~ o d i f y i n gcoatings is most convenientlyeffected at the expense of formation of Si -Si and Si- C bonds that are stable in acidic and weakly basic media. Additionof organic radicals with various functional groups to surface of silicas (in order to produce functional organosilicas) usually involves reactions between structural hydroxyl groups of the surface and appropriate alkoxy- or halogenorganosilanes: ESi

OH R,Si&-,

Si

Q

SiR,(X3-,)

HX

where R is the organic radical bearing a functional group,X is the alkoxy group or halogen, n 1-3. In this casetheaddition of a necessary functional group is determined in practice by the choice of a suitable organosilane that is added to the carrier through the electrophilic substitution reaction of a proton in a silanol group on the surface[3]. Most commonly appliedin the past few years was also the method for production of surface chemical compounds withSi C bonds directly silicas at the expense of reactions of solid-phase catalytic and thermal hydrosilylation of functional olefins:

where R is the organic radical bearing a functional group [18]. As is evident, the introduction of a necessary functional group is governed in this case by the choice of a suitable olefin that is attached to the carrier surface at the expense of the reaction of nucleophilic addition through asilicon atom of a silicon hydride group. Under favorable conditions the functional organosilicas can be used to immobilize various compounds without any additional activation of surface. In most cases, however, fortheir attac~ment,the modified silica surface should have groups that under mild conditions are ableto effect bonding between a biologically active compound and the matrix surface. Further in this section the discussion concerns the results achieved by us lately in developing methods foractivating the functional organosilica surface whose modifyinglayer bears amine, vinyl, mercapto, and silicon hydride groups.

Immobilization of biologically active compounds is mostcommonly effected throughapplication of amine~containingorganosilicas(amino-organosilicasj. In this situation the introduction of amine groups into a modifying organic coating is made through treatment of the surface with suitable alkoxyorganosilanes, for example with ~"aminopropyltriethoxysilane119-1211. Earlier, during studies by IR spectroscopy techniques itwas shown that the grafted amine groups formed strong hydrogen bonds involving nitrogen atoms and proton-co~tainingsites of the silica surface (residual silanol surface groups, adsorbed water molecules) [22-241. Later on, this inference was corroborated by systematic research into the grafted amine group state applying the 29Siand I3C NM spectroscopy [4,25--371. Although the chemical activity of amine groupsin complexes with hydrogen bondsdecreases, the functional groups of the surface quite readily enter into a reaction with vapors of acetic or methacrylic acid. Using IR spectroscopy it hasbeen found that in this case on the surface one can observe formation of surface chemical compoundsof the salt 2NHZ *-OOCRj thatdecompose at 220-250°C 121,241. In addition, studies carried out by many researchers also showed that the bonded functional groups of the surface underwent other chemical transformatio~scharacteristic of amine groups.

ining silicas without any additional activationis amenable to a direct imnlobilization of compounds whose composition includes aldehyde,isocyanate,isothiocyanate,carboxyl,epoxygroups, or labile atoms of halogens bound to carbon. It has been found that participation of aminopropyl groups substantially facilitates proceeding of reactions with organic compounds C- I or C Br bonds, for instance, with monoiodoacetic acid or oylpyroracemicacidmethylether 138-421. Formation of azomethine bonds between aminopropyl surface groups and aromatic aldehydes of various chemical constitutions (benzaldehyde, salicylaldehyde, p-oxybenzaldehyde, p-resorcy1 aldehyde,protacatechuicaldehyde, vanillin, ~"nitrobenzaldehyde,and pdim~thylaminobenzaldehyde)makesitpossible to immobilizecompounds that regarded as analogs of substrates and inhibitors of polyphenol oxidase After the reaction with aromatic aldehydes the infrared absor tion spectra disappearance of absorption bands at 3290 and 3360cm-'forvalence vibrations of N in primary amine groups of the modified silica and appearance of characteristic absorptionintheregion of1640-1650cm" assigned to azomethine bonds in the surface compounds formed. In the case of the silochrom surface this method makes it possible to provide attachment of 0.20-0.25 mm01 of benzaldehyde and its derivatives per 1 g of the carrier (80 m2/g). It is significant thatafterthereduction of Schiffs bases withsodiumborohydride or sodium hydrosulfitetheabsorptioninthe spectral regionof1640-1650cm-lis not observed. The surface cornpounds formed have the following constitution:

where R -H or -OH; R' -H, -OH or -OCH3; R" or -N(CH&. Such matrices with attached aromatic aldehydes allow one to effect successfully noncovalent immobilization of polyphenol oxidases of vegetable and microbial origin, which seems to be attained owing tobiospecific adsorption of the enzyme through sites that are close in their constitution to substrates or inhibitors of the biocatalyst. Aldehyde groups canbe introduced into the compositionof glycoproteids by the oxidation of their carbohydrate component. For example, it has been shown that the carbohydrate component of catalase ~ e ~ j c j Z Z i u ~ is efficiently subjected to oxidation under the action of solutions of periodic acid or potassium periodate:

It is characteristic that at the oxidation of 30% of carbohydratesthe enzyme completely retains its activity. Thermal stability, optimum pH, and pH-dependence of its stability do not practically change, but its storage stability at 4 4 ° C substantially increases. Thus, the catalase with the oxidized carbohydrate component is a stable active enzyme containing aldehyde groups that can be utilized for covalent immobilizationdirectly on the amino-organosilica surface owing to formation of azomethine bonds between the carrier and biocatalyst SiCH2CH2CH2Nx CH enzyme. In comparison with the native catalase the specific catalytic activity of the immobilized enzyme is higher, which indicates stabilization of the biocatalyst during its attachment to surface. the The immobilization through theoxidized carbohydrate Component permits oneto produce preparations with an activity of up to 2000 Becker units per 1 g of the carrier,which was impossible when applying other methods for bonding. These results give evidence for certain advantages of the immobilization through the oxidized carbohydrate component, namely that this method allows oneto attachhigher amountsof an enzyme as against its bonding through amine acid residues, and, besides, when stored in suspension the preparationretains its highactivity for morethan ayear. In addition, as to the optimum pH, temperature, and pH-dependence of its stability the immobilized catalase does not practically differ from the soluble one. Whenasolutioncontainsgalactoseoxidase anditssubstrate (galactose), a peculiar process of autoimmobilization of the enzyme on the amino-organosilica surface takes place. In this situation, under the action of a biocatalyst galactose is transformed into 2,4-galactohexodialdose,which serves as a cross-linking agent and provides bonding of the enzyme with the amine-containing silica matrix sur~usariu~ this method face. Whenusinggalactoseoxidasefrom

enables one to produce immobilized preparations withan activity of up to 90 units per l g of a carrier. Such preparations weresuccessfully appliedfor analytical determination of galactose-containing carbohydrates (galactose, lactose, raffinose) in complex mixtures [47,48]. lmmo~iiizationof nzymeson Activated Amino-~r~anosilicas ~ ~ t h o d s ~ ~ Since ~ 1970, the efforts of many researchers have resulted in the development of various methods for activation of the aminoorganosilica surface, which are suitable to provide bonding of biologically active compounds under mild conditions. Special note should be taken of the methods that are based on the application of glutaraldehyde, diazonium salts, water-soluble carbodiimides, isothiocyanates, acylazides, haloalkyls, and some other cornpounds (the literature is covered in detail in our review paper [7]). The method for activation of the amino-organosilica surface by glutaraldehyde has become mostwidespreaddue toits relative simplicity and availability of reagents. It isbelieved that various enzymes are attached to such an activated surface through amine groups of amino acid residues:

zCH2CHzNz CHCH~CHzCH2C It should be taken into account that glutaraldehyde preparations practically do not exist in a monomeric form, and, in fact, surface cornpounds have a more complex constitution. The azomethine bonds formed during immobilization of enzymes are stable in neutral and alkaline media, but in acidic media they undergo hydrolysis. In a number of cases they are efficiently reduced by sodium borohydride. pplication of the azocoupling reaction for attachment ofenzymes involves synthesis of cross-linkeddiazonium salts onthesurface at theexpense of the intro~uctionof nitro groups (with the aid of ~-nitrobenzoylchloride),their subsequent reduction by sodium hydrosulfite, and diazotization by nitrous acid. The surface compounds formed have the following constitution: At neutral or weakly basic values of pH the diazonium salt attached to the surface under mild conditions canreact with various groupsof an enzyme, in the first place with phenolic residues of tyrosine, imidazolic residues of histidine, with and amine groups of terminal amino acids, sulfhydryl groups of cysteine, indolic residues of tryptophan, guanidinic grouping of arginine, as well aswithhydroxyl groups of serine and treonine. ~ l t h o u g hthe activation of the amino-organosilica surface involves multistage processes, the method is universal enough and applicable to a wide range of biocatalysts. Some steps have been taken to decrease the number of stages by utilizing modified silicas whose amine groups are bound to solublecarbodiimidesrepresented by thegeneralformula N=C=NR’, the activation actually affects carboxylgroups of an enzyme which thereupon is bound to the amino-organosilica surface through formation of peptide bonds

2NHC(0)-enzyme

This method is very convenient for immobilizationof enzymes that are unstable in neutraland alkaline media, since thecondensationreactionproceeds at acidic values of pH. At the same time, when carbodiimides are used in this way, polycondensation of the enzyme itself possible. To all appearances, this difficulty can be avoided by applying carboxyl-containing organosilicain the capacity of a starting matrix. After activation of such organosilicas by an appropriate carbodiimide theybecome suitable forimmobilization of protein molecules throughamine groups with formationof surface structures whose constitution can be represented as E~iRC(O)NH -enzyme

We have made a comparison[7] between activities of immobilized preparations, amounts of attached enzymes, and degrees of retention of protein activity in the final product that were determined in cases various methods for bonding. It has been shown, in particular, that when choosing a bonding procedure one should allow for the structureof a concreteenzyme or other active compound intended for attachment to the surface as well as conditions underwhich the preparation will be applied afterits immobilization. At present, there is no universal method applicable for attachmentof all the biocatalysts. Moreover, generally speaking, such a method seems to be impossible. Therefore, subsequent development of efficient methods for producing activated inorganic matrices designed for i ~ ~ o b i l i z a t i oof n enzymes and other active compounds is a task of great scientific and technological importance. It is also necessary to carry out systematic research into propertiesof active compoundsimmobilized on a solid-state matrix in order to ascertain optimum conditions of their application for practical purposes. the The experience gained by us from producing activated matrices on the basis of functional organosilicas has shown broad prospects of the idea to use reagents whose groupsof the same composition differ substantially in their chemical activity. The use of such reagents permits oneto carry out aprocess of bonding intwo steps. First,areaction is broughtaboutforamino-organosilylsurface groups of the modified silica matrix, and then functional groups of compounds subjected to immobilization enter into the reaction. It is this principle that forms the basis of the methods proposed in our work 1,523 for activation of the aminoorganosilica surface employing tolylene 2,4-diisocyanate and 2,4,6-tricloro-l,3,5triazine. by tolylene Isocyanate groups tolylene 2,~-diisocyanateboundto benzene rings in position4 are more active than NCO group in position 2. Therefore, when amino-organosilica comes into contact with solutionsof tolylene 2,4-diisocyanate in carbon tetrachloride, it is isocyanate groups in position 4 that predominantly enter into the surface reaction. The modified silica matrices activated in this way contain grafted isocyanate groups in position 2 benzene rings. Using I R spectroscopy we have established that all thegraftedaminopropylsilylgroupsenter into the reactionwith tolylene 2,4-diisocyanate. As a result, the modifying coating formed contains isocyanate groups thatin the 1R spectrum are characterizedby an intensive absorption

ii

band with the peak maximum at 2290 cm-'. During the subsequent stage these isocyanategroupscan beused forsurfaceimmobilization of compoundswith amine groups undersufficiently mild conditions. The structure of the surface compounds formed can be represented by the following scheme: ESiCH2CH2CH2NHC(0)NH NHC(0)NH -Enzyme

The isocyanate groups attached to the amine-containillg silica surface with the aid of tolylene 2,4-diisocyanate exhibit a higher activity in reactions for bonding of enzymes than that of the isothiocyanate groups formed during treatment of aminoorganosilicas with thiophosgene:

Of importance is also the fact that when activating by tolylene 2,4-diisocyanate it is not necessary to utilize the activation reagent excess, and the groupings of the urea type are hydrolytically stable. Below we give concrete examples of effecting processesforactivation of thesurface and subsequentimmobilization of an enzyme in thecase of covalentbonding of a-amylasefrom oryzae 153,541. Amino-organosilica was prepared from a nonporous pyrogenic silica witha specific surface area of about 175 by boiling with ~-aminopropyltriet~oxysilane in a toluene solution for 2 h. The silica was preliminarily dried at 400°C for2h.Afterthewashout of themodifyingreagent excess withtoluenethe sampleobtained was driedundervacuum at 100°C. According to the data of potentiometrictitrationwithperchloricacidthecontent of aminopropylsilyl groups is equalto 0.7 mmol/g.Next, 2.5 g of aminopropyl-Aerosil was suspended in 25 mLcarbon tetrachloride, then 0.39 tolylene 2,4-diisocyanate dissolved in CC14 was added.Themixture was stirred atroomtemperature for 5 min and washedwithcarbon tetrachloride. Thesampleprepared was driedundervacuum at 100°C. The carrier activated in this way can be either used directly for bonding with an enzyme (or other active compound) or stored in sealedglass ampules. The enzyme immobilization on the activatedcarrier was effectedin the following way: 0.25 a-amylase with a protein contentof 96-98% and amylolytic activity of 280,000 units per gram was dissolved in 10 mL of 1 M phosphate buffer at pH7.0. Then the solutionwas slowly stirred while 1 g activated carrier was quickly added. The bonding of the enzyme with the matrix surfacewas effected at room temperature for 10 min. Further, the obtained preparation with the immobilized enzyme was washed in turn with l M phosphate buffer at pH 7.0 (200 mL) and 0.005 M calcium acetate at pH 5.5 (200 mL) to disappearance of amylolytic activity of the washing liquid. The product obtained was dried at 4°C and stored at the same temperature.Theproteincontent of driedpreparationswiththeimmobilized

enzyme amounted to 8-12%, and their amylolytic activity was equal to 800-1250 units per 1 g. min no-organosilicas activated by tolylene 2,4-diisocyanate were also applied for the anchorage of catalase [45,46], urease papain and pronase E enzymic preparation-amylosubtilin G10x and cells of methane-oxidizing bacteria [60]. ~ c t i v a t i o ~ anzi~~o-organosilicas by cyanuric c ~ l o r i ~ e .In a six-member triazine ring of cyanuric chloride in position 2,4,6 there are groups with highly labile chlorine atom. Interaction between one of these bonds (C-Cl) of the activating agent and a grafted aminopropyl group can result in formation of chemicalcompoundsattachedtothesurface by hydrolytically stable ECNH C bonds. spectra of aminopropylsilicas treated with cyanuric chloride do not display bands of absorption of valence and cm-') 1 S90 cm-') vibrations of N- H bonds in primary amine and deformation groups. Sideby side with this, the presence of absorption in a region of 16601400cm-'' (vCzN and characteristic bands of a triazine ring (1600 and 1560 cm") provides evidence for addition of cyanuric chloride to aminopropyl groups of the surface. It is significant that if one of the groups of cyanuric chloride molecule enters into the reaction, the activity of the rest of the C-Cl bonds in the surface chemical compoundsformed slightly decreases. Atthesame time, the mobility of chlorineatoms in such C-C1 bondsremains sufficiently high to allow bonding ofenzymes and other amine-containing biologically active compounds under mild conditions:

neof themajoradvantages of themethodforactivation of thearninoorganosilicasurfacewiththeaid of cyanuricchloride is that in this case process of immobilization of appropriate compounds can be effected at low values of a medium. It important, for example, for covalent proteinases thatare unstable in neutraland alkaline media. concrete examples of effecting processes for activation of the amine-containing silica surface by cyanuricchlorideandsubsequentimmobilization of an acid proteinase (pepsin). First, l 0 g of anlinopropylsilica prepared from macroporous Silochrom with a specific surface area of ca. 80 m2/g was suspended in 50 mL of dioxane. The S0 mL of 5% solution of 274,6-trichloro-1,3,5-triazine in dioxane was added,andthe mixture was agitated using a magnetic stirrer at room temperature for 30 min. The product was rinsed in dioxane to wash out the cyanuric chloride excess (50 mL 4) and dried under vacuumat 100°C. In the case of aminopropyl-~erosil,l g of the substance was suspended in S% solution of 2,476-tricloro-1,375~triazine in a dioxane-toluene mixture (4: 1 by volume, 15 mL). The mixture was agitated uisng a magnetic stirrer for 30 min. The productwas washed with dioxane(50 mL 4) and

dried under vacuumat 100°C. Before utilization the silica preparations activatedby cyanuric chloride were stored in sealed glass ampules. One gram of aminopropyl-~erosilactivated by cyanuric chloride was put in a solution (27 of pepsin (460 mg) in 0.1 M acetate buffer (pH and the mixture was stirred for 1 h at room temperature and for 18 h at 4°C. To remove physically adsorbed enzyme from the silica matrix the preparation obtained was washed withl M aqueoussolution of sodiumchlorideup to disappearance of absorption at 280 nm in washing water and then with 0.1 M acetate buffer (pH 4.5) up to absence of chlorine ions. The activity of the immobilized pepsin with respect to hemoglobin was 155,000 units per gram of the dry preparation. When using a~inopropyl-Silochromactivated by cyanuric chloride, 1 g of the carrier was put into a solution (7.5 of pepsin (46 mg) in 0.1 acetate buffer (pH and this step was followed by the procedures applied in the situation of the enzyme attachment to the activated aminopropyl-Aerosil surface. In the case of Silochrom the activityof the immobilized pepsin withrespect to hemogloblinwas 5 1,000 units per gram of the dry preparation. The preparations of immobilized pepsin were stored in 0.1 M acetate buffer (pH 4.5). min no-organosilicas activated by cyanuricchloride were also used forthe immobilization of trypsin [61-631, catalase [45,46], urease [55,64], andpapain [56,65,66]; enzymic preparation-amylosubtilin ClOx [59], imidazole, and histamine [38,39,42] includingforthefollowingco-ordinationattachment of hemin 167,681, acetylacetone [39-42], thiourea, thiosemicarbazide, aminothiazole, and sulfanilic acid 169-721, tiva vat ion of a~ino-or~anosili~as by ~ i ~ ~ l o a l k a n eIns . number of a cases it is convenient to activate the surface of amino-organosilicas by di~aloalkanesof the general formula X(CH2),X, where X stands for bromine or iodine and 2-10 hen the surface activationwas effected by dibromoet~anethe concentration of the grafted ES i ( ~ ~ 2 ) ~ N H ( C Hgroups ~ ) 2 ~was r 0.10 mmol/g. In the case of other dibromoalkanes (1,6-dibromohexane and 1 10-dibromodecane) the obtained concentrations of the grafted S i ( C H ~ ) ~ ~ H ( C Hgroups 2 ) , ~ ~were r approximately the same (0.12-0.14 mmol/g). If use was made of more active corres~ondingdiiododerivatives of alkanes (1,2-diiodoethane9 1,6-diiodohexane~ and 1,10-diiododetion was carried out at 20-60°C, the concentrations of the e groups attached to the same matrix surface amounted to 0.200.25 mmollg. The treatment of the surface of amino-organosilicas with dihaloalkanes made it possible to produce activated matrices that were suitable for efficient immobilization of acetylacetone and other diketones [38-41], of many amine-containing organic omp pounds including organic diamines[38,67,68], sulfur-containing ligands [71,72], and of some amino acids 1741. Instead of dihaloalkanes for activationof the surfaceof amino-organosilicas one can use other bifunctional compounds, for example, epichlo~ohydrin. Inthis case, however, two types of reactions (additionand substitution) proceed equally according to the schemes:

The grafted C-Cl groups and especially epoxy groups canbe successfully applied for covalent attachment of biologically active compounds.

In order to produce activated matrices one can use not only amino-organosilicas but also organosilicaswithotherfunctionalgroups. In this respect, of serious interest are, in particular, olefin-containing silicas that in the presence of organic peroxides and other initiators or under the actionof ionizing radiation are ableto add compounds having double bonds [24]. It is by copolymerization of vinylsilica and methacrylic acid that modified silicas with grafted carboxyl groups have been synthesized [75]. We have described [61] a method for producing activated silica matrices for immobilization of enzymes. This method is based on a reaction betweenvinyl groups and maleic anhydride. Although the control over proceeding of radiationinduced cograft polymerization is impeded, the general scheme of the chemisorption processes can be represented as follows: =Si

-CH= CH,

=Si

To carry out the synthesis, 20 g of Silochrom (dried at 400°C for 2 h) was suspended in 50 mL of 6% solution of vinyltrichlorosilane in carbon tetrachloride and kept at room temperature up to the complete evaporationof the solvent. The silica modified in this way was put in ampules and suspended in 50 mL of l5Y0 solution of maleic anhydride in acetic anhydride. Then the ampuleswere evacuated and sealed. These steps were followed by radiation-induced copolymerization. The samples obtained were washed with water up to the absence of absorption at 220 nm and dried in vacuum at 100°C. Immediately beforeenzyme immobilization the activated carrier was heated at 130°C for 1 h. Consider a specific example of applying the method. Immobilization of trypsin waseffected in the following way. First, 0.6 g of vinyl-Silochrom activated by maleic anhydride was put into solution (5 mL) of trypsin (20 mg/mL) in 0.1 phosphate buffer (pH 7.6). Then the mixture was shaken for h at 20°C. The produce was washedwith 0.05 tris-HC1 buffer (pH 8.2) containing 0.02 CaC12 up to disappearance of absorption at 280 nm in washing water and then with distilled water and 0.2 solution of CaC12. The immobilized enzyme was

stored in 0.02 M aqueous solution of CaClz. The amount of the enzyme attached to the carrier was calculated from the difference in the amount of the enzyme takenforimmobilizationandwashed outfromthe surface after bonding (in terms of absorptionat 280 nmj.Thepreparations of trypsin immobilized on matrices with grafted polymaleic anhydride exhibit the highest thermal stability and can function for a long time even at 65°C. Of special interest is the effect of a sharp increase in the specific activity of trypsin immobilized in this way in the presence of calcium ions. Thus, when observing this phenonlenon [61], we noticed that the activity the preparation immobilized in the presence of calcium ions increased by a factor of about 3, while in the case of the native biocatalyst the degree of activation was approximatelyonly 20%. This givesevidence fora substantial role of the activated carrier in increasing specific activity of the immobilized trypsin. The mechanism the effect observed has not yet been completely elucidated. Such a considerable activation islikely to be caused by a favorable arrangemelit of the immobilized trypsin on the surface, or related to the interaction between calcium ions and carboxyl groups of the carrier, or else attributed to a joint action of these factors.

Silicaswhose modifying layer containsmereaptogroupsarepromising in two major aspects. First of all, suchmatricescan be employedfor purification of proteins containing free groups by the covalent chromatography technique [76]. In addition, such carriers are suitable for producing i~mobilizedenzymes attached to surface through disulfide bridges [77]. The most simple method for introducing sulfhydryl groups into a modifying silica layer involves application of appropriate al~oxyorganosilanes, forinstance, ~-nlercaptopropyltrimethoxysilane 1781. Theauthors of Ref. 78 studiedthe kinetics of chemisorption of ( C ~ ~ O j ~ S i C ~ 2 C H z Cthrough silanol groups 2SH of the silica surface. Since the control over proceedingof reactions of electrophilic substitution of protons in s S i O H groups was exerted by spectral technique only in terms of decreases in the intensity of the band at 3?50 cm-', its results are ambiguous due to possible side reactions with participation of methanol molecules. The band of stretching vibrations of bonds in the range of 2600 cm"' has a rather low intensity, and, therefore, application of the spectral data to determine concentrations of the grafted functional groups difficult. At thesame time, identi~cationand numericalevaluation of bondedmercapto groups can be effected reliably enough with the aid of Ellman's reagent dithiobis~2-nitrobenzoicacid). The application of this reagent and availability of surface mercapto groups facilitate the process of the quantitative reaction of the thiol~isulfideexchange. As a result, in solution one can observe formation of pnitrothiobenzoate anions characterized by the intensive absorption in a range of 412 nm:

COOH =SiCH,CH,CH,S-S

H

The content of~-n~trothioben~oate anions in solution is completely consistentwith the concentration of sulfhydryl groups attached to the silica surface. Modification of silicas with ~-mercaptopropyltrimethoxysilane andsubsequent activation of thesurfacecan be effected in thefollowingway.First, 10 g of Aerosilwitha specific surfacearea of175 m2/g preliminarilydried at 400°C for 2 h was suspended in solutions of ~-mercaptopropyltrimethoxysilane in 100 mL of toluene. The mixture was stirred and boiled in a flat-bottomed flask with a backflow condenser. In doing so, periodical sampling of the Aerosil sus~ension was carried out. The samples taken were filtered to separate the silica. Then, in orderto removeuncombined ~-mercaptopropyltrimethoxysilane the modified Aerosil was washedwithtoluene and dried in vacuum at 100°C. Ina similar way onecan also effect chemicalmodification thesurface of Withotherconditionsbeingequal, in this casethestarting carrier mass is twice as high as when Aerosil is used. ~ e t e r ~ ~ i n a t iofo nconcentrationofmercaptogroupsattachedtosurface matrices (as well as of activation of such modified silicas) involved the following procedures. First, 0.5 g of Aerosil (1 g of Silochrom) with a chemically modified surface was suspended in 500 mL 0 M solution of 5,5'-ditllio~is(2-nitroben~oic 8.0). After1hthe carrier wasfiltered and acid) in 0.1 phosphatebuffer washed with 500 mL of the same buffer with an additive of 0.5 mol of sodium chloride. Then the filtrate and washing solution were mixed, and optical density of the absorption band with the peak maximum412 at nm was measured. The grafted group concentration was calculated with allowance for the molar extinction coefficient being equal to 14,140 1791. From the data gathereditfollows that carrying out reactions in boiling solutions of ~-mercaptopropyltrimethoxysilane in toluenefor 3-4 hmakesit possible to attach sulfhydrylgroups to the fumed silica surface, withthecon130 pmol/g. Silochrom modified underthe tent of suchgroupsamountingto sameconditionscontains 73 pm01 of graftedrnercaptogroupsper1g of the sorbent. The principle of determination of concentration of grafted mercapto groups with theaid of Ellman'sreagentformsthe basis of themethodfor activating SH-containing organosilicas designed to produce matrices for immobili~ationof thiol-containillg proteins and enzymes. The carriers activated in this way enable one to effect bonding of thecompoundsmentionedusingthe thiol-disulfide exchange reaction:

COOH

ESiCH2CH2CH2S S

E S i C H 2 C H 2 C H 2 SS -EnzymeH

HS-Enzyme

S

In particular, in order to immobilize urease on the thiol-containing matrix activated by Ellman’s reagent, 90-100 mg of the carrier was put into a column, then 10 of the enzyme solutionin 0.1 phosphate~itratebuffer was drawn through the column for 4 h, following which the column was washed with 40 of NaCl solution in the same buffer. The amount of the biocatalyst attached was calculated in terms decreases in the enzyme activity in solution.The attachedurease activity was determined in thesamecolumn by thepH-stat method. The pH dependence of urease activity was studied at 25°C. The kinetics of hydrolysis of urea was investigated at pH 7.0 and 25°C. It was established that the most active urease preparations from microorganisms phyticus L-l can be synthesized by carrying out the process of il~mobilization at 7-43. In the case of activated Silochrom with a particle size of 0.250-0.365 mm the process of the enzyme bonding is completed within 4-5 h. In order toincrease activity of preparations it is necessary to add disodium salt of ethylenediaminetetraacetic acid to solutions for immobilization (1 mmol/L). It has been also shown that the degree of retention of urease activity depends to a considerable extent on thelength of an organic radical which bindsa biocatalyst molecule to carrier surface. judgedfromthemethodfor its extraction,themicrobialurease belongs to enzymes that function in a livingcell in the dissolved state. In such a case removal of a biocatalyst from surface can be regarded as transer to its natural condition. With increasing spacer length, resemblance between conformations of native and immobilized enzymes becomes more and more close. Thus, when applying the same method for bonding urease to surface through disulfide bonds, under favorable conditions it is possible to succeed in producing preparations that completely retain the activity of a soluble enzyme Here it should be noted thatthe increase of the specific activity of agrafted enzyme with increasing distance from the surfacewas also observed for other enzymes (trypsin, amylo~lucosidase, chymotrypsin) that function in living organisms in solution. It is quitenaturalto assume thatduring immobilization, men~brane-bonded enzymes canincrease their specific activity not when their macromolecules are movingawayfromthematrixsurface but,on thecontrary,whentheyare approaching it. In the generalcase, it may be asserted that the degree of retention of the activity of a biocatalyst is determined not only by the method for bonding it to a carrier but also by the length of an organic radical between the enzyme and matrix surface.

In recent yearsit was shown that the solid-phase hydrosilylation reactions on silicas could be successfully employed for synthesizing hydrolytically stable surface chemical compounds with silicon-carbon bonds [18,80]. Of special interest is application of this method for attachmentof functional olefins, in particular of acrylic acid and acrolein, which allows one to produce activated silica matrices with carboxyl and aldehyde groups. Such matrices can be used for subsequent irnmobilizationof a wide range of amine-containing organic reagents and biologically important cornpounds. Using IR spectroscopy it has been found that solid-phase catalytic and thermal hydrosilylation of acrylic acid can beeffected at high pressures of the reagent, although the hydrosilylationmaybecomecomplicated by some side processes [81]. When carryingout athermal versionof hydrosilylation, attachment to surface SiH groups is feasible with participation of both C = C and C bonds of acrylic acid. Solid-phase hydrosilylationof acrolein proceeds more efficiently [82]. Silicawith silicon hydride groups(0.2 mmol/g) was put into aglass reactor containing acrolein and a catalyst (solution of H2PtC16 in isopropanol), and the mixture was kept at 50°C for h. Then the modified silica was filtered, washed multiply with isopropanol, and dried in air. Acrolein is known to combine properties of aldehydes and unsaturated compounds and so it can react with surface =Si participation of both C C bonds and carbonyl groups. The data gathered provide evidence for the complete involvementof =SiH groups of silica in th tion reaction. It has been established that the aldehyde groups(GEE S attached to the silica surface react with 2,4-dinitrophenylhydrazine ditions, A spectrophotometric method hasbeen developed for quantitative determination of concentrations of grafted carbonyl groups when utilizing theabovementioned reagent. The activated silica matrices with aldehyde groups aresuitable for bonding amine-containing compounds with the aidof azomethine bonds: SiCH2CH2CH0 H2N

enzyme"+ ESiCH2CH2CH N -enzyme

Revelation and optimization of major factors that have an effect on processes of immobilization call for kinetic research into attachment of enzymes to the surface of activated carriers. Description of kinetic curve is usually provided using empirical equations that allow one to calculate the terminal time of an immobilization process (or the time required to carry out such a process up to acertain degree of

completion) and extent of the maximum bonding for a given enzyme and chosen reaction conditions. Theauthors of Ref. 83 havemade studies of the trypsin immobilization kinetics on thesurface of amacroporous silica withgraftedaldehydegroups (specific surface area: 90 m2/g, total pore volume: 1.5 cm3/g, mean pore radius: 33 nm). A solution of trypsin in 0.1 M phosphatesitrate buffer (pH 7.0; 25°C) was circulated through a column with the activated carrier at a rate of 4.7 mL/ min with the aid of a peristaltic pump, and the rate of change of the absorption intensity for a range of 280 nm was determined by a spectrophotometer. It was shown that in most cases the process of attachment was governed not by the rate proper of the chemical reaction between protein molecules and functional groups of the activated carrier but by the rate of penetration of the enzyme into matrix pores. Thus, according to our data 153,541 the process of attachment of a-amylase to thesurfae of an activatednonporous silicais completedwithin 10-15 min, while for porous silica matrices activated in a similar way it takes from to 7 h (depending on the molecular weight of an enzyme) to complete iml~obilization. As was shown by way of example for bonding of trypsin [83,84], the description of the kinetics of immobilization of the enzyme on activated porous matrices with spherical particles can be given, applying the equations suggested earlier for the kinetics of adsorption of gases on porous spherical adsorbents from a carrier gas flow. These equations are based on analogous notions about the natureof a ratedeterminingfactor (~ntradiffusion resistance) andon theLangmuirtype of adsorption isotherms. It has been established (851 that the terminal time for an immobilization process doesnot practically depend on the initial concentration of an enzyme in solution. However, the kinetic curvesfor carriers withthesamesurface nature and pore sizes but with different particle sizes do not coincide. The terminal time for an immobilizationprocess varies in proportion to the square of themean radius of particles of the carrier used. Of great importance is also the relationship between sizes of pores and protein molecules. A plot of activity of immobilized enzyme preparations as a function of carrier pore sizes gives, as a rule, a curve with a maximum. The decrease in the activity of preparations at the mean pore radius exceeding the optimum one is attributed to a decrease in the specific surface area of the carrier. Insituationswherepore sizes aresmaller than the optimum values, the activity of preparations decreases because of reduction of the matrix surface accessible to biocatalyst macromolecules from steric considerations. An empirical rule has been formulated, according to which the activity of an immobilized preparation is at its maximum in the cases when the mean pore radius approximately twice as large as the enzyme molecule size. However, there are also substantial departures from this rule. In the caseof enzyme preparations immobilized on porous matrices one also can observe some intradiffusioll resistance for a substrate. For example, urease immobilized on Silochrom is characterized by two kinetically different forms.The enzyme located in anear-surface layer hastheMichaelisconstant that is close in the order of magnitude to its value for the soluble biocatalyst (3 and mm01 respectively). Urease situated in pores does not exhibit its activity at low

es

concentrations (0.5-5.0 mM) of urea. When the substrate concentrationis high l000 mM), the diffusion barrier is overcome, which leads to an increase in the reaction rate and Km value (21 mmol). In general, it can be stated that geometrical characteristics of carriers have a considerable effect upon the bonding rate of enzymes, capacity of a matrix for a protein, and manifestation of the immobilized preparationactivity, which is due to the intradiffusion resistance.

Activity of immobilized preparations depends to a great extent onof pH a medium, in which an enzyme bonding is effected. It has been established that the maximum degree of bonding of proteins on silica is observed at pH values close to the isoelectric point of a protein. It is believed that at such pH values globules of proteins have the mostrigid structure. As a result, the surface occupiedby a protein molecule under these conditions proves to be minimal. To all appearances, at pH values close to the isoelectric point, a macromoleculeis adsorbed in a statewhen its conformation bears a most strong resemblance to thatof the native protein and is distorted to a lesser degree under the action of the silica surface. Under favorable conditions, dependencies of the reaction rate on p native and immobilized enzymes concur. The same is true for the opt values for the manifestation of their activity. Such a concurrence seems to provide evidence for conservation of the conformation of a biocatalyst during its anchoring on a surface. In many cases, however, pH dependencies of activity for the free and immobilized enzymes are somewhat different. This can be related both to changes of properties of a biocatalyst during its attachment to the surface and to influenceof diffusion o other factors. In particular, when studyingtherate of hydrolysis of 1 urea by urease preparations as a function of p it has been established [64] that the activity profile for the immobilized enzyme differs rather strongly from that for the soluble urease. For the soluble enzy one can observe an ordinary bell-shaped dependence of the reaction rate on p with the pH optimum at about 5.8, and, in addition, the dependencies determined using bidistillate and 2.5 rnM phosphate-citrate buffer practically coincide. In the case of urease immobilized on Silochrom the rate of hydrolysis in the phosphatecitrate buffer decreases when pH of the solution varies from 4.5 to 6.5 (while in bidistillate it changes slightly) and irreversibly reduces practically to zero when pH lowers from 4.5 to 4.0. At pH values higher than 7.0 the rate of this hydrolysis agrees satisfactorily withthereactionrateforthesoluble enzyme. The irreversible loss of the activity withdecreasing pH to 4.0is also observedfor the soluble enzyme. The presence of a plateau for activity within the pH range from 4.5 to 6.5 in the case of bidistillate could be explained by the influence of the substrate diffusion. However, it has been shown that during hydrolysis of 1 urea the preparation is functioning in the kinetic region. The plateau observed is attributed to a shift of pH of the solution within the carrier pores caused by the urea hydrolysis products.

far as immobilized preparationsare concerned, itis of significanceto discern the part of an active compound that retains its main properties in the bound state. is evident from the data summarized in Table I, in a number of cases a method ed for activation of the silica matrix surface can be of prime importance. de by side with the initial activity of an immobilized enzyme there are other important characteristics, namely its storage stability and stability in operation conditions. It was established that the activity of trypsinimmobilized on CC-amino-Silochrom (amino-Silochrom activatedby cyanuric chloride) in the process of storage in a dried state at 4°C for 14 months almost did not change. When storing at room temperature in distilled water, in 74 days the activity of this preparation made up 78% of the initial one, thus the halflife of its activity conservation, was equal to 200 days. In order to gather numerical data on stability the of the preparations obtained, a study was made of performance of columns with immobilized trypsin at elevated te~peratures.The immobilized trypsin was drawn through a column, and at its outlet amounts of the formed p-nitroaniline were recorded. The results achieved

Specific Quantityand Activated Silica Matrices

Enzyme

Carrier

Specific quantity of bonded Degree Method enzyme binding conservation of activationa (mg/g)

cc

Silochrom Trypsin

MA TDI TDI

Silochrom Pepsin

cc

Fumed silica Pronase E Catalase Fumed silica

Activity ofEnzymesAnchoredonSurfaceof

Silochrom

TDI

cc cc

TDI

cc

Silochrom

TDI

cc

GA Oxid.

89 91 92 l00 100 45 42 97 20 20 40 40 40 40

of 4

of Degree activity

( 0 52 35 47 70 70 69 68 71 100 100 72 91 60 100

31 29 15 80 73 16 12 50 10 14 87 80 70 l00

CC: arninosilica activated by glutaraldehyde, tolylene 2,4-diisocyanate, cyanuric chloride respectively. MA: vinylsilica activated by maleic anhydride. Oxid.: catalase with oxidated carbohydrate component on aminosilica.

(Fig. 1) provide evidence for high stability of immobilized preparations.Even at 55°C it difficult to determine losses of activity within the studied periodsof time due to the low rate of the column inactivation. It should be noted that there is lower rate of trypsin inactivation at 65°C in the case of the matrix activated by maleic anhydride in comparison with the aminosilica activatedby cyanuric chloride. Under the assumption that the decreasein activity follows the exponential law, the rate constant for enzyme inactivation, Ki, was calculated by the formula:

where is the initial activity and E is the activity of the enzyme at instantof time t . For the total kinetic curve the Ki value was acquired by linearization of the gathered data in semilogarithmicco-ordinates(lnE,t)applyingthemethod of least squares. The obtained values of the inactivation rate constant formed the basis for calculating the halflife of activity conservation tl12:

or mean time of no failure

The values of the halflife of activity conservation of the column are given in Table 2, which lists also the dataon the mean time of no failure This factor, which is proportional to the rate constant for theenzyme inactivation, was offered to characterize stability of preparations in Ref. 87. Comparison between the values calculated from the decrease in the function in^ column activity and tl12 values calculated from the results of measuring the preparation activity before and after

Time,

Kinetics of inactivation of the trypsin preparations: (l, 2) soluble enzyme at 60 and 55°Crespectively; trypsinimmobilizedonaminosilica,activated by cyanuric chloride, at 65 and 55°C respectively;(4, 6) trypsin immobilized on vinylsilica, activated by maleic anhydride, at 65 and 55°C respectively. m

ii

Thermal Stability of Trypsin Preparationsa Tempera- constant column Rateof Time for ture inactivation (min"

operation (min)

tl/2

(h)

~unctioningcolumn CC-Amino-Silochrom-trypsin 7.72.8 (f0.5) 420 59.5 41.3 65 2.6 (f0.l) 7.8 230 loe3 6.4 4.4 ~~-Vinyl-Silochrom-trypsin 55 (310.4) 2.9 40. 65 1.5 (310.1) 340 11.1 7.7 In the absence of a substrate Soluble trypsin 1.39 0.96 1.2 (f0.2) 60 6.0 (f0.3) lop3 0.28 0.19 ~A-~inyl-Aerosil-trypsin 6.4 4.4 60 2.6 (310.3)

(h)

57.5

6.1

t1,2*

(h)

240

was calculated from the preparation activity before and after the column operation. CCamino-Silochrom-trypsin is trypsin immobilized on amino-Silochrom activatedby cyanuric chloride; ~A-viny~-Silochrom-trypsinand ~A-vinyi-Aerosil-trypsin correspond to the enzyme immobilized on vinyl-containing silica (Silochrom fumed silicarespectively) activated by maleic anhydride.

the column operation shows that the tllZ*values are substantially lower. our opinion, this may be explained by two causes. On the one hand, one can assume that the decreasein the immobilizedenzyme activity has an insigni~cant effect upon the column operation, and, on the other, the sharp decrease in the preparation activity may take place during the initial period of time, while the remaining active part of the enzyme has a higherstability, i.e., there are atleast two enzyme fractions that differ in their thermal stability. The validity of the last inference is corroborated by the data gathered when studying inactivation of MA-vinyl-~erosil-trypsin at 60°C. In this case the enzyme was incubated at this temperature without the substrate, and measurementsof the residual activity were made at 25°C. From the comparison beween the constants of inactivation at the same temperature it follows that the stability of MA-vinylsilicatrypsin at incubation without a substrateis approximately 20 times as high as that of the soluble enzyme (calculations of the inactivation constant for MA-vinylAerosil-trypsin were done starting with 50 min). When the inactivation constant was calculated from the data on the decrease in the functioning column activity, the stability of the immobilized trypsin preparations proved to be approximately 40 times as high asthat of the soluble enzyme. The immobilizedtrypsin preparations, inparticularthoseattachedtotheSilochrom surface, turned out to behighly suitable for investigations of their kinetic characteristics.

already mentioned, side by side with a high catalytic activity and selectivity characteristic of soluble enzymes, the immobilizedenzymes have also an important advantageovertheformer, namely theyare readily separatedfromareaction medium and can be used repeatedly. Combination of these properties opens up broad vistas for application of immobilized enzymes when analyzing various compounds in complexmixtures. Selectivity of biocatalysts makes it possibleto choose the necessary enzyme for agiven substance, and employmentof modern methods of physicochemical analysis allows oneto select an appropriate system for monitoring products of a given reaction. So far, biochemistsand microbiologists have extracted and derived over l000 enzymes, and, therefore, owing to their efforts the range of problems that can be successfully tackled with the aid of these enzymes is already quite wide. Thus, for instance, work in progress on a large scale intended for development of enzyme electrodes whose operation is based on combination of immobilized biocatalysts and special electrodes that are selective for the reaction products synthesized using a given enzyme. Substantial potentialities are also inherent in the application of immobilized enzymes in flow reactors. Such reactors enable one to effect synthesis of products of corresponding enzymaticreactions. Besides, when at the outlet of a reactor there is a flowing sensing device of the electrochemical, spectrophotometric, enthalpimetric, or other type, it becomes possible to automate analysis of substrates or products of reactions of an enzyme attached to a matrix. this In case the reactor can operate in the pulse mode (responses of a recording system are similar to chromatographic peaks) or in themode that allowsone to estimatethetotal levelof transformation of a substrate. The general a r r a n g e ~ e nof t an installation for carrying out analyses by employing immobilizedenzymes is relatively simple. It is shown in Fig. 2. corresponding buffer solution (1) is continuously circulated by means of a pump (2) in a column (4) packed with a carrier with an immobilized enzyme. With the help of a dosage system (3) a certain volume of a solution of the substance being analyzed can be

Flow-injection system for analyses with use of immobilized enzymes: buffer solution;(2) peristaltic pump;(3) dosing apparatus (injector); (4) flow reactor; (5) detector; (6) amplifier; recorder.

introduced into the buffer flow (loop feeders are usuallyused). The compounds that are substrates of the immobilized enzyme undergo transformations when they are drawn through the reactor column. Detector responds to changes of properties of the flowing solution. A. corresponding signal is transmitted through amplifier (6) to recorder (7). In the capacity of a sensing device it is possible to use a spectrophotometer, flowing oxygen electrode, or flowing enthalpimeter. In the cases when the enzymatic reaction processis accompanied by evolution of oxygen, use can be made of the Clark electrode. The flowing enthalpimeter makes a running recordof quantity of heat evolved in the processof transformation of enzyme substrates. When carrying out our research we applied all the three above-mentioned types of sensing devices. The results achieved with the helpof enzymes immobilized on organosilica surface which were used to analyze corresponding substrates in aqueous solutions are summarized in Table 3. The most universal (though not the most sensitive) sensing device used an enthalpil~etricdevice, since practically all the enzymatic reactionsare accompanied by thermal effect. The basis of the enthalpil~etric installation for determination of concentratiolls of enzyme substrates is glass adiabatic column packed with an immobilizedpreparation.Thedesign features of the flow enthalpimeter were

Enzymes Immobilized on Organosilicas for Analysis of Their Substrates Immobi~ized Analyzed enzyme Trypsin

Catalase Butyrylcholine esterase Urease Glucose oxidase Galactose oxidase

substrate a-benzoyl-l-arginine ethyl ester a-benzoyl-L-arginine amide Methyl ester D, L-lysine Hydrogen peroxide Butyrylcholine Acetylcholine Urea Glucose

Galactose Lactose Raffinose Polyphenol oxidase p-Cresol Pyrocatechol Adrenaline ATP Hexokinase Hydrogen peroxide Peroxidase

Analysis region (mmol/L) Type 0.5-10 0.5-10 0.5-10 0.1-3 0.5-7 0.5-7

0.1-10 0.1-1 0.1-1 0.1-1 2-20 00 0.1-2 0.1-1 0.1-1 0.1-5 0.5-10 0.1-1

of detector Enthalpimetric

Enthalpimetric Enthalpimetric Enthalpimetric Enthalpi~etric Spectrophotometric Electrochemical Electrochemical

Enthalpimetric Enthalpimetric

es

isor

described by us in Ref. 88. The sensing device sensitivity is not lower than 4.2 J. The baseline drift per hour makes up f 2 % of the whole scale. In the pulse mode of feeding of the measured substrate doses the reproducibility of signals observed was good of the mean valueof peak heights). consider the results of the analytical determination of some substrates of trypsin (0.5-10 mM),hydrogen peroxide (0.1-3 mM),urea (0.5-10 mM),and glucose (0.12.5 mM) by means of the flow enthalpimeter using enzymes immobilized on the surface of activated amino-Silochroms (see Fig. 3). A note should be made of the substrate specificity in terms of thermal effects of enzymatic reactions (Fig. 3a), as well as feasibility of constructing a system with several immobilized biocatalysts, where the product of an enzymatic reaction canserve as a substrate for subsequent transformations (Fig. 3d). Good potentialities latent in the enthalpimetric method for determination of substrate concentrations were also revealed in the case of

0.5

1

1.5

C, mmollL

E

g 150

50 0

2

4

6 8 C. mmollL

1

0

v

~eterminatiollof enzyme substrates in flowing enthalpirneter with use of ~ i o ~ ~ t a l y s t s i ~ ~ oon b ithe ~ i surface z e d of activated aminosilicas. (a) Trypsin, substrates: a-benzoyl-L-arginine ethyl ester (l), methyl ester D,L-lysine (2), N , benzoyl-l-arginine amide (3); (b) catalase, substrate hydrogen peroxide; (c) urease, substrate urea; (d) glucose oxidase (l), glucose oxidase with catalase substrate glucose.

immobilized cholinesterase, hexokinase, and peroxidase. The results achieved when determining urea in the flow enthalpimeter (Fig. 3c) correlate well with the data gathered during the spectrophotometric recording of pH changes at the outlet of the column using p-nitrophenol in the capacity of a pH indicator. The obtained data show that after 200 analysesthesynthesizedpreparations of immobilized urease did not change substantially their characteristics. Analysis of substrates of oxidoreductases that catalyze oxidation involving consumption of oxygen from a buffer solution canbe made in a number of ways. It is most convenient to carry outsuch an analysis by employing the Clark oxygen electrode. We have shown [B]that in order todetermine concentrationsof glucose (0.1-1.0 mM) in blood and blood serum at a rate of 60 analyses per hour it is possible to use columns with immobilized glucose oxidase.It has been established that one colunm with the immobilized enzyme (100 mg of carrier) enables one to make up to 1000 of reproducible analyses. The preparations ofglucose oxidase immobilized on amino-Silochroms by our methods have a high storage stability (up to three months) at room temperature without any substantial decrease in activity. It has been also found that the application of a flow oxygen electrode facilitates analyses of substrates of catalase (Fig. 3b), polyphenol oxidase [90,91], and galactose oxidase [47,48]. It has been found [90-943 that preparations of native and immobilized phenol oxidases can be employed for qualitative and quantitative analysesof phenols and aromatic amines, for example pyrocatechol, 3,4-dioxyphenyamine, adrenaline, nor adrenaline, 3,4-dioxyphenylethylal~ine,phenol, o-, ~-aminophenols,cresol, and a numberof other compounds. The detectors used for these analyseswere the Clark oxygen electrode, spectrophotometer with a flow-through cell, and ellthalpimetric sensing device. The reproducibility of analyses was better than 2% at a reliability when the volume levelof 95%. The sensitivity of analysis amounted to mol/L an analyzed sample was 0.05-0.10 mL and could be increased up to mol/L when the pulse modeof introduction of a substrate was replaced by its continuous drawingthroughacolumnwithimmobilizedpolyphenoloxidase.Theaverage production rate of analysis was 30-40 samples per hour. method has been proposed whereby it is possible to increasethe sensitivity of analysis of phenols through introductionof specificadditives (e.g., substances capableof nonenzymatic reduction of o-quinones to o-diphenols) into a reaction mixture. Thus, it has been established that the application offlow-type reactors with immobilized biocatalysts is very promising for analyzing substrates of enzymes in complex mixtures. The use of flow-sensing devices of various types (oxygen electrode, entllalpimeter, etc.) makes it possibleto automate the process of carrying out an analysis. The most versatile for analyzing substrates of enzymes have proved to be systems based on a flow enthalpimeter.

The high catalytic activity and exceptional selectivity of enzymes offer ample scope for their application not only for analysis but also for fine organic synthesis. The

broad prospects of their use have been considerably expanded owing to development of methods for immobilizing enzymes on surface of inso*lublematrices. At present, reactions of synthesis of peptides involvean increasingly wide use of proteolytic enzymes. This related to the fact that theenzymatic synthesis of peptides does not require anyspecial activation of carboxyl groups and protection of side functional groups of amine acids. Moreover, this synthesis is stereoselective and enables one to carry out reactions under mild conditions (room temperature, aqueous media with values of pH close to neutral ones). When utilizing heterogeneous preparations of proteolytic enzymes it is possible to effect continuous synthesis of peptides with the help of a flow reactor, which is of obvious practical interest. Special attention should be given to preparations of immobilized papain. This enzyme has a high substrate selectivity, sufficient esterase activity. characterized by a weak secondary hydrolysisat the o ~ t i m u mvalues of pH for the enzymatic synthesis of peptides.

In spite of clear advantages that can be gained while employing immobilized proteases for synthesis of peptides, at present there is no detailed research into corresponding processes. Since an immobilized enzyme acts as a heterogeneous catalyst, the most efficient approach seems to involve effecting an enzymatic synthesis under kinetic control conditions. The unquestionable advantagesof such an approach are rapidity of synthesis and ease of separation of products from an insoluble biocatalyst. In this situation there are no restrictions attributed to solubility of final products to their distribution in water and an organic solvent. Especially attractivefor synthesis of peptides is thiol proteasepapain. It is relatively inexpensive andhasahighsubstrate specificity. Kineticcontrol is known to have a substantial drawback caused by the secondary hydrolysis of a final product. Earlier it was found that in the pH range of 8.0-9.5 papain retains its considerable esterase activity, which enables one to use it for production of peptides under kinetic control conditions. At the sametime, under these conditions the rate of hydrolysis of peptide bonds is insignificant and does not affect substantially the final product yield. Theaforementionedproperties of papain make it suitable for a stepwise synthesis of peptides. References 56, 65, 66, and 96 are devoted to research into properties of papain (immobilized on organosilica) in reactions of formation of peptide bonds. By way of illustration, theinfluence of various factors on thefinal product yield was investigated using a reaction of synthesis of tripeptide:

BzAlaOCH3 H ValGZyOH

BzAla ValGlyOH

CH30H

where Bz,Ala, are benzoyl, alanine, valine, and glycine respectively. When employing immobilized enzyme, it is important for all the starting components and products tobe in a dissolved state. Alkyl ethersof amine acids which are used in the capacity of carboxyl components are slightly soluble in aqueous media. To ipcrease the solubility of the carboxyl component the above-mentioned reaction was effected at 40°C in an aqueous-organic mixture containing (by

volume) of methanol. As a rule, the ratio of carboxyl and amine components was 2:1. The carboxyl component was synthesized using I4C-alanine.In order to ascertain the yield, aliquots from this reaction mixture were separated by employing a thin-layer chromatography technique that made it possible to determine the distribution of radioactivity along theplate length. At pH 9.0 the yields of the synthesized tripeptide amounted to 49 f2%. Figure 4a illustrates the kinetics of the final product accumulation in the course of the reaction for synthesis of tripeptide. The maximum product yield is attained within 40 min. The subsequent incubation doesnot lead to any increase in thefinal product yield. The secondary hydrolysis of the final product is insignificant and does not exert a substantialinfluence on the yield value. After the incubationof the reaction mixture for 120 min the final product yield decreases by 6-8% in comparison with the m a ~ i m u mvalue attained.

Dependence of the yield of peptide BzAlaValGlyOH in the reaction catalyzedby immobilized papain on time (a), pH of the medium temperature (c). Reactionconditions:mixturevolume, 2 ml; volume fraction of MeOH, BzAlaOMe, 0.4 mmol; HYalGlyOH, 0.2 mmol; pH (a, b); 40 min (b, c); 40°C b).

The graph of the final product yield versus pH value has the form of a bellshaped curve (Fig. 4b). The maximum yields are observed at pH 9.0. This situation with the optimum pH is related to the fact that at pH 8 protonation of group NH2 of the amine component leads to a decrease in its nucleophility. The acylenzyme formed under these conditions from the carboxyl component and enzyme is subjected to a predominant hydrolysis and decomposes into the N-protected amine acid and enzyme. At pH 9.5 the decrease in the final product yield seems to be related to the immobilized enzyme inactivation. Thus, the range pH 8.0-9.5 is the optimum one for synthesizing peptides under the kinetic control conditions when employing il~mobilized papain. The initial concentration of starting componentsis of special importance for the synthesis of peptides under the kinetic control conditions. In the case of the immobilized-on-silica papain that catalyzes the reaction of synthesis of tripeptide, the optimum concentrations of the starting components are equal to 0.14.2 mol/L. When the concentrations are lower, the yields slightly decrease. The twofold excess of the carboxyl component leadsto the product yield increase by 15-20% as against the equimolar ratio of the reaction components. The temperature dependenceof the tripeptide yield in the reaction catalyzedby immobilized papainis shown in Fig. 4c. The maximum yields are observed at 45°C. Since the synthesis of tripeptide was carried out in an aqueous-organic mixture, gradual vaporization of methyl alcohol from the reaction medium resulted in that the carboxyl component passedinto aninsoluble sediment. Thus,at 60 and 70°C in 40 min the concentration of the unhydrolyzed carboxyl component was equal to 25-30%. As a consequence, at 50-60°C the final product yields decrease. The thermal stability of the papain immobilized on organosilica is higher than that of the soluble enzyme. This is evidenced by the fact that the immobilized papain exhibits its synthetic activity at 7O0C,whereas the solubleenzyme practically completely loses its activity at a temperature of 50°C. We have also effected synthesis of tripeptide using the same weighed sample of the immobilized papain. As a result, the yields attained were as follows: 51, 50, 46, 45%. The tendency towards some decrease in the yield is probably related to the action of the organic solvent on the enzyme. From the above-presented data a conclusion can be drawn that the low solubility of the carboxyl componentlimits utilization of immo~ilizedpapain for production of peptides. It is evident that using i~mobilizedenzymes will be most efficient in the case of soluble components involved in the synthesis. Addition of an organic solvent to a reaction medium, even if it allows one to tackle a problem partially, cannot be unlimited.Due to this fact it isof interest to searchforsuch ester derivatives of N-protected amine acids that are highly soluble in aqueous media, can serveas substratesof papain, and are capable of acting as carboxyl col~ponents for the enzymatic synthesis of peptides. The authors of Ref. 9'7 made use of glycolylglycinic esters of N-protected amine acids in the capacity of a caboxyl component. The synthesis of such compounds was effected from N-substituted amine acid and chloroacetyl glycine. ~lycolylglycinicesters are readily soluble in aqueous solutions; at pH 6 they are hydrolyzed by papain and can be used for production of peptides containing hydrophobicamineacids(phenylalanine,tryptophan, leucine). Whentheyare

ennployed in reactions for the enzymatic synthesis of peptides, the functionof amine components can be performed by the compounds that are usedin the case of alkyl ethers. Glycolylglycinic esters of N-protectedamine acids readily interactwith papain. The hydrolysis of such compounds by papain proceeds, as a rule, several times asfast as the hydrolysisof analogous alkyl ethers, which leads to a substantial decrease in the reaction time for formation of a peptide bond under the action of immobilized papain. We have studied the reaction of formation of a peptide bond by employing papain immobilized on silica as a catalyst and glycolylglycinic ester of benzoyl-lalanine in the capacity of a carboxyl component:

When use was made of glycolylglycinic ester of benzoylalanine, all the components for the synthesis were in a dissolved state. The reaction was carried out without addition of an organic solvent at room temperature. Thus, the conditions to use immobilized papain were optimum. At a ratio of carboxyl and amine Components of 2:l and pH 9.0 the reaction yieldof tripeptide amounted to 59 f 6%. The reactiontime was substantially shortened and made up 15 min.Themaximum yields of tripeptide were observedpat 8.0-9.5 is optimum the synthesizing for one peptides when capacity papain thein e of a catalyst. At p 8-10 a slight alkaline hydrolysis ofglycolylglycinic ester of benzoylalanine observed. This process, however, does not exert any substantial influence onthe final product yield,since theenzymatic synthesis rate exceeds considerably the spontaneous hydrolysis rate. Excess of one of the components leads to an increase in the yield. At the equimolar ratio of components the yieldis equal to 30%. n thecase ofexcessof carboxyl or amine Component (at a ratio of 2 1 ) the yield increases by 30 or 10% respectively. Therefore, the enzymatic synthesis of peptides is expedited by a twoto fourfold excess of one of the components.

pentapeptide Leu-enkephalin. The synthesis was catalyzed by papain il~mobilized on silica, and the reaction was effected according to the following scheme:

The carboxyl component was glycolylglycinic ester of ~~e~-butyloxycarbonyl tyrosylglycine. The amine component was glycylphenylalanylleucine. The components and products wereeasily soluble in buffersolutions that were utilized for this synthesis. The reaction time was 20-25 min. The reaction proceeded at o of the components thefinal product yield amounted f. 66 was to investigate the influence of a component ratio on the final product yield (Table 4). it was in the case of the synthesis of peptide

Reaction yieldsof Boc-Leu-enkephalin in the Synthesis Catalyzed by Papain Immobilized on Silochrom Content of components in the reaction mixture (mol/L)

Yield of Boc-Leu-enkephalin weight) by

Carboxyl component Amine component 0. 0.2 0.

60; 62; 57; 60 72; 70 81; 89

the yield of pentapeptide increased at an excess of one of the components used. A study was also made of the synthesis of Boc-leu-enkephalin catalyzedby immobilizedpapainundercontinuousprocessconditionsusing column.The reaction proceeded accordingto theabove-mentioned scheme. To reduce variations of pH values of themediumwiththereactionprogressthereactionmixture included M sodium carbonate buffer (pH 9.5). The column contained 400 mg of immobilized papain. The pumping rate for soluble substrates was adjusted so to ensure the complete enzymatic hydrolysis of the carboxyl component during the time of its residence in the column. The immobilized papain preparation was used multiply. The experimental results are summarized in Table 5. The final product yield is substantially affected by the pumping rateand volume of the solution that contains reaction components andpasses through the column. With increasing pumping rate the final product yield increases. Preference should

Yields of Boc-Leu-enkephalinDuringSynthesisin Containing Papain Immobilized on Silochrom Component content Pumping (mmol)

rate of reaction

Reaction nCarboxyl enine volume mixture component component (mL/min) l 5 5

5

0.1 0.1 0.5 0.5 0.5 0.5

Yield of Boc-Leu-

by weight) 1

23

a FlowReactor

0. 1 0.5 0.5 0.5 1.5

23

.6 0.7 0.9 1 0.9

20 28 43 54

be given to prolonged circulating of a solution of substrates through the column. In this case the final product yields are considerably higher than in the pulse mode of introduction of substrates. So, when the volume of the reaction mixture flowing through the column increases from 1 to 5 mL,the yield increases from 20 to 43%. In the mode of a continuous passing of a solution of starting components the conditions set up in the columnare optimum, and thefinal product yields approach the values achieved under static conditions. Thus, the synthesized preparations of papain immobilized on Silochrom for the first time were used for production of a biologically active pentapeptide ~ ~ c - ~ e ~ - e n k e p h It a lhas i n . been shown that the final product yields are notlower than in the caseof the native enzyme.In addition, for thefirst time it hasalso been established that immobilized papain canbe applied for production of peptides in a continuous flow reactor.

Adsorption of proteins and other biopolymers on silica surfaces takes place during their iml~obilization198-1021, application of silicas in processes of purification and concentration [103-1061, development of sensors (107-1121, determination of the body’s initial response to foreign objects [l 13,1141 and implantedbioglasses [I 15,1161, well as under natural conditions in soils [l 171. Knowledge of the major law-governed phenomena of interaction of biomolecules with a surface is necessary for purposeful variation of the surface chemistryof silicas and properties of solutions with the aim of varying adsorption equilibria in a desired direction. Depending on applications of silicas the range of the required affinity of surface to biomolecules can change from very great value characteristic of enzyme immobilization to a practically negligible value in the case of electrophoretic separation using quartz capillaries. The intermediate values of the affinity are necessary for variouschromatographyprocedures.Atpresentatthe qualitative level it is a d ~ i t t e dthat the main contribution to interactions between biopolymers and silica surfaces is made by electrostatic interactions [l 18-1241. In this case the silica surface charge is developed as a result of dissociation of surface silanol groups, while the biopolymer chargeis formed in consequenceof the appropriate dissociationof acid and/or base groups. The process of adsorption of proteins at levels over their isoelectric points calls for additional assumptions to be made. At the same time, when considering equilibria of interactions between polyelectrolytes and silica surface, this apparent contradiction can be accounted for without making any additional assumptions. According to Gibbs, adsorption takes place during formation of a surfaceexcess of solved moleculesat solid-liquid interfaces. In this case introductionof silica into in its conce~tration.Attachment of solution of protein leads to a decrease molecules to surface by chemisorption can also be effected without any decrease in their content in a solution, i.e., without their surface concentration. For a quantitative descriptionof interactions on a surfacelet us consider equilibria maintained during adsorption of biomolecules M on the silica surfxe with adsorption sites L. We have already investigated various equilibria involving sur-

es

face sites and molecules in solution [125]. Recent evidence suggests that adsorption of biopolymers on surfaces of silicas presupposes interactions of two polyelectrolytes. In this case on the silica surface there canbe acid silanol groups and attached groups of a different nature. At the sametime, biopolymers (proteins in particular) have numerous functional groups with various dissociation constants.In any case, reversal of charge signs both of silicas and of biomolecules in solution is accompanied by ionization of twogroupsthat specify transition. between these two charged forms. Thus, in afirst approxi~ationthe interactionof proteins with silica surface can be interpreted as an interaction between dibasic sites of the surface and dibasic molecules in solution. For a modified surface having amphoteric properties ionization of adsorption sites can be represented in the following way: K1

[LH] 0

23

+l

[H]

[L] -1

[H]

Similar equilibria are also established during ionization of molecules in a solution:

K3 [MHz] +l

K4

[MH] [H] 0

[H] -1

All the three forms occur in the solution in the state of a dynamic equilib~ium, and their ratios are determined by pH values and dissociation constants K3 and On. the assumption that the interaction involves only oppositely charged surface sites and adsorbed particles it is possible to write the following equilibria:

where

In thecase of sorption of a negatively charged form [M](Kdl isotherm equation can be written as follows:

K&) the adsorption

while for sorption of positively charged molecules [MH2] the equation

Kd2)we can get

is seen from these equations, when on the surface there are simultaneously both acid and base sites (ampholyte), the ascending and descending curve sections are more steep than in the presence of only acid or only base sites. Figures 5 and 6 show that for ahigh affinity (a low a significant sorption is observed also at pH p1 forcationexchangersand atpH p1 foranionexchangers.Thus,the equations of adsorption can be written in the following general form: [LM]

[Llo[Ml,q [MI,,

Since theadsorption sites arenotalwayshomogeneous in their affinity to adsorbed molecules, the adsorption isotherms observed in this case are described by a sum of several isotherms [125,126]. For example, for two typesof binding sites the experimental isotherms consist of a sum of two isotherms:

Adsorption of positively charged molecules with an isoelectric point of 6.5 on negatively chargedadsorption sites with pK 9, The curves were calculated from (l) using 10"' mol/L, mol/L, lo-' mol/L, mol/ L, 1 0 - ~m ~ i /PM: ~, log [M],,.

es

Adsorption of positively charged molecules with an isoelectric point of 6.5 onnegativelycharged adsorption siteswithanisoelectricpoint of pK 9. The curveswerecalculatedfrom Eq. (1) using K1 mol/L, 10-l~ IX~I/L, PM [rvrleS.

is thetotalamount of thebound molecules,is the overall concentration bindingsites ofthe first type, [L"Io is that of another type, and F' and and and are the corresponding acidity functions and association constants for sites of the first and second types. From the isotllerms reported one can easily determine the distribution coefficient, i.e., the ratio of concentrations of the sorbed and dissolved substances. In a general form, the distribution coefficient may be represented as follows:

where F i s the acidity function whose form is governed by the nature of the system studied. The observed constants of equilibria

and, correspondingly, the free energy of adsorption and concentration of other particles in the solution. From the equations for the equilibrium constants and adsorption isotherms it obvious that at thesame equilibrium constantand free energy of adsorption, therates of adsorption kl and desorption can be different. At a high affinity of molecules to a surface the observed desorption rate

may be sufficientlylow, and itwill give an opportunity for prolonged functioning of an immobilized compound even in the situation of the electrostatic mechanism of attachment. Taking into account that in case the of application of porous silicas the adsorption processes involve mass-transfer resistance it is possible to expect that this will give rise to an additional decelerationof desorption, i.e., to anincrease in the stability of immobilized preparations. On the other hand, under the conditions of covalent bindingat low rates of sorption and desorptioneven small equilibrium constants yield a sufficient stability of iml~obilizedmolecules. Therefore, it maybe inferred that in the first case the stability is determined by thermodynamic parameters, whilein the second case it is controlled by kinetic parameters. During covalent immobilization there is a possibility for a combined action of these two mechanisms.

As mentioned above, covalent attachment may not lead to formationof any surface excess of an adsorbate, i.e., the ther~odynamiccoefficient of distribution is equal to unity and the preparations produced will be stable. At the sametime, adsorptive attachment with such a coefficient of distribution will require the action of additional factors necessary for stabilization of attachment of biomolecules. At present there is a method that finds increasingly wideuse for physical introduction of enzymes and other proteins into the silica gel structure. The method is based on the sol-gel process [107,127-1321. The corresponding procedureconsists of formation of silica gels whose pore diameter is smaller than the dimensions of immobilized molecules. A adsorption onto colloidal silica sol and three polypyrrole-silica nanosites (untreated and amine- or carboxylic acid-functionalized) was investigated. The results suggest that DNA adsorptionis mainly governed by electrostatic and hydrophobic interactions. At neutral pH silica the sol has anet negative surface charge, which probably accounts for zero or very low amounts of DNA adsorbed se substrates [324]. research into adsorption of oligodeoxyribonucleotide onto aminopropylsilane-modified silica wafers it has been shown that the major role in the adsorption process is played by electrostatic interactionsbetween the negatively charged adsorbate and aminated groups on the silica wafers [133]. Other examples of DNA adsorption studies using silica and involving covalent immobilization of si~gle-stranded DNA on optical fibers designed for the development of biosensors are reported [134]. The adsorption of both plasmid and chromosomal duplex DNA to silica was investigated with the purpose of determining the dominant forces acting in the binding reaction111351. The adsorption isotherms indicate that three factors, namely: (l) shielded intermolecular electrostaticforces, (2) dehydration of the DNA and silica surfaces, and (3) intermolecular hydrogen-

bond formation in the DNA-silica contact layer, make the dominant contributions to the overall driving force for adsorption. Capillary zone electrophoresis in quartz tubes is often used for the separation and characterization of humic acids. It has been established that the adsorption of humic acids on an uncoated capillary wall is high (1361. Immobilization of purified Aldrich humic acid on aminopropylsilica and glutaraldehyde-activated arninopropylsilica has been also investigated. In general, the humic acidis bound to thesolid by both physical and chemical bonds. The physically adsorbed humic acid can be released to large extent at high pH. Tominimize physical adsorption, the samples were thoroughly washed using a 1 M sodium chloride solution of p free amino groups of the aminopropylsilica were end-capped 137). The surface complexation model (with no electrostatic term) has been applied in order to describe the mechanisms operatingat the oxide-water interface and determine apparent surface complexation constants for each species (cation, organic ligand). As a result, surface reactions with the hydroxyl groups of silica surface are suggested[138]. The sorption of humic acid hasbeen shown to be related to the nature of charges on the mineralsurfaces, with positively charged surfaces sorbing the humic acid to a greater extent than negatively charged surfaces. The sorption process involving positively charged mineral surfaces has been found to be ciated with c o n s u ~ p t i o nof protons as a consequence of the interaction, whereas that involving negatively charged mineralsurfaces does not consumeprotons 1171. It hasbeen also established that in thecase of the silica surface as substrate the immobilization of various cells [60,128,132,139] and organelles of albino rat liver microsomal fraction 11401, well immobilization of submitochondrial particles prepared from beefliver mitochondria leads to stabilization of their enzymatic activity [141].

Nowadays, researchers have at their disposal a wide range of sufficiently versatile and sophisticated methods for chemical attachment of enzymes and other active compounds to pristine and modified silica matrices, Worthy of a particular note is the store of methods for activating the surface of ami~e-containingsilicas. The gathered experimental data give evidence for the fact that the high efficiency of immobilization of enzymes on the organosilica surface is determined both by the surface chemistry and geometrical structure of a silica substrate (namely by the nature of functional groups of modifying layer, by particle sizes, and by the character of a matrix pore structure) and by properties of a high-molecular adsorbate and conditions for proceeding of adsorption from solution. For a large set of enzymes it has been shown that depending on a method for bonding of biocatalysts to the surface of a silica matrix it is possible to vary amounts of an immobilized enzymatic protein, its thermal stability, optimum pH value for attachment to the surface, static and operative stabi1ity of preparations, degree of conservation of their activity, and resistance to denaturating actions of various compounds. These factors are also of utmost importance for efficient use of flow reactors for fine organic synthesis involving catalysts immobilized on organosilicas and for analysis

of enzymatic substrates incomplex mixtures. Withreference to the use for practical purposes, of a particular significance is consideration of equilibria of interactions (adsorption and chemisorption) between surface sites of unmodified and modified silicas and enzymes or other biologically active compounds in solution. hysicochemical interpretation of such equilibria furnishes more favorable initial positions for addressingand tackling the problem of the most effective attachment and application of preparations immobilized on silica matrices.

The research described in this chapter was made possible in part by grant ffUN2438 from the U.S. Civilian Research and Development Foundation and from the Government of the Ukraine, and by grant N3.4/175 from the StateFoundation for Fundal~entalResearch of the Ministryof Scienceand Technologies of the Ukraine.

3. 4.

6. 7 8. 9. 10. 11. 12. 13. 14. 1

ed OH Uu. 1. Ermakov, V. A. Zakharov, and B. N. Kuznetsov, ~ n c ~ o r Complexes Oxide Carriers in Catalysis, Nauka, Novosibirsk, 1980 (in Russian). LisichkinandA.Ua.Yuffa, ~ e t e r o g e ~ e o~ ~e st a Contplex l himiya, Moscow, 1981 (in Russian). V. A. Tertykh and A. Belyakova, Chemical ~eactionsr~volving Surjace, Naukova Dumka, Kiev, 1991 (in Russian). E. F. Vansant, P. VanDerVoort,and K. C.Vrancken, C~aracterizatio~ and ~ o d ~ c a t ofi othe~ Silica Surface, Elsevier, Amsterdam, 199.5. mat, Konieczka, B. J. Tarbet, J. S. Bradshaw,and Separation and Purification Methods 23:77 (1994). N. Zaitsev, C o m ~ l e ~ ~ iSilicas: ng Synth of ~ o n d e dLayer and Surface C h e ~ i s t r yFolio, , Khar’kov, 1997 V. Tertykh and V. V. Yanishpolskii, in d ~ d s o r b e ~ tVol. s , 8 (V. A.Tertykh, ed.), Naukova Dumka, Kiev, 1980, pp. 3-27 (in Russian). Weetall. Appl. Biochem. Biotechnol. 41:1S7 (1993). V. K. Antonov, and K. Martinek, eds., Im~obilized Enzymes, ‘vols. University, 1976 (in Russian). d and Erzzymes, IRL Press, Oxford, d, ed., I ~ ~ o b i l i z e Cells Was~ington,1985. Weetall. Nature 223:969 (1969). Weetall. Science 166:61S (1969). H. H. Weetall. Biochim. Biophys. Acta 185:464 (1969). R. Haynes and K. A. Walsh. Biochem. Biophys. Res. Commun. 36:23S (1969). A. N.Emery, I. S. Hough, and I. M.Novais.Chem.Eng.(London)258:71

eetall and C. C. Detar. Biotechnol. Bioeng. 17:295 (1975). 16. 17. P. Monsan and G. Durand. FEBS Lett. 16:39 (1971).

Inorganic 18. V. A. Tertykh and L. A. Belyakova, in~ d ~ ~ o r p t i o n Sorbents (A. Dabrowskiand V. Tertykh, eds.), Elsevier, Amsterdam, 1996, pp. 147-189. 19. V. A. Tertykh, A. A. Chuiko, and I. E. Neimark. Teoret. Eksperim. Khirniya 1:400 (1965). 20. A. A. Chuiko, V. A. Tertykh, G. E. Pavlik, and I. E. Neimark. Kolloid. Zhurn. 27:903 (1965). 21. V. A. Tertykh, E. A. Chuiko, A. A. Chuiko, and I. E. Neimark. Kolloid. Zhurn. 28:278 (1966). 22. V. A. Tertykh, A. A. Chuiko, V. A. Khranovskii, and I. E. Neirnark. Zhurn. Fiz. Khimii 42: 1758 (1 968). 23. V, A. Tertykh, A. A. Agzamkhodzaev, A. A. Chuiko, and L. T. Zhuravlev. Izv. AN SSSR, Ser. khim., 1739 (1968). n 24. V. A. Tertykh,T. N.Burushkina, and A. Chuiko, in~ o d i ~ c ~ tofi oProperties P o l y ~ e rand ~ ~ P o ~ y ~ e r i c ~ a t e r i(K. u l , A. ~ Kornev, ed.), Naukova Dumka, Kiev, 1965, pp. 85-95 (in Russian). 25. IS. Tanaka, S. Shinoda, and Y. Saito. Chem. Lett. 179 (1979). 26. C.-H.Chiang, H. Ishida, and J. L. Koenig. J. ColloidInterface Sci.74:396 (l 980). 27. C.-H. Chiang and J. L. Koenig. J. Colloid Interfxe Sci. 83:361 (1981). 28. C.-H. Chiang, N.-I. Liu, and J. L. Koenig. J. Colloid Interface Sci. 86:26 (1982). 29. S. Shinoda and Y. Saito. J. Colloid Interface Sci. 89:293 (1982). 30. G. S. Caravajal, D. E. Leyden, C. R. Quinting, and G. E. Maciel. Anal. Chern. 60:1776 (1 988). 31. K. C. Vrancken, P. Van Der Voort, I. Gillis-D'Harners, P. Grobet, and E. F. Vansant. J. Chem. Soc., Faraday Trans. 88:3197 (1992). K. Posserniers,P.Van Der Voort, and E. F. 32. K. C. Vrancken,E.Carteleyn, Vansant. J. Chem. Soc., Faraday Trans. 89:2037 (1993). 33. K. M. R. Kallury, P. M.Macdonald, and M. Thompson. Langrnuir 10:492 (1994). 34. K. C. Vrancken, P. Van Der Voort, K. Possemiers, P. Grobet, and E. F. Vansant, in ~ o d ~ Surfaces e d Pesek and I. E. Leigh,eds.),Royal Chemical Society, Cambridge, 1994, pp. 46-57. 35. K. C. Vrancken, K. Possemiers, P. Van Der Voort, and E. F. Vansant. Colloids Surfaces A 98:235 (1995). 36. K. C Vrancken, P. Van Der Voort, K. Possemiers, and E. F. Vansant. J. Colloid Interface Sci. 17486 (1995). 37. K. C. Vrancken, P. Van Der Voort, K. Posserniers, P. Grobet, and E. F. Vansant. J. Colloid Interface Sei. 170:71 (1995). ion 38. L. A. Belyakova, T. P. Kolotusha, and V. A. Tertykh,in ~ ~ s o r ~ t and ~ d s o r b e n ~Vol. s , 12 (V. A. Tertykh, ed.), Naukova Dumka, Kiev, 1984, pp. 312 (in Russian). A. Belyakova, and T.P. Kolotusha,in Chemistry of 39. V.A.Tertykh,L. Heterogenized (A. Ya. Yuffa, ed.), Tyumenskii University, Tyumen', 1985, pp. 3140 (in Russian). 40. L. Belyakova, T. P.Kolotusha, V. A. Tertykh, and A. G. Maidannik. Teoret. EksDerim. Khirniva 22:631 (1985).

41. T. Kolotusha, L. A. Belyakova, and A. Tertykh. Ukr. Khim. Zhurn. 51:470 (1985). A. Tertykh, in 42. T. P. Kolotusha, A. G. Maidannik, L.A.Belyakova,and Syi~thesi~~ and F h y s i c o c ~ e ~ i ~ a l F ~ o p e rotfi eInorg~nic s and Carbon So~bents(V. A. Tertykh, ed.), Naukova Dumka, Kiev, 1986, pp. 42-47 (in Russian). A. 43. K. B. Yatsimirskii,A. A. Chuiko, A. P. Filippov, G. A. Karpenko, Tertykh, Yanishpolskii, S. B. Grinenko, and M.Belousov.Doklady AN SSR 2371 137 (1977). A. Tertykh. Teoret. Eksperim. Khimiya 23:237 (1987). 44. 1. A. Butovich and 45. L. Gudkova, N. Latyshko, R. G. Degtyar’, M. F. Gulyi, Yanishpolskii, and A. Tertykh. Ukr. Biokhim. Zhurn. 52:614 (1980). Yanishpolskii, A. Tertykh, L. Gudkova, and N. Latyshko. Ukr. 46. Biokhim. Zhurn. 53N4:53 (1981). Yanishpolskii, A. Tertykh, T. T. Buglova, and 0. 47. L. Kondakova, Koroleva. Ukr. Biokhim. Zhurn. 56:394 (1984). 48. L. Kondakova, Yanishpolskii, A. Tertykh, and T. Buglova. Ukr. Biokhim. Zhurn, 6ON6:34 (1988). 49. M. A. Marshall, H. A. Mottola. Anal. Chem. 55:2089 (1983). 50. M. A. Marshall, H. A. Mottola. Anal. Chirn. Acta 158:369 (1984). A. Tertykh, E. I. Fodorchenko, Yanishpolskii,A. S. 51. A. A. Chuiko, Tsyperovich, 1. Galich, and T. A.Koval’chuk,U.S.S.R. Patent 672,203 to Institute of PhysicalChemistryand Institute of BiochemistryAcademy of Sciences of Ukraine (1975). 52. Yanishpolskii, A. Tertykh, A. A. Chuiko, Yu. 1. Krylova, L. Kozlov, and K. Antonov, U.S.S.R. Patent 621,680 to Institute of Physical Chemistry Academy of Sciences of Ukraine and Institute of Bioorganic Chemistry Academy of Sciences of U.S.S.R. (1976). 53. Galich, A. S. A. Tertykh, Yanishpolskii,A.A. Chuiko, 1. Tsyperovich, and T. A. Koval’chuk. Doklady AN UkrSSR, B, N7:654 (1977). 54. I. Galich, A. S. Tsyperovich, A. Koval’chuk, Yanishpolskii, A. A. Chuiko, and A. Tertykh, Ukr. Biokhim. Zhurn. 49N6:44 (1977). Yanishpolskii, A. Tertykh, D. Yu. Yuodvalkite, and 55. G. Lyubinskii, A. A. Clemzha, Ukr. Biokhim. Zhurn. 59N4:35 (1987). 56. S. A. Ogii, Yanishpolskii,and A. Tertykh, Ukr. Biokhim. Zhurn., 62N6:48 (1990). 57. A. Bogatskii, T. 1.Davidenko, A. Chuenko, Yanishpolskii, A. Tertykh, and A. Chuiko. Doklady AN IJkrSSR, B (1978). 58. A. Bogatskii, T. I. Davidenko, A. Chuenko, V. Yanishpolskii, Tertykh, and A. A. Chuiko. Ukr. Biokhim. Zhurn. 51:315 (1979). 59. S. El’chits, Yanishpolskii, M. B. Abairnova, and L. A. Zatorskaya. Ukr. Biokhim. Zhurn. 51:374 (1979). 60. 1. Karpenko, Yu. R. Malashenko, Yanishpolskii, and A. Tertykh. Ukr. Biokhim. Zhurn. 51:387 (1979). 61. Yanishpolskii, A. Tertykh, and G. Lyubinskii. Ukr. Biokhirn. Zhurn. 51:324 (1979).

Yanishpolskii. Doklady AN UkrSSR, B 62. V. A. Tertykh, I. V. Klishchar, and N12:44 (1983). V.Yanishpolskii, G. V.Lyubinskii, and V. A. Tertykh. Teoret. Eksperirn. 63. Khimiya 25:113 (1989). Yanishpolskii, A. Tertykh, D. Yu. Yuodvalkite, and 64. G. Lyubinskii, V. A. Glemzha. Ukr. Biokhirn. Zhurn. 54:145 (1982). 65. Yu. V. Mitin, S. A. Ogii, and A. Tertykh. Bioorgan. Khirniya 11:1476 (1985). A. Tertykh, and Yu. Mitin. Doklady AN UkrSSR, B N9: 76 66. S. A. Ogii, (1987). 67. L.A.Belyakova, T.P. Kolotusha, T. E.Serova,V. A. Tertykh, and K. B. Yatsimirskii. Doklady AN SSR 288:1358 (1986). 68. T, P. Kolotusha, L. A. Belyakova, and V. A. Tertykh. React. Kinet. Cat. Lett. 46225 (1992). Tertykh, in Synthesis and 69. I. N. Polonskaya, L. A. Belyakova,and P l ~ ~ s i c o c ~ e ~Properties ical and Carbon Sorbents (V. A. Tertykh, ed.), Naukova Dumka, Kiev, 1986, pp. 9-16 (in Russian). A. Tertykh. Ukr. Khirn. Zhurn. 70. I. N. Polonskaya,L.A.Belyakova,and M1135 (1988). A. Tertykh. Ukr. Khirn. Zhurn. 71. I. N. Polonskaya, L. A.Belyakova,andV. 11 (1990). Tertykh. Ukr. Khirn. Zhurn. 72. I. N. Polonskaya,L. A. Be1yakova, and 62N4:85 1996). P. Kolotusha, I. N. Polonskaya, L. A. Belyakova, G. Ivanova, and A. 73 Tertykh, U.S.S.R. Patent 1,153,975to Institute of Physical Chemistry Academy of Sciences of the Ukraine (1982). 74. I. N. Polonskaya, L. A. Belyakova, and Tertykh, Ukr. Khim. Zhurn. 55:114.5 (1989). 75. A. A. Chuiko, G. E. Pavlik, A. Tertykh, E. A. Chuiko, V. A. Arternov, I. E. Neimark, and E. V. Tsinenyuk, Ukr. Khirn. Zhurn. 32:371 (1966). 76. I. Lozinskii, I. G. Tsoi, Yu. Davidovich, and S. V. Rogozhin. Izv. AN SSR Ser. Khirn. N2:417 (1981). 77. J. Carlsson, R. Axen, K. Brocklehurts, and E. M. Crook. Europ. J. Biochern. 44189 (1974). 78. G. V. Lyubinskii, V. V. Yanishpolskii, A. Tertykh, D. Yu. Yuodvalkite, and A. A. Glemzha. Ukr. Biokhim. Zhurn. 5624 (1984). 79. Yu. M. Torchinskii, Suyur in Proteins, Nauka, Moscow, 1977 (in Russian). 80. J. J. Pesek, in Che~ically Surfaces Pesek and I. E. Leigh, eds.), Royal Chemical Society, Cambridge, 1994, pp. 1-23. 81. A. M.Varvarin, L. A. Belyakova, A. Tertykh, R. Leboda, and Colloids Surfaces A 129 (1996). 82. L. V,Bereza, V. V. Yanishpolskii, and V.A. Tertykh, in Abstracts Second Int. Conf. Chemistry ~ i g h l y - o r g a n i ~ e dsubstance^^ and ScientiJic P ~ ~ n ~ ~ p l e ~ ~ N a n o ~ e c ~ n o l o St. g ~Petersburg, , 1998, p. 1.55. G. V. Lyubinskii, V. V. Yanishpolskii, and V. A. Tertykh. Ukr. Biokhirn. Zhurn. 55:499 (1983).

84. G. V. Lyubinskii,V.V.Yanishpolskii, and A. Tertykh. Teoret. Eksperirn. Khirniya 20:l 12 (l 984). 85. G. V. Lyubinskii and V. A. Tertykh. Teoret. Eksperim. Khirniya 26:201 (1990). 86. G. Lyubinskii. Ukr. Biokhirn. Zhurn. 56:390 (1984). 87. D. Varfolorneev and 1. V. Berezin. Molekulyarnaya Biologiya 10:818 (1976). 88. A. Tertykh, V. V. Yanishpolskii, and A. Ogiy, in~ ~ ~ o r p tand i o nA d s ~ r ~ e n t s , Vol.12(V.A. Tertykh, ed.), Naukova Durnka, Kiev,1984, pp. 68-71 (in Russian). V. Yanishpolskii, in Catalysis and Catalytic Processes in 89. V. A. Tertykh and P r o ~ ~ c t i oofn Chemico-Pharmaceutical ~reparations,Vol. l (K. K. Razikov, ed.), Moscow, 1985, pp. 137-140 (in Russian). 90. I. A,Butovich and V. V. Yanishpolskii, in~dsorptionand ~ d s o r ~ e n tVol. s , 12 (V. A. Tertykh, ed.), Naukova Durnka, Kiev, 1984, pp. 68-71 (in Russian). A. Tertykh, in Catalysis and Catalytic Processes in 91. 1. A.Butovich and Productio~ Chemico-Pharmaceutical Preparati~ns,Vol. l (K.K. Razikov, ed.), Moscow, 1985, pp. 141-144(in Russian). 92. V. D. Sernichaevskii, I. A. Butovich, and N. V. Tsarevich. Doklady AN UkrSSR, B NIO: 80 (1984). 93. I. A. Butovich and V. A. Tertykh. Ukr. Biokhirn. Zhurn., 56:437 (1984). 94. I. A. Butovich and V. A. Tertykh. Ukr. Biokhirn. Zhurn. 56:527 (1984). 95. A. Ogiy and V. A. Tertykh, in Catalysis and Catalytic Processes of ChemicoPhar~aceuticalProductions, Vol. l (IS. Razikov, ed.), Moscow, 1989, pp. 2124 (in Russian). 96. Yu. V. Mitin, N. P. Zapevalova, and E. Yu. Gorbunova. Int. J. Peptide Protein Res. 23528 (1984). 97. Yu. V. Mitin, E. Yu. Gorbunova, S. A. Ogiy, and P. Kul’. Bioorgan. Khirniya 21:709 (1985). 98. M. Arroyo, J. M. Sanchez-Montero, and J. Sinisterra.Enzyme Microb. Technol. 243 (1999). 99. R. Maeda, H. Tornida, M. Matsurnoto, and K. Kondo. J. Chern.Eng. Jap. 30:910 (1997). 100. E. Castillo, V. Dossat, A. Marty, J. S. Condoret, and D. Cornbes. J. Arner. Oil Chern. Soc. 7477 (1997). 101. K. Naka, K. Iwaki, N. Kitano, A. Ohki, and S. Maeda. Polymer 28:911 (1996). 102. B. Selrni and D. Thomas. J. Amer. Oil Chem. Soc. 75:691 (1998). 103. C. Stathakis, E. A. Arriaga, D. F. Lewis, and N. J. Dovichi. J. Chrornatogr. A 817:227 (1998). 104. G. Cartoni, F. Coccioli, R. Jasionowska, and D. Rarnires. Ann. Chirnica (Ital.) 88:319 (1998). 105. D. Reifsnyder, C. V. Olson, T. Etcheverry, H. Prashad, E. Stuart, and S. E. Builder. J. Chrornatogr. A 753:73 (1996). 106. Kaltenbrunner and A. Jungbauer. J. Chrornatogr. A 734:183 (1996). 107. J. Wang, D. S. Park, and P. V. A. Parnidi. Electroanal. Chem. 434: 185 (1997). 108. L. Chut, J, Li, and S. N. Tan. Analyst 122: 1431 (1997). 109. Kunzelrnann and H. Bottcher. Sensors Actuators, B-Chemical 39:222 (1997). 110. L. Coche-Guerente, Cosnier, and L. Labbe. Chern. Mater. 9:1348(1997).

L. Tiefenauer, S. Kossek, C. Padeste, and P. Thiebaud. Biosensors Bioelectronics Li, S . N. Tan, and H. L. Ge. Anal. Chirn. Acta B. Fubini, L. Mollo, S. Bodoardo, B. Onida, D. Oberson, and C. Lafurna. Langrnuir Fubini. Ann. Occup. Hyg. K. D. Lobel and L. L. Hench. J. Biorned. Mater. Res. K. Lobel and L. L. Hench. J. Sol-Gel Sci. Technol. K. M. Spark, J. D. Wells, and B. B. Johnson. Austral. J. SoilRes.

H. N. and E. Yeung. Science T. J. Su, J. R. Lu, R. K. Thomas, Z. F. Cui, and J. Penfold. .J. Colloid Interface Sci. L. Wannerberger, M. Wahlgren, and A. C. Eliasson. Cereal Chem. A.Kazakova,V. Gun’ko, E. F. Voronin, S. S. Silchen~o,and A. Chuiko. Kolloid. Zhurn. M. Gun’ko, TLKOV, V. I. Zarko, Dudni V. A. Tischennko, 0. A. Kazakova, E. F. Voronin, S. S. Silchenko, V. N. arvinchenko,and A. A. Chuiko. J. Colloid Interface Sci. E. Killrnann and Reiner. Tenside Surfactants Detergents B. Saoudi, N.Jarnrnul, M. Chehirni, G. P. McCarthy, and S. P. Armes. J. Interface Sci. V. A. Tertykh andV. Yanishpolskii,in ~ d ~ ~ o on r ~New ~ ~ o n ~ n o ~ g a n Sorbents ~c (A. Dabrowski and A. Tertykh, eds.), Elsevier, Amsterdam, pp. G. Sajonz, P. Sajonz, G. Zhong, and G. Guiocon. Biotechnol. Prog. unn, J.

Miller, B. C. Dave, J. S. Valentine, and J. I. Zink. Acta eteille, C. Roux, M. Chatry, and P. Davidson.Acta

C. Roux, J. Livage, K. Farhati, and L. Monjour. J. Sol-Gel Sci. Technol. rowiec-Bialon, Bielecki, B. Jarzebski, J. J. ~alinowski,A. d E. Galas. Biotechnol. Techniques S. A. Yarnanaka, N. P. Nguyen, B. Dunn, J. S. Valentine, and J. Zink. J. mer. Chern. Soc. B. andr rand.

ColloidInterfaceSci.

L. Henke, P. A. E. Piunno, A. C. McClure, and U. J. Krull. Anal. Chirn. Acta K. A. Melzak, C. S. Sherwood, R. F. B. Turner, and C. A. Haynes. J. Colloid Interrace Sci. D. Fetsch, M. Fetsch, E. M. P. Mendez, and J. Havel. Electrophoresis

K. Koopal, Y. Yang, A. J. Minnaard, P. L. M. Theunissen, and VanRiernsdijk. Colloids Surfaces A 141:385 (1998). 138. N. Labonne-Wall, V. Moulin, and J. P. Vilarern. Radiochirn. Acta 79:37 (1997). 139. Livage, C. Roux, J. M. DaCosta, I. Desportes, and J. F. Quinson. Sol-Gel Sei. Technol. 7:45 (1996). 140. A. V. Bogatskii, 0. V. Sevastyanov, T. I. Davidenko, V. V. Yanishpols~ii,V. A. Tertykh, and A. A. Chuiko. Doklady AN N8:656 (1979). 141. M. Habibi Rezaei and M. Nernat Gorgani. Int. J. Biochern.Cell 30:417 998).

137.L.

Institut fur Anorganische Chemie und Analytische Chemie, Johannes GutenbergUniversitat, Mainz, Germany

I. Introduction 11. Fundamentals A. Thermodynamic aspects of separation processes B. Adsorption isotherms C. Chromatographic separation processes D. Kinetic aspects E. How to achieve a chromatographic separation? 111. Characteristics of Native and Surface-Modified Silicas Employed for Chromatographic A.Properties of native and surface-modified silicas as bulkpowders properties of columns packed with native and surface-modified silicas C. Surface ~odificationof silica (chemical bonding of functional at groups the silica surface) D. Chromatographic processes on silanized silicas (reversed-phase CLC)

565 568 568 57 1 572 573 578 579 579 585 587 59 l

and IV. Conclusion

593

References

594

Silica is a well-known technical product in powdered form manufactured as both pyrogenic and precipitated silica The world production is estimated to be just over one million metric tons. The products are used as filler, lubricant, and adsorbent. During the classical period of chro~atographybetween the 1920s and the

1960s silica xerogels, made from waterglass and sulfuric acid, found widespread application as adsorbents. These lumpsof xerogel were milledand size classified to a mean particle size between 40 and 500 The columns filled with these irregularly shaped particles were operated in chromatographic separations by hydrostatic pressure or low pressure at linear velocities at about 0.1 mm/s. Due to the relatively large particle size the efficiency of these columns was rather low, which resulted in rather poor separations. etween theyearsl965 and 1975 theoretical concepts were worked out by various research scientists to lay the foundations for high-performance (column) liquid chromatography (HPLC). Thekey points developed were to use microparticulate silicas of xerogel type and of 5-10 average particle diameter, to pack them into stainless steel columns and to runthe colums at high flow rates about 1-5 1nm/s of linear velocity. During the primary stages of HPLC the coarse particles were milled to particles with the desired mean average particle diameter and sized into a narrow fraction by using novel technology based on air elutriation. Next, special procedures were worked out, called slurry techniques, to filter the suspension of the particles at high pressure and high flow rate through a stainless steel column in order to pack the column. The high flow rates were achieved by using high-pressure pistonor membrane pumps. Silica was superior to other adsorbents dueto its high mechanical strength, its flexibility in adjusting theporosity, the average pore diameter and the specific surface area. uring the decade1970-80 irregularly shaped silicas were increasingly displaced by spherical particles which were processed by specific procedures. However, sizing was still required. Silica beads were seen to generate a higher column permeability and more stable columns than those withirregularly shaped silica. Columns packed with 5-10 @m silica particles generated about 50,000 theoretical plates per meter column length. The use of native silica as applied in classical column liquid chromatography imposedanumber ofsevere limitations in terms of practical application. The equilibration of columns with nonpolar organic solvent mixtures took a relatively long time. When working with a nonpolarsolvent, for example n-hexane, the water content of the eluent and the silica had to be carefully adjusted and controlled. Furthermore, many sample mixtures appliedwere in aqueous solution and had to be extracted into an organic eluent for separation. The way to overcometheproblem andto handleaqueoussamples was to employ hydrophobic silicas. The latter were made by the silanization of native silica with reactive long-chain n-alkyl silanes, prefera~lyn-octadecyl silanes. The bonded n-octadecyl groups rendered the surface hydrophobic. These silica derivatives were called reversed-phase packings. Water-organic mixtures were employed as eluents. In reversed-phase HPLC, column equilibration was fast (about five to six column dead column volumes) and water was a component of the eluent. Chemical surfacem o d i f i ~ ~ t i o nsilica by silanization was studied in great depth Apart from hydrophobic ligands other typeswere also bonded (for example, polar, ionic, and biospecific to name a few). The separation of low-molecular-weight sample mixtures was performed with silicasof approximately l0 nm average pore sizebeing equivalent to a specific

This chapteris devoted to physisorption phenomena onsilica from solutions, under the aspect of adsorptive (chromatographic) separation processes. The solutes of interest comprise low molecular weight compounds between 500 and 2000 daltons of widely differing polarity, from hydrophobic to polar, and to ionic species, as well high-molecular-weight compounds, for example,synthetic polymers and biopolymers, respectively. The solvents employed span the full scale from nonpolar to polar aprotic, protic to buffers applied as single solvent, binary, ternary, and quaternary mixtures, respectively. The concentration of solutes ranges from picomolar (10"~ M)to molar. The separation processes are not performed under static conditions in a batch process. In contrast the solvent is pumped through a column bed containing the adsorbent in the form of an array of particles as a continuousbed at a constant flow velocity, The solution containing the solutes to be separated is injected discontinuously into the solvent stream applied as a continuous feed. The success of the separation in other words the chromatographic resolution is governed by the thermodyna~ics andkinetics as well as by mass transfer phenomena. Even at a high speed of the solvent solution through the column bed one will have quasi-equilibrium conditions. The major difference between a single batch operation and a chromatographic column operation istlnat equilibrium is repeated thousandfold in the latter. This results in a tremendous enhancement of separation efficiency and speed.

[l A brief review will be given of the thermodynamic aspects. The distribution of a solute i between two immiscible liquid phases y and x is designated as i,

il,

and is described by the equilibrium distributioncoefficients and The definition and the relationship between these is given by

g"

xx

where xx and

respectively.

(2) are the molefractions of solute i in the phasesx and y , respectively,

KC CY

where and are the molar concentrations respectively, and

of solute i in the phases x and y ,

(4)

where V" and are the molar volumesof compounds x and y comprising the two immiscible phases. The truly thermodynamic distribution constant is defined as

where and are the activity of solute i in phases and y , respectively. KO is related to the change of the chemical potential A p 0 as In

E;'

pot-'

where R is thegasconstant, theabsolutetemperature,andand chemical potentials of the solute in phases x and y , respectively. KO is a function of the chosen standard state and is related to

the as follows:

are the activity coefficientsof solute i in the phases where y x and respectively, and so

x and y

For the reference state of infinite dilution of each phase one obtains the relationship yx and 2"io

(9)

The change in the standard free energy AGO for adsorption is AGO

In9

(10)

A,XO

(1 1) where Aho and Aso are the changesin the partial molar enthalpy and entropy of the standard state. KCis related to a measurable quantity in liquid column chromatography known as the chromatographic retention coefficient k. The latter is defined as

where ti is the retention time (retention volume) of solute i and t, (v,) is the column dead time (column dead volume): ki where and Vm are the volume of the stationary phase and the mobile phase of the chromatographic system and is the phase ratio. The retention coefficient related to the ther~odynamicquantities as

where and ASo are the changes in the standard enthalpy and the standard entropy of the solute i.

The ratio of the retention coefficient of two solutes is defined as the selectivity coefficient

is directly related to the changes in the standard chemical potential of solutes and j :

where and are thedifferences in the standardchemical potentialof solutes and j between the two phases and y . great deal of interest was directed towards the estimation of the distribution coefficients, however, a detailed discussion of this subject would unduly expand the scope of this chapter. One of the classical approaches in liquid chromatography on native silica is the Snyder equation 121:

where, v, is the volume of the adsorbed phase as a monolayer per unit weight of silica, is the adsorbent activity parameter, So is a measure of the adsorption energy of solute on a standard adsorbentsurface with A, is the molecular cross-sectional area of the solute and is the solvent strength. Another classical concept is Hildebrand's solubility parameter theory, which was also applied in liquid chromatography on silica and its chemically bonded of a given solute is given by derivatives [13j. The distribution coefficient

where iji is the molar volumeof solute Si is the solubility parameter of solute i, S, and S, are solubility parameters of the stationary phase and the mobile liquid phase, respectively. For simple organic solutes the prediction of retention coefficients in column liquid chromatography can be performed by using the software package Chromsword from Merck KGaA, Darmstadt. This software package uses experimental data from known solutes on a distinct number of different columns and mobile phase compositions and applies increments of the molecular structure of a compound to calculate its chromatographic retention. number of models were proposed to describetheadsorption interactions between solutes and silica surfacesamongstthese are the Snyder-Soczewinski model, the Scott and Kucera model and the JaronieE model. These models differ, forexample, in thetype of underlying interactions, localized and delocalized adsorption, homogeneous and heterogeneous surfaces, and displacement phenomena. The reader is advised to consult the corresponding literature [14].

The adsorption of a solute from the liquid phase on an adsorbent at constant temperature characterized by an adsorption isotherm. with gas adsorption the amount of solute adsorbed is plotted against its equilibriul~ concentrationin solution and iscalled an apparent isotherm. In the case of strong interaction between thesurface and the solute aLangrnuirtype of isotherm is the result. During the processof separation when two solutes have to be resolved the composite isotherm is more informative that the individual isotherm. Adsorption isotherms are usually reported in the literature in the form of the surface excess isotherms where is plotted against xi, the latter being the liquid mole fraction of the solute Axi is the decrease in mole fraction of solute when moles of the original solution are equilibrated with m grams of the adsorbent. The resulting isotherms are termedcompositeisotherms,from which the individual isotherms of the single component can be computed. Figure 1 shows the types of composite isotherms accordingto the classification of Schay et al. 1151. A negative value of nOAxi/vnmeans that the solvent is preferentially adsorbed. Surface excess isotherms of binary solvents on adsorbents have been measured by chromatographictechniques [16-211. A comparison of various static and dynamic methods to assess adsorption isotherms from solution was reported. by Seidel-Morgenstern From the isotherm the distributioncoefficient, of a solute is derived as the slope of the isotherm at agiven equilibrium concentrationin the solution.By fitting theexperimental data of themeasuredisotherm with amodelequation,the

Composite isotherms according to the classification

Schay et

Langmuir equation, the dissociation constant and the association constant KA as well as the maximum adsorption capacity amaxare calculated. Figure 2 shows an experimental isothermof a synthetic peptide on aspecial type of silanized silica ( ~ i c ~ r o s p hADS e r C18, Merck KCaA, Darmstadt)

Chromato~raphicseparationprocessesare classified into the analytical and preparative. In analytical column liquid chromatography the aim is to achieve reasonable a resolution of the component of interest within the samplemixture atan acceptablepressure drop of thecolumn in a short enough time period,namelya few minutes. Analytical HPLC stems from the linear range of the isotherm where the distribution coefficient, KC,and the retention coefficient k are independent of the solute concentration:

K!

constant

Thus, thedilute sample is injected and resolved. If the isotherm takesa concave or a convex course towards the abscissa the slope of the isotherm changes as does the distribution coefficient This nonlinear regime is called the overload regime of

Bi-Lang~uir-~odeli Model: TwoSite Bind

concentration

(~~rnl)

Adsorptionisotherm of a syntheticpeptideonLiChrospher C18 (Merck, KCaA, Darmstadt).The experimentaldata were fitted by the two-site binding Langrnuir model. Solvent: phosphoric acid,pH 3.5, triethylamine 0.1 w/w

the system. It is in this regime that preparative and processliquid chromatography are preferably performed. The aimof this mode is totally different from the analytical mode. The targets are a high mass throughput per unit time with a high yield and high purity at reasonable cost. To achieve these goals, eluents with a high degree of solubility for the isolate and samples are chosen and injected at relatively high concentrations. When a concentrated solution of two compounds i and j is injected on a column, specific phenomena take place during the separationand are known as displacement and tag-along effects [8,22]. In the first case the elution profile of the first eluting compound is compressed by displacement of the strongeradsorbed late eluting compound (see Fig. 3). Due to the tag-along phenomena the first eluting compounds smear along through the whole elution profile of the second compound (see Fig. 4). The second phenomenon is not a favorable remedy for isolation, whereas displacement is very beneficial to maximize the throughput, In order to visualize the different situations in analytical and preparative column liquid chromatography two chromatograms are displayed representing typical solvations. Figure 5 shows an analytical chromatogram with a baseline separationof some adjacent compounds. Figure6 demonstrates the situationof a continuous chrornatographic process by simulated moving-bed technology; the two profiles heavily overlap. Only at the beginning and at the end of the profile are pure substances A to be seen, which are isolated by cutting fractions.

The dynamics of chromatographic separations are best described in the classical book of Giddings We will give a brief and simplified summary. The reader therefore, referred to the extended literature. When a solute travels or migrates

(mio)

Comparison of elution profiles of a binary system with Competitiveadsorption (thick lines) and without competitiveadsorption (thin lines) at a concentration ratio of component l to component 2 of (From Ref. 22 with permission of the publisher.)

57

0.05

4 time

Comparison of elution profiles of a binary system with competitive adsorption (thick lines) and without competitive adsorption (thin lines) at a concentration ratio of component 1 to component 2 of 9:1. (From Ref. 22 with permission of the publisher.)

AmWflSSER: %B= 1 6 - 0 . 9 ” 1 8 - 7 . 5 ‘ - 9 5 - 4 . 9 ’ - 9 5 -TEST

n

2

Separation of a synthetic mixture on ananalytical column. Test compounds: thiourea, benzamide, aniline, phenol, phenyl ethanol, acetophenone, nitrobenzene, benzoic acid methyl ester, benzene, toluene, ethyl benzene, propyl benzene, anthracene, n-butyl benzene, n-pentyl benzene, n-hexyl benzene, n-octyl benzene, n-nonyl benzene, n-decyl benzene,and n-dodecyl benzene.(From Ref. 26, with permission of GIT Verlag, Darmstadt.)

Elution profiles of a two component mixture A and B on a 16 column simulatedmovingbed chromatograph. (Courtesy of Dr. Schulte,Merck KCaA, Darmstadt.)

through a column with a dense bed of particles a dispersion of the initial profile takes place through several phenomena (Fig. 7):

Eddy i.e., convective mixing in the interstitial voids of the packed bed L o ~ g i t ~ d i ~ ~ Z i.e., diffusion along the column axis transfer of the solute between the mobile phase and the stationary phase. The peals dispersion of a column is expressed in terms of the height of the theoretical plate H as

curve

Mass transfer

Eddy diffusion Longitudinal diffusion U

Dependency of the total plate height, H , on eddy diffusion, longitudinal diffusion, and mass transfer.

where is the standard deviation of the peak in length units and L is the column length. In other words,H equals the increaseof the variance of the peak dividedby the column length. can be calculated by different methods from the peakprofile. To obtain a dimensionless quantity, the reduced plate height h, the theoretical plate height H is divided by the average particle diameter

The plate height concept stems from classical unit operations, such as distillation, etc. With the known column length the number of theoretical plates can be calculated:

H N is given in number of plates per unit column length. The objective in chromatographic separation (analytical mode) to achieve a high plate number and correspondingly alow plate height. The total plate height of the column Htotal is assumed to be a sum of the three contributions, as discussed before:

B

CU

U

where U is the linear velocity of the eluent. H is a hyperbolic functionof U with the three terms A, B, and C reflecting the single contributions. Figure 8 presents a measured curve It observed that the diffusion term B dominates the column efficiency at low linear velocities, while at higher U the C term becomes dominant. The minimumH is controlled by the eddy diffusion termA. The latter is equivalent to about in other words inHPLC columns them i n i l ~ u nplate ~ height H cannot be smaller than twice the averageparticle diameter. It is also apparent fromFig. 8 that the H versus U dependency is independent of the value of the retention coefficient in the range between 1 k 10. To increase theanalysis time the linear velocity has to be increased aswell. With higher linear velocities, however,thecolumn efficiencyis impaireddue to the limitations of the mass transfer kinetics. circumvent these disadvantages silica columns with a continuousbed were manufactured and exhibited a higher porosity with respectto packed columns dueto the presenceof large macropores of 1-3 pm. Thesecolumns exhibit a flat versus U curve even at hi her rates. The commercialcolumns of monolithic type, e.g.,Silica Rod"of Merck KGaA, Darmstadt generate a column performance equivalent to a 5 p m silica, however, at five times lower pressure drop at comparable conditions and scarcely a loss in column efficiency at linear velocities above 1 mm/s

1

r-----"

I

5

2

8

8

9

H vs. U curves of five solutes on a Purospher@RP1 8 column (MerckKGaA, ~armstadt).(From Ref. 25.)

In analytical HPLC spherical silica beads of about 5 p m average particle diameter are conveniently employed. The column length varies between 100 and 200 mm with an inner diameter of 4 (4.6) mm. Such columns generate about 10,000 theoretical plates per column length, which is sufficient to resolve most of the sample mixtures. To shortenthe analysis time even further, particles ofsize were employed in short columns of 40-70 mm. The reduction in column length is mandatory toavoid a too high column pressuredrop. These short columns allow oneto separate sample mixturesin a few minutes [26]. further reduction in the particle diameter to about 2 p m has already been accomplished. The column length in this case is about 30 mm or less. The number of theoretical plates generated is in the order of several thousands. The speed of analysis isvery fast and in theorder of 1 min or less. Inordertooperate suchcolumns, specifically designed HPLC equipment is needed, which has minimum dead volume contributions and a short detector response time as well a high data sampling rate [S]. The injection volume should be about 0.5 The particles can be porous or nonporous. In preparative and process HPLC efficiencyis a less importantparameter. Typical particle sizes for preparative HPLC are between 15 and 30 pm for silica. For continuous chromatography, suchas simulated moving-bed separations the total number of required theoretical plates amounts to several hundred.

In analytical HPLC, optimum chromatographic resolution is the main goal, i.e., baseline separation between two adjacent elution bands. Chromatographic resolution is defined as:

where ti and ti are the retention time of solute j and i, respectively, and oli is the standard deviation of the first eluting peak in time units. baseline separation (2% overlap of the two peaks) is equivalent to 6otiassuming a Gaussian profile and the same peak area o f j and i. The relationship between and the chromatographic parameters is essential:

where 1 is the selectivity term, k i / ( l ki) is the retardation term, and is the efficiency term. The most important term with regardto resolution is the selectivity term. must be larger than unity to achieve separation. When k varies in the range 1 k 10 the second termis nearly constant. Itbecomes dominant when the retention falls to k 1. Higher k values of 10 do not improve the resolution. The column efficiency plays a secondary role. For example, by increasing by a factor of two the resolution improves only by a factor of 1.4. The general procedure to attain a chromatographic separation is as follows: 1. Chooseacolumnwith an appropriatepackingandstationaryphase. It is recommended to start with a high quality n-octadecyl-bonded silica. Choose a binary eluent, e.g., water-acetonitrile and adjust the composition in sucha way that the components of thesamplemixture elute within 2-1 1 column dead volumes (1 k 10). In the case of acidic and basic cornpounds use buffer instead of water and adjust the pH in such a way that for acids the pH pk, 2 and for bases the pH pkb 2. 4. After retention control, the mainobjective is to achieve resolution. The shift of the peaks at the chromatogram willbe affected by the eluent composition. Ternary quaternary solvent compositions shouldbe employed pH shifts. It may be the case that some peakselute with tailing. For basic compounds use specifically base-deactivated silica columns. Add a mM amount of detergents, e.g., N-heptane sulfonic acid for peak shape improvement.

In preparative HPLC a substantial degree of information is needed to select a proper variant of column liquid chromatography: What is the solubility of the product to be purified? Usually a solubility larger than 20 mg/mL is required. 2. Is the solvent identical with the eluent to be employed at isolation? 3. Select a column and eluent which enable the maximum degree of selectivity.

4. What is the loadability of the column with the eluent chosen? Howlarge is the linear range of the isotherm? 5. Use volatile eluents. 6. What is the best mode of operation with respect to the sample throughput, purity, and yield: repeated injections of small amounts, recycling with peak cutting, prefractionation of impurities, or by-products?

Porous silicas manufactured for analytical, preparative, and process-scale separations by HPLC are high-purity products as compared to technical-grade silicas being applied as fillers, lubricants, andadsorbents.Thetermpuritycomprises inorganic constituents as well as organic residues. For chromatographic applications the metal content is most decisive because it affects the adsorption of complexing analytes at the silica surface. The characteristics of silicas are usually discussed under twoaspects: firstly, the properties which characterize the bulk powder of the native or surface modified silica; secondly, the propertieswhich characterize the column packed withsilica, as well as its chromatographic parameters. The former group of properties are of more interest in terms of the producer of the silica and the silica column manufacturer. The latterset of properties are of vital interest for the user, thus, in other words the chrol~atographer’s judgment on the performance of the column for a given separation task. We will follow this discril~ination forpractical reasons.

Physical P r o ~ e r t i e Particle ~: ~ o r ~ h o l o Average ~y, Particle Size, and Particle Today, spherical particles are exclusively used,whicharemade by specifically designed and elaborated processes. While in 1980, 5-10 silicas were employed as packings, therewas a clear tendency in 1990 to preferably use 5 p m packings (see Fig. 9). The trend is now toward 3-4 particles; 3 particles are the smallest which can be size fractionated by the current sizing technology at economically viable atches with smaller particles contain submicron particles which strongly adhere to larger particles and which are difficult to remove or separate. When quoting the particle diameter of a packing it is recommended to give the average particle diameter as the volume average and not as the weight number or surfaceareaaverage,becausethevolume of a chromatographic column is the decisive parameter. It is also essential that the particle size distribution covers a certain range or width. The size distribution is usually characterizedby the ratio of the to the value. The (dplo) value corresponds to the averageparticle

5

tion: left

Scanning electron micrograph of 5 right 1 0 0 0 ~ ) .

(clp5@

6.0

silica (magnifica-

diameter at 90% (10%) of the cumulative particle size distribution. The ratio should range between 1.5 and as shown in Fig. 10. roader particle size distributions impair the column performance in two ways. Particles at the lower end of the distribution, so-called fines, lead to an increase of the column pressure or might even cause a blocking the outlet sieve. To detect fine particles of a packing, particle analysis should be carried out with an instrument which coversthe size range between 0.1 and 5 The size distribution should be plotted based on the number average for better visibility. The presence of large particles at the higher end of the distribution impairs the column performance because the column plate number is inversely proportional to at otherwise constant conditions. Theaverage particle diameter of apackingcontrolstwoimportantcolumn performance parameters: 1. The pressure drop Ap across the column as

where is the viscosity of the eluent, U is the linear velocity of the eluent, is the length of the column, and is the column resistance factor, usually between 500 and 1000 for well-packed columns. The pressure limit of common HPLC equipment is about 400 bar. However, HPLC columns cannot be operated constantly under such high pressures. After a certain period the column bed is compressed, voids and channels are formed, and the column efficiencyis impaired or the column bedeven collapses. The usual column working pressure is between 50 and 100 bar. Whenreducingthe particle diameterfrom 5 to 3 pm thecolumnlength is simultaneously reduced from 100-150 mm to about 50mm to avoid too high a back pressure.

Differential and cumulativeparticlesize (method: Coulter counter, volume average).

distribution of native silica

2. The plate number of the column The column efficiency expressed in terms of number of plates is related to the particle diameter

Most separation problemsin HPLC can be solved with column generating about 5000 plates, for example employing columns of l00 4 (4.6) mm packed with5 p m particles. Columns with 3-4 p m particles are used for fast separations with analysis times of less than 5 min. To eliminate the tedious and laborious sizing and column packing procedures monolithic columns havebeen developed with continuous bed Such columns consist of porous silica rod which is cast in a plastic column. The rods have through pores of about 1-2 p m and diffusion pores in the order of about 10 nm. Thus, silica rods have the samespecific surface as microparticulate 5 p m silicas and show the same retention behavior,when they are chemically modified, to reversed-

surface area, (BET), of about 300 m2/g. Due to the accumulated knowledge in chemical bonding of biocompatible and biospecific ligands to the silica surface in thebeginning of the 1980s, such species becameattractiveadsorbentsforthe separation of biopolymers [4]. avoid the hindered diffusion of dissolved biopolymers in the chromatographic separation and toenable full access to the stationary phase within the pore system, silicas with pore diameters larger than l0 nm were required. This led to the l~allufactureof wide-pore silicas with pore diameter (pd) 50 nm. The corresponding reductionin the specific surface areawas not a critical parameter for biopolymer separation. It could even be demonstrated that approximately 2 pm nonporous silica beads with an appropriate surface modification were able to resolve biopolymers with short analysis times of less than 1 min During thelast decade the qualityof silicaadsorbents was further enhanced. The improvements consisted of processing silicas with a higher purity and stability as well as a more reproducibleand stable bonding of ligands. Today thetypical batchto-batch reproducibility for a typical silica employed in HPLC is as follows [6]: Specific surface area, Specific pore volume, diameter, Average pore

330 f 10 m2 g 1.1 f 0.2 cn1 /g 12 f 1 nm

i

Joyner, and Halenda)

The average particle size of silicas today is about 5 pm, which results in a column efficiency of 100,000 plates per meter column length. number of suppliers offer short 4-7 cm columns with 3-4 km silicas for fast HPLC separations [7]. Besides analytical HPLC, which is now a mature separation technique, preparative HPLC became a purification technique of interest for the manufacture of reference compoundsand the isolation impurities and by-products. The progress made in the manufacture bulk quantities of selective silica adsorbents, in the assessment of individual and composite isotherms, in modeling and Simulation of the therl~odynamicsof ligand-solid adsorption systems and masstransfer kinetics, as well as reinvented continuous procedures, such as simulated moving-bed technology, led to an enormous growth in the application of HPLC in process chromatography a result, highly efficient systems are commerciallyavailable to purify value-added compounds in 10-100 g quantities daily at a reasonable cost. One striking example is the production of enantiomers from racemates on chiral selective paclcings at process scale. During the last decade a hybrid method emerged called capillary electroendosmotic chromatography (CEC), known more concisely as electrochromatography. Columns are fused silicaof100 to 500 mm length and about 100 p m inner reversed-phase silicas. Instead of using H diameter packed with 3 rnent with a pump, modified capillary electrophoresis equipment is used where a voltage up to 30 is applied between theends of the capillary column.The analytes injected aredriven by an electroendosmotic flow (EOF).Resolution takesplace and depends onthenature of theanalyte(neutral,polar, ac basic), on the partition solute-surface interactions and on electrornigration. to the absence of a hydraulically generated flow and the corresponding column pressure drop, particles of 3 p m and smaller canbe used, which generate morethan 200,000 theoretical plates per column length [lo].

phasepackings.Therodscommercialized by Merck CaA, Darmstadt, showa similar efficiency as 5 p m particulate columns with tw major differences 1251: the column performance remains unchangedat higher flow rates up to about7-8 cm3/ min where a microparticulate 5 column cannot be operated. Furthermore, the column pressure drop on silica rod columns is about five times less than on microparticulate columns at otherwise constant conditions. Silica packingswith 1-2 p m and smalleraverage particle diameterare employed in packed fusedsilica capillaries in capillary electrochromatography (CEC), Instead of apressure-driven flow an electroendosmotic flow (EOF) is a voltage up to 30 kV [28,29]. CEC is a hybrid technique between micro and capillary electrophoresis (CE) and combinestheadvantageous features of both the high column efficiencyof CE and the o~~tstanding selectivity of HPLC.

ace

Pore ~tructura~

The pore structure of a solid is generally characterized by its porosity, the pore diameter, the poreshape,thepore connectivity, thepore size distribution,the specific porevolume, and the specific surfacearea.Theseparametersare obtained from sorption measurements with nitrogen, intrusion techniques, radiang, pycnometry, thermoporometry, calorimetric measurements, spectroscopy, andothertechniques [30]. Thecalculationmethods are based on certain assumptions regarding the pore shape and the mechanisms of pore filling. The term porosityrefers to anopen porosity being accessible to molecules in the gas phase or liquid phase. With respect to silicas the framework density is used to differentiate between porous and nonporous solids. According to the I ~ P A C classification 1301 porousmaterials are divided into microporous,mesoporous, and macroporous according to the value of the pore diameter: icroporous solids: porediameter esoporous solids: porediameter2 acroporous solids: pore diameter

2nm 50nm 50 nm.

Pd

Pd

Silica supports in HPLC have mesopores rangingbetween 6 to 50 nm. The presence of micropores leads to slow desorption kinetics and a lower loading capacity than for supports with mesopores. The specific surface area and the specific pore volume of mesoporous solids, such as silica, were calculated from adsorption isotherm measurements. The adsorption isotherms, which result mainly from physical adsorption of nitrogen, may convenientlybe grouped into five classes,namely types Ito V of the classification originally proposed by Brunauer et al. 1311. A typical isotherm of a mesoporous silica is shown in Fig. 11. It is of type IV. A characteristic feature of a mesoporous solid (type IV isotherm) is its hysteresis loop. The type of a hysteresis loop depends on the pore type 1321. For the calculation of the poresize distribution with the assumption of cylindrical pores the Kelvin equation [33,34] is suitable. A pore size distribution is shown in Fig. 12.

Isotherm of

mesoporoussilicasample.

Monomodal pore size d~stributionof a mesoporous silica sample.

Typical values

mesoporous silica are

50-500 m2/g vp(G) pd(BJH)

0.5-1.2 cm3/g 4-20

During the last decade a new family of high-purity silicas has replaced the traditional silicas which were mainly synthesized from waterglass solutions and acids. The latter contained large amounts of metal impurities, such as sodium, magnesium, calcium, iron, titanium, and aluminum [35]. Typical values are sodium 1000 ppm, aluminum 200 ppm, and iron 150 ppm. The high-purity silicas are mainly synthesized from tetraethoxysilane, high-purity colloidal silicas, or silicas which were acid-washed after synthesis. The metal content of these silicas is reported to be lower than 20 ppm. Analytical methods to assess the metal contentat these trace levels are: neutron activation analysis, atomic adsorption spectroscopy, voltammetry, and optical emission spectroscopy. Metal impurities in the bulk and at the surface of silica give rise to a chemical surface heterog~neity,which leads to irreversible adsorption of basic compounds ortailing of the peaks.

Native silica is a hydrophilic adsorbent and adsorbs about to 10 (w/w) physisorbed water when exposed to humid air. The silica surface bears hydroxyl groups which can be divided into three types according to their co-ordination to surface silicon atoms: free or isolated, geminal, and vicinal (hydrogen-bonded)groups. Theycan be identified by IR spectroscopy, 2oSi CP MAS NMR spectroscopy, and other physicochemical methods [36-381. At a fully hydroxylated state [36,39] the surface concentration of hydroxyl groups, aOH,amounts to about8 to 9 ,umol/ m2. Physisorbed water is usually removed from the silica surface by subjecting the silica to heat treatment under vacuum at 150°C. With increasing calcination temperature the surface hydroxyl groups condense to siloxane groups, whereby is wat evolved. A drastic decreasein aoH is observed in the range between 200 and 400°C. Above 600°C mainly free or isolated hydroxyl groups condense. Rehydroxylation of the silica surface by exposingthedehydroxylatedsurface to watervapor is reversible up to acalcination temperature of about 600°C. Above 600°C rehydroxylation becomes a slow process or remains even incomplete [40]. Exposingporous silica to water or watervapor at an elevatedtemperature causes changes in the pore structure of the silica via hydrothermal treatment. As a result the specific surface area diminishes, the average pore diameter increases; however, the specific pore volume remains constant. The magnitude of the effect depends on the temperature and on the duration of hydrothermal treatment The surface hydroxyl groups are weak Broensted acids with apk, value of about 7 However, depending on the method of determination the plc, value has been reported to vary between and 10 [38]. Another reason which might cause the wide variation of the pk, value originates from metal impuritiesin the silica. Silicas made from waterglass or from colloidal silica often contain significant amounts of elements, such as sodium, aluminum, and titanium. In the presence of aluminum strong Broensted acid sites of the type Si OH -A1 can be formed, as encountered in crystalline MFI type zeolites [41]. Acidic impurities might also be present which are left behind from the postsynthetic treatmentor from theparticle shaping

processes. The pH of zero charge of the silica particles, pH,,, is related to the plc, value(s) according to

The pH of zero charge of silica varies between 1.5 and 3 [42]. Above the pH,, silica acts as a weak cationexchanger, below this valueas an anion exchanger.The following ion exchange processes have been postulated by Allen and Matijevic [43]: =Si-OH

OH"

=Si-O"

I-

pH

8

By increasing thepH of a silica suspension from 3 to 8 the surface becomes increasingly negatively charged. This is evidenced by the dependence of the electroendosmotic velocity on function a of pH of packed silica capillaries in electrochromatography [44]. While deprotonation of the surface hydroxyl groups is a fast process, the protonation is rather slow, for example, a pronounced hysteresis is observed in the titration curve of silica with sodium hydroxide and hydrochloric acid, respectively. Above pH 8 silica starts to dissolve as soluble silicates. Above pH 10 the solubility increases exponentially [45]. Previously used and chemically modified porous silicas are subjected to an activation processwhich consists of either a thermal treatmentand/or a treatment with a concentrated solutionof an acid (except hydrofluoric acid) to decrease the surface heterogeneity and to remove inorganic impurities from the surface [46].

Prope~~es The hydrodynamic properties of a separation column are usually characterizedby the column permeability, the columnresistance factor, Q, and the dependencyof the column pressuredrop as a function of the volume flow rate, F, which is related to the linear velocity, by

l.

being the free cross-sectional area of the column. It is preferable to use rather than F, because U is independent of the cross-section of the column. The column per~eabilityis defined as Ef

where is the viscosity of the eluent, L the column length, and dp the mean particle diameter of the packing. usually amounts to 500-1000 for well-packed columns. It should be noted that the column "permeability is equal to the square of the average particle diameter. The permeability of the column can change during column operation due to fracture the particles, formation fines, changes and

rearrangement of the column bed, and particle blocking of the frits at the column top and end. As indicated by Eq. (26) the column pressuredrop, Ap,is linear function of the flow rate, this is shown in Fig. 13, whichcompares the Ap versus curves of two types of silica columns. Thefirst one is packed with 5 pm silica particles, the second one is a monolithic silica column. At a flow rate of l cm'/min the latter shows a fivefold decrease in the pressure drop than the particulate column. The reason for this is the higher porosity, of the monolithic column as compared to the microparticulate column (0.78 versus 0.45).

2. Column efficiency is characterized by the theoretical plate height, H , and the nurnber of theoretical plates, The relationship between and H is L H

(3 1)

where is the column length. It is expressed in length units, preferably in is related to the unit length of the column. The plate height, H , is dependent on the linear velocityof the eluent. The dependency has a hyperbolic shape. At low values H decreases with increasing linear velocities until a minimum is reached. H increases at higher linear velocities.

6

flow

F

Column back pressure, as a ic columns and a microparticulate c ment (without a column), 18e column. (From Ref. 25.)

e flow rate, F, on three

7

The total plate height, H , is assumed to be a sum of the contributions

H

B

cu

where a term reflecting the contribution of eddy diffusion, the term B characterizes the contributionby longitudinal diffusion, and C is the mass transfer term being dependent 011 the kinetics of mass transfer of solute between the stationary and mobile phases. For practical use it is advisable to operate the column at a linear velocity where H is at its smallest. This corresponds to about 1-2 mmjs. At the optimum linear velocity, uOpt, the total plate height corresponds to about two to three times the mean particle diameter of the silica packing.Thisenablesone to calculate the expected number of theoretical plates at a given and L value at otherwise constant conditions. For dp and L 100 mm, N is calculated to be about 10,000 and 7000, respectively. With this column performance the majorityof sample mixtures canbe resolved. The analysis time, t R , being equal to the retention time of a given analyte is given by the following equation:

where is the retention coefficient and D, the diffusion coefficient of the analyte in the eluent. According to the equation the analysis time is equal to the plate number and the square of the mean particle diameter of the packing.

Column Ret~ntion

Selectivity

Retention expressed by the retention coefficient, should be kept in the range between and 10. Small coefficients impair the resolution, whereas larger h. values enhance the analysis time. The retention behavior of the column is assessed by applying test mixtures and standardized conditions. A numberof tests are proposed to monitor certain properties of the column, such as its hydrophobicand silanophilic character and its shape selectivity towards test solutes. Not only the retention but also the peak symmetry assessed. The same test mixtures are also used to manifest the selectivity of a given column and to highlight the differences between colurnm prepared by different manufacturers. It should be emphasized that columns canbe grouped into those for general applications and those servingto resolve certain groups of compounds or even single analytes.

Among the chemically bonded silicas, silanized silicas have achieved major importance as lipophilic adsorbents in HPLC. Thus, the focus of this section is directed to this type of adsorbent.

1. of Silanizationmeansthe useof reactiveorganosilanestointroduceorganosilyl groups to reagent. In the case of silica the surface hydroxyl group and adsorbed water at the surface,respectively, react with the organosilane formingeSi-OSi- R bonds to make the hydrophilic surface hydrophobic. The objectives of thesurfacereaction of silica with an organosilaneare follows: To change the original hydrophilic surface into hydrophobic one To fine-tune and to tailor the hydrophobicity of the surface by the type of the organo-group and its surface concentration To create chemically stable bonded layer To minimize matrix effects by effectively shielding the parent silica surface. The hydrophobicity of the silica surface proportional to the carbon content after modification. The carbon content canbe converted into ligand densityof bonded groups per unit mass, unit surface area,or unit volumeof the bonded silica. Carbon moieties are n-alkyl or aromatic groups, for example, phenyl. Silanized silicas with hydrophobiccharacterare employed as packings in reversed-phase HPLC with solvents ranging from mixtures of organic solvents, mixtures of water or buffer with organic solvents and buffers. One should keep in mind that the term reversed-phase packings does not only refer to silanized silicas. Reversed-phase packings can be: Silicas with a coating of a hydrophobic polymer, for example, polystyrene Oxides other than silica with hydrophobic polymeric coating, for example, alumina, titania, zirconia Cross-linked hydrophobic organic polymers, for example, polystyrene divinylbenzene Porous graphitized carbon. Bonded hydrophobic silicas representonegroup of reversed-phasepackingsin HPLC. They are named according to the bonded hydrophobic groups such as Ealkyl- and phenyl-bonded silicas. The abbreviations employed are C4-, C8, and C18-bonded silica. The n-alkyl groups such n-octyl and n-octadecyl are called ligands. The most common ones are n-octadecyl ligands. The bonding of these ligands to the silica surface can occur via different reactions and to varying degrees. The latter is measured by the carbon content through carbon elemental analysis. Typically, the carbon content of E”alkyl-bonded silicas varies between 10 and 20% (w/w) carbon. The carbon contentis simple measure of the hydrophobicity of the bonded silica; the higher the carbon content of a silanized material then the higher willbe the retention of hydrophobic analytes in the reversed-phase column. At a known stoichiometry of the surface reaction one can calculate the ligand concentration or the ligand density of silanized silica; however, the molecular weight of the n-alkyl ligand and the carbon contentof the bonded silica needs to be known.

The ligand density, a, is then calculated according to the equation

a=

P, lo6 1200n,

l

n,)

where a is expressed in pmol/m2, P, is the measured carbon content % C (w/w), n, is the number of carbon atomsin the bonded silane,A4 is the molecular of the silane, n, is thenumber of functionalgroupsubstituents in thesilane molecule, and is the specific surface area (BET). The maximum ligand density achieved with n-alkyl groups was reported [48j to be about 4.2 pmol/m2. When silica surface is fully covered with hydroxyl roups thesurfaceconcentration aOH amountstoapproximately 8 pmol/m This means that only half the original hydroxyl groups have reacted under otherwise optimum conditions. In other words, the ligand densities of the n-alkyl groups and the r e ~ a i n i n ghydroxyl groups are in the same orderof magnitude. Commercial nalkyl silicas exhibit ligand densities in the order of about 3 pmol/m2. In order to the remaining hydroxyl groups the silanized silicas are subjected to second silanization with reactive short-chain organosilanes. This procedure is known endcapping. Some commercially available silanized silicas are even double endcapped. For a clearer picture silanized silicas have been classified into several families. Brush-type bonded silicas are those which resemble monolayer of bonded n-alkyl groupsanchoredtothesurface. Incontrastto thistype,the polymeric group possess more or less dense polysiloxane layer with terminating n-alkyl groups. The brush type and polymeric type of silica-based reversed-phase packing behave differently in shape selectivity towards hydrophobic analytes. It has been shown that long flexible n-alkyl groups are an essential requirement for chromatographic separation with respect to retention, selectivity, and fast transfer kinetics. This principlehasalsofoundapplicationinthe design of cross-linkedhydrophobic organic polymers, for example, in commercial packings in which the n-octadecyl ligands are attached to polystyrene backbone. To overcometheintrinsically limited pH stability of bonded silicas andto improve the elution pattern of basic analytes in reversed-phase HPLC the classical n-alkyl-bonded silicas have been replaced during the last decade by the improved so-called high-quality silicas. These include stable bond packings and base deactivated silicas. Improvement is based on the development of high purity and more stable parent silicas and on more effective bonding chemistry. Parallel to the advancement in materials science and technology the quality of silicas and bonded silicas for HPLC will be further improved in the future.

F

anes which are employed in the surface reaction with the silica bear two important constituents:

A reactive group, such chloro or alkoxy, which is capable of reacting with the hydroxyl groups of the silica surface

functional group R, such as n-octadecyl, which mainly controls the hydrophobic properties. In principle the n-octadecyl group can be introduced by the use of one of the following silanes: I

I

-Si-R

R'

x

I

X-Si-R X-Si-R X

onofunctional Bifunctional Trifunctional where X is a chloro alkoxy group and R' is a methyl or n-octadecyl group. hen synthesizing the corresponding three packings the hydrophobic chromatographic properties will be about the same at constant carbon content; however, theretentiontowardsthepolaranalytesunderreversed-phaseconditionsmay differ. The bonding of the three aforementioned silanes at the silica surface is performed under different reaction conditions taking into account the differences in the reactivity towards the hydroxyl groupsand water. A typical reaction procedure is as follows. Silani~ationwith n-octadecyldi~ethylchlorosilane. First, 100 g of rehydroxylated, dry silica (l 50"G, 12 hPa, 15 h) is placed in a previously evacuated and nitrogen flushed L three-necked, round-bottomed flask. The latter is equipped imroth condenser with a drying tube, a 500 dropping funnel, a stirrer bearing, a stirring shaft and astirring motor. Then400 cm3 of dry toluene is added whilst stirring (speed: 3.5). The flask is immersed in a preheated oil bath (electric thermometer: heater Z O O T ) . The toluene is heated to reflux and 16.4 (0.24 moles) of imidazole is added while stirring. After 10 min a solution of 86.4 (0.24 moles) of melted n-octadecyldimethylchlorosilane in 200 mL of toluene is added to the silica (while stirring) over approximately 15 min, using a dropping funnel. The solution is stirred under reflux for 1 h, after which the oil bath is removed and the suspension is allowed to cool for 30 min while stirring (in most cases a cluster of imidazole hydrochloride is formed; the suspended silica is decanted off from the cluster and filtered through a 1000 cm3 glass filter). The silanized filtered silica is sequentially washed with 200 mL of toluene, 300 mL of methanol, 300 mL of methanol-water (50/50 v/v), and finally with 250 mL of methanol. The residue consequently subjected to tetrallydrofuran (THF) treat nt by pouring it into a 1 three-necked, round-bottomed flask equipped with a mroth condenser, a stirrer bearing, astirring shaft and astirring motor. Then 500 mL of tetrahydrofuran is added and the suspension is stirred (speed: 3), under reflux, for 20 min (electric thermometer: 80"C, heater: 150OC). After removal of the oil bath, the hot suspension is filtered and washed once with 300 mL tetrahydrofuran and twice with 300 mL methanol. Finally the product is dried under vacuum at 40°C. A decisive feature in the bonding chemistry of an n-alkyl-bonded silane is the reactive group of the silane, be it chloro, methoxy or ethoxy, respectively [49].

-N(CH3)2

-0COCF3 >C1 >>-OCH~X-OC~H~

The kinetics of silanization are enhanced by catalysts and scavengers for acids in the case of reactive chlorosilanes. Most importantly the water contentof the silica and the reactants including that of the solvents must be controlled. If this is not the case then nonbonded products are formed, such as disiloxanesand oligomer siloxanes, which are different to extract and which might slowly be removed.

[50] Silanized silicas are characterized by: Carbon content Ligand density Hydrophobic selectivity Silanophilic selectivity Shape selectivity Stability under defined elution conditions.

In reversed-phase systems, adsorptionof solutes takes place from a polar solvent on a lipophilic surface. The typical, classical system is the adsorption of organic compounds from water on activated carbon adsorbents, as still performed in water purification processes. Activated carbons tend to be applied in adsorptive highresolution separation processes in which two main disadvantageous features are combined. The micropores create a high adsorption potential with a short linear range at the initial part of the isotherm. Due to the presence of micropores the desorption kinetics are rather slow and severe tailing of the chromatographic peaks occurs. The most hydrophobic surface known together with a high degree of orientation is graphitized carbon black. However, such graphitic materials, even when agglomerated to larger microsized particles, do not attain sufficient stability. An elegant way to process porous graphitizedcarbon was elaborated by Knox and his coworkers. Theyused a porous silica as aprecursor and completely loaded it with a polymeric resin. After activation the resin was converted into a silica-entrapped activated carbon. Dissolution then removed the silica and the product was subjected to a graphitization process. Thus, in this manner the silica was employed as an imprint for the graphitized carbon. The material went on to become a commercial product manufactured by Hypersil Ltd., Cheshire, IJK. Based on the rich knowledge on silanization toprotect hydroxylgroups in synthetic organic chemistry, the silanization ofsilica surfaces was developed in the mid 1970s. The most popular silanes were the reactive n-alkylsilanes containing n-octadecyl (C18) or n-octyl (C8) ligands. The bonding chemistry of these silanes with the surface hydroxyl groups was studied in great detail. This information is collected in a series of monographs. The reaction consists of the formation of one or two surface siloxane bonds during which hydrochloric acid isreleased when chlorosilane reagents are used.Specific recipes weredeveloped for monofunctional, bifunctional, and trifunctional silanes. The reaction is correctly named a silaniza-

tion reaction because an organosilyl moiety is bonded to the silica surface. bonded silicas are the most widelyused adsorbents in reversed-phase (RP)CLC today, followed by C8-bonded silicas. The interaction between the solutes in an aqueousoraqueous-organicsolvent and the lipophilic surfaceare collectively called solvophobic orhydrophobicinteractions [51]. More precisely, there is a driving force of an organic solute from its water solvent structure to be released for attraction to alipophilic surface, whereby an association is formed between the solute S and the hydrocarbon ligands L:

The strength of association is determined by the total hydrocarbon surface areaof the solute as well as the total hydrocarbon surface area of the n-alkyl ligand and the surface tension of the solvent. Quite simply one can say that the interactions increase with the hydrophobicity of the solute and the carbon content of the silanized silica. Thus, elution increases with enhanced hydrophobicity of solutes for an homologous series for a given silanized silica. There has been on ongoing debate for a long period of time as to whether the mechanismis one of adsorption or of a liquid-li~uid partitioning type. In the latter case the solvated n-alkyl layer of the surface of a silanized silica is considered as a liquid layer. Evidence for partitioning hasbeen given by the agreement between theretention coefficientof the solutes in RPCLC and partition coefficients derived from n-octanol-water systems. thorough treatment of this subject is published by Tijssen [52]. Commercial silicas usually exhibit a ligand density of about 3 pnol/m2. The a maximul~would be about 4 pmol/m2. This impliesthat when one starts with fully hydroxylated silica with an hydroxyl group concentration of about 8 to 9 pmol/rn2 less than half of the original hydroxyl groups have reacted. This means that a notable part of the original silica surface is still present. Better shielding of the silica surface can be accomplished in several ways:

l. Endcapping via an additional reaction with a reactive short-chain silane, for example, hexamethyldisilazane the reactionof a silane with bulky protectiveside chains with thesilica atom the silane, e.g., ~-octyldipropylsilane binding a protective hydrophiliclayer close to the original silica surface and n anchorillg the organosilylgroup by a carbamate -CO-NHlinkage. The named procedures led to substantially inlproved silanized silica adsorbents for chromatographic separation. During the course of these attempts both the reproducibility and the chemicalstability of the silanized silica was drastically improved, Solvents used are typically binary mixtures of water-methanol or water-acetonitrile. Acetonitrile is preferred to methanol owing to itslower viscosity and better suitability for UV detection. The water content can be increased to 95/5 vjv. small amount of organic solvent in the solvent mixture is required to wet the surface. To be able to work in a reproducible manner a buffer should be employed, instead of just water. Buffers are essentially needed for the separatio acidic solutes. The pH of the buffer should be one which is about 2

than the pk, of the acidic solute and about 2 pH units higher than the pk, value of the basicsolute. Studies carried out onthe pH stability of silanized silicas at pH 7 indicate that phosphate buffers are inferior in terms of deterioration than other buffers (tris, borate, etc.). For the acidic range, trifluoroacetic acid, phosphoric acid, and acetic acid were employed. When working with organic/buffered mixtures the miscibility limits should be considered. Mixtures with a buffer concentrationof up to 50 mM can be employed without precipitation of salts. It should be noted that the dissolution is a function of temperature. Mixtures of water and higher alcohols exhibit a relatively high viscosity and are less suited as solvents. 2-Propanol is often employed as a solvent mediatorbetween organic solvents and water in column cleaning and washing procedures. The concept of solvent strength is also applicable to RPCLC. In the case of silanized silicas water the weakestsolvent. It is used in organic-water mixtures to control the retention coefficient. There is nearly a linear dependency between the logarithm of the retention coefficient and the water content of the eluent. Thus, retention of asamplemixture in RPCLC is adjusted to 1 k 10by controlling the water-acetollitrile (methanol) composition. With respect to solvent strength, acetonitrile has a somewhat higher elution strength than methanol. Thus, the solvent strength increases as the solvent polarity decreases. Cl8-silanized silicas are the adsorbents of choice in RPCLC. These adsorbents not only offer the highest selectivity for hydrophobic solutes but also exhibit the highest chemical stability towards hydrolytic cleavage. C8-silanized silicas are superior to C18 when a selectivity for more hydrophilic solutes is considered. Solvent selectivity is controlled by the choice between basic, acidic, and dipolar solvents. Tetrahydrofuran, methanol, andacetonitrile fall into this scheme. Often a change in the composition of these mixtures can leadto retention changes. The p of the eluent has the most pronounced effect on the resolution of acidic and basic compounds. A major problem arises with the resolution of basic organic solutes. called silanophilic interactions excessive retention and peak tailing occ case special base-deactivated silanized silicas should be employed. Another solution is to add a base, for example triethylamine, at low concentration to the solvent mixture. The amine blocks the acidic centers and basic analytes are eluted with more symmetrical peaks. Another alternative is to add an ion-pair reagent in 5 to 10 mM concentration to the eluent, for example n-heptane sulfonic acid. Ion-pair formation with the positively charged amine improves the peak profile. RPCLC is apowerfulseparationtechnique to resolve nearly all ki rganic solutes. Resolution is best for polar and phobic compounds. ion of solutes must be avoided by adjustingthethesolventmixture.

Native silica and its chemically bonded products have acquired a dominant position in separation processes in the liquid phase. These applications range from analytical ch~ol~atographic separations to preparative isolation and purification

ofcomplexmixtures inthepharmaceutical and chemical industry. Chemically tailored silicas arealso employed in process chromatography wherehigh-value products, for example racemates, are resolved into enantiomers in large quantities. Advanced sol-gel chemistry, the availability of refined polymers as constituents in composite structures, and the powerful tools of high-resolution analytical techniques in materials science provide a basis for novel innovative products in thefield of silica adsorbents. One promising area is the synthesis of defined nanoparticles with desired chemical functionalities. Another field is the control of pore structure and surface properties of micron and submicron silica spheres. Both these materials are candidates in the engineer in^ of miniaturized devices applied as chemical reactors, optical devices, analytical systems, etc. The control of material morphology is not only restricted to spheres of various sizes from nanometer to micrometer dimensions. Through advanced technologies silica fibers, foams, and defined agglomerates are synthesized. Putting these facts into perspective, silica will be an extremely powerful hightech adsorbent for future applications.

1. H. K. Ferch, in The Colloid Chemistry Silica (H. Bergua, ed.), Advances in Chemistry Series 234, American Chemical Society, Washington D.C., 1994, p. 481. 2. K. K. Unger, ed., Pac~ingsand Stationary Phases in Chro~atogruphic ~ec~niques, Marcel Dekker, New York, 1990. 3. F.Vansant, P. van der Voort, and K. C. Vrancken, eds., Characterization and C ~ e ~ i c a l ~ o d ~ cofa tthe i o Silica n ~ ~ r f a cin e ,Stud. Surf. Sci. Catal. (1995). 4. M. T. W.Hearn, ed., HPLC of Proteins, Peptides and Polynucleotides, VCH Publishers, Cambridge, 1991. 5. M. Hanson, K. K. Unger, C. Mant, and R. S. Hodges, Trends Anal. Chem. 25:102 (1996). 6. Ehwald, PhD. Thesis, Johannes Gutenberg Universitat, Mainz, 1999. 7. K.K. Ungerand Weber, eds, A Guide to Practical GitVerlag, armstadt, in press. 8. G. Guiochon, S. G. Shirazi, and A. M. Katti, eds., ~ ~ n d u ~ e n t a lPreparat~ve s and Nonlinear Chro~atography,Academic Press, New York, 1994. 9. G. Subramanian, ed., Preparative and Process Scale Liq~id Chro~atography, Ellis Horwood, London, 199 1. 10. S. Lurie, S. Couver, and V. L. Ford. Anal. Chem. 70:4563 (1998). 11. L. Karger, L. R. Snyder,andC. Horvath, An ~ntr~ductionto ~eparation cience, John Wiley and Sons, New York, 1973. 12. R.Snyder, Principles ~ d ~ ~ o r p tC~z~omatograp~zy, ion Marcel Dekker, New York, 1968. 13. L. Karger, L. R. Snyder, and C. Horvath, I n t ~ o ~ ~ c t to i o Sepuration n Science, Wiley Interscience, New York, 1973, p. 11. 14. Oscik, A ~ ~ o r ~ tEllis ~ o nHorwood, , Chichester, 1982.

15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31.

G. Schay, L. Nagy, and T. Szekremesy. Polytechnica 4:95 (1960). F. Huber and R. G. Gerritse. J. Chromatogr. 58:137 (1976). H, L. Wang, J. L. Duda, and C.J. Radke. J. Colloid Interface Sci. 66:1277 (1978). F. Riedo and E. Kovats. J. Chromatogr. 239:1 (1982). N. Le Ha, J. Ungvaral, and E. Kovats. Anal. Chem. 54:2410 (1982). F. Koster and G. H. Findenegg. Cllromatographia 15:743 (1982). W. Markowski, K. Czapinska, and H. Poppe. Chromatographia 17221 (1983). A. Seidel-Morgenstein, in ~ a t l z e ~ a t i s c h e~odellierung der p r a ~ a r ~ t i ~ e n Fl~~sigchrovnatographie, Deutscher Universitatsverlag, Wiesbaden, 1995, p. 67. M. Huber, Diploma Thesis, Johannes Gutenberg Universitat, Mainz, 1998. J. C. Giddings, Dynamics Chrovn~~tography, Part Marcel Dekker, New 1965. B. Bidlingmaier, K. K.Unger, and N. von Doehren. Chrornatogr. 832:11 (1999). K. Ungerand E. Weber,eds., ~ a n d b u c hder €€PLC, Teil l , GIT Verlag, Darmstadt (1989). K. Nakanishi. Materials 467 (1997). M. M.Dittmann, K. Wienaud, I. Bele, and G.P. Rozing, LCGC International 13:800 (1995). S. Ludtke, Th. Adam, K. K. Unger. J. Chromatogr. 786229 (1997). J. Rouquerol, Avnir, C. Fairbridge, D. H. Everett, M. Haynes, N. Pernicone, J. D. F. Ranisay, K. S. W. Sing, and K. K. Unger. Pure Appl. Chern. 1739 (1994). S. Deming, W. S. Derning, and E. Teller, J. Amer. Chem. Soc.

32. J. H. de Boer, in The Structure 33.

Properties Porous ater rials H. Everett ths, London, 1958, p. 68. d ~Porous r ' ~ J. Rouquerol, ~dsorptionby ~ o ~ ~and

34. W. T. Thornson (Lord Kelvin). Phil. of Size clus us ion C~romatography,Chromatographic Science Series 69, 35. Marcel Dekker, New York, 1990. 36. K. K. Unger, Porous Silica, Elsevier, Amsterdam, 1979. Nawrocki. Chromatograp~ia31:177, 193 (1991). 37. 38. 39. 40. 41. 42. 43. 44. 45. 46.

sitat, Mainz, 1998. to Zeolite Science and Practice (H. van Bekkum, E. M. Flanigan and J. C. Jansen, eds.), Amsterdam, Elsevier, 1991, pp. 268-278. P. G. Dietrich, K. H. Lerche, J. Reusch, andR. Nitzsche. C~ronlatograpllia 44363 (1997). L. H. Allen and E. Matijevic. J. Colloid Interface Sci. 33:420 (1970). T. Adam, P h D . Thesis, Gutenberg Universitat, Mainz, 1998. R. IC. Iler, The C h e ~ i ~ t r ySilica, Wiley Interscience, 1977. K. Unger, K. D. Lork, B. Pfleiderer, K. Albert, E. Bayer. J. Chromatogr. 556703 199 1.

47.B.Buszewski. Chromatographia 28:574(1989). 48. K. Szabo, N. Le Ha, P. Schneider, P. Aeltner, and E. Kovats. Helv. Chim. Acta 67:2128 (1984). 49. K. D. Lork, K. K. Unger, and J. N. Kinkel. Chromatogr. 352:199(1986). 50. H. Engelhardt, M. Arangio, and T. Lobert. LC CC Int. 10:803 (1997). i o ~ in Phase Liquid C ~ r o ~ f f t o g ~ aChrornatogr. ~~y, 51. The ~ e t e ~ t Process 656 2) (1993). 52. R.Tijssen, P.J. Schoenmakers, M. Bohmer, L. K. Koopal, and H. A. H. Billiet. J. Chromatogr. 656: 135 (1993).

Laboratoire de Physique des Fluides Organisks, Collhge de France, Paris, France Rhodia-Silicone, Rhodia, Saint Fons, France

I. Introduction

597

11. Molecular Parameters Which Fix the Surface Excess at Saturation A. Materials and technique used to form the layers B. Dry thickness of the layers C. Molecular parameters controlling the surfaceexcess

598 598 600 600

Internal Structure of the Layers Role of the OH extremities in the nonreversibility of the chain anchoring B. Concentration profile and swollen thickness of the layers in good solvent

604

AdsorptionKinetics

610

Role of the Chain Extremities on the Adsorption Kinetics

61 1

Exchange Between the Adsorbed Layer and a Bulk Polymer

612

Modification of the Silica Surface in Order to Form End-Grafted PDMS Layers

616

Conclusions

617

References

619

604 606

Silica surfaces are naturally highly active withrespect to polydimethylsiloxane (PDMS) polymer chains, which spontaneously form hydrogen bonds between the silanol sites of the surface and the oxygen atoms of the backbone the polymer

chains. number of practical applications of silicon-based polymers rely on such interactions. This is thecasefor siliconeseals routinely used to preventwater flowing around kitchen sinks, forexample, in which small silica beadsact as cross-linkers between the chains. This also the case for more sophisticated and technical applicationssuch as thebondingwith silicone adhesives of the glass windows in the“pyramid du L o u w ~ , ~bonding ’ which had to be invisible, but nevertheless strong. The characterization of the interactions between PDMS chains and silica, and the understanding of how polymer chains interacting with a solid surface organize themselves are two important steps towards the control and the optimization of practical systems based on silica-anchored BDMS layers. We shall report here series of experiments on PDMS adsorption andend grafting on plane silica surfaces. Even though adsorptionis more difficult to characterize plane surfwes than on poroussilica or on colloidal silica beads, because of the small specific area, plane surfaces are nevertheless interesting: they allow several complementary methods to be used to analyze in detail the organiz~tionof the polymerchains inside theadsorbed layers (for example,ellipsometry,x-ray or neutron reflectivity, etc.), as well as the design of well-defined adhesion or friction tests. full characterization of surface-anchored polymerlayers and of their role in adhesion and friction can thus be conducted. The aim of the present chapter is to report on such an investigation for the system BDMS-silica. The chapter is organized as follows: in Section the molecular parameters which impose the surface excess, i.e., the saturation of the adsorption or grafting process will be identified. Section will be devoted to theanalysis of the organization of the polymer chains inside the surface layer, with a special attention paid to the question of the reversibility of the adsorption, in order to delineate the relevance of the theoretical description of such layers adsorbedfromconcentratedsolutionsormeltsin terms of “pseudobrushes.” Then, in Section IV, the characterization of adsorption kinetics will be presented, while the effect of the natureof the chains extremities on the adsorption processwill be discussed in Section To answer more precisely the question of the possible reversibility of the adsorption, experiments c~aracterizing the degree of exchange between chains pertaining to the layer and chains from the bulk polymer melt will be described in Section VI. Finally, the possibility of modifying the silica surface in order to prevent PDMS from adsorbing at the silica surface while keeping the possibility of end grafting the PDMS chains to form grafted polymer brushesat thesilica surface will be described in Section before drawing some conclusions.

The polymer samples used were narrow fractions of OH-termillated PDMS, obtained from com~erciallyavailable products ( ~ h ~ n e - P o u l e nby c ) fractionated precipitation techniques. few samples withtrimethyl- or vinyl-terminated PDMS

chains have also been prepared in order to test the role of chain extremities in the adsorption process. The molecularweights of all those polymers arein the range 5500 kg/mol, withtypical polydispersity indices between 1.05 and 1.15. The choiceof polymer chains with OH extremities was dictated by the idea that such extremities could form double hydrogen bonds or, athigh enough temperature, covalent bonds with the silanol sites of the surface, and thus prevent chain desorption when the adsorbed layer was put into contact with bulk chains. One can thus imagine the adsorbed layer as being formed with chains which are attached to the surface by many monomers, forming loops and tails, and possibly also by the chain extremities, with an interaction between thesurface andtheextremitiesstronger than between the surface and any monomer in the chain. For large enough molecular weights, however, one can expect that this stronger interaction of the extremities does not affect the structure of the adsorbed layer. In order to test these ideas, layers formed at different adsorption temperature have been systel~aticallycompared, and their behavior compared to that of chains having extremities leading to weaker interactions with the surface than OH groups. The detailed procedure used to form the adsorbed layers is described elsewhere [1,2], and we only recall here the essential steps, necessary to understand how the results presented below are obtained. Different solutions of these polymers were prepared in a good solvent, heptane or octane, with polymer volume fractions in the range 5 to 1. small volume of the solution (typically 0.2cn13)was put ozone into contact with the silica surface of a silicon wafer freshly cleaned by techniques 131 and enclosed in a specially designed stainless steel reactor, minimizing dead volumes,and equipped with O-ringseals in order to prevent from solvent evaporation during the incubation time. This reactor is schematically presented in Fig. 1. The reactor, with the sample, was placed in a thermally controlled ovenat a temperature during a time l,,,,. After the silicon wafer was extracted from

Schematic representation of the reactor used to form the adsorbed layers from semidilute solutions in hexane, octane, or toluene, and allowing one to perform the adsorption at high temperature while keeping the concentration in the incubation solution constant.

the reactor, rinsed under flux a of heptane andtoluene, left in a toluenebath during the time trinsing,and finally vacuum dried at room temperature. All the silica surfaces used in the present investigation are thus the surfaces of the native oxide of silicon wafers. We found that such surfaceswere highly reproducible both from the point of view of the surface chemistry and of their geometrical characteristics: for example, silicon wafers cut along the [l i l] plane of the silicon monocrystal have, after a UV-ozone cleaning procedure (this cleaning treatment was not observed to affect the surface roughness, contrary to any other chemical cleaning process), a roughness with an RMS amplitude of 0.45 mm as characterized by x-ray reflectivity techniques [4]. All the work presented here has thus been conducted on the particular amorphous silica resulting from the oxidation of silicon monocrystals. The typical characteristics (number and nature of the silanol sites and of the silanic bonds) of such a silica surface can be inferred from the information available for plane silica layers deposited on silicon wafers [5-71.

e After rinsing and drying, the thickness, h, of the polymer film remaining on the surface of the silicon wafer was measured by ellipsometry or x-ray reflectivity. We have observed that:

1. W ~ ~ t e vthe e r concentrations in the reaction bath and the molecular weights in the investigated range, h became independent of t,, for t,,, 20 h. 2. When large molecular weightswere used, h decreased with increasing trinsing, and became stationaryfor t,insing 24 h. For smallmolecularweights lo5), flushing the surface with toluene was enough to provide stationary h. This dry thickness of the layer at saturation of the adsorption, h,, has been observed to be independent of the temperature at which the adsorption was conducted. It gives an estimation of the number of monomers pertaining to chains attached to the surface, as the monomers are densely packed in the dry layer, and thus provides a direct measurement of the surface excess. It appears tobe a simple characteristic of the surface-anchored layer.

The evolution of the final thickness of the dry layers, with the polymer molecular weights M,, and the volume fraction in the reaction bath, Q?,is reported in Figs. 2 and 3 respectively. All the results for high enough Q? can be gathered on unique master curve as shown in Fig. 4, demonstrating that obeys the scaling law h, with N the polymerization index of the polymer, and a length comparable to the sizeof a monomer. Such a scaling law can beeasily interpreted, as first noticed by P. Auroy et al. by assuming that during the reaction time all the chains located in the slab with a thickness R(@) from the surface [R(@)is the radius of the chains in the reaction bath at the volume fraction Q?]are able to rapidly attach to the surface by either adsorption or end grafting (for trimethyl-terminated PDMS the only possibility is adsorption of the

h ,

1

(kg/mole)

Evolution of the dry thickness of the adsorbed layer obtained after a long incubation time, i.e., when theadsorption process has saturated, h,, as a function of the molecular weight of the polymer, for incubations from a melt. All data correspond to OH- terminated chains.Filledcircles, 1;opencircles 0.1.

Evolution of h, as a function of the polymer volume fraction in the incubation bath CD, for twodifferentmolecularweights,respectively 410 kg/mol(filled circles) and 140 kg/mol (open circles), both OH terminated.

40

35

15 10 5

0 0

Evolution of the final thickness of the dry layers h,, as a function of the scaling variable with N the polymerization index of the chains and the polymer volume fraction in the incubationbath. Filled symbols =-l; open symbols d, 1. Each symbol corresponds to a given molecular weight in the range 29.6 kg/ mol to 740 kg/mol. monomers in the chain,while both adsorption and end attachment through double hydrogen bonds or covalent bonds are possible for OH-terminated chains). Since in a good solvent semidilute solution R(@) [9], this leads to a surface excess in the film: r and thickness of the dry film in which all the monomers are close packed:

h The line in Fig. 4 is the best fit of the data to Eq. and yields 0.55 nm, a reasonablevalueforthemonomer size. Indeed,neutron-scatteringexperiments have shown that the average end-to-end vector of PDMS chains in the melt was given a function of the index of polyl~erization by with 0.56 nm [lo]. The scaling law of Eq. (1) has been observed for adsorption of PDMS on porous silica beads [S], with a valueof smaller than in the present experimentsby a factor of 2. This may indicate that the density of adsorption sites on the pretreated silica beads used in Ref. 8 was significantly smaller than for the freshly cleaned silicon wafer surface used here, and that this reduced number of active surface sites was limiting the number of chains able to attach to the surface. For small a departure fromEq. (1) is visible in Figs 2 and 3 and the measured h, values are larger than expected by the simple arguments leading to Eq. This can be accountedfor by noticing that as the silica surface attracts the PDMS monomers, the argument used to produce Eq. is oversimplified: the attractive interaction between the surface and the monomers tends to build a polymer concentration larger than close to the solid surface, over a distance from the surface

comparable to the average distance, 6,between entanglements in the solution (6 is the correlation length in the bulk solution [l l]). One can reevaluate h, with arguments similar to those used to get Eq. i.e., countingthetotalnumber of monomers in a slab of thickness R(@) from the surface, and taking into account this concentration profile close to the surface [12]. One obtains

Equation (2) obviously matches Eq. (1) for going to and also matches the expected thickness of a layer adsorbed from a dilute solution when Equation (2), with 0.55 nm accoants quitewell for the experimental data [1,2]. One can noticethat for high the thickness of the dry layer can be expected to follow exactly the same scaling law whatever the nature of the anchoring between the chains and the surface: the limitation to the attachment of chains to the surface comes from the accessibility of the surface binding sites. With the arguments used to produce Eq. (l), this limitation is assumed to be only due to thefact that as soon as the chains located in the first layer with a thickness comparable to the chain radius in the reaction bath have attached to the surface, they fill the space close to the surface with a concentration equal to that of the bulk solution. Any attachment of an additional chain is then impeded, because it would imply an overconcentration of polymer close to thesurface, i.e., an osmotic barrier which repels the chains. We shall see in Section VI1 that for an end-grafted brush (no adsorption, only end grafting), provided the interactions between the surface and the chain extremities are not large enough to enrich the slab of thickness R(@) close to the surfxe in chains extremities during the surfaceexcess obeys exactly the same scaling law, with the same prefactor as for the adsorbed layers discussed above. The thickness of the dry layers, h,, is then an easy measurement of the surface excess, but not a parameter sensitive to the internal structure of the layer. It is also worthwhile noting that the surface layers presented above are quite different from what is obtained in the case of reversible ~onomer-surfaceinteractions. For reversible interactions, the equilibrium concentrated profile within the adsorbed layers has been worked out by de Gennes in terms of self-similar profile 121. The overall surface excess in reversibly adsorbed layers formed from dilute solutions remainssmall, corresponding to a dry thickness of the orderof a monomer size, i.e., much smaller than in the layers formed from concentrated solutions as discussed above. The second important difference between a reversibly adsorbed layer and those discussed in the present paper is that the layers resulting from the interaction between a semidilute PDMS solution and a cleansilica surface are able to resist a changeof their environment without any appreciable change in the surface excess: the layers can be dried, swollen with a pure solvent, dried again, put into contact with another PDMS solution, and dried again, yet h, remains essentially constant. In fact, the chains which are ableto reach the surface during the adsorption or reaction step are trapped by the surface and stay there. The interaction between. the silica and the P D M S chains is strong enough so that the surfaceexcess is essentially fixed bythe conditions under which the adsorption has been conducted. The layers are able to support achange of their environment, adapting the conformation of the chains to this change, but keeping the interaction pointsbetween the

chains and the surface, at least on the average. These layers can be imagined as presented schemati~~lly in Fig. 5: the chains which have touched the surface during the adsorption time have a certain number of monomers anchored to the surface, forming loops and tails, with a large distribution of loop sizes. This a rather complicated internal structure. Is it possibleto describe it in a systematic way, and to relate the distribution of loops and tails to the statistical properties of the chains in the adsorption bath?The next part will help answer such questions.

In order to describe in detail the molecular organization inside the surface layer, one needs to answer twoimportant questions. First, one hasto characterize to what degree the adsorption. is nonreversible. If such is the case, the layer should retain a memory of the statistics of the chains intheir adsorption bath, a problem which has been solved theoretically by Guiselin The layer should then followdefinite laws for the concentrationprofile when swollen in a goodsolvent, and for the evolution of its swollen thickness with the molecular weightof the chains or the volume fraction in the reaction bath. This can be checked experimentally. Conversely, if the adsorption at least partly reversible, reorganization of themonomers in contact with the surface should proceed during the whole lifeof the layer, and be affected by what is in contact with the layer. Such a reorganization may occur at constant surface excess, but it may also imply an evolution of the surface excess when the environment of the layer is changed. We presentnow several experiments performed with the aim of answering such questions.

The first question we have tried to address is related to the role of the OH extremities of the U OH-terminated chainsin the anchoringprocess: does end graft-

Schematic representation of an irreversiblyadsorbedlayer,with distribution of loops and tails, named “pseudo-brush.”

a large

ing occur simultaneously with adsorption, which would imply that whatever the further changes in the environment of the layers the surface excess remains constant, or not. We have not been able to observe directly, by infrared spectroscopy, evidence for the formation of a covalent bond between the chain extremities and the silanol sites of the surface by OH condensation, for the adsorbed layers formed at high temperature.But several indirect experimentshave been conductedinorder to characterize the role of the OH extremities of the chains, and to test how far one could assume irreversible surface anchoring. First, systematic desorption experiments have been performed. A good solvent such as toluene or octane not able to desorb anoticeable amount of chains of the surface layer formed by theproceduredescribed in Section 11: several days of immersion in pure good solvent only lead to a decrease of the dry thickness of a few angstroms, i.e., at the limit of sensitivity of the thickness measurements (ellipsometry or x-ray reflectivity), whatever the initial thickness of the layer, in the whole range investigated. A similar result was obtained with layers exposed to dilute ammonia or hexamethyldisila~ane solutions. We have observed, however, that trimethylchlorosilane (TMSCII) in dilute solutions was able to almost completely desorb the layers. In Fig. 6, the desorption kinetics by TMSCl are reported for various layers of PDMS with a molecular weight 412 kg/mol in contact with a 0.5% TMSCl solutionin toluene. The layers were formed at 110°C. The dry thick-

Effect on the dry thickness of the layer, h, of an immersion in a 0.5% solution of trimethylchlorosilane (TMSC1) in octane, for different immersion times (5 mn, circles; 20 mn, squares; 80 mn, triangles; 12 h diamonds), for surface layers formed at 110°Cwith chains havinga molecular weight410 kg/mol (various The results are reported as a function of the thickness of the layers before immersion in the TMSCl solution, h,.

ness of the layer after an immersion time t in the TMSCl solution, hd, normalized by the initial thickness ho, is reported as a function of t in Fig. 7 . The thickness of all layers decreases exponentially withtime, and the kinetics is almost insensitive to the temperature at which the layer has been formed. This could be taken as a possible indication that the chain extremities are not chemically bonded to the surface (high temperatures should favor bond formation). It is not clear, however, that the origin of the observed decrease of the layer thickness is only desorption. When TMSCl grafts to the silica surface, HCl is produced which may cut a PDMS molecule yet present at the surface and thus lead to a decrease in the layer thickness. We have checked by gel permeation chromatography (GPC) that the molecular weight distribution of a PDMS sample to which TMSC1-toluene solution was added was not affected even after several days. It may well be, however, that when in contact with the surface, the local concentration of HCl becomes locally larger than on the average in the solution, catalyzing the chain cutting. the available surface of our plane system is quite small, we have not been able to put into evidence the presence of HCl at the surface.

Keeping in mind the fact that the surface excess of the PDMS adsorbed layers remains essentially fixed when the layer is swollen by a good solvent, we analyze

0.8

Evolution of the ratio of the dry thickness of the layer after immersion in 0.5% TMSCI solution in octaneto the initial thickness,h, a function of the time of immersion, t, for layers formed at two different temperatures, respectively 110°C (filled symbols) and 20°C (open symbols), and different ratio r between the initial thickness of the layer, h, and the maximum initial thickness obtainable with that particular molecular weight, h,.

now how the layer reacts when immersed in a good solvent. Two characteristics of the layer will be discussed: the shapeof the concentrationprofile closeto the surface which should be sensitive to the way the polymer molecules are attached to the surface (this concentration profile is expected to be very different for chains which are only attached to the surface by’one extremity or for adsorbed chains [14]) and the swollen thickness of these layers. To measure any of these characteristics, concentration profile or swollen thickness, one needs a technique which gives a contrast between thepolymer and the solvent. Neutron reflectivity can be used,taking advantage of thecontrastforneutrons between hydrogen and deuterium. Neutron reflectivity allowsthewholeconcentration profile inside the swollen layer to be characterized. Such experiments have already been reported 134,151. The reflectivity data were adjusted with the calculated reflectivity curve of a layer made of 10 or 20 thin slabs, each having a constant concentration, the adjustable parameters being the concentrations in the different slabs. The adjustment procedure is efficient to determine the concentration profile, provided one fixes the surface excess, which for our layers is known independently. Such a procedure allows to extract the coilcentration profile from the reflectivity data without assuming a priori the shapeof this concentration profile. typical result is shown in Fig. 8. The full line gives the experimental concentration profile, while the dotted line is the predicted Guiselin’s profile [l31 for a layer having the same surface excess. The

.o 0.8

0.0

0

20

100

Example of a concentration profile of an adsorbed layer swollen in octane, for chains with a molecular weight 92 kg/rnol and a surface density of adsorbed chains X 0.02. The full line is the profile determined by neutron reflectivity. The dashedlinecorresponds to theprofilepredicted by theGuiselin’sdescription of “pseudobruslnes,”imposing to thepseudobrushthesamesurfaceexcessasthe layer. Cuiselin’s description has been correctedto take into account the finite length of the polymer chains 131.

agreement is satisfactory enough to establish that the layers formed by the procedure described in Section are Cuiselin’s “pseudobrushes.” We have also investigated, in a systematic manner, how theswollen thickness of these layers, L, was varying with the molecular weight and the volume fraction in the reaction bath used to form the layers. The teclmique used for these measurements was to compress the layers with a modified atomic force microscopetip, and to take as anestimate of the swollen thickness the distancebetween the tip and the surface of the silicon wafer at which a detectable repulsive force appeared. The details of such experiments are reported elsewhere 161. We only summarize here the main results of these experiments. In Fig. 9, the evolution of the swollen thickness of the layer, with the molecularweight of the chainsis reported for layers formed at 1. A. scaling law is observed over almost two decadesin molecular weights, with an exponent 0.85 rfr: 0.04. This exponent is in good agreement with what has been predicted for either end-grafted brushes or Cuiselin’s pseudobrushes [18], L FY if one introduces the fact that the surface density of chains in these layers 2; varies withthepolymerizationindexas C leading to L n ~ ’ / ~ It is interesting to notice that, as the surface excess, the swollen thickness of these layers, which is an averaged characteristic, is expected to follow exactly the samescaling law for either end-grafted or adsorbed layers. This is not a quantity sensitive to the internal organization of thechains inside the surfaceanchored layer. Whenthe layers areobtainedfrom solutions, withvolume fractions in the reaction bath smaller than one,the swollen thickness of the layers no longer follows the expected scaling laws with the surface densityin chains attached to thesurface:

l00

M

Swollen thickness of layers obtained from melt and swollen in octane, L, as a function of the molecular weight. The observed power law an hasexponent 0.85, close to the predicted value for Guiselin’s pseudobrushes,0.83.

the volume fraction in the reaction bath fixes the surface density, C [see Eq. and notice that the surface density is related to the dry thickness by (h,/a)/N)]. Scaling laws are observed, but with an exponent close to 0.6, larger than the expected value 0.33 [16]. This is quite surprising: Why do the layers formedfrom meltfollowthe behavior of “pseudobrushes,” while the layers formed in almost the same conditions, except for the concentration in the incubation bath, do not? We believe that this fact could be taken as an indication that during the timelayer the was kept in the dry state, some reorganizationof the attachments pointsbetween the chains in the layer and the surface could have takenplace: layer formed with volume fraction in the reaction bath smaller than one has dry thickness smaller than the radius of the chains. When such layer is in the dry state, its chains have to flatten down on the surface. This may favor the reorganization of the loops and tails, changing the loop size distribution. This is not the case for a layer formed from melt where the chains keep their Gaussian conformation in the dry layer. It is, however, quite difficult to test in more detail this possible reorganization of the interaction pointsbetweenthechains andthe surface. If somereorganization occurs, it should take time, and thus the swelling of the layer should depend on how long the layer has been maintained in the dry state, fact that we have not investigated in systematic way. The picture of the adsorbed layer which can be built from all the above results is the following. When the surface is put into contact with the solution to form the adsorbed layer, the chains located close to the surface, inside layer with thickness comparable to the radius of the chains inside the solution attach to the surface rapidly, and then prevent any further enrichment of the adsorbed layer. One chain has many attachment points to the surface, with possibly stronger interaction between the surface and the chain extremity than between the surface and any monomer in the chain, for the a-u O~-terminatedchains. The fact that the concentration profile inside the layer when swollen in good solvent decreases rapidly with the distance from the surface is sign of the fact that the chains in the layer haveadopted conformationwithmanyloopsand few tails. The concentration profilereflects thedistribution of loop sizes; monodisperse layer of loopsshould give profile close to the parabolic profile expectedfor end-graftedmonodisperse chains. Theconcentration profileof ouradsorbed layers can be described by Guiselin’s approach, demonstrating that the distribution of loop sizes is governed by the statistics of the chains in the incubation bath. The layer can be visualized formed by dense layer of small loops and few large loops and tails which are the only parts of the chains able to extend far from the surface. It seems that the swelling forces exerted on the loops when the layer is in contactwith goodsolvent arenotabletodetachthemonomers anchored to the silica surface, otherwise the concentration profile would become close to the parabolic profile observed for end-grafted chains. It is not obvious, however, that theattachmentpoints between thechains and thesurface are totally fixed, and that the elementary adsorption process totally nonreversible; the fact that the swollen thickness of the layers formed in solutions versus the surface density of chains does not follow the expected law may be taken an

indication that some degree of reorganization of the attachment points may take place while the layer is in dry state.

In Fig. 10, the evolution of the dry thickness of the layers, h, as a function of the incubation time is reported for two different incubation temperatures, respectively 110°C and 20"C, for polymer samples with threedifferent molecular weights, terminated, and for a volume fractionin the incubation bath 1. In all cases, the dry thicknessof the layers first increases with the incubationtime, and then saturates. The time necessary to reach the saturation appears to only weakly depend on the polymer molecular weight. It strongly dependson the temperature(it more than ten timeslarger at room temperature than at 1 10"C>, whilethe thickness at saturation, h,, appears to be independent of this temperature, as it canbe inferred from the Fig. 11 where the data of Fig. 10 are gathered in term of the ratio ~ / h The ~ .fact that the adsorption kinetics appears independent of the molecular weight is coherent with the picture we have built of the adsorption process: the polymer chains which adsorb do not have to diffuse towards the surface; they are brought into contact with the surface when the sample cell presented in Fig. 1 is filled; the adsorption only requires local motions of chain segments, in order to allow the monomers to form the hydrogen bonds with the silanol sites of the surface. The temperature dependence of the kinetics appears as an important way to better characterize how the chains interact with the surface, but it has not been investigated in a systematic manner up to now.

A

A A A

Adsorption kinetics, as seen through the evolution of the dry thickness of the layers a function of the incubation time treat, for surfaces immersed in melts, and two incubation temperatures, 110°C (filled symbols) and 20°C (open symbols). Three different molecular weights of ~H"terminatedchains havebeen investigated: 50 kg/mol (circles); 282 kg/mol (squares); and 410 kglrnol (triangles).

1.o 0.8

0 O

A

A

10

tion, h,,

Ratio of the dry thickness of the layers, h, to the dry thickness at saturaas a function of the incubation time, for the data of Fig. 10.

In order to better understand the role of the chain extremities in the formation of the surfacelayers, we have compared the behaviorof trimethyl-terminated and OH-terminated chains discussed above. vinyl-terminated PDMS to the Typical adsorption kinetics for OH-terminated chains have been reported in Fig. 10. The behavior of the trimethyl-terminated chains is very similar to that of the OHterminated PDMS,as canbe seen in Fig. 12 which displays the evolutionof the dry thicknessof the layers as afunction of the adsorptiontime. The dependence of thethickness at saturation, h, versusthe scaling variablefor layers formed with a volume fractionin the reaction bath dt> l , follows the same scaling as for OH terminated chains and the slopeof the straight line, a 0.47 nrn, is also Comparable. The behavior of the vinyl-terminated chains is, however, quite surprising: these chains hardly adsorb on a clean silica surface, and form layers with thicknesses muchsmaller than those of theequivalentOH-terminated (see Fig. 13). We believe that this is a clear indication that indeed the OH e ities strongly influence the adsorption process. In Fig. 14, we have reported the kinetics of formation of layers from a melt of vinyl-ter~inatedPDMS chains (M, 34 kg/rnol) mixed with traces of OH-terminated (M, 410 kg/mol). c hat ever the i~cubationtemperature, the thicknessesof the layers formed from the mixture are larger than those obtained for the pure vinyl-terminated even though only -terminated PDMS were used. We believe that this demonstrates the efficiency of the OH-terrninated chains to incorporate in the surface

et 25

20

'5 v

10

5

10

Adsorption kinetics, as seen through the dry thickness of the layers, for tri~ethyl-terminatedPDMS chains, adsorbed at 110°C. Data have been taken for four different molecular weights: 32, 57, 101, and 160 kg/mol.

layer. A possible explanation of the fact that the methyl-terminated PDMS gave results almost similar to those obtained with OH-terminated chains is that even though assumed trimethl terminated, the commercial products used contained a noticeable amount of OH-terminated chains. The vinyl-terminatedPDMS samples, anionically polymerized a shorttime before use, were better controlled, leading to a lower adsorbed amount. Additional evidence that the chain extremities are important in the process of the layer formation is given by the simple experiment in which mixtures of two different molecular weights, both OH terminated, with various proportions of the two components areused to form thelayers. The results of such an experiment are presented in Fig, 15, for 282 kglmol. The fact that the 34 kg/mol and thickness doesnot obey a simple linear mixture law and is below the linear mixture curve, is a strong indication that the surface layer is enriched with the shorter chains, i.e., in chains having more OH extremities.

In order to furthertest the way a given layer reacts to a change of its environment, another experiment for which no chain cutting had to be expected (contrary to what has been presented with TMSCl in Section 111) has been performed. Layers formed with one molecular weight, M I at , avolume fractionin the reaction bath were furtherincubatedwithamelt of PDMS withamolecular weight MI; M2 410 kglmol). When was lower than one, the dry thickness of

12 10

m

m

cl

2

m

10

m m

cf

m 2

1000

100

10

10,000

Adsorption kineticsfor vinyl-terminated chains, with different molecular weights: 20, 50, and kg/mol. For comparison, the kinetics for a 50 molecular weight, O~-terminatedsample are also reported. (a) adsorption at 100°C; (b) adsorption at 20°C.

the layers increased by an amount Ah. Ah normalized by the dry thickness of the initial layer is reported in. Fig. 16 for layers formedwith different molecular weights, and different volume fractions in thereactionbath.Thisadditional adsorption has been conducted at two different temperatures, 20°C and 110°C. The initial molecular weight of the chains in the surface layer seems to be unimportant, but two clear tendencies emerge fromFig. 16: Ah is much larger for layers

0

10,000

l

Evolution of the dry thickness with incubation time, for layers formed from a mixture of a-u vinyl-terminated PDMS with a molecular weight 36 kg/mol and traces of a a-u 0H"terminated PDMS of large molecular weight (410 kg/mol) at volume fraction smaller than 1 YO.The thickness is always larger than that of layers formed with pure vinyl-terminated polymer, by more than lo%, whatever the conditions under which the layersare formed, even if the total amount of OH~ter~inated chains is kept low.

35

20 15

10

0 1

Final dry thickness of layers formed from mixtures of two a-u OH-termielts with molecular weightsM Iand ( M l 34 kg/rnol and 282 kg/ mol; incubation temperature 110°C) as a function of the volume fraction of low molecular weight in the mixture. The factthat the dry thickness doesnot obey a simple linear law isan evidence that the OH extremities favor theadsorption of the lighter species.

1

0

0

0.8

l

Evolution of the increment of the dry thicknessof the layers whenput into contact with a melt of high molecular weight, 410 kglmol, Ah, for different layers formed from solutions (volume fraction smaller than one), and different molecular weights. To demonstratethat suradsorption is quite efficient for layers formed with a thickness much smaller than the maximum thickness for that particular molecular weight, Ah normalized by the maximum thicknessof a layer made with pure410 kg/ mol in the melt, Ah/hm,,, (410 K) is reported as a function of the ratio of the initial thickness of each layer divided by the maximum thickness of that particular molecular weight, h ~ ( ~ ) / ~The ~ ~suradsorption ~ ~ ( M ) is.more efficient for layers formed , ( from ~ ) semidilute solutions with smaller further from their h ~ ~ ~(i.e.,

formed at low and the additional adsorption of chains with the molecular weight is favored by an incubation at higher temperature. The fact that the two sets of e~periments, incubation at20°C and at 110"C, both give no increase in the layer thickness when the initial layer is formed with a volume fraction in the reaction bath close to one is in support of a thickness increaseAh due to suradsorption, and not to chain exchange (desorption of the surface chains by the long bulk chains). Then the parameters govering are the accessibility of the silica surface, i.e., the surface densityof the chainsinitially present inside the surfacelayer, and the degree of interpenetration between the surface and bulk polymer chainswhen the layer is put into contact with the melt of large chains. The interpenetration between a surface-anchored polymer layer and melt has been analyzed theoretically by de Gennes [l91 forend-graftedbrushes and by Aubouy et al. [20] for irreversibly adsorbed layers. Thesepredictions andthe existing experimental investigations are discussed in a review article [14]. The general tendency is that dense layers tend to expel the chains from the melt, and this occurs more rapidly as the surface density inside thesurface layer isprogressively increased, whenthemolecular weight in the melt is larger than that in the surface layer. These tendencies are consistent with the fact that no supplementa~adsorption takes place when the

surface layers are formed from a melt, i.e., at their maximum surface density. In order to go further in the characterization of these processes, and independently follow the evolutionof the surface excess of the chains at M Iand one needs to use a system in which there is a contrast between both types of chains. Series of such experiments havebeen conducted using deuteratedand hydrogenated and infrared spectroscopy. These experiments are, however, difficult to interpret, because preferential interactions between the surface and the hydrogenated chains (compared to deuterated) do exist [14,15], which affect the process.

All the surface layers discussed above have to some extent the rather complicated internal structureof “pseudobrushes,” i.e., are madeof loops andtails, with a large distribution of loop sizes. This size distribution makesComplicated the modelingof important properties of these layers, such as their ability to promote adhesion, which is governed by the aptitude of chains from an elastomer put into contact with the layer to interpenetrate into the layer. Attempts have been made toward such modeling,but they remain complicated,and this is an intrinsic consequence of the internal organization of the chains in the surface layers [21,22,14]. It may thus be desirable to build simpler surface layers, in which the size distribution of chain portions attached to the surface is narrower. The easiest thing to imagine, at least for the models,is end-grafted chains, with a narrow molecularweight distribution. The desirable situation is that of chains attached tothe surface onlyby one of their extremities, and with no particular interactionsbetween any monomers in the chain and the surface. form such end-graftedlayers with the PDMS-silica system, one needs to modify the silica surface in order to prevent direct interactions between the monomers and the surface, while keeping the possibility of end grafting. To do so we have developed a surface treatment consistingof grafting on the silica surface a self-assembled layer of a short oligomerof PDMS, containing four monomer units, and terminated by a chlorine at one extremity and a Si-H group at the other extremity. On such self-assernbled monolayers, PDMS hardly adsorbs: under incubation conditions similar to those described in Section for the adsorbed layers, onlyadry layer withathickness of l nmremains fixed on the surface, to be compared to the 20-50 nm fixed in the case of bare silica. On the functionalized extremities of the molecules in the monolayer, monovinyl-terminated PDMS canbe grafted by a hydrosililation reaction in the presenceof platinum as catalyst [23,24]. In Fig. 17, the dry thickness of end-grafted layers thus formed, for a variety of molecular weights and of volume fractions in the reaction bath are reported as a function of the same scaling variable as in Fig. 4, The remarkable feature is that the same scaling law, with the same prefactor, 0.5 mm is observed, while the concentration profile of these layers when swollen in a good solvent has been observed to be parabolic, as predicted [14,25,26]. These two facts together strongly support the argument used to produce the scaling Eq. (l), and to establish the fact that what limits the number of chains anchored to the silica surface in the two situations described inthis chapter is not the number of available interaction sites

35 30

E

5

0 0

30

50

70

Final drythickness of end-graftedlayersas a function of thescaling variable N1/2@7/8. The behavior of the dry thickness of these layers is identical to that of adsorbed layers. The surface excess is governed by the polymer molecular weight and the volume fraction in the reaction bath.

on the silica surfae, but the number of chains which haveaccess to the surface. For thetwo situations, adsorptionandend grafting, thenature of the interaction between the chains and the surface is not a relevant parameter to fix the surface excess, but it imposes the internal organization of the chains inside the surface layer.

We have studiedin a systematicway how adsorbed PDMS layers form on the plane silica surface of silicon wafers. On bare clean silica, we have shown that surface-anchored layers form spontaneously when a polymer solutionor a polymer melt put into contact with the surface, and that the total number of chains (per unit area) able to attach to the surface was fixed by both the molecular weight of the polymer chainsand by the polymer volume fraction in the solutionused to form the sur~~ce-anchored layer. These chains attach rather strongly to the surface, so that the surface excess, which dependson the adsorption conditions, remains essentially fixed when. the layer is taken out of the adsorption bath, dried, or swollen again in a pure solvent or put into contactwith a polymer solution. To that sense, the adsorption can be considered as nonreversible. The silica surface has a large number of silanol sites, which are potential adsorption sites for the PDMS chains, the monomers interact in^ with the surface through the formation of hydrogen bonds between some silanols of the surface and the oxygen of the backbone of the chain. Any monomer in a particular chain is able to attach tothe surface, and

we have shown that the concentrationprofile of the adsorbed layers, when swollen in a good solvent, could be described in terms of the profile predicted by Guiselin for irreversible adsorbed layers. The internal structure of the layer is thus rather complicated: each chain has many contact points with thesurface, and the layer is made of many loops andfew tails, with a large distribution of loop sizes, which are determined by the statistics of the chains in the adsorption bath. The fact that the concentration profile decreases rapidly close to the surface reflects this distribution of loop sizes, and means that the layer is essentially made of many short loops densely packed close to the surface, and of a few large loops and tails, able to extend far from the surface under the effect of excluded volumeinteractions. In the presence of pure good solvent, it seems that the stretching force acting on the loops is not able to detach the monomers interacting with the surface: we have seen no evidence for an evolution of the concentration profile with time after swelling. The adsorption thus looksessentially irreversible. We have, however, gatheredevidence that rearrangement of the attachment points could take place to some extent: when the layers are in the dry state, the chains flatten down on the surfaceif the layer has been formed from a solution, and thus does not haveits maximum thickness. This seems to favor a reorganization of the attachment points; at least this is a plausible explanation of the fact that the swollen thickness of these layers versus the surface density of chains no longer follows the expected scaling law. We have also been able to observe that those incomplete layers could be saturated by a further suradsorption of large-molecular-weight chains. The adsorption of PDMS on silica thus appears to be either irreversible partly reversible, depending on what is put into contact with the layer, and the layers formed at their maximum surface density behave asif the adsorptionwas nonreversible. Layers formedat low surface density are much more sensitive to reorganization and suradsorption than layers formed from a melt. The adsorption process is, however, a complex one, in which the nature of the chain extremities plays a major role. We have investigated in detail how the adsorptionkinetics were influenced by the nature of the chain extremities. It seems that an OH termination strongly favors adsorption, butwe have not been able to establish whether or notthese extremities were able to form acovalent bond by condensation onthe silanols. Wehave also developedsurfacemodification techniques of the silica surfaceinorder toobtainend-grafted PDMS chains, with no adsorption of their monomers on the surface. It is remarkable to notice that with the systemsilica-PDMS, thetwokinds of layers, adsorbed and end grafted can thus easilybe formed with the same surface excess (in both cases, this surface excess can be adjusted by choosing the molecular weight of the chains and the polymer concentration in thereactionbath),withthesamemolecular weight of the surface-anchored chains, but the two categories of layers having a very different internal structure: brushes for the end-grafted chains, and pseudobrushes for the adsorbed layers. A major interest of producing surface-anchoredlayers on plane surfacesis to use these layers as model systems to try to understand the molecular mechanisms of adhesion and friction. The silica-PDMS system discussed here is a perfect tool to do so, because the layers are controlled from the point of view of their internal structure and of the surface density of anchored chains. We have used such layers

to test in a systematic way the recent models of adhesion promotion between surface and an elastomer in terms of connector molecules, and also, through a recently developed technique of near-field laser anemometry, to characterize how surface-anchored polymer chains could be used to adjust friction (see Ref. 13 and references therein). These investigations open the way to thedesign of surfaces with adjustable mechanical properties, and the silica-PDMS system should certainly play an important role in this domain.

1. M. Deruelle, R. Ober, €%.Hervet,and L. LCger, submitted to Macromolecules. 2. M. Deruelle, Thesis University Paris VI, 1995. Clean Surface Technology (K. L. Mittal, ed.),Plenum 3. J. R. Vig,in Treatise Press, New York, 1987, pp. 1-26. 4. J. Daillant, J. J . Benattar, and L. Leger. Phys. Rev. A 41:1963 (1990). 5. D. L. Angst and G. W. Simmons. Langmuir 7:2236 (1991). 6. Y. J. Chabal and S. B. Christman. Phys. Rev. B 29:6974 (1984). 7. P. Gupta, A. C. Dillon, A. S. Bracker, and S. M. George.SurfaceSci.245:360 (1991). 8. P. Auroy, L. Auvray, and L. LCger. Macromolecules 245158 (1991). 9. P. G.de Gennes, in Scaling Concepts in Polymer Physics, Cornell University Press, 1979, 82. 10. L. Fetters, D. J. Lohse, D. Richter, A. Witten, and A. Zirtel. Macromolecules 27:17 (1994). 11. P. G. de Gennes, in Scaling Concepts in Polymer Physics, Cornell University Press, 1979, p. 80. 12. P. G. de Gennes. ~acromolecules141637 (1981). 13. 0. Guiselin. Europhys. Lett, 17:225 (1992). 14. L. LCger, E. Raphael, and H. Hervet. Adv. Polymer Sci. 138:185 (1999). 15. C. Marzolin, PhD Thesis, University Paris V, 1995. 16. M. Deruelle, Ondarphu, and L. Lkger, submitted to Langmuir. 17. S. Alexander. J. Phys. Paris 38:977 (1977). 18. P. G. de Gennes. Macromolecules 13:1069 (1980). 19. M. Aubouy, G. H. Frederickson, P. Pincus,and E. Raphael. Macromolecules 28:2979 (1995). 20. M. Aubouy and E. Raphael. Macromolecules 275182 (1994). 21. F. Brochard-Wyart, P. G. de Gennes, L. LCger, U.Marciano, and E. Raphael. Phys. Chem. 98:9405 (1994). 22. C. Ligoure. Macromolecules 29:5459 (1996). 23. J. P. Folkers,E. Durliat, M. Deruelle, H. Hervet,and L. LCger, submitted to Macromolecules. 24. E. Durliat, H. Hervet, and L. LCger. Europhys. Lett, 38:383 (1997). 25. S. Milner, A. Witten, and M. E. Cates. Macromolecules 21:2610 (1988). 26. E. B.Zhulina, 0. V.Borisov, and V.A.Pryamitsyn. J. ColloidInterface Sei. 137:495 (1990).

This Page Intentionally Left Blank

Department of Physics, University Joseph Fourier, Grenoble, France

I. Introduction

TI, Silica Aggregates Law of Chain Adsorption A.Purepolymer B. Simple modellayer C, Bimodalpolymer

TV, Initial State of Adsorption

622 623 623 624 624 625 626

Formation of Silica-PDMS Networks A. Networkstructure Density of active hydrogen bonds C. Tightly bound water molecules D. Frozen monomeric units

627 627 629 629 629

VI. Kinetics of Chain Adsorption A. Release of water molecules B. Molecular weight dependence C. Effect of silica concentration variations

630 63 63 632

VII. Intermediate States of Gelation: Swelling Effect

634

VIII. Surface Treatment of Silica

63 5

Conclusion

638

Appendix A

640

Appendix B

64 1

List of Symbols

64 1

References

642

The purpose of this contribution is to show that the characterization of the adsorption of polydimethylsiloxane (PDMS) chains onto fumed silica provides an insight into properties of the surfaceof aggregates. Twoof the most experimentally important reinforcing fillers are silica and carbon black. The understanding of the effect of reinforcement of rubbers is partly based on the description of the mechanism of adsorption of macromolecules on the surface of mineral aggregates. The adsorption plays an essential role in the formation of physical systems which are experimentally found toexhibit a gel-like behavior. For suitable concentrations of the mineral filler associated withsuitable chain lengths, loaded polymers canbe swollen, using a good solvent; fL~rthermore,they behave as cross-linked long chains when they are weakly strained. The adsorptionof polymer chains from the melt usuallygives rise to the formation of loops and trains on the surface of the filler; aggregates are then connected to one another by loops. This stresses the need to characterizethe process of adsorption of thepolymer onthe filler.owever, one of themain difficulties encountered in describingphysicalproperties of loadedpolymers is deterlniningthe ele ntarymechanism of interaction of monomeric units with the mineral surface. espite accurate investigations into properties of the surface of carbon black, usi atomic force microscopy and large number of approaches based on chemicalphysics, the characterization of the polymer-surface interaction is still a baffling problem; both the activity and the morphology of the surface are involved in this interaction Furthermore, thereexists a broad distributionof the energy of adsorption of small molecules; considering cyclohexane molecules, for example, the width of the distribution is about 7 1;Jimol and the peak is around 30 kJimol. The silica surface hasbeen widely studied [2-4]. The mechanismof interaction of chains is contrasted with that of most polymers mixed with carbon black; in mechanism of interaction of PDMS is through hydrogen bonding of the polymer to the filler through the silanol groups on the silica surface. The enthalpy of interaction of one mononleric unit with the surface is about l0 kJ/mol and the in~eractionmay be considered as punctual. The well-defined interaction allows a molecular approach to the description of properties of the adsorbed layer; this interaction will be evoked in Section Attention will be focused on properties specific to the full inlmersion of silica particles in molten S. Theformation of polymeric layers on solid or liquid surfaces has received a lot of attention in recent years mostly due to the use of scaling concepts and of a continuous chain description[5-91. Layers are described fromtheprobabilitydistributionfunction of loopsand tails per unit surface. Considering layers which originate from chains in solutions or in a liquid state, a thorough theoretical analysis has been developed in order to relate the distribution of loops to the thickness of the layer and to the concentration of monomeric units near the solid surface 10,1 Sections and IV are devoted to the description of the adsorbed layer. The formation of networks can be described in terms of a bond percolation process that takes place when percolation units (mineral aggregates) can be connected to one anotherby adsorbed polymerchains. The functionality of percolation

units is determined by the mean total numberof chains that canbe adsorbed on one aggregate. The estimate of the functionality requires one to know the specific area of the mineral filler and the law chain adsorption which applies to aggregates fully immersedintothe melt. Thebondpercolationproperty is conveniently detected from the observation of the swelling effect; it allows a reasonable determination of the threshold of percolation. The bond percolation must be distinguishedfromthepercolation effect which results fromthecontact of coated aggregates, as long there is more polymer than available surface. The formation of networks is briefly outlined in Section

Primary spherical particles l4 mn diameter) are fused together to form aggregates. Theseaggregates are supposed to be theprimarystructure of the filler; clusters of aggregates are called agglomerates, The weight fraction of adsorbed water is about 2%. It has been shown that water molecules which are hydrogen-bonded to silica are easily removed from free silanol groups upon heating, whereas water molecules which interact with vicinal silanol groups are tightly bound through two hydrogen bonds [12,13]. They give rise to some kind of bidimensional lattice (Fig. 1). The polymer and filler were mixedin themelt, using either a two-rollmill or roller sigma blades; the appropriate amount ofsilicawas added to the melt to achieve the desired filling level and the mixture was stirred for a p p r o x i ~ ~ t e 2l yh. Bound rubberis usually analyzedby repeated washingof the sample with a good solvent; the sample is then dried to a constant weight and the amount of rubber is measured from a microchemical analysis.

The adsorption of PDMS chains occurs through the formationof hydrogen bonds between silanol groups located on the silica surface and oxygen atoms attached to chain skeletons. The free enthalpy of monomeric desorptionis thus higher than the

4.4A

Schematic representation of free geminal and vicinal silanol groups.

thermal energy at room temperature. Consequently, monomers adsorbirreversibly when aggregates are fully immersed in the melt.

Several mainfeatures are associated with the polymeric adsorption which occurs in PDMS-silica mixtures. The distribution of loops in the layer arises from the distribution of loops in the melt; this distribution is frozen when macromolecules are in contact to thesilica surface. Correspondingly, it has been already shown that the ~DMS adsorption on silica aggregates obeys Gaussian statistics in spite of the geometrical nature of the surface which cannot be considered as an infinite plane 114,151. However, the surface of any aggregate can be pictured as an ensemble of finite planes, the area of which is usually much larger than the square the radius of gyration of chains which are in contact to the surface. The statistical framework of description is Gaussian because polymer chains are in a liquid state prior to the adsorption process. It has been shown both experimentally and theoretically that as a consequence of the Gaussian statistics, the mean number of contact points, of one chain with thesilica surface is proportional to the square root of the number of skeletal bonds,

The resulting amount of PDMS, adsorbed per gramof silica, for aggregates fully coated and well separated froE one another, is expressed as a functionof the number average molecular weight according to the equation

with

in which M m is the molar weight of one monomeric unit and is the Avogadro number; is the mean area associated with one hydrogen bond participating in the adsorption process [14]. For the adsorption of PDMS on fumed silica, the typical order of magnitude pm is5.0 (g/mol)"", with 150 m2/g.The principle of the measurement of the factor is presented in Section It may be worth noting that the numerical value of depends on the nature of the mechanical mixing using roller sigma blades or a two-roll mill.

The usual description of a random flight, considered in the presence of a reflecting wall, provides a simple expression for the distribution of contact points of one Ga~lssian chain,with adsorbed ends

ion

The mean valueof the number of contact points, calculated accordingto Eq. (4) is equal to 1.2 0 ; this is also the mean number of loops formed above the surface of adsorption, in one chain with fixed ends. The contact points are supposed to be randomly distributed on the surface. The maximum number of contact points of one chainis N/2 l. Equation (4) describes the formationof loops; the presenceof two tails is accounted for by dividing one chaininto three parts which consist of n l , and bonds. The number of configurations of one chain with contact points thus written as

the probability that after takingn1 steps the chainwill arrive at any point above the surface without ever having touched or crossed the surface, at any earlier step, is proportionalto l / A . Thenormalization of nl, is considered in Appendix A; it is related to the total numberof active silanol groups. This equation gives a simple picture of any Gaussian chain,in a melt. Monomers adsorb irreversibly when aggregates arefully immersed in the molten polymer and configurations are frozen. More sophisticated approaches have been developed to derive the loop density profile of the polymeric layer, i.e., the monomeric unit density, the extension fo the layer, and the adsorbance [lo, 1 l].

Experimental results concerning mixtures prepared from a bimodal polymer lead to the following interpretationof the adsorption process. It is first considered that the elementary step in the chain adsorption the hydrogen-bonding of one monomeric unit to the silica surface; the probability of adsorption of any monomeric unit is an uniform function. It is now supposed that in the case of a bimodal polymer, this probability is weighted using the molar fraction of each chain species. Then, the total number of available sites of adsorption of long chains is @L cAT/(cefi) per gram of silica; NL is the number of monomeric units per long chains. The total number of available sites of adsorption of short chains is similarly equal to

(1

cMT/(cefls)

(7)

The total amount of polymer, adsorbed per gram of silica, is thus expressed as

in agreement with experimental results. An effective molecular weight is defined according to the equation:

With regard to the adsorption of a bimodal polymer, the silica surface is occupied according to the weight fraction of each chain species. The weighting factors apply also during the process of adsorption.

The characterization of the initial state of PDMS adsorption on fumed silica provides a method for theanalysis of silica-water molecule interactions. In this section, it is shown that tightly bound water molecules are not released from the silica surface during the mechanical mixing. This feature concerns the specific amount (208, measured right after the end of the mechanical mixing and after eliminating free chains; has been found to vary as the square root of the chain molecular weight according to the equation 1) The linear variation of the fraction of bound rubber is depictedin Fig. 2; Qi is the initial amount of polymer per gram of silica in the mixture. Such a result

0

50

200

2-3)

300

Fraction of bound polymer. Pure polymer; specific surface area: 300 m2/g and 50 m2/g (the scale of wasdivided by 3). bimodal polymer (specific area: 50 rn2/g; the scale of was divided by 3). Curve A, fraction measured right after the mechanical mixing of PDMS with silica. Curve B, fraction measured at equilibrium. 50 m2/g

shows that solidly bound chains have a mean number of contact points with the silica surface which obeys the Gaussian law of adsorption represented by Eq. (l). The chain molecular weight was varied over the range 2 lo3 to 3 lo5 g/mol, while the amount of silica was varied over the fraction range2.5 to 40% gig. Usin roller sigma blades, the experimental valuesof are equal to l .4 and 0.48 10(g/mol)"/2 when the specific surface areas of silica are l50 and 50 m2/g, respectively, while the experimental values of are 1.0 and 0.4 (g/mol)"/2, respectively when a two-roll mill is used [16,17]. The ratio of specific surface areas is 3; it must be compared to the ratio of factors. Thisis equal to 2.9 for roller blades and to 2.5 for a two-rollmill. The small difference between these two values showsthat the mixing procedure, i.e., the shear applied to the mixture, may have a slight influence on the adsorption process. It must be worthnotingthatthenumerical fityields pm 5.8 (g/ too when the PDMS adsorption occurs on fumed silica with a specific surfaceareaequal to 300 m2/g, using roller sigma blades. Sucha result shows clearly that macromolecules do not occupy fully the silica surface involved in the determillation of the specific area, probed from the adsorption cf small molecules. Equation (12) applies also to a bimodal polymer, provided M nis replaced with the effective molecular weight M:, defined according to Eq. (l 1). Typical bimodal polymers that have been used consisted of long chains 3 lo5 g/mol) and short ones (43 lo3 gimol); the concentrationsof long chains were 0.5 or 0.75 w/w and the concentrations of silica were or 40 phr It will be shown in the next section that the comparison of the numerical values of with the numerical values of givesaccess to the detection of the two po~ulationsof water molecules adsorbed on fumed silica.

In most loaded systems prepared from a mechanical mixing, the specific a ~ o u nof t adsorbed polymer measured at equilibrium,is smaller than that obeys the following equation:

is the initial specific amount of polymer in the mixture 181. The difference accounts for theeffect of bridging; it is negligible as long as the numerical value of / 3 , ~ / 4 ( 2 ; is smaller than about 0.05. Equation (12) has been derived from thesimple numbering of chains which are adsorbed on one aggregatewhich or bridge two aggregates.

where

According to the simple modelpresented here, any site of chain adsorption is picturedasachemicalfunction; silica aggregatesaremultifunctional units. Furthermore, any polymer chain is pictured as a ,bifunctional unit because it is supposed to bridgetwoaggregates, only. Letdenotethetotalnumber of

en

sites of polymer adsorption,also called silica functions, available to polymer chains on aggregates; any site of adsorption corresponds to the formation of hydrogen bonds between the surface and one chain. is the total number of so-called chain functions that can "react" with silica, i.e., that can be adsorbed on silica; contact points are involved in the definition of one chain function. The number, is twice the total number of chains, U:, per gram of silica in the filled polymer. Let cjjsi denote the fraction of sites of adsorption occupied by polymer chains; is the fraction of chain functions in contact with the silica surface. Then, the numbersof functions, and dt,,obey the equation

or

with (2isi when the silica surface is saturated, i.e., fully occupied by the adsorbed polymer. The fraction, is thus expressed in a simple way e m

Applying now a mean field approach to the description of the PDMS-silica mixture, the fraction is also considered as the probability that onechain function has one contact with one silica function on one aggregate. The probability that two chain functions, located on one chain, have reacted with two aggregatesis Starting from theinitial amount of polymer, Qi, per gramof silica, the amount of polymer, PFolywhich connects two aggregates is thus written

Similarly, the a ~ o u n of t polymer, PLoly, adsorbed without connecting two aggregates is

mount of Adsorbed Polymer Conseq~lentl~, the total amount of adsorbed polymer is expressed as the sum of p;o,y and p;o,y, (18) or

Qi@c(2 Equation (12) is obtained by combining Eqs. (15) and (19).

The fit of Eq. (12) to experimental results provides the numerical value of the factor. The chain molecularweight was varied over the range 2 lo3 to lo4 mol, while the amount of silica was varied over the fraction range to 40% It may be worth emphasizing that this numerical valueslightly depends on the nature of the mechanical 'mixing. For example, pm loR3(g/mol)-'j2, with AT 150 m2/g and 2.6 (g/mol)"/2, with AT 50 m2/g when roller sigmablades are used. Corresponding numerical values of pm are 4.3 and l , l mol)"/2, when a two-roll mill is used. The numerical values of the mean area, are 0.36 nm2 and 0.27 nm2, in the first case and 0.50 nm2 and 0.47 nm2, in the second case. A reasonable estimate of the density of active hydrogen bonds is 2.5 per nm2. It must be noted that the numerical fityields also 5.8 (g/mol)-1/2, when the PDMS adsorption occurs on fumedsilica with a specific area equal to300 m2/g, using roller sigma blades. Such aresult shows clearly that macromolecules do not occupy fully the silica surface involvedin the determinationof the specific area, probed from the adsorption of small molecules.

Numerical values of the ratio are in the range 0.2 to 0.27; the reasonable estimate of this ratio is 0.25. It shows that the surface occupied by strongly linked chains is only a fraction (about 25%)of the surface of silica. This property implies the existence of two types of sites of adsorption of monomeric units on aggregates [12,13]. One type correspondsto free silanol groups, which havebeen shown experimentally to cover about 25% of the silica surface, while the other type corresponds to vicinal silanol groups (Fig. 1). It may be asserted that the PDMS adsorption is closely related to the release of water molecules, which play the role of a poison, hindering the binding of monomeric units.

It may be of interest to observe directly monomeric units whichare adsorbed on the silica surface. The magnetic relaxationof protons attached to thefixed units canbe discriminatedfromtherelaxation of protons attached to polymer loops. More precisely, nuclear magnetic properties of any spin system submitted to a strong magnetic field 1 present an axial symmetry: the relaxation along thedirection of the magnetic field may be very different from the relaxation which occurs along the direction perpendicularto the magneticfield direction. For example, theproton transverse relaxation, observed in solids, usually occurs over the range 0 to 20 while the longitudinal relaxation occurs over the range 0 to 1 It has been established that the transverse relaxation of protons is sensitive to the statistical structures that are formed in polymers observed above the glass transition temperature

(191. Withregard toadsorbedpolymers,thetransverserelaxation of protons, attachedto fixed units, exhibits amagneticrelaxation specific to spinsystems embedded in a solid; the corresponding relaxation time is about 10 The relaxation of protons attached to loops exhibits a behavior specific to polymeric gels, wherein segmentalfluctuations are strongly hinderedby the presence of cross-links; in the caseof adsorption, adsorbed unitsplay the role of cross-links. The relaxation associated with loops occurs over the range 0 to ms; it is well contrasted with the relaxation associated with adsorbed units. Considering a silica filled PDMS,it is possible to determine the fraction of the proton magnetization, which belongs to fixed units; the fraction providesin turn the fraction of units, which participate in the chain adsorption. Such an experimental approach has been applied to the study of loaded PDMS;the fraction, was expressed as

where y, accounts for the presenceof monomeric units which cannot move because they are directly attachedtoadsorbed units. According to Eq. (l), M, is also expressed as

The molecular weight was varied over the range 1.8 lo3 to 257 lo3 g/mol and the fractionof fumed silica 150 m2/g) was 40 phr. NMR measurements were performed after extracting all free chains. The experimental value of y, was 3.8; such a result shows that due tothe adsorption effect, the random motions of about 4monomeric units, on average, are fully hindered [20]. The similar approach, applied to end-hydroxylated chains, lead to 7.4,while 8 (g/ More end~hydroxylatedpolymer is adsorbed on silica and the number of monomeric units fully affected by the adsorptioneffect is about eight. This value is probably related to the density of adsorption higher for end-hydroxylated chains than for end-methylated ones.

~ t t e n t i o nto experimental values of and pm shows that after the mechanical mixing, about 110 m2/g of the silica surface must be covered by the polymer to 150 m2/g. reach the equilibrium state of adsorption when the specific area is The surfaceto be covered by the polymeris about 40 m2/g when the specific surface area is 50 m2/g. One of the main features which characterize the PDMS interaction with silica concerns the time dependence of the progressive adsorption process on the surface offered to the polymer; this process occurs over a time interval about equal to two months at 343 K. It has been established both experimentally and theoretically that the kinetics of PDMS adsorption strongly depends on the chain molecular weight. In contrast with most dynamic p~enomenaobserved in molten polymers, an increasing rate of adsorption is measured when the chain molecular

weightis increased. This property shows clearly that the chain diffusion in the moltenpolymer is nottherelevantprocess which governsthemechanism of PDMS adsorption on silica aggregates.

The adsorptionprocess, observed after the mechanical mixing, has been interpreted by considering that it isclosely related to the slowreleaseof water molecules, tightly bound to vicinal silanol groups. The probability n(t)that one poison molecule occupiesoneadsorption site onthe silica surfacehas been shown to be expressed as

n(t) exp(-jt/z) is

time constant specific to the diffusion of water molecules through PDMS More precisely, the release of water molecules is supposed to occur in two steps. During the first step, hydrogen bonds are broken and very near the surface, released molecules are in equilibrium with still bound ones; the equilibrium constant is called During the second step, water molecules have to diffuse through the polymer although they are not compatible with PDMS; they evaporate when they react the surface of the sample. The diffusion coefficient is called D,:

r=-

L)&’

a is the thickness of the interface formed by adsorbed monomeric units.

The molecularweight dependence of the adsorption processis closely related to the Gaussian statistics of one chain; the number of contact points that a given chain can offer for adsorptionis equal to N is the number of monomeric units in one chain. Then, the probability that one chain cannot bind is expressed as &(t)

(n(t))O

and the fraction of chains that can be bound at a time t is A4Bo~~~d(Nt/z)

This fraction is represented by the experimental value of the ratio

QB(t)is the fraction of polymer adsorbed at time t , after the mechanical mixing. The time constant associated with AgBoUnd(t)is predicted from Eqs.(24) and (25) to depend on N Considering eachsilica-filled polymer, experimentalresults about AqBound(t) are first plotted as a function of the reduced variable Two typical curves of kinetics of adsorption are plottedin Fig. 3; curve was obtained from a mixture prepared from an initial amount of polymer, equal to 5.5 g per

I---"-

1.2

0

o 0

0

s.000

25,000

30000

(g/llkol*

inetics polymer adsorption. Curve specific area 150 m2/g, M n 73,000 g/mol, and 0.18 g/g. (0) Curve B, specific area 50 m2/g, M, 300,000 g/mol and QC1 0.20 g/g.

gram of silica, the molecular weight was 73,000 g/mol. and the specific surfkce area was 150 m2/g. Curve B was obtained from Qi 4.93 g/g, 300,000 g/mol and a specific surface area equalto 50 m2/g. Curve Ais nearly in coincidence with curve B althoughthe specific areashave different values; thereducedvariable is relevant. For this concentration of silica 20 phr), the order of magnitude of the time constant, as derived fromFig. 3, is lo4 [g/(m~l.h)]'/~; is thus equal to 10" The estimateof the diffusion constant of water molecules leaving the surface, L), cm2/s, gives K 10"' forthe fraction of free water molFcules, nearthe surface and in equilibriu~with adsorbed water molecules 3 A).

The rateof adsorption is increased upon additionof silica. For each system, an empirical curve can then be drawn through experimental points to serve as a visual guide, The system prepared from theinitial weight fraction of polymer, 5.5 g/g, is arbitrarily chosen as a reference (specific surface area of silica, 150 m2/g andM n 43,000 g/rnol). All other kinetics curves that are drawn canbe put intocoincidencewiththereferencecurve by multiplyingthe scale of each curve by a suitable factor the result is illustrated in Fig. 4a. The sameanalysis applies to the bimodal polymer, using the effective molecular weight, already defined in Section the result is shown in Fig. 4b.

Adsorption kinetics: superposition propertyof curves. c area 150 m2/g, 73 lo3 g/mol, and g/g) area 150 m2/g: 73 lo3 g/mol, Qi 16.6 gig (0); 43 lo3 g/mol, Qi and 10.0 gig M, 3 lo5 g/mol, Qi 5.0 Specific surface area 50 m2/g: 73 lo3 3 lo5 g/mo1, 3.3 and 2.5g/ cific area 50 m2/g). Pure polymer: and Q i 5.0 gig. (A) Bimodal polymer; weight concentrationsof long chains:c 0.5 gig, Qi 2.6 and 5.0 gig; c 0.75g/g, Qi 2.4 and g/g.

Variations of the factors(Qi)are plotted in Fig. 5; S is proportional to the square root of the initial weight fraction of polymer, 5.5 g/g; the experimental value is 2.3 instead of 2.35. Taking measurement uncertainties into consideration, this result shows that any kinetics curve can be derived reasonably from the reference curve accordingto the simple relationship

with defined from the ratio QF/Qi. The mathematical expression for the AqSound function can be considered invariant. According to Eq. (27), the time constant depends linearly on the initial specific amount of polymer. Considering the definition of givenby (23), it is supposed that the only parameter which may depend on the silica concentration is the diffusion coefficient it is assumed that depends on the free volume of the system, which in turn is closely related to the texture of the mixture. No high pressure treatment was applied to the samples; consequently, the compactness of these systems results only from the mechanical

0.5

0

l

0.2

Scale factor s(Qi) used to put ce curve;specificarea 150 m2/g

0.4

0.5

kineticscurves into coincidencewiththe andspecificsurfacearea 50 m2/g

mixing. The compactness, determined from the spacial arrangement of hard objects (silica aggregates), decreases upon silica addition and L), is supposed to increase.

Different states of gelation of the filled polymers can be prepared by varying the time of observation defined from the end of the mechanical mixing. The states of gelation are determined from different states of adsorption of the polymer. The slow adsorption process occursat 343 K; it can be interrupted and all observations are made at 300 K. At this temperature the adsorption effect is nearly negligible. The purpose of this section is to show that filled polymers behave like covalent polymeric gels when they are swollen by a good solvent. It has been already shown that the state of maximum swelling of filledPDMS can be described within a mean field approximation 181. Such a description starts from the estimate of the mean number of strands which connect two aggregates, at a time t. This estimate is obtained in the following way. The probability to have one chain in contact with one aggregate, &.(t), is expressed as a function of the fraction of bound polymer fr(t) ~ ~ ( ~ ) Equation / ~ i . (19) applies at any time: (28) and

M

)

1

dl =.M)

(29)

The number of chains which connect two aggregates is determined by

Then, in accordancewithEq. thenumber of strandsper unit volume of adsorbed polymer which connect two aggregates is determined by

where

is the density of the pure polymer, or,

and

Assuming that deformations of strands are in affinity with the macroscopic volume of the swollen polymeric part of the mixture, theswelling ratio Q,, is expressed as

(QXA

(34)

UB(~>

The dependence of (Q%: Q,,/2)" function aas of r ( t ) = f r ( t ) (1 is illustrated in Fig. 6, for 43 lo3andfor several silica weight fractions (0.20, 0.15, and 0.1 w/w; the specific surface area is 150 m2/ g). The swelling effect is seen to be independent of the silica concentration. It only depends on the fraction of bound rubber. The dependenceof the swelling effecton the chain molecular weight is illustrated in Fig. 7, whereinthe silica concentrationhas also several values [1'7,23]. Experimentalvalues of the slo esof thethreestraight lines in this figure are found to be roportional to l/&; the swelling quantity is plotted as a function of r(t)/&in Fig. 8. For both figures the molecular weights are 43 lo3, 73 lo3, and 3 lo5 g/ mol, respectively. Finally, considering two specific surface areas, 50 and 150 m2/g, the swelling effect is shown experimentally to be independent of these two values (Fig. 9). All these experimental results show that strongly linked chains form a network structure simply described by assuming that the law of chain adsorption, described by Eq. (l), applies also during the progressive saturation of the silica surface.

K

The fomation of hydrogen bonds between silanol groups located on the silica surface is not questioned nowadays. It has been shown that hydrogen bonds can

0.02

centration:

states of gelation, swelling effect; variations of the silica con150 1/Qi 0.20; l/Qi 0.15;

an 43,000

be broken whensilica-filled is immersedintolueneunder an ammonia atmosphere [24]. The amount of polymer left bound to silica is a strongly decreasing function of immersion time in the presenceof ammonia. Further investigations into properties of the silica surface can be based on the competitive adsorption 15,000 g/mol) and short endobserved between long end-methylated chains hydroxylated chains 350 glmol). The weight ratio of hydroxyl functions is 5 % (w/w); the mass ratioof short chains per gramof silica, @OH, is varied from0 to 0.8 g/g. Bothlongend-methylatedchains andshortend-hydroxylatedones are introduced in the mixer chamber without any premixing because these two polymersareincompatible.Thecompetitiveadsorption is studied right afterthe mechanical mixing, using silica aggregates characterized by a specific area equal to 150 m2/g. The total amountof adsorbed polymer Qi(0) is due to the contribution of long chains, Qi(O), in addition to the contribution of hydroxylated chains, (0). Theamount Qb(0) is derivedfrom by consideringthetwofollowing assumptions: the silica surface is uniformly occupied by short chains at any concentration; (2) short chains present in the mixture are fully adsorbed on silica aggregates. Therefore, Qf;(O) obeys the relationship

(K

@OH

(K

Intermediate states of gelation, swelling effect; variations of the molecular weight (AT 150 m2/g): l@, 43,000 g 1/Qi 0.10; Gn ’73,000 l/Qi 0.10; Gn 300,000 g/mol, ( X ) l/Qi 0.19; 1/Qi 0.10.

Variations of thenormalizedratio Qk(O)/Q: are illustrated in Fig. 10 as a function of the weight ratio A straight line can be drawn through experimental points with aslope equal to -4.7. This result is analyzed in the followingway. As a consequence of the uniform occupancy of the silica surface, the number of free silanol groups replaced with hydroxylated ends of short chains

where is the molar weight of an hydroxyl group; the numberof chains which cannot be adsorbed is thus equal to n s i o H / n on average and the normalized amount of adsorbed long chains is expressed as

K

The numerical value of the slope is -4.9 for 115,000 g/mol, in reasonable agreement with the experimental value. It is worth emphasizing that starting from the estimate of the total number of silanol groups which participate in the chain adsorption process, 4 lo2’,

Intermediate states of gelation, swelling effect; variations of the molecular weight and the silicaconcentration. The reduced variable, l " ( t ) / a ,is used; key to points as in Fig. 7.

per gram of silica; the weight ratio of end-hydroxylated chains required to occupy all sites can be derived from the relation

Thenumericalvalue ofis0.23 g/g; it is in agreementwiththevalue of the intersection of thestraight line withthe axis (Fig. 10). No long-chainadsorption can occur when the specific amount of end-hydroxylated chains per gram of silica higher than or equal to 0.21 g/g. This result shows clearly that the occupancy of the silica surface by short chains is uniform whatever the concentrationof this hydroxylatedpolymer. Inotherwords, it cannot be considered thatshort chains first adsorb on free silanol groups.

In this chapter, it is shown that the polydimethylsiloxane adsorption on fumed silica provides an insight into properties of the surface of aggregates. The most striking feature concerns the quantitative aspectof the descriptionof the bindingof

0.0002

Intermediate states of gelation, sidered. 50 m2/g, lo5 g/mol: l/Qi 0.30. AT mol, l/@ 0.18; /Qi ( X ) 1/Qi 0.19; 1/Qi 0.10.

0.00w

ecific areas are con0.18. AT

150 m’/g,

lo5 g/mol:

the polymer to the surface; the interpretation relies on the well-defined elementary interaction of PDMS monomeric units with silanol groupsthroughhydrogen bonds. The PDMS adsorption provides the mean number of silanol groups participating in the adsorption; this number is revealed to be equal to fourper nm’, for specific surface areas equalto 50 or 150 m2/g. Thisresult is in agreement with other investigations f251. Free silanol groups are discriminated from vicinal groups; the estimate of the fraction of free silanols is 0.25. Furthermore, surprisingly enough, the amounts of polymer adsorbed on silicas with specific areas equal to 150 or to 300 m2/g are equal. Structures of networks that are formed are easily handled, using a mean field approach to the analysis of the effect of polymer bridgingbetween aggregates; both the square root of the polymer molecular weight and the silica concentration are involved in a quantitative way in a master curve which determines the fraction of bound rubber. The threshold of this bond percolation is well detected fromswelling effects. Thepolymerbridging between aggregatescannot be confusedwith thepercolation of hard objects, whichis usuallydetectedfrom viscoelastic ~easurements.

Competitive adsorption of end-methylated chains and end-hydroFraction of adsorbed end~methylatedchains measured ri ht after the mechanical mixing as a function of the fraction of hydroxyl groups. tion of adsorbed end-methylated chains measured from mixtures in the equili~riumstate of adsorption as a function of the fraction of hydroxyl groups.

Finally, the kinetics of polymer adsorption reveals the role of poison played by water molecules previously adsorbed on silica. The slow release of water molecules, from the surfae, governs the temperature dependent kinetics of polymer binding; the rate FDMS adsorption is a function both of the square root of the polymer molecular weight and of the concentration of silica. C o n s i d e r ~ nany ~ silica-filled FDMS,the glass-like interface formed by monomeric units embedded in the surface, and the interface formed by polymer loops and free-movingchainscanserveasamodelforthestructures resulting from polymer adsorption on mineral aggregates, provided the elementary interaction between monomeric units and the surface is reasonably characterized.

The normalization constantof this function, A,is calculated from thefit of the total number of contact points to the number of available silanol groups, per gram of silica: AT rcn1n2

The probability that one polymer chain is adsorbed on the surface is r,

Y E ,Inz)[1

(n(t))fic]

The computation of Eq. (A3)yields agreement with experimental curves.

~olecular-weight-dependent function, in

The fraction of chain functions, which have reacted is considered as the probability that one chain function has a contact point with the surface. The probability, that one bridge has been formed between two aggregates is expressed

( W

4c46sifSi

where f s i is the mean functionality of one aggregate. The total number of chains connecting two aggregates is p2/2 times the numberof aggregates per gramof silica, while the functionality is determined from the number of sites of adsorption located on one aggregate:

taking Eq. (12) into consideration,

in accordance with Eq. (17). The same description applies to dangling chains; replaced with p1 such that (1 This equation leads then to Eq. (18).

I

Avogadro number specific area of silica prefactor of the lawof chainadsorption,for

fully coatedaggregates

is

he

prefactor of the law of chain adsorption,right after the mechanicalmixing diffusion coefficient of water molecules in PDMS fraction of bound polymer at time t mean number of sites of chain adsorption, per gram of silica fraction of chain functions in contact with the silica surface mass fraction of end-hydroxylated chains per gram of silica fraction of sites of chain adsorption on silica occupied by the polymer molecular weight of a monomeric unit of FDMS number average molecular weight of PDMS effective molecular weight of a bimodal polymer molar weight of hydroxyl group fraction of proton magnetization which belongs to adsorbed units mean number of aggregates per gram of silica number of strands per unit volume which connect two aggregates number of chains which connect two aggregates, per gram of silica mean number of monomeric units in one chain probability that one chain adsorb on one aggregate, only probability that one bridge has been formed between two aggregates amount of adsorbed dangling chains, per gram of silica amount of polymer, per gram of silica, bridging two aggregates amount of polymer adsorbed per gram of silica, at time t amount of polymer adsorbed per gram of silica, right after the mechanical mixing amount of hydroxylated polymer adsorbed per gram of silica, right after the mechanical mixing amount of polymer adsorbed per gram of silica, at time t amount of polymer adsorbed per gram of silica, at time t amount of polymer adsorbed per gram of silica when the surface is fully occupied initial amount of polymer per gram of silica, in a mixture amount of polymer adsorbed per gram of silica, on fully coated aggregates swelling ratio of the polymer in a silica-filled PDMS AqBoundfraction of polymer adsorbed at time t , after the mechanical mixing meannumber of contactpointsperchain s(Qi) factor of the time scaling of the kinetics of adsorption meanarea ofsilica associatedwithonehydrogenbond participating in the adsorption timeconstant of the kinetics of adsorption fraction of units per gram of silica which are fixed

1. Meng-Jiao Wang and S. Wolff, in Black, 2nd ed., (J. B. Donnet and R. C. Bansal, eds.), Science Technology, Marcel Dekker, New York, 1993, Chapter 6. 2. G. Berrod, Vidal, E. Papirer, and J. B. Donnet. J. Appl. Polymer Sci. (1981).

3. E. Papirer, H. Balard, and A. Vidal. Eur. Polym. 24:783(1988). abstracts, 4.E.Papirer and H. Balard,EurofillersMeeting,Mulhouse,Extended 1995, p. 135. 5. P.-G. DeGennes.Macromolecules14:1637(1981). 6. P,-G. De Gennes and P. Pincus, J. Phys. Lett. France 44241 (1983). 7. T. Miller.Science251:905(1991). 8. 0 . Guiselin.Europhys. Lett. 17:225(1992). 9. P.-G. DeGennes,in Scaling Concepts in Polymer Physics, Cornel1University Press, Ithaca, 1979, p. 29-53, 10. M. Aubouy, 0. Guiselin, and E. Raphael. Macromolecules29:7261(1996). 11. M. Aubouy, Ph.D Thesis of the 'University of Paris VI, 1995. 12. G. Armistead, A. J. Tyler, F. H, Hambleton, S. A. Mitchell, and J. A. Hockey. Chem. Phys. 73:3947 (1969). 13. Mathias, G., Wannemacher. J. Colloid Interface Sci. 125:61 (1988). 14. P. Cohen Addad, Polymer 30:1820(1989). 15. P. Cohen Addad, and Touzet. Polymer 34:3490 (1993). 16. P. Gohen Addad, and L. Dujourdy. Polymer Bull., in press. 17. P. Cohen Addad, and N. Morel. Phys. 6:267 (1996). 18. P. Cohen Addad, Polymer 33:2762 (1992). 19. P. Cohen Addad, inProgress in Spectroscopy Eesley, J. Feeney, and L. H. Sutcliffe, eds.), Pergamon Press, Oxford, 1993, p. 303-31 20. P. Cohen Addad, and E. Ebengou. Polymer 33:379 (1992). 21. P. Cohen Addad, and P. G. De Gennes. C. R. Acad. Sci. Paris 319 (S&. 11):25 (1 994). 22. P. Cohen Addad, and N. Morel. C. R. Acad. Sci. Paris 320 (Six. TIb):455 (1995). 23. L. Dujourdy, Thesis,Universityof Grenoble, 1996. 24. P. Vondracek, and M. Schatz. J. Appl.Polym.Sci.,23:2861(1979). 25. P.Gallas, J. C. Lavalley, A. Burneau, and 0. Barres. Langmuir 7:1235 (1991).

This Page Intentionally Left Blank

Department of Inorganic, Physico-Chemical, andMaterial Chemistry, University of Torino, Torino, Italy

US.National Institute for Occupational Safety and Health, Morgantown, West Virginia

I. Introductio~

645

11. Biological Responses 646 to Silica A. Silica-related health effects 646 Fateinhaled of an particle lung in the 646 Mechanisms of silica toxicity and chemical functionalities involved 648 111. Adsorption ofMaterials Endogenous adsorption A. Particle proteins Silica adsorption of pulmonary surfactant and in~ibition of toxicity Surfactant removal and restoration of toxicity D, Surface-surfacta~t interactionsandtheunique toxicity of silica

650 650 652 653 655

IV. ~odificationof Silica Surface Properties and Toxicities by Exogenous Materials 657 A. ~ o d u l a t i o nof toxicity by adsorption organics 657 B. ~ o d u l a t i o nof toxicity by surface inorganic materials 659 References

66

Lung disease associated with occupational mineral dust exposure has been recognized in thewritten record for over two millennia, at least since the time of Hippocrates. Crystallirie silica is probably thefirst material tohave been recognized as particulate toxicant, being responsible for the developmentof long-term disease in people exposed to respirable-sized silica dusts. However, its mechanismof action at the dolecular levelis still obscure. This is due in great part to the extreme

variability in surface properties among quartz dusts arising from different sources, so that any classification of crystalline silica dust as single substance is somehow cumbersome. This is reflected in the conclusions of most present literature in the field, such as the IARC monograph on silica carcinogenicity (IARC, 1997) and a few publications which have followed (Donaldson and Borm, 1998; Fubini, 1998b) describing crystalline silica dustsasavariableentity.Much of this variability resides in surface processes taking place when the dust generated, stored, airborne, and inhaled. In this respect, adsorption processes on silica dusts are of paramount importance to understanding the pathogenic mechanisms.

Exposure to some kinds of silica dusts adversely affects the lungs, causing silicosis as recently established by the International Agency for esearch on Cancer C), bronchogeniccarcinoma(IARC, 1997). Recently videncewas also reported of the association between several autoiml~unediseases and exposure to silica (Steenland and Goldsmith, 1995). C ~ r o l ~ i c s i l i c (nodular o ~ ~ i . ~ pulmonary fibrosis) has been recognized since ancient times as an occupational disease which afflicts people chronically exposed to dusts containing crystalline silica. Acute silicosis (alveolar proteinosis) usually occurs in occupations where silica is fractured or ground into fine powders by mechanical processes (drilling, sandblasting, etc.). In contrast to chronicsilicosis, acute silicosis becomes clinically apparent within a few years of exposure and is a serious, often fatal disease, resulting from acute injury to alveolar lining cells. ~ r o ~ c ~ ~carg e ~ i is lung cancer which can occur in experimental animals exposed to silica dusts and is suspected to occur preferentially with patients with silicosis (IARC, 1997). It also is associated with smoking, which may act synergistically with silica or which may confound epidel~iologicalevidence. Silica-related autoimmulle diseases are typically systemic sclerosis, r ~ e u ~ a t o art~riti,s, id and c ~ r o ~ renal ic ~i,sease. When discussing chemical models it has to be stressed that different mechanisms and thus different surface functionalities may be involved in the various diseases provoked by silica inhalation, although some interplay may occur between these events. In a11 cases the primary event is inflam~ation and recruitl~ent of defence cells in the alveoli (macrophages, neutrophils, etc.).

espirable particles which penetrate deep into the lung andsettle onto the respiratory bronchioles or the pulmonary alveoli will first contact a surfactant coating on the thin fluid coating of the epithelium. That hypophase contains surfactants consisting largely of lipids and proteins synthesizedand recycled by the alveolar typeTI epithelial cells. Those surfactants exist a surface film on the air-liquid interface and as micellar dispersions within the aqueous lining layer. A primary function of this surfactant coating is the decrease and regulation of the pulmonary surftax

tension (Clements et al., 1970; Bourbon, 1991). It appears that another functionis the prompt suppression of otherwise direct membranolytic action of some mineral dusts on the alveolar epithelium. Lavaged pulmonary surfactant or primary components of that surfactant have been shown to rapidly adsorb to the silica surface and to transiently suppress cytotoxicactivity (Emerson and Davis, 1983; Wallace et al., 1985). The subsequent fateof the particle and nature of the disease process dependsat least in part on interactions of the so-conditioned particle with pulmonary alveolar macrophages free on the alveolar surface. Macrophages can phagocytose asurfactant-coated quartz particle and subsequent cellular digestive processes may modify or remove the prophylactic surfactant coating with consequent restoration of cytotoxic activity. A particle may be cleared to the ciliated airways and out of the lung (Fig. 1, path a). Conversely the macrophage may die and the particle, perhaps associated with cellular residue; willbe available again in the alveolar space for possible recoating and possible reingestion by other recruited macrophages, establishing a continuous ingestion-reingestion cycle with accumulation of the free particles in the lung (Fig. 1, path During these cycles the activated macrophagewill secrete a large amount of substances such as cytokines, reactive oxygenspecies (ROS),reactive nitrogen intermediates (RNI),arachidonic acid metabolites and growth factors, which cause persistent in~ammation and may severely damage epithelial target cells (Fig. 1, path c). In the case of chronic silicosis and some other mixed-dust pneumoconioses, the evidence of animal model studies and of histopathology examination of human pulmonary tissue is that some particles do enter the interalveolar septa; this may involve transport by the alveolar macrophage or other mechanisms of penetration

Events following phagocytosisby alveolar macrophages: (a) clearance to the ciliated airways and out of the lung; (b) ingestion-reingestion cycle with accumulation of the free particles in the lung; (c) damage to epithelial target cells following persistent i n f l a ~ ~ a t i ocaused n by macrophage activation.

of the alveolar epitheliuminto theinterstitium. Interaction is then possible with the pulmonary macrophage, fibroblasts, or other cells within the septa. studies of pulmonary response ofmice to intratracheal instillation ofsilica dust using irradiation to suppress the initial inflammatory influxof alveolarmacrophages andpol~morphonuclear leukocytesindicatedtranslocation ofsilica particles through the pulmonary alveolar epithelium with subsequent formation of granulomas and fibrosis there (Bowden et al., 1989). It is not clear as to which disease processes result from interactions of particles with the free surfixe alveolar macrophages prior to this distribution of the particles, or which follow processes begun within the septa adjacent to the pulmonary fibroblasts located there. Research has indicated that the path followed by a deposited particle to benign clearance or to pathogenic interaction with alveolar macrophages or with cells behind the alveolar epithelium, may be dependent on the state of the particle surface. In particular, the path taken may be determined by endogenous biochemical constituents of the lung or exogenous organic or mineral materials adsorbed onto the particle surface.

Aproposedgeneralmechanism of toxicity of silica causing silicosis andlung cancer(Kane, 1996; IARC, 1997; DonaldsonandBorm, 1998)is schematized in Fig. 2. On the basis of literature data, surface functionalities have been associated withthe biological steps in which such functionalities appearinvolved (Fubini, 1998b). Cytotoxitymay be related to thedistributionandabundance of silanols (SiOH) groups at the surface (Pandurangi et al., 1990; Hemenway et al., 1993; Fubini et al., submitted) and/or to silanol groups dissociated in water (Nolan et al., 1981). It has been longhypothesized that cytotoxicity was originated by strongadsorption of membranecomponentsontothe silica particle (Nashet al., 1966). A role for silanol activity in themembranolyticaction of quartz was supported by the more recent observation of the diminution of hemolytic activity by quartz dust with dehydroxylation of the quartz particle surface by calcination (Razzaboni et al., 1990). It has been suggested that toxic membranolytic activity is based upon hydrogen-bonding of silica surface silanol groups with nitrogen or oxygen moieties of macromolecules of biological membranes or upon interaction of dissociated anionic surface silanols with ionic lipids or proteins of biological membranes (Nolan et al., 1981); ,or radicals on the surface of freshly fractured silica causing damage of membrane lipids by direct peroxidation or by Fenton reaction generation of hydroxyl radicals (Castranova et al., 1997). One or more of these reactive particle surface sites might then interact with a pulmonary macrophage, triggering toxic events in the cell and pathogenic effects in the pulmonary alveolus. Anactivatedmacrophage cycle may be established, and particle-derived ROS (Fubini et al., 1989; Giamello et al., 1990) and cell-derived ROS (Vallyathan et al., 1992) will both contribute to state of oxidative stress, persisting as long as the inflammation persists. Cells will also release nitric oxide, which contributes to the oxidativestress and in the presence of the superoxide ion forms the dangerous compound peroxonitrite. Particle-derived ROS, such as free

silica particle in the lung form, crystal structure Si0H Si0

surface radicals, charges

inhibition

clearance

all~miniumions other metal ions hydro~hobicsurface

iron particle derived *free rudicals surface r a ~ ~ c a l s

silicosis

Thepossiblerole of physico-chemical factors in thesequence of events leading to the pathologies associated with inhalation of crystalline silica reported by Fubini (1998a).

radicals or peroxides, may also be implied in direct damage to the epithelial cells. Several particle-derived ROS have been reported, such as hydroxyl radical, superoxide anion, and peroxides (Shi et al., 1995). The production of sili~~-originated free radicals is much higher on freshly ground materials, where surface peroxide or hydroperoxides are formed (Fubini et al., 1990; Giamello et al., 1990), therefore such a step is rnore relevant in the case of freshly ground than of aged silicas. If some iron, even in trace amounts, is present or has been adsorbed at the silica surface-which is a very common situation withminerals-Fenton chemistry may be activated, withconsequentprolonged releaseof radicals, which maycause DNA damage andtransformationjn target cells. Free radial generationdoes not usually relate to the actual amount of iron but to small fractions of iron with a particular redox and co-ordination state (Fubini et al., 1995b; Gilmour et al., 1995).As aconsequence of theabove effects, mutations. and proliferation in epithelial cellsmay initiate a neoplastictransformation. It appearstherefore that the potential of the inhaled particles to catalyze ROS release and to persistently activate macrophages would determine the pathogenicity of a given dust. Surface processes act to inactivate these processes, e.g., removal of radical generating sites such as adsorbed iron or others, should also inhibit the pathogenic response. Any surface property, moreover, favoring path a instead of b, by lowering the extent of accumulation of the dust in the lungs and consequent inflammation, will also lower the pathogenic potential.

However, theories of particle surface functional groupsand interactions responsible for the disease-inducing potentials of silica must consider the possible modulation of silica pathogenicity by endogenous or exogenous processes of particle surface modification. Several studies of exogenous pretreatmentof dust surfaces have been performed. If a respirable dust is covered by polymers (Mao et al., 199S), has been chemically modified (Wiessner et al., 1990), is hydrophobic (Hemenway etal., 1993; Fubini et al., submitted), or has been treated with aluminum salts (Brown and Donaldson, 1996), the effect of silanols is much reduced or even eliminated. Under these circumstances theparticle may followpath a inFig. l , i.e., will be cleared out from the lung in the upper airways or to lymph nodes by macrophages. Alternatively, following path b, clearance will be inhibited and phagocytosis will eventually cause cell death following disruption of the phagolysosome membrane. There are two broad concerns bearing on the role of endogenous modificationof silica surface in the determination of pathogenic activity in vivo. One is the question of the prompt modulation of silica particle surface activity from conditioning by the mucus and surfactant protective layers on the lung surface, e.g., the prompt masking or neutralization by surfactantadsorption of particle surface silanol groups or surface free radical species (Marks, 1957; Emerson and Davis, 1983). The other is the question posed by the comparable in vitro cytotoxic activities of silica and some aluminosilicate dusts, despite the relatively weak disease-inducing potentials of thealuminosilicatedusts in vivo (Brownet al., 1980; Daniel and LeBouffant, 1980). That is, what differences in surfacepropertiesprevent clays but not quartz from expressing comparable toxicities in vivo? Tests of postulated mechanisms of silicosis typically do not test such comparably cytotoxic but relatively weakly pathogenic mineral silicate clay dusts as a negative control. However, measures of the in vitro cytotoxicity of quartz and alu~inosilicateclay respirable dustshavefoundthattheyare Comparably cytotoxic to lavagedmacrophages (Vallyathan et al., 1988). Thissuggests thattheoutstandinghazard of quartz dust is not due to the presence of unique surface toxic functionalities and interactions, but rather is due to a lack on the quartz surface of surface functionalities and interactions which are responsible for prophylacticeffects on other silicate mineral surfaces.

Proteins are large amphipatic molecules which, being intrinsically surface active, tend to adsorb toall surfaces. The enthalpy of adsorption varies over a wide range and depends upon the chemicalnature of both surface and proteins. In some cases the enthalpy of adsorption may even be positive and the process of adsorption is entropically driven (Brash and Horbett, 1995). Entropic factors are always relevant because of the displacement of water molecules from both surface and protein well as limited unfolding of the protein at the surface. The physicochemical nature of the solid surface determines the kind of protein which preferentially adsorbs, as well as the strength of the bond. a general rule

the more hydrophobic thesurface, the greater the extent of adsorption (Brash and Horbett, 1995). This rule has several exceptions.Low-densitylipoproteinsare adsorbed preferentially on hydrophilic silica surfaces (H0 and Hlady, 1995); often adsorption on hydrophilic surfaces is underestimated because the adsorbate is more easily removed(Brash andHorbett, 1995). Serumalbumin is strongly adsorbed at silica or glass surfaces; heterogeneity in adsorption sites, namely the presence of calcium ions acting as electron acceptors, attracts the electron donor moieties of the protein (Van 1994). EDTA in fact decreases adsorption. Silica does adsorb proteins, likely viahydrogen bondingof their polar parts onto silanols (Van Oss, 1994). One of the hypotheses for thelytic action of silica on some cell membranes was in fact a strong adsorption of the external part of membrane protein onto silanols ( S u ~ m e r t o net al., 19’7’7). The adsorption of several proteins was thoroughly investigated by Kozin et al. (1982) in the course of an il~vestigationof the l~embranolyticpotential of several form of silicas. The following orderof affinity for adsorption by quartz was found when results were expressed as percent of protein bound: fibrinogen lysozyme IgG ribonuclease ovalbumin ,k”ctoglobulin bovineserumalbumin. Differences were also noted in the affinity of various silicas for proteins. Too much speculation on these differences could be pointless because the authors used “synthetic” and mineral samples. The mineralcristobalite adsorbed two ordersof magnitudemore than the synthetic one,probablybecausethe synthetic onehad experienced a high-temperature treatment which inactivates surface functionalities (Fubini, 1998b). The hemolytic potentialof the silica particles correlated with their capacity to adsorb IgG orlysozyme (LYS), both positively charged, but not bovine serum albumin(BSA),negatively charged. Expressedas a molar ratio of adsorption SA, theresults indicated that only particles with such a ratioof 5.0 or below were membranolytic, suggesting a requirement of an optimum negative surface charge for l ~ e ~ b r a n o l y s iAdsorbed s. proteins partly inhibited hemolysis. Adsorption mayalso stern from hydrophobicinteractions. The adsorption bebavior of human plasma fibronectin was investigated on silica substrates of different surfaceenergy,onehydrophilic and onehydrophobic(Jonsson et al., 1982). Fibronectins are high-molecular-weight glycoproteins present on many cell surface, in connective tissues, and in extracellular fluids. They are involved in many cell adhesion phenomena. Similarlyto what happens with humanfibrinogen, there was an increased amount of protein adsorbed at the plateau on a hydrophobic surface as compared to a hydrophilic one. Furthermore, adsorption was nearly irreversible onahydrophobicsurfacebut partially reversible on thehydrophilicone. Interaction of antibodies with preadsorbed fibronectin suggests that ~bronectin adsorbs in different conformations and/or arrangements on the two types of surfaces. Thesurface of crystalline silica exhibits hydrophilic andhydrophobicareas (Bolis et al., 1992; Fubini et al., 1992, 1993). Proteins will therefore be adsorbed at different sites depending on their characteristics. The possibility that particular arrange~entsof hydrophilic and hydrophobic sites on a crystalline surface may strongly and modify a proteinis one of the hypotheses put forward years ago to explainthepeculiar biological properties of most crystalline silica polymorphs

(Langer, 1978) and has neverbeen fully discarded. Irreversible modifications of biomolecules at the surface may activate the immune system which will regard them as nonself, hence the autoimmune silica-related diseases and a possible “immune” pathogenesis of silicosis as postulated by Pernis and Vigliani (1982). It has to be pointed out that plasma proteins adsorbed on silica dusts acquired antigenic properties as reported long ago (Scheel et al., 1954).

Pulmonary surfactant is a multicomponent mixture of lipids, proteins, and carbohydrates (Clements et al., 1970; Kingand Clements, 1972; King et al., 1973; Haagsman and van Golde, 1985; Bourbon, 1991). Lipids are the niajor component of pulmonary surfactant; the major fraction is phospholipids, principally diacyl(palmitoy1)phosphatidylcholine (PPPC) (Gilfillan et al., 1983). DPPC in physiological saline can reproduce almost all of the surface-tension-altering behavior of pulmonary surfactant on the air-liquid interface of the alveolar surface. Therefore DPPC in saline frequently is used as amodel for pulmonary Surfactant. Adsorption of complete (lavaged) pulmonary surfactant or of PPPC can suppress the in vitro cytotosicity of quartz (Marks, 1957). This suggests a surfactant role in preconditioning mineral particlesurfaces to prevent such interaction from damaging alveolar epithelial cell membranes. Decrease in hemolytic activity versus surface-normalized amounts of adsorbed PPPC has been measured for quartz and kaolin respirable dusts (Keane et al., 1990; Wallace et al., 1985). Isothermsat 37°C foradsorption of PPPC from dispersion in physiological saline were measured in two ways. The DPPC remaining in dispersion indicated that approximately 15 mg PPPC adsorbed per square meter quartz surface area (the surface area was measured by BET nitrogen gas adsorption) (Wallace et al., 1988). However, multiple saline rinsing of the quartz dust removed all but about 4 mg DPPC/m* as determined by organic solvent elution and wet phosphate assay quantificatio~(Wallace et al., 1992). The kinetics of enzymatic removal of this tightly adherent DPPC suggest that this coating may be a bilayer with ionic phosphatidylcholine headgroups of the DPPC oriented toward the quartz surface for the inner molecular layer and oriented toward the outeraqueousmediaontheouter molecular layer, with thefatty acids tails between (Fig. 3). Such DPPC pretreatment and rinsing of quartz dust fully inhibits the hemolytic activity of quartz againsterythrocytes invitro andagainst lavaged rat pulmonary macrophages as measured in vitro by lactate dehydrogenase, beta glucuronidase, and beta N-acetyl glucosaminidase release assays (Wallace et al., 1985). Prompt and total suppressionof quartz membranolysis by DPPC ph~spholipid surfactant adsorption has been observed for freshly fractured quartz dustas well as for’standard quartzdusts. It is not clear that surface free-radical or~highlyreactive sites would survive a surfactant adsorption~igestioncycle of events. Perhaps the question to be asked of a role for free-radical sites is: Under what conditions could highly reactive silica surface sites produce significantlevels of toxic intermediates by reaction with lung-lining organics? And, are suchhighly reactive surface sites also

Adsorption of DPPC surfactant on a quartz surface silanol. The cationic trimethyla~nloniu~ end of the DPPC interacts with an acidic silanol surface site. Silicate basic aluminol surface site interactions with the acidic phosphate may alter this DPPC conformation. Sites of phospholipase A2 and C hydrolysis are shown. This molecule would be backed by an outer oppositely directed DPPC molecule.

present on freshly broken mineral dusts which are not strongly pathogenic? With such information one could then reasonably speculate about involvement of such interactions in acute or accelerated silicosis.

The lack of persistence of prophylactic effectsof surfactant on quartz and the consequent restoration of particle toxicity have been measured in cell-free system models of enzymaticdigestion processes, for cellular processes and in limited animal model studies. In cell-free enzymatic digestion e~periments, phospholipase A2 (PLA2) (Fig. digestion of DPPC-treat that enzyme could hydrolyze the fatty acids from the ads sequent loss of lysolecithin digestion product from the parti digestion rate data used a kinetics model for e n z y ~ a t i cdigestion of an adsorbed bilayer considering steric hindrance both of the outer molecular layer to the inner particle-surfac~-associatedDPPC layer, and of hindrance by the enzyme-substra~e complex to enzyme access to adjacent DPPC molecules. The analyses suggested adsorbed as a tightly held bilayer, with the outer part the bilayer igested by PLA2, and the inner part of the bilayer digested much more slowly, and with mineral specific rates. Hemolytic activities of parallel samples showed a correlation of restoration of membranolytic activity with removal of the particle-surface-ads or be^ layer of surfactant(Wallace et al., 1992). These cell-free analysesindicated that phospholipasecould digestively remove from silica and consequently restore particle cytotoxicity; but cell-free experiments cannot predict absolute values of rates of such restoration for cellular or processes (Wallace, et al., 1994).

rates of cellular digestive removal of radiolabeled DPPC from respirable quartz dust using the macrophage-like P388D1cell line found that two-thirds of the mineral surface-adsorbed DPPC was digested in a one-week period. Tests used radiolabeled DPPC adsorbed to quartz to challenge a cell line Over the one-week periodafterdust challenge, columnchromatographyandsubsequent liquid scintillation counting wereused at one- to three-day intervals to deternline the radioactivity which remainedwiththe PPCfraction of lipid extractedfrom particles after cell disruption;that is, tdeterminethe rates of digestion of particle-adsorbed DPPC. This found that approximately two-thirds of the adsorbed DPPC was digested in a six- to nine-day period by these cells with no significant mineral specificity. Lysosomes of these cells were acidic. phospholipase r digestive activi seen (Hill et al., onary macrophage viability and damage induced C-treated dust challenge was measured in parallel with loss of fluorpid surfactant from phagocytosed particles overaoneof digestion of a commercial organoboron fluorescent-labeled challenge of lavaged macrophages was measured by fluorescenceconfocalmicroscopy. Parallel systemsmeasured damage by the single-cellgel “comet” assay, and viability by trypanbluedye exclusion and by adualfluorescencecommercial “live-dead” assay. Surfactant adsorption fully suppressed quartz toxicity at onedayafter challenge; approxil~ately 60% of thelabeled-phospholipidfluorescencewas lost overa seven-day period, and viability was restored to three-quarters to full activity of native quartz over that period. DNA damage activity also was initially suppressed; but activity began to appear again after several days (Liu et al., 1997, 1998). Using Chinese hamster lung fibroblast-derived V79 cells, challenge with silica dusts was found to induce micronucleus formation at high dust doses. C pretreatment of the dusts suppressed this activity over a five-day post silica enge time period (Liu et al., 1996). commercial bovine-based pulmonary surfactant treatment of silica was used to challenge lavaged rat alveolar macrophage Cell viability as measured by trypan blue dye exclusion was approximately 10% of controls for l h challenge withnative silica, and 90% forsurfactant-treated silica. After 24 b of challenge by surfactant-treated silica, viability droppedtoabout 70% of the his bovine surfactant also was used by tracheal instillation in the rat, with subsequent assay of bronchiolar lavage assays at 1 and 14 days fortotalprotein,betaglucuronidase activity, and neutrophil influx. ~ i g n i ~ c a n t differences indicative of a protective function were seen between the native and sur~~ctant-treated challenges at one day for protein and enzyme assays, but not forneutrophil i~filtration.The effects haddiminished to nonsignificant levels at 14 days(Antoniniand 1994). Another set of experimentssuggested that the loss of surfactan laxisis duetopbospholipaseenzymatic rior to and following intratracheal instillation of silica, rats were treated orally with amiodarone, a drug known to induce phospholipidosis in the lung, presumably by inhibiting phospholipase activity. At 60 days after silica instillation, lung weight, hydroxyproline content as an indicator of fibrosis, neutrophil

percentage of lavaged cells, and lavage fluid content of albumin, /?-glucuronidase, and lactate dehydrogenase were measured. In all the assays the values for silicachallengedbutnon-drug-treatedanimals weresignificantly increasedoverthe negativecontrolnon-silica-challengedanimal values; forthe silica-challenged andarniodarone-treatedanimalsthevalues were signi~cantlygreater thanthe negative controls, butwere significantly lower than the values for silica-challenged but non-drug-treated animals (Blake et al., 1996).

In order to attempt to find the mineral-specific surfaceproperties and interactions responsiblefortheuniquelystrongpathogenicpotential of quartzdust, explicit direct comparisons of quartzdust with kaolinite aluminosilicate clay respirable dustshave been madeforsome of thesurfactantinteractions discussed above. This has shown the two dusts to have comparable cytotoxicity in vitro as measured byhemolysis or by macrophage damage, to have equivalent adsorption capacities per unit specific surface area for the primary phospholipid component of pulmonary s~~rfactant, and tohave their cytotoxicities fully suppressed by surfacecoveragewith DPPCsurfactant.Itappearsto befirmly established that silanol groupsaremembranolytic, andalthough transiently suppressed by adsorption of surfactant,the activity can be restoredthrough a cycle of prophylactic surfactant adsorption and removal; that is, the potential membranolytic activity survivesthe initial defensive interaction in the lung. A corollary to this is thatit is nota lack of toxic silanol groups which makes aluminosilicate clays relatively nonpathogenic; rather it is the presence of adjacentprotectivesurfacefunctionalgroups, e.g., aluminolgroups or heteroatoms on claywhich counterthe silanol toxicity (Wallace et al., 1989; Keane et al., 1990). The differing pathogenic potentials appear not to be due to differences in initial surfactantinteractionsasbothkaolinandquartzadsorbsurfactantand are promptly inhibited in their direct membranolytic action. The question arises: Are there significant differences in restoration of toxicity dueto prophylactic s u r ~ ~ c t a nbeing t removedmore fully or morequicklyfromquartz, such. that quartz pathogenicity isexpressed andother silicate activity remains inhibited foraperiod that is longcompared to particle clearance or retention in the lullgs? Cell-free phospholipasedigestionremoved DPPCmorerapidlyfromquartz than from kaolin: half the tightly adherent DPPC was digested from both dusts at a comparable rapid rate, while the remaininghalf, postulated to be the molecules in contact with mineral surfaces, showed a fourfold decrease in the kinetic rate constantforremovalfromkaolincornparedtothatforremovalfromquartz (Wallace et al., 1992). Those cell-freesystem studies used aporcinepancreatic exudate phospholipase A2, which was optimally active at neutral pH. However, the P388D1 cell in vitro studies did not see a mineral specificity for intracellular surfactant digestion or restoration of cytotoxicity. Preliminary data do not show mineral-specific differences in the rate or restoration of cytotoxicty of DPPC-treated quartz and kaolin to lavaged rat macrophage in vitro (Gao et al., 1999). The P-

388Dl cell-line studies observed an extracellular phospholipase released by the cells in this phospholipase was a pH-neutral-acting enzymewhich didshowa mineral specificity for DPPC removal from quartz and kaolin. Thus extracellular phospholipase appears to be more strongly active against quartz-adsorbed DPPC, while the intracellular phospholipase of phagocytic cells with acidic lysosomes and -active phospholipase do not appear to show a mineral specificity in restoration of dust toxicity. A hypothesis emerges from these considerations on the effects of endogenous surfactant in determining the exceptionally strong pathogenicity of silica. Silica or silicate particles depositing in the deep lung adsorb components of pulmonary surfactant and are thereby promptly inhibited in their cytoxicity. Such inhibited particles are phagocytosed by alveolar macrophages or make their way into the alveolar epithelium. The surfactant coating on the particles is subject to attack by lysosomal enzyme in the alveolar macrophage phagolysosome or in the cellular lysosomes of interstitial cells, orattack by extracellular enzyme in the interstitium. Therate ofsuccessof this enzymaticdigestiondependsonthe detailed structure of the particle surface, different surfacefunctionalgroups having different affinities for surfactant functional groups, affecting the strength of adsorptionandtheconformation of theadsorbedsurfactant molecule and thereby affecting the rate or degree to which the digestive enzymes can remove theprophylacticsurfactant. A molecularmodel is that theconfornlation of kaolin-adsorbed DPPC providesgreaterhindrance to phospholipase enzyme access to the phosphate and carbonyl ester bonds which the enzymes hydrolyze alongthe glycerol moiety of thephosphopholipidmolecule.Both quartzand kaolinhave acidic silanol surface sites whichcan interact withthe positively charged choline trimethylammonium group at the head of the however, clay and not quartz would have basic or amphoteric aluminol surface sites tointeractwiththe acidic phosphateorcarbonylgroups. T ~ L Ithe S clay surface interactions with the hydrolysis-susceptible region of the molecule might result in an adsorbedDPPCconformation which retards enzymeaccess and activity there. Thusbothmineralswouldhave potentially toxic surfacefunctions, but quartz would lack adjacent “protective” functional groups. The particular type ofcell interacting withthe particle may also be critical in this process, asalveolarmacrophageshave acidic phagolysosomal processes, while epithelial and mesothelial cells have neutral pH lysosomes (Johnson and Maples, 1994). Thus molecular conformations and susceptibility to enzymatic action may y between alveolar and interstitial locations of the particle. In thecase of PC sur~dctant, the available data suggest a nonmineral specific restoration of particle toxicity by alveolarmacrophages and the possibility of morerapid ofsilica particle toxicity compared to clay for particles subject to pho~pholipaseactivity, e.g., sequestered in the interstitium behind epithelial surface. However, the question is openas to the effects and kinetics of other components of pulmonary surfactant. esearch is needed on the comparative rates of loss of pulmonary surfactants and restoration of toxicity for quartz and clay particles within the alveolar interstitium to determine if intracellular or extracellular enzymaticprocesses these

could distinguish the longer term dusts.

in

response for different respired mineral

uartz on the development of lung disease was investigated by treating quartz with various organosilanes (Wiessner et al., 1990). The samples havebeen tested in on a mouse model measuringparameterslinkedtolunginflammationand fibrosis. Red-blood-cell lysiswas also measured as ameans to compare the reactivity towardsthe cell membrane with the pathogenic response. The samples studiedwere well characterized and they were administered to mice by intratracheal injection with each crystalline material at constant surface area, which implies that it is the surface of the particles whichelicits thepathogenicresponse investigated. Several parameters related to the inflammatory and fibrotic response were investigated, such as lung index, cell number, lavage protein concentration, and hydroxyproline functional groups attached to the quartz surfaces were (-CN), NHJ, and (-N(CH&). The crystals showing the highest degree of biological activity were nativequartz,and N(CH3)3 and CN-modifiedquartz. Conversely, the -NH2-modified one was as unreactive as the crystal preparation modified with the hydrophobic group -CH3. The authors conclude that electrostatic interactions may be more important indetermining effective biological activities than are hydrogell-bonding interactions, as the -NH2 group, which can give hydrogel1 bonds, was as unreactive as the hydrophobic "-CH3 terminal group. This very nice set of data deserves probably further interpretatioll as it contains information still to be drawn on the effect of the chemical modifications in the various steps of the pathogenic process. For instance, the possibility that the terminal groups of the chain might be involved in interaction with some other surface sites has not been considered. All silane-l~odifiedsurfaces were less hemolytic than pure quartz. This is likely due to the eliminatioll of surface hydroxyls in the silanization andtothe hydrophobicitybrought about by thehydrocarbonchain. ~ispersionof the quartz dusts in the pulmonary surfactant PPC inhibited hemolysis but had little or no effect in i~flammation and fibrosis. I n ~ a m ~ a t and io~ fibrosis were higher for pure quartz than for the modified ones, the -N(CH3)3 and -CN ones was higherthan those of the ones. The effect of a commercial organosilane reagent for treatment of laboratory glassware, Prosil@,was tested for its effect on silica toxicity in The organosilane was applied to silica dust by incubation at 100°C for 10 min. This treatment reduced the hemolytic activity of silica by 78% from the untreated silica-induced value. In challenge of lavaged rat macrophage showed a decrease in darnage to cellular membrane integrity as measured by a fluorescent propidium uptake assay

ini

!lac

at 5 hafter cell challengefororganosilane-treated quartz compared to native quartz. Organosilane treatment of silica also lowered by 83% the oxidant release from challenged lavaged macrophages in vitro as measured by chemiluminescence with a luminol indicator. Intratracheal instillation of Prosil@-treated silica in the rat resulted in no apparentsignificant difference in the viability of cells lavaged one day after dust challenge between native and coated silica dusts, and no apparent significant difference in amounts of protein in the lavage fluid. A decrease, not tested for statistical significance, is seen for the induction of beta glucuronidase in lavage fluid for the treated dust (Castronova et al., 1997). It has been observed that somepreparatoryconditionsfor in vitro toxicity studies ofsilica dustscaninadvertently alter themeasurements,apparently by modification of the quartz particle surface. In the course of sterilizing dusts for long-term in vitro testing it was observed that boiling quartz dustin glass test tubes suppressed the cytotoxic and membranolytic activity of the quartz (Wallace et al., 1990b). It is not clear if this is due to contamination of the quartz dust by organosilane or other laboratory glassware treatment materials partially released by the boiling conditions,or due to contaminationby a low-solubility silicon or aluminum or boron or other compound released by the glassware glass surface itself under boiling conditions. The suppression did not occurwhen quartz was boiled in polycarbonate tubes. However, it did occur when quartz was boiled in polycarbonate tubes containing glass beads or ground up glass cover slips. The latter result suggested that the prophylactic effect was not due to release of some organic coating on the test tubes. An alternative hypothesis is that some slightly soluble form of silica in aqueous media has a greater solubility from glass than from crystalline quartz surface in boiling water, resulting in silicic acid or some polymerized derivative being adsorbed on the quartz surface.

One of the most used silicosis inhibitors is (P~PN~), which is much more effective against quartz-induced fibrosis than against asbestosis. The -NO groups of PVPNO provide a periodic point for strong attachment The mechanismof action is still partially unclear. It is likelydue to the coating of the surfaceof the dustby the polymer (Schlipkoterand Brockhaus, 1960) but it has n found to act as a scavenger of free radicals (Gulurnian and van Wyk, inding to the surfaceis certainly important andmay explain thedifferences tween the effect of PVPNO on quartz and asbestos on the basis of their different H-bonding potential. Investigating the effects of surface-modifying agents on the production of reactive oxygen metabolites by polymorphonuclear leucocytes, Klockars et al. (1990) foundremarkable differences between quartzand various asbestos types. PVPNO inhibited oxygen metabolite production by quartz but had little effect on asbestos. Conversely, carboxymethylcellulose only reduced chrysotile activity, but wasineffective with other particulates, including quartz. There is clearly asort of sur~dce-inhibitorspecificity in polymer which is determined by the surface properties of the particle surface charges, etc.).

The toxicity of quartz to fetal rat lung epithelial cell line was investigated by in-U-Si1 as received and precovered by twoforms of yl), one of which (2-vinyl) only was active in blocking is (Nolan et al., 1981). The toxicity was evaluated as colony-forming efficiency (proportion of cells that survived to form colonies). inding of both forms of PVPNO effectively inhibited the toxicity of quartz to the cell line investigated, mostly at the same level. The authors suggest that quartz toxicity to epithelial cells may be caused by hydrogell-bonding interaction of the silanol group withcellular components (different from those involved in hemolysis) and/or formation of free radicals at the silica surface.

oluble aluminum salts, administe d during the exposure to silica dusts, blunt the adverse effect of silica. Long ago aldane (1917) hypothesized that the nonfibrogenicity of some quartz-co~taining dustscould be related to the presence some components-l~ainly claysilicates-capableof blunting silica toxicity. years later it was reported that theinhalation of aluminumdustsinh experimental animals the pathogenic response to crystalline silica dust. ons side ring that most probably the inhibition is carried out by soluble aluminum ouffant and associates experimented with different salts and ad~~inistration protocols (Le et al., 1977). They reported that among various alumi nu^^ salts, most alum the appro~riate,because of its high stability in aqueous tion of aerosols prevented the develo~ment of pulmonary fibrosis pure to crystalline silica dusts or silica-contsmineral taining tudies then followed by gin and associates m lactate inhibits t stratedthatsurface si~nificalltly reduced the biological activity of quartz and increased its clearance with no detectable particle retention in the lung 10 months after exposure. Further studies demonstrated that the in~ammatoryresponse to quartz in the ameliorated even when aluminum was administered post exposure 1988). The ~ e c h a n i s of l ~ action is not fully clarified, but it is noteworthy that aluminum is active if deposited on the silica particle and if admini few days after administrationof silica to experimental an genic nlechanism had proceeded, aluminum wasineffective. action of aluminum ion on some surface fullctionalities ofsilica implied in the ce modi~cationsbrought about by aluminum lactate, the measured on a pure quartz dust b e and after a treattate,followingtheprotocol used b &ginand associates (198’7) to prevent silicosis (Fubini et al., 1995a). The most remarkable difference between the two samples was found in the irreversible reaction with ammonia on

the partially dehydrated samples, which evidences a much higher Lewis acidity on the aluminum-treated sample. Aluminum decreases silica solubility. Aluminumin a silica framework substitutes forsilicon in a tetrahedral position, and acts asLewis a acid (electron acceptor), but also enhances Brransted acidity (proton donor), by facilitating donation from nearby silanols. A higher aciditymay affect surface affinity for biomolecules, membranolytic potential and the overall reactivity the surface. The presence of aluminum at the silica surface likely favors path a (clearance) in Fig. 1 relative to path b (macrophage activation and death). ~pidemiologicalstudies of fibrogenic pulmonary lung disease in some occupations with exposure to quartz in mixed mineral dusts havefailed to find correlations of disease risk with the quartz component of cumulative respirable dust exposure. This is a well-known phenomenon in epidemiological studiesof pneumoconiosis or progressive massive fibrosis in coal workers (Walton et al., 1971; Hurley et al., 1982; Robock and Bauer, 1990; Attfield and Morring, 1992). Innate aluminosilicate surface contamination of respirable quartz particlesin the mixed dusts may explain the seeming anomaly. Submicroscopic layers, e.g., few a atomic layersto tenths of a micrometer thick, of aluminosilicate clay or other minerals of low fibrogenicity coveringarespirable-sizedquartzparticlewould be detected only as low-level bulk c o n t a ~ i n a n tof theexposure by conventionalindustrial hygiene exposure characterization methods, i.e., infrared spectroscopy or x-ray diffraction of collected dust samples. But if such clay coatings of the quartz substrate particles exist and remain adherent following depositionin the lung, then they would represent a qualitative difference in the nature of the exposure: the particle would have the toxicological properties of the occluding mineral coating rather than that of the un~erlyil~g mineral particle identified by the conventional exposure analysis. Auger spectroscopic analysis and thermoluminescence analysis of silica particles from German coalmine dusts found evidence of aluminum contamination on all silica particles studied (Kriegseis and Scharmann, 1982). In vivo toxicity studies challenging the rat with silica added to coalmine dust of low silica content versus a coalmine dust of equal inherent silica content indicated that the silica found inherent in the coal dust did not possess the strong pathogenic activity of pure ouffmt et al., 1982). Laser microprobe mass spectrometric ( L A ~ ~ S ) study of quartz particles in German coalmine dusts found evidence of admixtures of quartz with clay minerals as a surface contamination of quartz particles, From this it was postulated that most quartz particles are neutralized by such coatings in the coalmine dusts, and proposed that the in vivo toxicities of the dusts are functions of non-quartz-related factors (Tourmann and Kaufmann, 1994). The use of scanning electron microscopy with x-ray spectroscopic analysis was modified to detect submicro~eter-thickcoatings of aluminosili~ate onrespirable-sized quartz particles by performing the analyses at successively lowered electron beam accelerating voltages, thus generating elemental x-ray spectra from decreasing depths into a particle (Wallace et al., 1990a). This revealed the presence of aluminosilicate occlusion of some fraction of quartz particles in dusts from clay works and from coalmines. This analysis of dusts from a limited number of U.S. coalmines suggested that the fraction of quartz particles surface occluded increases with decreasing coal rank, coincidentwith the epidemiological observationthat disease risk per

ici

unit cumulative respirable coalmine dust exposure decreases with decreasing coal rank (Harrison et al., 1997).

Antonini, J. M. and Reasor, M.J. (1994) Toxicol. Env. Hlth 43535. Attfield, M. D. and Morring, K. (1992) Am. Ind~st. Hyg.Assoc. .53:486. Bkgin, R., Massk, S., Skbastien, P.,Martel, M., Bossk, J., Dubois, F., Geoffroy, M,, and LabbC, J. (1987) Exp. Res. 13:205. Blake, T. L., DiMatto, M., Antonini, J. M,, Mccloud, C. and Reasor, Lung Res. 22:l 13. Bolis, V., Fubini, B., Marchese, L., Martra, G., and Costa, D. (1992) Chem. Soc. Faraday Trans. 87:497. Bourbon, J. R. (1991) in Pulmonary Surfactant, ~unctional,regulator^, Chemical Concepts, (J. Bourbon, ed) Chap. 1, CRC Press, Boca Raton, FL. Bowden, D. H., Hedgecock, C., and Adamson, L. Y. R. (1989) Pathol. 2.58:73. Brash, L. J. and Horbett, A. (1995) in Proteins at I n t e ~ f f c e s Fundamentals and Applicatio~s(T. A. Horbett and L. J. Brash, eds.), ACS Symposium series 602, San Diego, CA, pp. 1-23. Brown, G. M. and Donaldson, K. (1996) in Silica Silica-In~ucedLung Diseases (V. Castranova, W, Wallace, and V. Vallyathan, eds.), CRC Press, Boca Raton, FL, Chap. IV.4. Brown, R. C., Chamberlain, M., Davies, R., Morgan, D. C., Pooley, F. D., and Richards, T. (1980) in The Vitro E’yeeets ~ i n e r a Dusts l (R. C. Brown, et al., eds.), Academic Press, London, pp. 47-52. Castranova, V., Dalal, N. S., and Vallyathan, V.(l 997) in Silica and S i l i c a - I n d ~ Lung ~~e~ Diseases (V. Castranova, W. Wallace, and V. Vallyathan, eds.), CRC Press, Boca Raton, FL, Chap. 11.3. Clements, J. Nellenboyer, J., and Trahan, H. (1970) Science 169:603. Daniel, H. and Le Bouffant, L. (1980) in The Vitro Effects ~ i n e r a Dusts l (R. C. Brown et al., eds.) Academic Press, London, pp. 33-39. Donaldson, K. and Borrn, P. J. A. (1998) Ann. Hyg. 42:287. Dubois, F.,Bkgin, R., Cantin, R., Massk, S., Martel, M., Bilodeau Dufresne, A., LabbC, J., and Skbastien, P. (1988) Am. Rev. Resp. Dis 137:1172. Emerson, R. and Davis, G. S. (1983) Environ. Health Perspect. 51:81. Fubini, B. (1994) in Cellular ~ o l ~ c ~Eflects l a r of ~ i n e ~ a l Synthetic Dusts and Fibers (J. M.Davis and M. C. Jaurand, eds.), Springer-‘Verlag, Berlin-Heidelberg, NATO AS1 series, sub. H, Vol. 85, pp. 34’7-358. Fubini, B. (1998a) Ann. Hyg. 42:521. Fubini, B. (1998b)in The Surface Properties Silicas (A. P. Legrand,ed.),Wiley, Chichester, pp. 415-464. Fubini, B., Giarnello, E., and Volante, M.(1989) Inorg. Chim. Acta 163:187. Fubini, B., Giarnello, E., Volante, M,, and Bolis, V. (1990) Toxicol. Ind. 6.571. Fubini, B. (1994) in Cellulur and Effects ~ i n e r a l Synthetic Dusts and Fibres (J. M.Davis and M. C. Jaurand, eds.), Springer-Verlag, Berlin-Heidelberg, NATO AS1 series, sub. H, Vol. 85, pp. 347-358.

Cavenago, A., andUgliengo, P. (1992)

Faraday

Cavenago, A., and Volante, M. (199%) Scand. Work Environ. ~ e a l t / z , (suppl. 1):9. and Giamello, E. (1995b) Free Rad. Res. 23:593. Ong, T.,and Wallace, W. E. (1999) in E n v i r o n ~ e n t a l ~ u t a g e n Society # ~ tAnnual /~ Meeting. Giamello, E., Fubini, B., Volante, M., and Costa, D. (1990) ColEoids #5:155. Gil~llan,A. M.,Chu, A. J., Smart, D. A., and Roonet, S. A. (1983) Lipid Gilmour, P. S., Beswick, P. H., Brown, D. M., and Donaladson, K. (1995) ~arcinogenesi's 16:2973. and van Wyk, A. (1987). Lav. 78:124. P. and van Golde, L. M. G. (1985) Lung 163275. (1917) Trans. Inst. Min. Eng. 55:264. Harrison, J. C., Brower, S., Attfield, M. D., Doak, C. B., Keane, J., Grayson, R. L., and Wallace, W. E, (1997) Aerosol Sei. 28:689. ~ e m e n ~ aD. y ,R., Absher, M. P., Fubini, B., and Bolis, V. (1993)Arch. En~iron.~ e a / t h #8:343. Hill, C. A., Wallace, W. E., Keane, M. J., and Mike, P. S. (1995) Cell Bid. Toxicol. 11:119. Ho, C. and Hlady, V. (1995) inProteins at Interfbces F u n d a ~ ~ n t aand l s ~ppl~cations A. Horbett and L. J. Brash, eds.), ACS Symposium series 602, San Diego, CA, pp. 371-384. Hurley, J. F., Burns, J,, Copland, L., Dodgson, J., and Jacobsen, M.(1982) Brit. 39: 120. IARC Monographs on theEvaluation the CarcinogenicRiskofChemicals H ~ ~ m a(1997) ns Silica, Silicates, Dusts and Organic Fibers, Lyon, Vol. 68. Johnson, N. F. and Maples, K. R. (1994) in Cellular and ~ o l e c u l a r.Efects Mineral G. Davis and C. Jaurand, eds.), Springer-Verlag, Berlin, Dusts and Fibers Heidelberg, pp. 23Jonsson, U.,Ivarsson, Lundstrom, I., and Berghern, L. (1982) C o l l o i ~Interface Sci. 90:148. (1996) in Silica and Si/ica-Induced Lung Diseases: Current Concepts (V. Castranova, W. Wallace, and Vallyathan, eds.), CRG Press, Boca Raton, FL. Keane, Wallace, W., Seehra, M., Hill, C., Vallyathan, V., Raghootama, P., and Mike,P.(1990)in Proceedings I n t e r ~ a t ~ o n a /P n e ~ ~ o c o n i ~Cs iosn ~ ~ r e ~ c e , Pittsbur h 1988, U.S. Department of Health and Human Services DHHS -108 Part 1, U.S. GovernmentPrintingOffice, ~ashington, King, R. J. and Clements, J. A. (1972) Am. Plzysiol. 223:707. J., D. J., Gikas, E. C., and Clements, J. A.(1973) Am. P / ~ ~ ~ s i o l , 22#:788. Klockars, M., Hedenborg, M., and Vanhala, E. (1990) Arclziv. Emir. ~ e a l t h Kozin, F., Millsten, B., Mandel, G., and Mandel, N.(1982) Colloid I n ~ e ~ ~ aSei. ce 883326.

Kriegseis, W. and Scharmann A. (1982) Hyg. 26:625. Langer, A. M. (1978) 11:534. Le Bouffant, L., Daniel, H., and Martin, C. (1977) in Particles (W. H. Walton, ed.), Pergamon Press, Oxford, pp. 389-400. Le Bouffknt, L., Daniel, H., Martin, J. C., and Bruyerer, (1982) 26:625. ~ ~ ~ e ~ - D u s t LeBouffant, L., Daniel, H., and Martin, (1987)in (L. Le Bouffant, ed.), Colloque de 1’1NSERM pp. 481-492. Liu, X., Keane, M. J., Zhong, B. Z., Ong, and Wallace, W. E. (1996) 361539. Liu, X., Keane, M. Ong, Antonini, J. M,, and Wallace, W. E. (1997) 41 (suppl. 1):415. Liu, Keane M. J., Harrison, J. C., Cilento, E. Ong, and Wallace, W. E. (1998) 96:77. Mao, Y., Daniel, L. N., Knapton, A. D., Shi, and Saffiotti, U. (1995) Appl. Hdyg. l Marks, (1957) Br. J. Ind. Med. 14:81. Allison, A. C., and Harington, S. (1966) 210:259. olan, R. P., Langer, M,,Harington, J. S., Oster, G., and Selikoff, (1981) 26:503. Pandurangi, R. Seera, M. S., Razzaboni, L., and Bolsaitis, P. (1990) Envir. 86:327. and Vigliani, E. C. (1982) Am. Ind. Med. 3:13 i, B. L., Bolsaitis, P., Wallace, W. E.,and Health and Human Services D HS (NIOSH) Publ. 90-108 Part 1, U.S. Government Printing Office, Washington, D.C., pp. 21 5-230. VII Robock, K. and Bauer H. D. (1990) in Pittsburgh 1988, Department of Health and Human Services DHHS(~I0SH) Publ. 90-108 Part I, U.S. Government print in^ Office, Washington, D.C., pp. 208-283. Scheel, L. D., Smith, B., Van Riper, and Fleisher, E. (1954) 929. Schlipkoter, H. W. and Brockaus, A. (1960) Ditsch. ~ o c ~ e n s c h85:920. r. Shi, X., Mao, Y., Daniel, L. N., Saffiotti, U., Dalal, N. S., and Vallyathan, V. (1995) up. Hyg. 10:1138. and Goldsmith, D. F. (1995) Ind. 28:603. Summerton, J., Hoenig,Cooleyutler, M.,and Chvapil M. (1977) 26:123. Tourmann, L. and Kaufmann, R. (1994) Hyg. 38 (suppl 1):455. Vallyathan, V., Schwegler, D., Reasor, M., Stettler, L., and Green, F.H. Y. (1988) 32 (suppl. 1): 279. Vallyathan, V., Mega, J. F., Shi, X., and Dalal, N. (1992) Am. Respir. Bid. 6:404. Van C. (1994) Forces Aqueous Dekker, New York, pp. 336350.

Ila

Wallace, W. E., Vallyathan,V.,Keane, M. J., and Robinson,V.(1985) En~iron.Health 16:415. Wallace, W. E., Keane,M. J., Vallaythan, V., Hathaway, P., Regad, E. D., Castranova, V., and Green, F. H. Y. (1988) Ann. Hyg. 32 (suppl. l): 291. Wallace, W. E., Keane, J., Mike, P. S., Hill, C. A., and Vallyathan, V. (1989) in E’ecst of ~ i n e r a Dusts l on Cells, (B. Mossman and R. 0 . Begin, eds.), NATO AS1 Series Vol. H30, Springer-Verlag, Berlin, pp. 49-56. Wallace, W. E., Harrison, J., Keane, M. J., Bolsaitis, P., Eppelsheirner, D., Poston, J., and Page, J. (1990a) Ann. Hyg. 34:195. Wallace, W. E., Hill, C. A., Keane, M. J., Page, S. Bolsaitis, P., Razzaboni, B, L., Vallyathan, V., and Mike, P. (1990b) In Proceedings rnternational P ~ e ~ ~ o c Conference, o ~ ~ ~ s Pittsburgh ~ s 1988, Department of Health and Human Services DHHS(~1OSH)Publ. 90-108 Part l, U.S.GovernmentPrintingOffice, Washington, pp. 755-764. Wallace, W. E., Keane, M. J., Mike, P. S., Hill, C. A., Vallyathan, V., and Regad, E. D. n. 37:391. (1992) Toxicol. ~ 7 ~ v i r oHealth Wallace, W. E., Keane, J., Harrison, J. C., Stephens, J. W., Brower, P. S., Grayson, R. L., Vallyathan, and Attfield, M. D.(1994) in Cell~larand Eflects of NATO AS1 Series, Vol. H85(J. G. Davis ~ i n e ~and a l Synthetic Dusts and and M. C. Jaurand, eds.), Springer-Verlag, ~erlin-Heidelberg, 369-3’79. Walton, W. H., Dodgson, J., Hadden, G. G., and Jacobsen, M. (1971)in I ~ ~ a l e d Particles W , Vol. (W. H. Walton, ed.), Pergarnon Press, Oxford, pp. 669-689. Wiessner, J. H., Mandel, N. S., Sohnle, P. G.,Wasegawa, A., and f an del, G. S. (1990) Rev. Respir. Dis. 1.

Department of Material Sciences, Pacific Northwest National Laboratory, Richland, Washington

I. Introduction

665

11. Mesoporous Silica

666

Self-Assembly

667

IV. Direct Inorganic Functiol~alizationof MCM-41 A. Cocondensation B. Surfactantdisplacement C. ~ostcalcinationmetalization

VI.

669 669 67 1 672

Direct Organic functionalization of MCM-41 A. Cocondensation B. Surfactantextraction

674 674 676

MonolayerChemistry of MCM-41 A. State of the calcined surface B. Directsilanation C. Hydration of the calcined interface D. Drivingcondensation equilibria E. Fullydensemonolayercoverage

678 678 679 68 l 68 1 682

VII. Conclusions References

683 685

Silica has long history of service to humankind in the making of glass, abrasives, reinforced composite materials, sorbents, and catalysts. In these applications, the

interfacial chemistry of the silica surface, as well as the surface area provided by the particle morphology, are critical to the success or failure of the final product. Over the years, much effort has been expended in exploring the interfacial chemistry of silica, as well as manipulating particle size and morphology of the silica in an effort to enhance the surface-area-to-mass ratio. ecent advances in the areas of nanostructured ceramics and molecular self-assembly have provided thekey components to what is, perhaps, the ultimate refinementin catalyst support and sorbentdesign. This chapter will summarize the synthetic methods used to prepare interfacially functionalized mesoporous ceramics.

In 1992, scientists fromMobilTechnologyCenterpublishedalandmarkpaper describingthe synthesis and characterization of mesoporous silica [1,2]. In this work, a unique form of silica with a massively parallel pore structure, arranged in a hexagonal lattice was unveiled and name$ MCM-41. The pores in this honeycomb-like struct3re were approximately 30 in diameter and the ceramic walls were roughly thick (see Fig. 1). This highly porous morphology results in extremely high surface areas, typically 1000 m2/g or more. In addition, because of the parallel orientation of the pores and the high degree of order, all of the pores are open-ended, leaving all of the internal pore surface accessible to gas-phase or solution-borne chemicalspecies. These properties make MCM-41ideally suited for catalysis and sorbent applications. The origin of the template responsible for the formation of this intriguing superstructure is summarized in Fig. 2. Under well-established conditions, surfactant molecules self-assemble into long rod-shaped micelles. In a very specific portion of the oil- at er-surfactant phase diagram, these rod-shaped micelles aggregate into a hexagonal dense-packed phase. In the presence of silicate anions (or silicic

EhAron micrograph structure.

MCM-41 showing orderedhexagonalpore

Oil-water-surfactantphasediagram:theorigin used to template the synthesis of MCM-41.

of theordered structure

acid), the quaternary ammonium terminus of the surfactant molecule undergoes anionic metathesis with the silicate species from solution, forming a dense-pac~ed hexagonal ceramic precursorthat eventually precipitates out of solution. Calcining this mesostruct~redammonium silicate at 540°C results in the fusion of the aggregated silicic acid units and the combustion of the surfactant molecules, leaving the honeycombed ceramic superstructure. This basic strategy has since been found to be quite general. In various modifications it has been applied towards the synthesis of many varied ceramic materials with widely varying pore sizes and lattice-packing symmetries 131.

The spontaneous aggregation of molecules into an organized, ordered, and coherent matrix is referred to as “self-assembly.” Chemistry and biology are rife with examples of molecular self-assembly, ranging from the formation of micelles by simple surfactant molecules (as described above) to the elegant complexity of the double-stranded a-helix of DNA. From the soap we use to wash our hands to the genetic coding we pass on from generation to generation, molecular self-assembly plays a central role in both nature and society. Perhaps the best-studied example of this pervasive phenomenon is the area of self-assembled monolayers, which was unveiled in the modern era in a report by Sagiv in 1980 These are composedof a single layer of molecules in which the headgroup of the molecule interactswith an active interface andthependant hydrocarbonchains align withone another in parallel fashion. The interfaces upon which these can be built include solids (e.g., silica), liquids (e.g., mercury)

and the air-water interface. Of themanytypes of monolayers that have been preparedandcharacterized,undoubtedlythemostthoroughlystudied are the long-chain alkylthiols on gold (see Fig. [7]. Immersion of a clean gold surface into a solution of alkylthiol results in the adsorption of the thiol moiety to the gold surface, and eventual formationof a fully dense monolayer coveringvirtually all of the gold surface (small pinhole defects are commonly encountered). These monolayers are characterized by a very uniform film thickness and highly reproducible interfacial wettabilities, and typically have the hydrocarbon chains reclining at an angle of approximately 30". While the thiol monolayers are easily prepared and studied, they are limited in terms of their thermal and chemical stability. Monolayers constructed of organosilanes on silica offer a robust alternative [S]. Thesemonolayersarecovalently anchored and cross-linked, forming a solid foundation upon which to build the desired chemical interface. The recent literature offers several examples of studies that have synthetically elaborated terminally functionalized monolayers to create complex and useful chemical interfaces (an exampleof which isshown in Fig. 4) [S161. This ability to manipulate the chemicalidentity of a functionalized interface is key to the design and development of functionalized mesoporous materials. The increasing complexity of the monolayer interface has naturally led to the design and construction of elaborate monolayer superlattices 1171. The next step in the evolution of synthetic superlattice complexity is the construction of free-standing superlattices that can serve as hosts, templates, or foundations for even more elaborate hybrid materialdesigns 1 ringing all of these concepts together allows for the construction of complex assemblages, withintricate, repeating, and ordered molecular structure that canbe created easily and called upon to do highly specific tasks. Thus, the concepts and

Self-assembled monolayers of long-chain alkylthiols on gold.

An example of interfacial chemistry in silane-based self-assembled monolayers (see Ref. 9).

tools of supramolecular architecture [l91 have given birth to interfacially functionalized mesoporous ceramics.

The internal pore surfaces of mesoporous silica have been functionalized with wide variety of inorganic species by three fundamentally different methods. One of the most widely used methodologies involves the addition of an inorganic dopant (typically metal salt or alkoxide) to the original sol-gel mixture. This cocondensation route is, in fact, simply variation on the original Mobil recipe [1,2]. More recently, a clever method has been reported in which the cationic surfactant is replaced within the mesostructured greenbodyby a metal salt via cationic metathesis. A third, and very popular strategy, is to prepare the mesoporous silica using the traditional hydrothermal synthesis, and then simply incorporate the desired inorganic species postcalcination via condensation chemistry, proton exchange, or simple adsorption.

It has been shown that, by simply altering the solution composition of the initial sol-gel mixture, it is possible to influence the composition of the resulting mesoporous product [20]. While rather general, this approach is not completely universal, as somesystems fail to achievethemesostructuredmorphology, orthe mesostructure collapses upon removal of the surfactant template. This strategy particularly successfulwhen the primary structural component issilica and the

dopant species is a 3, 4, or 5 cation that can replace silicon in the oxide lattice and is added as a relatively minor constituent. When successful, this constitutes an expedient route for the incorporation of catalytically active species into the internal pore surfaces of mesoporous silica (see Fig. 5). A variety of different species have been incorporated into a mesoporous matrix by cocondensation. For example, boron has been incorporated using either boric acid or sodium borate as the boron source [21]. The B - ~ C ~ - materials 41 obtained via calcination were found to have surface areasof 900 m2/g, whereas the material obtained by ion-exchange removal of the surfactant template was found to have a surface area of 500 m2/g or less. This reduction in surface area was attributed to partial extractionof the boron during the ion-exchange step and partial collapseof the mesostructure. Detailed characterization revealed that the hydrolysis of the framework boron and its coordination state was highly dependent on the counterion, being stable when the counterionwas either sodium or ammonium cation, and unstable in the acid (protonated) form. Manyresearchershaveincorporatedaluminuminto their recipe formaking mesoporous ceramics. A recent study of the cocondensation of aluminum into mesoporous silica providesa detailed characterization of the interfacial active sites In this work, 2-D solid-state NMR was employed, using the ~ ~ T C O R pulse sequence, to examine the structure of both the as-synthesized and calcined mesoporous aluminosilicate materials. These 2-D HETCOR experiments were carried out in order to unambiguously determine the incorporation and structure of aluminum species within the inorganic mesophase framework, as well as the spatial proximity of thealuminum centers relative to thepore surface. These studies unambiguously established the presenceof aluminum in the polymeric oxidelattice of both the hydrothermally as-synthesized materials and the calcined aluminosilicate MCM-41 product. In the as-synthesized greenbody,it was shown (via dipoledipole coupling to the protons in the surfactant molecules) that both tetrahedral and octahedral aluminum atomswere present at the inorganic-or~anicinterface. In addition, their HETCOR m~asurements confirmed the presenceof both tetrahedral and octahedral aluminum in the framework of the calcined mesoporous aluminosilicate. Similarconclusions were reachedforthetetrahedralaluminum species found in similar alu~inosilicatematerials prepared under ambient conditions (no

B(OH)3

CTAC base

calcine

Inorganic cocondensation to prepare metal containing MCN-41.

octahedral aluminumwas observed in samples prepared under ambientconditions). Well-ordered gallium-containing mesoporous materials have also been prepared by cocondensation, and examined in catalysis studies (231. Thorough characterization of these gallosilicate materials revealed that thegallium was 4-coordinated and that part of the gallium was lost during calcination if the Si/Ga ratiofell below 30. Partial structural collapsewas observed in the Si/Ga 20 sample upon calcination, which can be attributed to the loss of gallium in this phase of the synthesis. The synthesis of Ga-MCM-41 was found to be highly sensitiveto the pHof the gelation mixture, with the optimum pH being 1 1 (at pH 1.S,1 lamellar phasewas formed, which collapsed upon removal of the surfactant template). Interestingly, no contraction of the rnesophaseoupon calcinationwas observed, which was attributed to the thick porewalls (14.4 and the suggestionthat these mesoporousgallosilicates were highly polymerized and ordered. Vanadium has also been entrained in a mesoporous silica matrix via cocondensation [24,25]. These studies were motivated to find an environmentally friendly, and yet selective, oxidation catalyst. In the first report [24], an optimized synthetic procedure is described, along with complete characterizationof V-MC rials with Si/V ratios ranging from 13 to 304. It was found that the vanadium is present in these materials as two tetrahedral 5) species and two square pyramidal VO( 2) species. That each of the two oxidation states was found in two environments suggests that one is bound within the silica framework and the other is on the pore surface. Through a series of spectroscopic and colorimetric cycles, it was shown that only those vanadium species present on the walls of the mesoporous structure were able to participate in catalytic cycles. These conclusions were supported by the second research group [25], which went on to demonstrate that the interfacial vanadium species are somewhat mobile.In addition, after calcination it was found that the oxidized vanadium species were in fact aggregated and these vanadium oligomers were nonuniformly distributed a thin film across the pore surface. In the examples outlined above, thecentral theme that of a mesoporous silica matrix used to support a dopant species in order to exploit its useful chemical properties. Generallyspeaking, if the dopant levelexceeds a certain threshold, then the mesostructure collapses due to an inadequate structural backbone. In a unique twist on this theme, Elder and coworkers have employed aZr02 dopant in orderto stabilize amesoporousTiOzmatrix that wouldhaveotherwise been unstable [26]. Since mesoporous Ti02 is unstable towards calcination (dueto facile crystallization, grain growth,and subsequent mesostructuralcollapse), it was necessary to stabilize the nanoscale anatasecrystallites in the intermediate mesostructure so that the morphology could be retained throughout the calcination stage of the synthesis. This wassuccessfully accomplished by adding 25 mol% Zr02 asa dopant. It was shown that the 25 anatase crystallites did not change in size during calcination. These are the smallest stable Ti02 crystallites reported to date.

Recently, two separate groups have reported their efforts involving the replacement of the cationic surfactant template within the mesotructured greenbody with tran-

sition metal cations. This strategy allowsfortheincorporation of thedesired metallic species before calcination, in the presence of silicic acid units with their easily exchangeable protons (instead of the functionally devastated postcalcination silica interface). a result, the incorporated metal is found in the form of a covalently bound metal silicate, lining the interior pore surfacesof the final mesoporous material (as depicted in Fig. 6). Calcination burns off the rest of the metal’s ligands and solidly anchors the metal center in the silica matrix. In the first of these reports, the researchers describe “planting” manganese(11) cations on the internal pore surfaces of MCM-41 throughcationic metathesis of the quaternary al~moniumsurfactant ion template, followedby calcination in air Following this simple procedure, they were able to prepare a range of manganese loadings in the final adduct simply by varying the initial manganese acetate concentrations. Calcination of the manganese adducts at 600°C for 6 h resulted in the manganese being strongly bound to the Mn-MCM”41 and resistant to aqueous leaching. The manganese cation was shown to be octahedrally coordinated and bound tightly to the silica framework (although it is not clear whether the Mn is bound within the framework or simply on the surface). More recently, cobalt ethylenediamine complexes have been employed in similar fashion [28]. This seems to be a very facile process as the displacementis complete in 20 min at 47°C. These cobalt complexes were steadfastly adsorbed and resisted leaching, even with repeated washing. However, theywere found to be sensitive to cation exchange in the presence of ammonium chloride (note that these cobalt ethylenediamine adducts were not calcined).

The literature is replete with numerous examples of metallic species being introduced into the pore structure of MCM-41 after the silica has been calcined to remove the surfactant template. A wide variety of different chemistries have been exploited to effect this metallic incorporation, including oxide-alkoxide condensation, acid-base chemistry, ring openingof strained molecules, ion exchange, hydrogen bonding, and simple adsorption. An excellent review nicely summari~esmuch of the effort in this area [29]. Many of these cases involve simple washing or soaking of the MCM-41 material in a solution of metal salt(s) or alkoxide(s), and allow adsorption or con-

Displacement of surfactants by metal cations.

densation to take place with the pore wall. The chemistry of these depositions is straightforward and will not be further discussed here. However, a few examples of somewhat more exotic chemical interactions will be briefly mentioned to illustrate some of the more novel applications of this synthetic route and the adducts that result from it. Neutral metal carbonyl complexes are typically relatively nonpolar species, typified by their solubility in organic solvents and their moderately high vapor pressures. In one interesting study Mn2(CO)rowas found to undergoreaction with the interfacial silanols of MCM-41, resulting in a tethered manganese species that could be further manipulated. Silanation of the silanols prevented this binding. Subsequent air oxidation for 2 h at 300°C resulted in a manganese-oxide-de~ivatized mesopore. This procedureresulted in the formation of ultrasmall l~anganesecontaining particles grafted to the pore surfaces and uniformly distributed throughout the mesostructure. A recent report details the attachmentof ferrocenophane 1 1”ferrocenediyl)dimethylsilane] via a ring opening of the strained cyclophane by the surface silanols of the MCM-41 silica surface [31]. Thernetallocenestructure was found to be retained in the product and the presence of covalent siloxane tethers were confirmed.Incorporation of theferrocenemoiety was foundto beefficient, with approximately 95% being bound in the mesoporous silica. In addition, these tethered ferrocenes were found to be kinetically quite accessible, as they underwent virtually instantaneous oxidation by iodine. A uniquevariation of the“aqueousmetal salt wash method” was recently employed to study the coordination chemistry of an adsorbed metal species. The tetrapyridine copper complex was prepared by first exposing MCM-41 to a solution of cupric ions, drying, and then exposing the Cu-MCM-41 to pyridine vapor at roomtemperature [32]. Thetetrapyridinecomplex was foundto be strongly adsorbed to the pore walls of MCM-41 silica. Studies using 2 - 0 proton and deuterium HYSCORE electron spin resonance spectroscopy revealed that the complex was rigid on the ESR timeframe and was somewhat distorted relative to its ground state conformation. The pyridine ligandwas found to be significantly flattened (i.e., more closely parallel with the planeof the complex) dueto adsorption to the adjacent silica surface, although some disorderwas apparent from the spectra. Polypyridine ligands, such as 2,2’-bipyridine (bpy) and phenanthroline (phen), are excellent ligands for transition metals, and they form strong, brightly colored complexes with thesecations. Their high affinity for transition metal cationsis due to the ideal chelating stereochemistryof the nitrogen atoms and then-acidity of the aromatic backbone, which sets the stage for n-backbonding between the metal center and the n* orbital of the arene. Examples of the inclusion of this lrind of complex into MCM-41 include Mn(bpy),[33] and Fe(phen)3 [34]. In the case of the Mn(bpy), addition [33], the initial manganese complex has only two bipyridine ligands, thereby leaving two coordination sites open for reaction with the silanols of the pore walls. This material not only demonstrated superior catalytic activity (relative to homogeneous catalysts), but also demonstrated good recyclability and stability to leaching. TheFe(phen)3adduct was also active in theoxidationof aromaticcompounds [34], butbeingcoordinativelysaturated,there is onlya

weak Coulombic interaction between the metal center and the silica, leaving these materials susceptible to pH changes, ion exchange, and leaching. More complex ligands, such as porphyrins, have also been incorporated into a mesoporous framework. In one recent study 1351, ~~e'~~-tetraphenylporphyri~l was incorporated intoseveral different mesoporous ceramicsand the photoionizationof the resultant adducts was examined. The MCM-41 silica-stabilized porphyrin radical cation was found to be long-lived and stable at 77 K, and decaying only 10% over 48 h at room temperature. The photoyieldof radical cation was correlated to the co~positionof the ceramic phase;a l u n l i n u ~incorporation reduced the photoyield, while titanium incorporation increased it. Loss of vibrational fine structure in theSoretbands of the adsorbed porphyrin's UV-Vis spectra was attributed to interaction of the porphyrin's n-system with the surface silanols in the pore walls. While the bipyridine, phellanthroline, and porphyrin complexes canbe thought of as a transition metal cation wrapped up in a stable, basically inert organic shell, resulting in a large coordinatively inert cation, the heteropolyacids can be thought of as just the opposite; an anion (typically phosphate or silicate) wrapped up in a stable, basically inert metal oxide shell, resulting in a large, nonbasic anion. This metal oxide cage mightbe predicted to sorb fizvorably to the silica walls of MCM41 materials. Researchers at Mobil were the first to recognize the value of these adducts in catalysis, and their patent describes the preparation and use of these materials [36]. Similarly, phosphomolybdicacid and silicotungstic acidadducts CM-41have been studied by several othergroupsas well 137-391. The ed materials haveall shown higheractivity than the solution-borne analogs. Finally, complex metalclusters have been incorporated into the pore channels of 41 withretention of the cluster structure.For example [40], a novel ~ ~ C ~ ( C O ) dianionic ~ * C l complex was prepared and incorporated into the porestructure o f MCM-41.Detailedcharacterization revealed that the cluster structure was maintained postdeposition. This adduct was then subjected to thermolytic decomposition to create bimetallic silver-ruthenium clusters, which were studied as hydrogenation catalysts.

Just as with the inorganic derivatization of M C -41 materials, organic functionality can be introducedduringthe sol-gelassemblyof themesophase, by ion exchange of theas-synthesizedgreenbody or by subsequent elaboration of the preformed mesoporous ceramic. Each of these strategies has its advantages, to be discussed below. Dueto thewealth of chemistry that has been elucidated, as well as some of the subtle requirements for effective coverage, the organosilane functionalization of calcined MCM-41 will be discussed separately.

One of the problems associated with silanation of mesoporous silica is that the finale of the traditional synthesis is calcination, typically at 540*C. This is incompatible with almost all organic functionality, and leaves the silica surface severely

dehydrated and almost totally devoid of surface silanols (see below). developed a clever method around this problem by incorporating fun silanes into the sol-gel mixture and then removing the surfactant template via acid extraction instead of calcination (depicted in Fig. 7) 141-431. Well-ordered, hexagonal mesostructures were observed at the lower loadings of functionalized silane (10-20%), but not at higher levels of incorporation [41]. It was suggested that an upper limit of 40 mol% of functionalized silane (or less) was imposed due to the need for adequate Q4 levels for a stable silica mesostructure. An impressive list of organic functionality has thus been incorporated into asilica mesophase, including amines, thiols, epoxides, alkenes, arenes, and heterocycles (although the degree of functional silane incorporation was not always a direct reflection of the original solution Composition, suggesting thepossibility of leaching during work-up)[42]. In addition, they used this strategy to make MCM-41 containing a pendant 2,4-dinitroaniline moiety, thereby constituting the first direct synthesis of a mesoporous silica with a tethered organic chromophore (431. Since the functionalized silane represents a de ct in the condensation chemistry of these cocondensed mesoporous materials, and their orientation is as yet undetermined, there may be some fraction of the functionalized termini that is kinetically unaccessible in these products. There is little or no driving force to align the functionalized molecules in any given direction, as the van der Waals interactions between the alkyl chains will be swamped out by the surfactant molecules. Thus, there is no guarantee of an exofacial orientation and it is possible that some fraction of the functionalized alkyl chains may be trapped within thesilica framework. The observation that cocondensation results in the formation of irregularly sized pores and less ordered coatings than monolayer methods lends credence to this conclusion [44]. Thiol-based materials, similar to those prepared by Mann, were prepared by Rhijn and coworkers[44]. Subsequent oxidationof the thiol was easily effectedwith hydrogen peroxide and resulted in the formation of 1.0-1,5*mmol/g of tethered sulfonic acid groups. Aseries ofdifferent mesoporous sulfonic acid derivatives were prepared and their catalytic activity compared. The material preparedby installing a mercaptopropyltrimethoxysilane monolayer on a calcined and -41 substrate was found to havethehighestfunctionalloading

Si(OEt)4

(RO)3Si-R

1) CTAC base

2) acid wash or surfactant extraction Alkyl-MCM-41 alkyl chain)

7 Organiccocondensationstrategy.

density and produce the highest degree of conversion and selectivity in FriedelCrafts alkylation chemistry. Cocolldensation has alsobeen used to prepare acid catalysts [45]. In this work, the organic component was terminated with a phenyl ring, which allowed for subsequent gas-phase sulfonation, affording the corresponding benzenesulfonic acid. This route provided material with a loading densityof approximately 0.14 mmol of acid sites per gramof material. The resulting mesoporous acid was found to be a shape- and size-selective catalyst for ketalization chemistry. Other groups have also employed this method to prepare other types of functionalized MCM-41 materials. Stein and coworkers [46]used cocondensation of tetraethylorthosilicate (TEOS) with vinyltriethoxysilane to make mesoporous silica with vinyl groups lining the pore walls. It was shown via subsequent bromination that most of these vinyl groups were accessible to solution phase species. However, a significant kinetic retardation relative to solution bromination of typical alkenes was observed,which was attributed torestricted diffusion into and outof the pores. Given the facile diffusion into mesoporous systems reported elsewhere (e.g., the replacement of surfactant molecu from a filled pore by cobalt ethylenediamine complexes [28], rapid reaction of MCM-41-bound ferrocene with iodine [31], metalization of MCM-41 porphyrins [47], this explanation may not be accurate. An alternate explanation might be that because of the small size of the vinyl group, it might be held tightly against the silica interface, or possibly even be partially buried. Since bromination of alkenes proceeds through cyclic a bromonium ion and requires ~ a c ~ s attack i ~ e of the bromide ion on this intermediate cationic species, anything that h.inders approach of the bromide ion to the backside of the bromoniumion will kinetically retard the reaction. Since thebromoniumion willbe formed on the kinetically most accessible face of the vinyl group, this leaves the kinetically least accessible face of the olefin as the reaction partner for the bromide ion. Since the vinyl group bound directly to the silica, it is rigidly held and has very little conformat~onalflexibility to facilitate these trajectory requirements. Richer and Mercier reported a useful modification of Mann’s method in the cocodensation of TEOS with MPTMS to make a thiol-modified MCM-41 1481. Theyreplacedthetraditionalquaternaryammoniumsurfactantswithnonionic polyethyleneoxide(PEO)surfactants, and carried outthe sol-gel chemistry under conditions of neutral pH. The use of the nonionic P E 0 surfactants eliminated the need for the final acid wash, and these surfactantswere readily extracted by Soxhlet extraction with ethanol. The researchers noted more disorder in their mesoporous products (relative to 100% TEOS MCM-41), as well as shorter lattice spacings and reduced surface area. However, they still observed significant thiol incorporation into their materials (up to 1.1 mmol/g). There was no determination of the kinetic accessibility of these thiols to solution phase species.

Recently, Pinnavaia has championed the approach of using long-chainalkyl amines as the surfactant template and removing them via acid extraction to avoid calcination in the preparation of MPTMS-derivatized mesoporous silica (Fig. 8) [49]. By side-stepping calcination, the silica surface remains highly hydroxylated, thereby

Surfactant extraction strategy.

allowing for ready incorporationof the functionalized silane. However, removal of surfactants via an acid wash was shown to result in a material with greater disorder than does calcination [49,50]. In addition, elemental analysis in co~junctiol~ with solid-state 29SiNMR reveals that half of the silicon atoms are hydroxylated, suggesting that perhaps the silicic acid units may not be fully condensed into a rigid silica backbone by this route. onet the less, they were able to incorporate as much as 1.5 mmol thiol per gram of sorbent material using this strategy, depending on pore diameter. In addition, all of the thiols were shown to be accetsible for binding mercury when the amine template contained 12 carbon atoms(27 A pore d i a ~ e t e r ) , but only 61% of the thiols wer? able to bind mercury when the template contained only eight carbon atoms (15 A pore). This reduced efficiencywas attributed to steric congestion induced by the grafted ligands within the smaller pore. Being a stepwise treatmentof a preformed mesostructure,there is verylittle chance that this synthetic approach will result in the functionalized chains ending up trapped within the silica framework (the possibility raised with the cocondensed ~aterials). clever functionalization scheme directly compared the incorporationof metal porphyrin complexes into MCM-41 via cocondensation, surfactant displacement, and surfactant extraction with acid and subsequent ion exchange [47]. With these large bulky molecules, cocondensation was not found to be overly effective, with only small amounts of the metal porphyrin Complexes being entrained within the pores. Both acid extraction and direct surfactant displacement were found to be effective methods for binding the porphyrin materials into the mesoporous matrix. The sharpness of the Soret bands indicated that, in this case, the n-system of the porphyrin molecule was notinteracting with the silica surface, which would not be expected dueto the presenceof the quaternary ammoniumsubstituents. Treatment of these M~M-41-supportedporphyrins with metal acetate solutions resulted in the facile formation of the corresponding metal porphyrin complexes, indicating that the pores were still oeen for diffusion of solution-borne species. Furthermore, dye molecules (about l 2 A in size) and their oxidation products were free to diffuse in and outof the porphyrinated pores,an importantconsideration in the designof an effective photocatalytic system. It is also important to note that the metal cations did not displace the quaternary-salt-anchored porphyrins from the pores during the metalization reactions.

The unique utility of mesoporous silica is its extremely high surfacearea. However, partialorincompletecoverage, inessence, wastesa 11 of this valuableinterface. effectively exploit the unique advantagesof install a fully dense self-assernbled monolayerea entire silica surface. irect silanation of calcined has been shown to result in very poor levels of silane incorporation (approximately 10% of full nxonolayer coverage) [49,51]. This can be directly attributed to the severe state of dehydration and the general lack of surface silanols in the as-calcined material. To understand the silanation chemistry of the calcined MCM-41 surface, we must first understand the state of the surface, and then recognize what must be done to make it receptive to the installation of organic functionality.

The calcined mesoporous silica surface is depleted of surface silanols relative to a “normal” amorphous silica surface (Fig. 9) 1521. The exact degree of desiccation seems to vary somewhat from laboratory to laboratory, and is probably related to the length of time the sample held at the final calcining temperature. The chemca interface is known to be dependent on the fraction. of hydroxylated silicon atoms on the to be anywhere from 4% to as much consensus seems to fall in the 2 0 4 0 % range The exact number is likely to be unil~portantas it iswidely recognized that not all of the hydroxyl groups are kinetica~lyaccessible to silanation. In fact, it has been s11own that only approximately 10% of the hydroxyls a available for silanation [SZ], in good agreement with other observations [49,51]. ne explanation that has been given for this observation is that most of the silano are adjacent to one another (this has been shown to be true 1541; for a more detailed picture, see also ef. and that hydrogen bonding between the silanols attenuates their 1Iucleop~ilicitysufficiently that they

State of the calcined

interface.

fail to react with the chlorosilane [52]. This explanation overlooks the fact that on any othersilica surface, hydrogen-bonded silanols have been routinely silanated for decades. Perhaps a more accurate vision has been presented recently. It suggested that these unreactive silanols are in fact within the silica walls or within kinetically inaccessible “cracks” in the walls In this research it was shown that 21 of the silanols are unreactive. Other place this fraction much higher, at approximately 65% [52], or 35-70% depending on silane Again, the exact numbers may not be that importanthere, other than to recognize that not all surface silanols are able to undergo silanation. Also, is it likely that these numberswill vary from one labto the next depending on the exact procedure used to make their silica foundation. The important conclusion hereis that calcined MCM-41 has notablyfewer surface silanols than most silica surfaces, and a significant number of those silanols present are kinetically inaccessible for silanation. Thisis the first major stumbling block to the formation of a fully dense self-assembled monolayer.

However, fully dense monolayer coverage maynot be needed, or even desirable for certain applications of functionalized mesoporous materials. Most notable in this sense would be catalysis. Thehighturnover at the catalyst centermakeshigh loading density a secondary consideration. However, diffusion in and out of the pore structure is critical to support the catalytic cycle and this diffusion may be hindered if the catalytic site is bulky (e.g., metal complex, porphyrin, etc.) and present in large numbers on the pore walls [47]. There are a number of examples to be found in the literature where silane coupling agents have been appended to the M C ~ - 4 1interface, almost always with the as-calcined material one that has been dried under vacuum. This a very simple and direct strategy to limit the surfkce coverage to modest levels. Many of these efforts have been summarized in an excellent review [29]. recent example of the direct silanation strategy was reported by Jaroniec and coworkers [573. In this work, avariety of simple silanes were appended toa predried sample of MCM-41 by refluxing it in excess chlorosilane and pyridine. Using this method, they were able to bind 0.76-1.96mm01of silane per gram of product, which corresponds to approximately surface coverage inthis case (relative to an ideal monolayer). Detailed analysis revealed that the resulting pore diameter varied mol~otonicallywith the steric bulk of the starting silane (see Fig. 10). In addition, the hydrophilic/hydrophobic character of the adduct couldbe adjusted by changing the identity of the appended silane. The Combination of these factors provides a powerful tool in the synthesis of tailor-made porous nanostructures. The previously mentioned synthesis of thiols used to make mesoporous acid catalysts provides an insightful comparison [44]. In this CommLlnication, the authors compare cocondensation, direct silanation and full monolayer formation as a means for making mesoporous acid catalysts. All three strategies were found to beeffective methodsforfunctionalizingthernesoporousmatrix. monolayercoverage was foundto give materialwiththe lowest area-to-mass ratio, smallest pore diameter, greatest pore uniformity and smallest pore volume.In addition, theacid catalyst derived from thefull monolayer material

Pore size dependence on silane

provided the highest degreeof reaction conversion and selectivity, all of which are consistent with the greatest degree of organosilane incorporation. Direct silanation also been used to introduce more complex ligands into mesoporous materials. For example,ethylenediamine-,diethylenetriamine- and silanes were grafted onto the pore walls of 40 MCM-41 These ligating interfaces were used to chelate cobalt (111) for electrochemicalandoxidation studies. Onceagain,onlypartialcoverage was obtained, which varied corresponding to the steric bulk of the silane. The assumption was made that only14% of the silicon atoms in theMCM-41contained accessible silanols. Factoring this assumption into the ratios obtained fromelemental analysis indicated that only 30-65% of the available silanols underwent silanation. Cyclic voltammetry revealed some interesting electroactivity. Inaddition, these cobalt complexeswere found to reversibly form a high-spin dioxygen adduct, suggesting that perhaps someof the cobalt is only chelated to two ethylenediamine ligands, or that one et~y~enediamine ligand thermally dissociates allowing formation of the dioxygen adduct. Similarly, an amine-terminated silane has been used to incorporate a ruthenium porphyrin complex for selective catalytic epoxidations 1581. The first step in this work was to install an aminopropylsilane interface on the pore walls of the silica, followed by binding the amine terminus to one of the axial co-ordination sites on the ruthenium complex. Using this approach, it was possible to incorporate as much as 0.8 mm01 ruthenium per gram of adduct. This material was found to be an effective epoxidation catalyst, producing high yields and turnover frequencies. ~pproximately5% of the ruthenium was found to leach out during the catalytic runs. The need for a basic catalyst entrained within a mesopore has been previously noted and pursued 1591. One way that this has been attained is to derivatize the pore wall with an alkyl silane terminated with. a primary chloride, and subsequent SN2 displacement of the chloride with piperidine 1601. Subsequent treatment with hexamethyldisilazane (HMDS) to “cap” theresidual silanols was found toimprove the

catalytic properties, presumably through the elimination of “inner salt” formation between the amine and silanol. These piperidine-substituted MCM-41 materials were effective catalysts for the formation of monoglyceride esters.

The second major stumbling blockto monolayer formation is that ~ ~ l c i n a t i oalso n drives off all physisorbedwater.This interfacial hydration is essential forthe hydrolysis and subsequent cross-linking of silane-based monolayers. Fortunately, this deficiency is easily remedied by any of a variety of methods. Initially, we felt that it was necessary to hydrolyze the interfacial siloxane (Si -0- Si) bridges to fully populate the interface with silanols. Thus, we took up the MC water and boiled it for 4 h, as this is known to hydroxylate other types of silica. After air drying for five days the silica still contained a large amount of water. Treatment of this material with a large excess of MPTMS in the usual manner resulted in approximately 25% surface coverage (relative to an ideal monolayer) [51]. If a similarly hydrated MCM-41 sample was taken up in toluene and had the surplus water removedvia azeotropic distillation, it was possible to increase surface coverage up to76%. Various attempts to improve this level of coverage W but all resulted in ’75-80% coverage. It was then learned that boiling in water did not,in fact, hydrolyze the siloxane bridges to form the su but rather this step simply resulted in a hydrated(wet) surface. This is supported by other observations [54]. Clearly, this interfacial hydration wassufficient for the hydrolysis and condensation chemistrynecessary for the majorityof self-assembled monolayerformation,and yet somehowinadequatetodrive this silanation to completion. Due to theextremely high surface area of the mesoporous substrate, itis possible to effect the desired level of interfacial hydration in a very simple and expedient fashion. The desired amount of water is determined (2-3 monolayers) and simply added to the silica-toluene slurry, and stirred for 1-2 h,allowingthewater to equilibrate and diffuse throughout the mesoporous matrix. This method is quick, easy, and provides approxi ely surface coverage when followed by routine 4-6 hperiod ofreflux). silanation (treatmentwith TMS in toluenewitha Contrary to one erroneous report [49], this hydration protocol requires only a single hydration step and a single s i l a ~ a t i ostep ~ and is quite easily carried out as a one-pot synthesis.

Evenusing

our optimizeddepositionprotocol,

silanation of

MPT~S was still limited to 80% of an ideal monolayer (an ide composed of approximately five silanes per nm2). This was trou limitatioll due to defects within the mesoporous silica? Was it due to steric compression at the inner radius (i.e., thiol terminus) of the monolayer? Was it due to competing pore blockage? Orwas it some unforeseen factor? Monolayer formation of thiols on gold is known to be an equilibrium process. Hydrolysis and condensation chemistry are also known to be equilibriu~processes. Thinking of silanation

as anequilibrium processwas the key to effecting 100% surface coverageof MCM41. Synthetic organic chemists havebeen exploiting Le Chatelier’s principle [61] by driving equilibrium processes through the removal of either the product or byproducts from the reaction as they are formed, thereby preventing the reaction from ever achieving true equilibrium conditions and forcing the reaction to go completely to one side of the equation. We chose totest this theorem by removing both the water and methanol from the silanation reaction mixture. Both of these materials form azeotropes with tolueneand their removal by distillation is straightforward (Fig. 11). The resulting monolayer-coated nlesoporous material indeed has lOOO/b fully dense monolayer coverage (see Fig. 13). Once again, this hydrationsilanation-distillation protocol is quite easily effected as a one-pot process, and requires no repetitive iterations.

By combining the optimized hydration procedure of simply adding two or three monolayers worth of water to the MC~-4l-toluene suspension, with the postdeposition distillation of methanol and water to drive the condensationequilibria to completion, we have been able to make fully dense (i.e., loo%, or about 5 silanes/ elf-assembled monolayers of MPTMS, aminopropyltrimet~oxysilane S), and (aminoethy1)aminopropyltrimetho~ysilane(ethylenediamine terminated; E ~ A - T M (see ~ ) Fig. 12). Using MCM-41 with a surface area of 900 m2/g, this optimized protocol results in monolayer-coated mesoporous materials containing 4.0 mm01 of functionalized silane per gram of final adduct. This class of heavy metalsorbentmaterialsprovidesunprecedentedloadingcapacityandbinding kinetics 1,621. By combining self-assembled monolayers with mesoporous silica, three generations of self-assembly have been embodied in a single supramolecular matrix; that of the micelle, that of the mesostructured ceramic precursor,and that of the monolayer interface. The union of these three generations ofself-assembly creates a highly ordered hybrid material that contains a chemically specific interface, extremely high surface area, and geometrically precise dimensions. S ~ represents ~ ~ the most highly ordered, highest capacity and most efficient heavy metal sorbent material ever made (Fig. 13). In fact, given the fact that SAMMS was designed and engineered entirely at the molecular level with very little wasted “bulk phase,” it is difficult to imagine how any substantive improvement could be made in capacity, efficiency, and order.

H20 OH

Anexample of drivingsiloxanecondensationequilibria removal of water.

removed azeotrope

by azeotropic

H 0 , l -Si-0-Si,

Ho,ki,oH Si-OH

I

-Si-0-Si-0-Si-0-Si-

I

A

Fully dense (100%)monolayerfornnation monolayer defects.

via azeotropicannealing of

Recent advances in thefield of derivatizing mesoporous ceramics haveresulted in a variety of means by which to functionalize these extremely high-surface-area materials, These strategies fall into three broad classes; cocondensation, surfactant displace~ent/extractioll and postcalcination derivatization. In cocondensation,the functiollalization reagent (either inorganic or organic) is included in the original “soup” of the MGM-41 sol-gel synthesis and gets incorporated into the mesoporous silica framework. This is a very quick and easy way of functionalizing a mesoporous interface, but may be amenable to only limited degrees of functionalization since, in this case, each functional site essentially is a defect in the silica lattice. In

Schematic of self-assernbled monolayers (S~~MS).

on mesoporous supports

surfactant displacement/extraction, the surfactant is removed from the mesostructured greenbody either by direct displacement with a cationic reagent, or by acid extraction(ion exchange) followed by functionalizationwitha silane coupling agent. This route avoids calcination, and thus has a receptive, hydroxyl-rich interface with which to bind the target species. Postcalcination derivatization involves reaction of the preformed MCM-41 materials with the desired functionalized molecule, either directly, or after some specified hydration protocol. These differing approaches each provide different levelsof functionalization, stabilityandorder.Cocondensation seems to be well-suited to bindingcertain transition metal species into the silica framework, as well as moderate incorporation of small organic molecules, but does not work well at all for large bulky molecules (e.g., porphyrins). Surfactant displacement is a very simple and direct method to quickly incorporate transition metal cationiccomplexes, or quaternized ammonium salts into a mesoporous matrix. While fast and easy, these adducts may be susceptible to subsequent leaching through ion exchange. However, their stability may be enhanced by calcination if the functional moiety has sufficient thermal stability. Surfactant extraction (i.e., acid washing) is a good way to remove the surfactant template while leaving a hydroxylated pore wall. This method leaves a somewhat disordered substrate andis limited to partial coverage and an incompletely fused silica backbone. Postcalcination silanation allows for systematically varied surfacecoverage, since thechemistcandictatethe degree of interfacial hydration due to the initial lack of surface silanols and physisorbed water in the as-calcined material.Inaddition,postcalcinationprovidesforamore highly ordered final product and is the only method to date that has demonstrated the ability to install a fully dense, functionalized self-assembled monolayer on a mesoporous ceramic. Some applications (e.g., catalysis) may notneed high levels of functionalization, and in fact high levels may be detrimental due to steric congestion of the reaction center, and/or diffusional limitations on turnover rate. These low-loading applications can be well served by all three of the synthetic strategies outlined above. However, some applications (e.g., environmental sorbent materials) require high degrees of surface functionalization and order for maximized performance. These applications are best served by the stepwise deposition of a fully dense selfassembled monolayer on a preformed mesoporous silica substrate. In the formation of fully dense monolayers on mesoporous silica, the level of surface hydration has been found to play a central role in the quality of the final coating. This level of interfacial hydration is very easily achieved through the addition of two or three monolayers’ worth of water at the beginning of the reaction and allowing it to disperseacrossthe silica surface.Alsoimportant is thedriving of thevarious silanation condensation equilibria via azeotropic removal alcohol and water. The forcing of these equilibria is easily accomplished by simply replacing the reflux condenser of the original reaction apparatus with a still head and distilling off the by-products of silanation. The functionalization mesoporous silica is a field still in its infancy. Many, many highly useful materials are likely to arise from efforts in this area in the coming years. Different functionalization levels are readily dictated and a wide

variety of functionality has been installed into this unique interface, and new catalyst systems are being designed daily. ~ u l t i p l egenerations of self-assembly are capable of creating exquisitely elaborate and highly ordered molecular superstructures and have created the highest degree of order in the final functionalized inteface, well unprecedented loading capacity in highly efficient environmental sorbent materials,

1. C. T. Kresge, M. E. Leonowicz, W. J. Roth, J. C. Vartuli, and 359:710 (1992). Beck, J. C.Vartuli, W. J. Roth, M. E. Leonowicz, C. T. Kresge, K. D. 2. J. Schmitt, C. T.-W.Chu, D. H. Olson, E. W. Sheppard, S. B.McCuXlen, J. B. Higgins, and J. L. Schlenker. J. Am. Chem. Soc. 114:10834 (1992). Zhao, P. Yang, Q. Huo, B. F. Chmelka, and G. D. Stucky. Curr. Op. Solid 3. State Mater. Sei. 3:ll 1 (1998). 4. J. Sagiv. J. Am. Chern. Soc. 102:92 (1980). Bigelow, D. L. Pickett, and W. A. Zisman. J. Colloid Interface Sci. 1:513 (1 946). 6. W. A. Zisman. Adv. Chem. Ser. 43:l (1964). 7. A. Ulman. Chem. Rev. 96:1533 (1996). to Organic Thin Films From ~ u n ~ ~ ~ i r - ~ l o d g e t t 8. A. Ulman, in An Introd~~ction ~ e ~ - ~ Academic ~ s e ~ Press, ~ l New ~ , York, 1991,pp. 245-269. 9. G. E. Fryxell, P. C. Rieke, L. L. Wood, M. H. Engelhard, R. E. Williford, G. L. Graff, A. A. Campbell,R. J. Wiacek, L. Lee,and A. Halverson. Langmuir 125064 (1996). Kato and Pac. Langmuir 14:2372 (1998). 10. Pan, D. G. Castner, and B. Ratner. Langmuir 14:3545 (1998). 11. E. Cleland, Jr., and C. L.Hussey. 12. R. C. Sabapathy, S. Bhattacharyya, Langmuir 14:3797 (1998). 13. K. Motesharei and D. C. Myles. J. Am. Chem. Soc. 120:7328 (1998). D. Tidwell, S. I. Ertel, B. D. Ratner, B. J. Tarasevich, S. Atre, and D. L. Allara. 14. Langmuir 13:3404 (1997). 15. D. A. Hutt and G. J. Leggett. Langmuir 13:2740 (1997). 16. M. V. Baker and J. D. Watling. Langmuir 13:2027 (1997). 17. Yitzchaik and T. J. Marks. Accts. Chem. Res. 29:197 (1996). 18. P. J. Stang and B. Olenyuk. Accts. Chem. Res. 30502 (1997). 19. Liu, A. Y. Kim, L. Q. Wang, B. J. Palmer, Y. L. Chen, P. Bruinsrna, Bunker, G. J. Exarhos, G. L. Graff, P. C. Rieke, G. E. Fryxwell, J. W. Virden, B. J. Tarasevich, and L. A. Chick. Adv. Colloid Interface Sci. 69:131 (1996). M. Antonelli and J. V. Ying. Curr. Op. Colloid Interface Sci. 1:523 (1996). 20. 21. D. Trong On, P. N. Joshi, and S. Kaliaguine. J. Phys. Chern. 100:6743 (1996). 22. M. T.Janicke, C. C. Landry, S. C. Christiansen, D. Kumar, G. D. Stucky, and B. F. Chmelka. J. Am.Chem. Soc. 120:6940 (1998). 23. C. F. Cheng, H. He, W. Zhou, J. Klinowski, J. A. S. Conclaves, and L. F. Gladden. J. Phvs. Chem. 100:390 (1996).

24. Z. Luan, J. Xu, H. He, J. Klinovvski, and L. Kevan. J. Phys.Chern.200:19595 (1996). 25. K. J. Chao, C. N. Wu, H. Chang, L. J. Lee, and S. Hu. J. Phys. Chern. 202:6341 (1997). 26. Elder, Y. Gao, J. Liu, D. E.McCready, and C. F. Windisch, Jr. Chern. Mater. 20:3140 (1998). 27. M.Yonemitsu, Y. Tanaka, and Iwarnoto. Chem. Mater. 9:2679 (1997). 28. R. Badiei and L. Bonneviot. Inorg. Chern. 37:4142 (1998). 29 Moller and T. Bein. Chem. Mater. 20:2950 (1998). 30. e, D. Gleeson, and S. C. Tsang. Chern. Cornrn. 951 (1996). 31. r, J. Barlow, J. Drevvitt, S. J. Heyes, and D. O'Hare. Chern. 32.

36.

ahr, ~artlnann,W. Bohlmann, and R. Bottcher. J. Phys. 1027752 (1998). Kim, W. Zhang, and T. J. Pinnavaia. Catal. Lett. 43149 (1997). Lett. 36:263 (1996). Kevan. Phys. Chern. 202:10455 (1997). Rav, and B. H. Rose, U.S. Patent 5,366,945 to

37

ikov, A. Sinnenna, R. J. Jansen, K. Parnin, and H. van Bekkurn.

Chern. 33. 34. 35.

38.

a, A. Sinnerna, H. W. Zandbergen, and H. van kkum. J. Mol. Catal. A: Chern. 224:287 (1996). 39. Chu, X. Uang, Y. Sham, X. Ye, and U.Wu. Catal. Lett. 42201 (1996). 40. Sheppard, Maschmeyer, B. F. G. Johnson, J. Thomas, G. Sankar, D. Ozkaya, W. R. D. Oldroyd, and R. G. Bell.Ang.Chem. Int. Eng.Ed. 41. Sims, and S. Mann. Chem. 1367 (1996). 42. Burkett, and S. Mann. Chern. Cornm. 1769 (1997). 43. C. E. Fowler, B. Lebeau, and S. Mann. Chern. Comrn. 1825 (1998). 44. Van Rhijn, D. E.DeVos,B. F. Sels, W. D. ossaert, and P. A. Jacobs. Chem. Comnn. 317 (1998). 45 46. 47 48. 49. er and T. Pinnavaia. Envir. Sci. Technol. 322749 (1998). 50. etroff, F. Scl-ruth,and G. D. Stucky. Nature 368:317 (1994). 51. X. Feng, G. E. Fryxell, L. Q. Wang, A. U.Kim, J. Liu, and K. Kemner. Science 6:923 (1997). 52. S. Zhao, G. Q. Lu, A. K. Whittaker, G. J. Millar, and H. Y. Zhu. J. Phys. 53. Y. Long, T. Xu, U. Sun, and W. Dong. Langmuir 146173 (1998). 54. A. Cauvel, Brunel, F. Di Renzo, E. Garrone, and Fubini. Langmuir 23:2773 (1997).

57. 58. 59. 60. 61.

J. Chen, Q. Li, R. Xu, and F. Xiao. Ang. Chem. Int. Eng. Ed. 343694 (1995). F. Diaz, IS. J. Balkus, Jr., F. Bedioui, Kurshev, and L. Kevan. Chem.Mater. 997). C. P. Jaroniec, M. ISruk, M. Jaroniec, and A. Sayari. J. Phys. Chem. B 102303 998). C. Liu, W. Y. Yu, S. G. Li, and C. M. Che. J. Org. Chem. 63:7364 (1998). K. R. Kloetstra, J. van den Broek, and H. van Beekum. Catal. Lett. 47235 (1997). A. Cauvel, G. Renard, and D. Brunel. J. Org. Chem. 62:749 (1997). L. Brown and H. E. LeMay, Jr., C ~ e ~ i ~thet Central r ~ , Science, Prentice-Hall, Englewood Cliffs, NJ, 1977, pp. 412-415. J. Liu, X. Feng, G. E. Fryxell, L. Q. Wang, A. Y. Kim, and M. Gong. Adv. Mat. 10:161 (1998).

This Page Intentionally Left Blank

Department of Physics, University of Helsinki, Helsinki, Finland

I.

690

11. Aerosol Silica in the Atmosphere A. Particles of natural origin B. Particles of anthropogenic origin C. Particles of extraterrestrial origin

693 694 694 695

111. Morphology Surface andProperties

of Fumed Silica 695

IV. Silica-Water Interface A. Similarity between water and liquid 698 silica Interaction between water and silica surfaces C. Effect of acids on thewatervaporuptake by fumed silica 706 V, Large Surface-to-volul~e Ratio Problems A.description General problem of the Large surface-to-volume ratio problems in the ice of supercooled nucleation from water

697 699 709 710

classical theory 714

VI. Silica-Aqueous Solution Systems at Low Temperature A. Freezing of adsorbed water pure on717 silica Heterogeneous mechanism ofice formation on partly 726surface hydrophobic silica of

716

VII. Implication for the Formation of Polar Stratospheric Clouds and A. ozone Stratospheric Freezing behavior of aqueous solutions adsorbed on fumed silica erethe C.forImplication

728 728

s

730 736

an D. Possible role of the silica surface in the photodecomposition of chloro~uorocarbons troposphere the in VIII.

737 738

References

739

This chapter willbe devoted to discussing how the unique surface properties of fumed silica can be used for obtaining aqueous systems with a large surface-tovolume ratio. Thermodynamic properties and phase transformations in such systems differ from thoseof bulk material and,therefore, their study very important for fundamental and many appliedinterests, in particular, for cloud physics, since in clouds the liquid and solid phases of water are highly dispersed. On the other hand, in the atmosphere, the silicon dioxide (Si02) surfacecan be regardedas representative of many natural dusts, products of combustion processes, and particles formed during evaporation of meteorites. In the atmosphere, such aerosol particles can serve as cloud condensation andice-freezing nuclei. The specific interaction of water with the surface of fumed silica,which can be considered as a laboratory counterpart of aerosol silica, allows one to obtain the populations of pure water and solution microdroplets which canserve as a model for the studyof freezing and melting processes taking place in the atmosphere. The knowledge in this field is also very useful for better understanding the formation mechanism of polar stratospheric clouds which are supposed to play a crucial role in stratospheric ozone depletion. For the background on the problems discussed in the chapter, the reader may wish to refer to the material covered in Refs. 10, 12, 27-29, 41, 62, and 68 (Vols. 5 and 7).

In the earth’s crust, silicon (Si) constitutes about 27.6% of the mass and, therefore, its different polymorphs (quartz, cristobalite, tridymite, etc.) play an important role in the evolution of diverse earth processes [l]. On the other hand, since in the exposed earth’s crust quartz alone is estimated to constitute about20% by volume [2] and mineral silica is a widespread constituentin fossil fuel, the particulate matter containing siliconis also widely present in theatmosphere. For instance, the amount of natural dust alone, which consists mainly of a mixture of silicates, such as clay minerals, feldspars, and quartz emitted annually into the atmosphere due towind erosion in aridand semiarid areas, is estimated to be from 1000 to 3000 Tg [3,4]. In the troposphere,particles containing large amounts of silicon are widely present in all size ranges, from ultrafine 0.06 pm) to coarse particles 1 pm) [5]. In the Arctic high troposphere and lower stratosphere, silicon together with sulfur (S) and chlorine (Cl) were found tobe the most abundant elements measured in size-separated aerosol samples [5-81. In this context, it is reasonable to expect that aerosol particles containing silicon in large amounts can play an appreciable role also in different physicochemical processes taking placein the atmosphere.

Enhanced amounts of particulate matter in the atmosphereresult in a variety of environmental problems ranging from reduced visibility through urban air to climate change l]. Aerosol particles can influence the climateof the earth in two main ways: (1) directly by changing the flux of solar radiation reaching theearth’s surface by absorbing and scattering the incoming radiation, thereby affecting the energy balance of the earth; (2) indirectly serving as cloud condensation and icefreezing nuclei and playing in such away a significant role in the processesof cloud formation and precipitation (10-121. In the highlatitudes a large portion of the atmosphericprecipitation is due to the formation of ice, which strongly influences the release of snow, rain, and hail 131. In the caseof absence of any foreign particles in supercooled droplets theice phase is formed homogeneously.But this mechanism is rare and usuallyice is formed due to a heterogeneous mechanism on the suspendedparticles serving as ice nuclei. Ice crystals also impact the radiative and physicochemical processes within the clouds. For instance, cirrus clouds, which cover about 35% of the earth’s surface, have an important influence on climate through their effect on the radiation budget [14]. The effect on radiative transfer depends to a large extent on size and form of ice crystals, which are closely linked to thenucleationmechanism and nucleating efficiency of ice nuclei. Aerosol silica and particles rich in silicon can serve as centers for condensation of cloud droplets and the formation of the ice phase both in the troposphere and stratosphere. Indeed, it was found long ago that particles of many compounds containing silicon (clays, sand, soil, etc.) are effective in the initiating ofice in the atmosphere [12]. In the upper troposphere and lower stratosphere, the crustal silicon-dominated particles were found to be among the most common types of ice nuclei [15]. Laboratory experiments showed that heterogeneous nucleation of ice on some modified surfaces of precipitated silica powders(dehydroxylated by heating at elevatedtemperatures) waseffective(16,171.Silicaceous materials, which are relatively cheap and ecologically acceptable materials have been considered as a possible substitute for the expensive silver iodide (AgI) in cloud seeding [l8,19]. Besides the suspended particulate matter, in the atmosphere a large number of gaseous pollutant species are present. Atmospheric acids such as HN03, H2S04 are involved in a number of environmental problems, ranging from the ecosystem acidification to stratospheric ozone depletion. Interaction between aerosol particles and these strong acids can modify surface properties of particles that, in turn,can influence theheterogeneousmechanisms of cloudformation. Atmospheric acids can also strongly affect the size of cloud droplets, whichdetermines the optical properties of clouds, and thereby theradiative transfer in a cloudy atmosphere at both solar and terrestrial wavelengths. For instance, it was shown that an enhanced concentration of nitric acid vapor decreases the mean sizeof cloud droplets, whichcan be stable even at relative humidity less than 100% [20,21]. Sinceaerosol silica represents a significant fraction of the particulate matter it may scavenge acids during transport through the atmosphere. A practical and vitally important problem, which is strongly connected with the formation of theice phase and presence of acids in the atmosphere, that concern-

ing the formation of polar stratospheric clouds Now it is well established that the heterogeneous chemical reactions which take place on/in the particles are central to the process of chlorine activation and subsequent depletion of the stratospheric ozone [22,23]. Unfortunately, the formation mechanism and phase state of are not basically understood and remain still a topic of considerable debate [24]. Knowledge of thephasestate is important, since thecharacter of heterogeneous reactions depends on whether the cloud particles are liquid, solid, amorphous,orsolidembeddedinaliquid-likelayer.Thephasestateinturn depends on the chemical composition and size of the particles. odel ling of the PSCs’ onset (and clouds in general) is complicated in that the responses to supercooling and the physicochemical properties(surfacetension, enthalpy of fusion, density, specific heat, etc.) of very small aqueous systems are not fully understood even in the case of pure water, let alone aqueous solution droplets 125,261. Since cloud particles represent thehighly dispersed liquid andsolid phases of pure water and aqueous solutions, the phase transitions in them can differ from those of bulk liquids. Therefore, for better understandingof the mechanisms of cloud formation it is necessary to improve our knowledge about the aqueous systems with a large surface-to-volume ratio. In other words, an answer should be found to the question of whether intensive thermodynamic parameters characterizing pure water and aqueous solutions change with size or not. Knowledge of this could help to resolve the so-called capillarity approsimation problemin the classicalnucleationtheory, which, at present,remainsthemaintoolfortheoretical simulations of theformation of the ice and liquidphasesintheatmosphere 12,2?-29]. That theproperties ofvery smallmetalparticlesdifferfromthose of bulk material has been reported in the literature [30-321. ~sperimentaland theoretical studies have shown that depression of the melting temperatures, decreasing of the enthalpy of fusion and existence of a quasiliquid layer are the main characteristic features of particles with a large surface-to-volu~e ratio.But in the case of pure water and aqueous solution microdroplets the situation is totally different. Since microscopic volume of water easily undergoeslargesupercooling,theoretical (molecular dynamic) simulations run up against the lack of fundamental understanding of the structure and the low-temperature behaviorof water. In the caseof solutions the situation is complicated even more by the presence of the solute ions. The esperi~entalinvestigation of very small aqueous dropletsmeets the problemof measuringtheheatevolved/consumedduringthephasetransformations in the separate particles, which is too miniscule to be detected using the existing techniques. Therefore, a population of relatively “compacted” finely divided aqueous droplets, the totalweight of which should be sufficient to detect the heat participating in the phase transitions, should be used [25]. Unfortunately, it is verydifficult to obtain a stable population of very small pure wateror aqueous solution droplets.In the unlimited vapor phase, or formed on a substrate, such a system would be unstable due to the Kelvin effect. The aim of this chapter is to show that the morphology and unique physicochemical surface properties of fumed silica allow one to circumvent the described above obstaclesand create a stableor quasistable population of microdroplets. The

very small aqueous systems obtained can serve as a model for investigation of phase transitions taking place in the atmosphere and for resolving the large surface-tovolume ratio problems that are important both for applied and fundamental interests. Due to the fact that fine aerosol silica particles from evaporating meteorites and combustion processes and fumed silica are both of high-temperature origin, their size, chemical composition, and physicochemical surface properties are similar. Therefore, fumed silica particles can be considered as a laboratory counterpart of aerosol silica. In this chapter attention will also be given to the study of the possible role of aerosol silica in the heterogeneous freezing of the background stratospheric sulfuric acidaerosols and nitric hydrates in the stratosphere. The possible removal mechanism of the atmospheric pollutant gases by aerosoil silica particles is discussed.

In the atmosphere, the concentration of solid aerosol particles varies greatly with time and location, and depends strongly on the proximity of sources, the rate of emission, the meteorological parameters which affect the vertical and horizontal distributions, and theefficiency of the various removal mechanisms 112,331. Aerosol particles are of different sizes and chemicalcompositions.Coarse particles are influenced by gravity and, therefore, have a short airborne lifetime. They do not ascend to high altitudes and do not travel large distances. But the submicrometer, fine aerosol particles pm) are not appreciably influenced by gravity and have long atmospheric residencelifetimes. Therefore, they cantravel very long distances around the globe and can reach even the stratosphere [7,15]. Elemental composition of the stratospheric aerosolsis very broad Almost two-dozenelementshave been determinedwithproton-inducedx-rayemission (PIXIE) [34]. Earlier it was assumed that the volcanic activity constitutes a major source of particulate matter in the upper troposphere and lower stratosphere. The presence of small particles containing high percentages of minerals characteristic of the earth’s crust suggested their crustal origin and served as a corroborationof this assumption [7,15,35]. owever, recent studies have shown that besides the volcanic activity, the desert storms in theSaharaand in Asia,extraterrestrialmaterial (meteorites), and high-~ying aircraft, together with industrial activity, also play an appreciable role in the increasingof the total stratospheric aerosol concentration [5,15,35-371. Large fractions offine aerosol particles are silica particles, which originate from combustion processes and evaporating meteorites. In the stratosphericaerosol layer between 15-20 km,wherethe intensive ozonedepletion takes place, theconcentration of silicon wasfoundto be about 0.1 1 pg/ m3. The most abundallt compound in the layer, sulfate, was found to be 6.83 pg/m3 Morerecentmeasurementsperformed in theArcticshowed that from high troposphere to low stratosphere, the concentration of silicon was even about three timeslarger than sulfur and comparable with chlorine [5]. Sincein the stratosphere fine aerosol particles have a residence timeof the order of several years aerosol silica and particles rich in silicon can play a considerable role in the~hysicochemical processesoccurringduringtheformation of high-altitude

clouds (cirrus and PSCs) serving as cloud condensation and ice-freezing nuclei. Aerosol silica particles enter the atmosphere due to the following main sources.

These are windblown dispersed soil particles from arid regions, particles resulting from natural weathering and dispersion of average rock materials, and particles formed during eruptions of volcanos. The elemental composition of particles of natural origin resembles the average crustal materials with the AljSi ratio equals approximately 0.3. (Si and A1 togetherwith oxygen constitutethe three most abundant elements the earth’s crust 1381.) Though the particles from the natural sources fall mainly in thecoarse size range and, therefore, precipitate near the sources, small but nonnegligible amounts of the dust particles, presumably due to wind erosion, occur in the smallest ultrafine size range, i.e., 0.25 pm. Therefore, such crustal particles can travel very long distances; for instance, an abundance of resuspended dust particles, found in the Arctic winter atmosphere had originated from the Asiandesert storms 1391. Crustal particles were found to be abundant even in the stratosphere (7,151.

Silica particles emitted from combustion of carbonaceous fuels have been foundto be present in all size ranges, with a maximum mass concentrationat 0.25 p m [39]. These particles are formed by condensation of the vaporized mineral silica during the col~bustionof high-rank coals (mainly bituminoLls) during power generation silica particles are also formed during incomplete combustion of diesel he characteristic feature of these particles is that they have an AljSi ratio much smaller (typically below 0.07) than the value of 0.3 expected for earth’s crust and soil mineral matter [5,39,40]. can be directly vaporized at extremely high temperatures t when a reducing agent [usually carbon (C)] is present the much lower temperatures, i.e., atabout 1000-1500°C. uring coal combustion, inclusions of mineral silica in the parent coal ash are reduced by carbon to relatively high-volatile silicon monoxide (SiO) according to the reaction (401 C

Si0 (g)

CO (g)

This chemical conversion of §io2 to the reduced-state species Si0 in the burning char particle strongly augments vaporizationof silica. The vaporization dependson the nature and distribution of other mineral inclusions (i.e., depends on type of coal) becausereactions between acidic and basic constituents can influence the activity of thevaporizingcompounds [42]. The alkaline metals,therefractory oxides of which are dominant in the low-rank coals, suppress silica vaporization, probably by reacting with §io2 to formsilicates in which the activity coefficient of the silica is reduced. In the high-rank coals, where the amount of alkaline metal is smaller, such reactions cannot occurreadily and the large activity of §ioz increases its vaporization [42]. The process of vaporization is followed by the subsequent

ati

recondensation of high-temperature vapor. The gas-phase reoxidationof Si0 back to Si02 occurs when Si0 vapor diffuses into the oxidizing gas atmosphere sufficiently far from the surfaceof the burning char particle. Since the oxidation of Si0 vapor is extremely exothermic(4 lo5 J/mol),the oxidation reaction is very rapid, which results in a large supersaturation of the refractory oxide, leading to nucleation of a very large concentration of nanometer-scale globular amorphous silica particles. Measurements by cascade impactor have shown the extremely narrow size distribution of particles with largest amounts ofSi in the ultrafine particle size range, i.e., smaller than 0.25 p m Since thenumberconcentration of the flame-generated particles is very large, on the order of 1014 particles/m3, collisions between the particles result in their subsequent growth by coagulation [43]. This mechanism leads to the formation of aggregates of particles of different size and morphology.

These particles are formed by recondensation of ablated vapor from evaporating meteorites entering the atmosphere of earth. Aglobal average annual mass influx of meteoritic objects smaller than 1 tonneis about 16,000 metric tonnes (44 tonnes per day) [44]. Approximately70% of the total incoming extraterrestrial material comp-letely vaporized during the entry into the atmosphere producing a global annual m i n i ~ u minput of about 2000 tonnes of meteoriticvapor at altitudes between 80 and l10 km 1451. The large number of “smoke particles” with initial diameters between 0.2 and l0 nrn results from spontaneous recondensation of the ablated vapor,which then accumulatesin the mesosphereand stratosphere. During transport in the stratosphere these particles can coagulate to form aggregates of part-icles, with subsequent sedimentation to lower altitudes. It is assumed that in the stratosphere these ultrafine smoke particles mainly consist of the oxides of the refractive elements iron and silicon, i.e., Fe304 and Si02[46]. In the stratosphere the number densityof smoke articles was estimated to be about 100 per cm3 with a total surface area up to 5 per cm3. The surface area density of smoke particles is comparable to the surface area of the background stratospheric sulfuric acid aerosols, the sizeofwhichis approximately 0.07 [47] and numberdensity typically 5-20 per cm3. Anthropogenic and extraterrestrial sources contribute the main fraction of ultrafine and fine silica particles in the stratosphere. Larger silica particles can be both of natural and anthropogenic origin and they are present basically in the low troposphere.

Particles of fumed silica powder can be considered as a laboratory counterpart of the atmospheric aerosoil silica originating from the last two sources described above.Indeed,theyhavethesamechemicalcomposition, and their size and morphologyare similar. Furthermore, since fumed silica and aerosol silica

particles areboth of high-temperature origin, their physicochemicalsurface propertiescan be similar too. These similarities allow one to use fumed silica, the physicochemical surface properties ofwhich have been studied most intensively, as a proxy for aerosol silica in laboratory experiments. A short review of the lnorphology and surface properties of fumed silica is given below. For more materialonthe subject thereadercan refer, for instance, to work by Barthel et al. [48]. Fumed silica, otherwise referred toas “silica smoke,” is avoluminous lowdensity snow-white powderwhich consists of aggregates and agglomerates of finely divided silica particles. It is characterized by highpurity and low energywith respect to water adsorption. Fumed silicais usually produced by the hydrolysis ofsilicon tetrachloride inan oxygen-hydrogenflame at elevatedtemperature [17,41,48]. The globular primary particles with sizeof about 10 nm are formed by collisions and coalescence of protoparticles of only 1-2 nm in size, which are formed by nucleation from the vapor phase at high temperatures. As the temperature lowers, the collisions lead to partial fusion of the primary particles, which results in the formation of the stable chainlike highly branched aggregates with size about 100 nm. As temperature lowers further, the collisions between the aggregates, in turn, lead to the agglomeration of aggregates, i.e., the f o r ~ a t i o nof the three-dimensional loose networks where the aggregates are linked together by physicochemical interaction between particle surfaces. The agglornerates are extremely weak,so they are very fragile and can easily be disrupted as the powder is subject to mechanical action. The size of the agglomerates determined with laser diffraction was found to be larger than 5 p m [48] and can reach a few tenths of a micron long 1491. The physicoche~icalsurface properties of fumed silica depend on the technological process [16,41], The siloxane Si) surface consists mainly of oxygen atoms,eachbondedtoadjacentsurface silicon atoms.Usuallyasmall fraction of the surface is hydrophilic, i.e., covered with surface hydroxyl groups, Approxilnately every other surface silicon atom bears hydroxyl groups. The silanol number (number of silanol groups SiOH per nm2) of fumed silica is about 2 116,503, whereas for fully hydroxylated faces of sols, gels, or precipitated silica, where each surface Si atom has groups, the silanol number is app~oximately5 [51,52] or even larger [16], Since silica particles have amorphous st~ucture,silica atomsarenot in regulargeometrical ~rrallgenlent.Therefore, hydroxyl groups attached to the surface silicon atoms are not in uniform arrangement butrather statistically distributed onthehydrophobicsiloxanesurface ue to the small silanol number the distance between SiOH on the surface of fumed silica is sufficiently large in order to form hydrogen-bonded silanol clusters. Therefore, a large fraction of SiOH groups are isolated 1171. everth he less, part of the hydroxyl groups can be hydrogen bonded. The number in the hydrogen-bonded hydroxyl clusters on the unlnodified and modified surface of Aerosil200 (Degussa) hasbeen determined from the adsorption of GO molecules [543, It was found that the average number of hydroxyl groups in the such clusters (chains) was very small and represented a value of about 2.4 for silanol numbers between 1.2 and 2.1. The length of the chains increased up to 3.8

hydroxyls as the degree of rehydro~ylationof siloxane surfwe has been increased by the autoclave method. There remains uncertainty concerning the surface density of hydroxyl groups, i.e., where concentration of -OH is larger on the concave surface (points where primary silica particles touch each otherin the aggregates)or on theconvex surface (with positive curvature) of primary particles. The knowledge of this is important for better understanding of the structure of adsorbed pure water and aqueous solutions on partly hydrophilic surface of fumed silica (see below). There are two opposite points of view. First it has been supposed that the density of the surface hydroxyls is larger on surfaces with negative curvature [25,4l]. This viewpoint is based on the assumption that negative curvature makes the isolated hydroxyls closer together, which then can become hydrogen bonded. The second viewpoint has been proposed by Balard who maintains that the concentration of the surface -OH is larger on surfaceswith positive curvature.Thisviewpoint is based on the assumption that the heat evolved during collision and partial fusion of primary particles prevents the formation of surface hydroxyl groupsat near the points of connection of particles. Another question, which remains to be answered, is whether primary particles have microporous structure or they are free of micropores, i.e., exhibit a smooth surface. Unfortunately, there is no agreement on the subject. Some authors suppose that the primary particles have smooth surfaces The absenceof hysteresis in the desorption branchesof argon and nitrogen isothermsseems to corroborate this point of view [16]. On the other hand, thereexists opinion that primary particles of pyrogenic silicas which have been condensed from aflame contain micropores [41]. It has been shown by infrared studies that there is a small but significant volume of micropores ranging to 0.01 cm3/g which hold water strongly. Such a reasoning can serve in favor of this viewpoint: the diameter of the primary particles is usually in the size range 10-20 nm. Packing of protoparticles, of which the primary particles are formed, can be different depending on the technologicalprocess. In the case of close packing, the size of the micropores would be smaller than in the case when packing is more loose. Unfortunately, itis difficult to determine not only thesize of the micropores but even their presence. Indeed, the diameter may be too small, rendering the micropores imperviousto nitrogen and argon adsorption[41]. On the other hand, the microporesmay be impervious also to water molecules becauseof the absence of surPdce hydroxyls if the hypothesis of Balard described above is correct.

As water is the most abundant liquid on earth, so different polymorphs of silicon dioxide, SiOz, are among the most abundant minerals in the earth’s crust. starting a discussion of the silica-water interface itself it consider some similarities between water and liquid silica. can be considered as the progenitors of the two most important natural classes of liquid solutions, the thermodynamic and transport properties of which have determined in considerable degree the picture of the organic and inorganic natureof the

present world. The properties of these solutions in turn are strongly dependent on the thermodynamic and dynamic properties of the pure solvents themselves [57]. Such a consideration would beuseful not only for better understanding of the silica-water system. It can show that the problem may also represent a significant fundamental interest.

remarkable degree of similarity, observed experimentally and calculated from molecular dynamic simulations, exists between water and liquid silica [57-601. The main featureis the similarity in the structure,which determines the unusual properties of these unique liquids. Their structures are similar in the sense that they both consist of large oxygen atoms with the smaller hydrogen and silicon atoms in the interstices [41]. Therefore,themolecularvolume of water and silicais mainly determined by oxygen atoms which are more-or-less close packed with a packing density depending on the temperature. The structureof silica isdetermined by the way in which large oxygen atoms are co-ordinated around smaller, more electropositive silicon atoms. Silicon and oxygen atoms have a radii ratio which permits a co-ordination number equal to four, i.e., silicon exists in tetrahedral co-ordination with four oxygen atoms [38]. Liquid and amorphous silicas have short-range ordering whichis described in terms of tridymite and cristobalite structuralmodels.Purewaterhas also atetrahedral crystalline structure which arises fromthearrangement of molecules in shortrange ordering due to hydrogen bonds. The structure of water is often described by the so-called ““mixture” model [61], in the framework of which at temperatures above O”C, water structure represents the simultaneous existence of two crystalline or pseudocrystallineforms.One of themrepresentsthearrangement of water molecules in crystalline clusters resembling an open tridymite-like structure, which are distributed in a quartz-like lattice incorporating all theremaining watermolecules. Below 0°C and withincreasingsupercoolingthestructure of water is altered and becomes moreice-like, i.e., the molecular arrangementbecomes moreopen and acquiresthestructure of hexagonal ice,which hasastructure similar to that of tridymite [61,62]. The tridymite structure of ice exists at atmospheric pressure and temperatures between153 K and 273 K. Below153 K ice exists in a cubic, cristobalite-like structure. Another similarity is that both liquids exhibit an anomalous density-temperature behavior. The “sharpness” of this behavior (i.e., the manner in which volume varies around the maximum density as a function of temperature) and the temf l~aximumdensity (TMD) depend on the pressure [57,58,63]. In water, is 4°C and it occurs in the stable region at normal pressure. The maximum density of water can be attributed to the competitionbetween two opposing effects: the gradual breaking downof the rather open ice-like structure to a somewhat closer-packed structure(as indicated by the increaseof the average numberof nearest neighbors with temperature)and a simultaneous increase with temperature of the average center-to-center distance [64]. In liquid silica, the maximum density occurs in the supercooled region 1571. Water and liquid silica also exhibit a related

anomaly in the particle diffusivity, which was found to increase with increasing pressure (density); their viscosities have negative pressure dependence. The origin of the parallels in the anomalous behaviorarises from the tendencyof water and liquid silica to establish their local molecular structure to the expanded open,four-co-ordinated(tetrahedral)arrangement in which internal energy, entropy, and density decrease with decreasing temperature[65]. The presence and character of the anomalous behaviorin these “tetrahedral” liquids depends on the value of average bridge-bond angles, .H 0 and Si 0 -Si. The differences and the melting point and the larger sharpness of the density maximum in water in comparisonwiththerather flat density silica are due to the fact that in water the average angle large, ~ 1 8 0 ”at , normal pressure, whereas the average Si mined for vitreous glass (661) is only about 144”. The structural compatibilities for ordered water near the fully hydrated surface (bound water) and amorphous silica results in the unique interfacial region, which willbereviewed in the next section. For better understanding of the interaction mechanism between water and silica it is useful to know the structure and cooperativeproperties of bulk liquid water. nowledgeof these isnecessary in order to distinguish between the structure, dynamic behavior, andthernlodyna~ic properties of bulk and bound water. The properties of water havebeen investigated for more than a century. A numberof models, the majority of which are based on the short-range ordering of molecules, have been created to explain the structure and unusual properties of liquid water. The existence of variation in the models, each of which taken separately explains the separate properties of water, indicates that the water structure is not fully understood, especially in the supercooled state. At present itis generally acceptedthat the distinctive thermodynamic and dynamic properties of water are determinedby making and breaking of hydrogen bonds1671. Though ltnowledge of the structure and properties of water is very important for understanding the water-silica interaction we cannot review the whole range of them in this chapter. The enormous volume of the e~perimental andtheoretical material concerning the subject has been reviewed, for instance, in the books by Franks [68] and the reader can refer to them.

esides the basic scientific interest involved (for instance, the unusual behavior of water adsorbed from the vapor phase silica surfaces) the study of water-silica interaction also has important practical applications. An improved understanding of the interaction between silica surfaces and water is of considerable technical, industrial, and environment importance, which arises from the widespread use of silica as an adsorbent andcatalyst support and from thefact that aerosol silica and particle ’conareawidespread solid particulatepollutant in theatmosphere. of the binding of water molecules to silica surfaces has been well do om infrared spectroscopy 1691. Most of the early investigations of the silica-water interface have been summarized in the reviewby Clifford 1701, and

In spite of the intensive study for well over 50 years, much yet remains to be studied of the physicochemistry of silica surface, of the structure and thermodynamic properties of water in the vicinity of the surface, and of the extent to which the surface-induced perturbations propagate. A full understanding of the role of silica surfaces in the modificationof the structure and thermodynamic properties of water requires answers to the following questions: (l) How much does the structure of the perturbed water differ from the structure of bulk water? (2) How does this structural difference depend on the physicochemical properties of the surface in terms of its hydrophilicity and microporosity? (3) How much time does a water molecule spend in the perturbed region near the surface beforediffuses it away into the bulk water? (4) ow do the concentration of the surface hydroxyl groups and the morphology of ica influence the adsorbed structure of water; in other words, is the structure layered or clustered?(5) Is it the material obtainedby dissolution of silica during the adsorption of water from the vapor phase that determines the modified properties of the adsorbed water? If so, then: (6) Why is it that such materialcannot be obtained when liquidwater is drawninto silica? From the variety of the silica-water systems we will mainly be interested in the one which in the allows us to obtain the systems with a large surface-to-volume ratio, i.e., system obtained from the adsorptionof pure water vapor and vapor mixtures of the atmospheric acids on partly hydrophilic surfaces of fumed silica. Below, a short review ofthe recent experimental data regarding thewater-silica surface interaction will be given. This review is not intended to be complete, but only indicatesthat the questions under consideration are very interesting and, in spite of the fact that the level of knowledge of the water-silica interaction is high, the subject still remains far from completely~nderstood.First, thecase of a fully hydroxylated silica surface will be considered. Then the interaction between water and a partly hydroxylated silica surface of fumedsilica and the formationof finely divided water droplets will be given. Effects of nitric and hydrochloric acids on the water vapor uptake and possible scavenging of acids by silica particles will be discussed. 1,53,71-74], neutron scattering measurement 1521, orption studies (16-18,75,76], and IR spectroscopy [77], have been used in investigation of thestructureanddynamicpropertiesof water in the vicinit surfaces of a number of silicas. The general picture which emerges from these studies is that the mutualspecific silica-water interaction results in a significant ~odificationof the structure and transport properties of the “vieinal’, (bound) water in comparison with those of bulk water. urfkce hydroxyl groups -OH are the primary sites for the adsorptionof water molecules [41,78]. The adsorption mechanism is determined by the field-dipole interaction of polar water molecules with the surface -OH groups. Numerous studies indicate that, on the oxide surfaces, water molecules exhibit strong bipolar behavior due to the electronegative and electropositive character of oxygen and hydrogen ions, respectively. In an applied electrical field, which arises at the interface between water and oxide surface, water structure and its dynamic properties change in comparison with the case when such fields are absent. On fully hydroxylated surfaces of silica, water molecules can adsorb on two hydroxyl groups,

accompanied by increasingtheheat of adsorption. In this case the adsorption structure is layered, i.e., water molecules tend to occupy all surface before formation of the subsequent layers has started. It is the difference between the structure and dynamic properties of bulk and bound water that gives information about the strength and character of interaction between water and silica surface. Factors such as concentration of thesurface hydroxylgroups, porosity, concentration of impurity ions, and surface acidity strongly influence the strength of the interaction which determines the structure and transport properties of bound water [51,52]. Therefore, the distance to which the surface-induced perturbation propagates still remains a controversial issue. The useof NMR spectroscopy in the study of the structure and diffusional mobility of bound water on different silica surfaces has found considerable application for many years [71,79]. The main feature of the method is based on the ability of the nuclear magnetic momentof hydrogen ('H, spin to reflect very subtle effects occurring on the molecular level, i.e., relaxation times are very sensitive to the changes in the arrangement of water molecules and to their mobility. In water adsorbed on silica, both the longitudinal and the transverse proton magnetizations decay in a nonexponential way[72,73]. The anisotropic interaction of water molecules with surface hydroxyl groups constitutes a mo on the otherwiseisotropic orientation of molecules in bulk liquid. tion components havebeen accounted forby three different states of water protons: (1) protons of the surface hydroxyl groups;(2) protons of the bound water; and(3) protons in water above the perturbed layer. The influence of silica surfaces on the longitudinalrelaxationtimehas been assumed to be limited to onlythe first adsorbedmonolayer.The existenceof rapidexchange between theboundand bulk water has been suggested [737. The interaction between water and silica surfaces manifests itself in increasing viscosity and decreasing diffusion coefficients. The first two or three monolayers of water subjected to the surface ordering effects represent so-called high-viscosity boundary layer in which entropy is lower than that in bulk water. The connotation of entropy is order-disorder, and the decrease in entropy is interpreted to mean that bound water isless random in its structural arrangement, Such a highly structured layer can only be formed on the polar surface with a large concentration of ordered silanol groups. Inelastic neutron-scatteringmeasurements showed that water molecules adsorbed on fully hydroxylated surface are double hydrogen bonded [52] and, therefore, have almost no freedomof motion [.51,73]. In contrast, molecules adsorbed on the isolated silanol groups (in the case of a partly hydrophobic surface) possess a certain freedom of motion. The extent to which the surface-induced water structure exists is still uncertain but it isbelieved that the disturbance amounts to no more than two or three molecular diameters [70]. The influence of silica surfaces is restricted to a layer with thickness comparable to the distance over which structuring occurs in bulk water, a 1 nm. Therefore, it is expected that the water structure is sensitive to the presence of the substrate or even solute molecules (ions) within this range. Within this perturbed layer clear differences fromtheproperties of bulkwater are observed. Beyond this layer, water is assumed to exist as bulk liquid.

of silica, hydroxyl groups -OH are not uniformly distributed throughout the surface but rather statistically. Surface -OH groups are of two types: isolated, and arranged in small patches where they can be hydrogenbonded.Thedistances between the isolated freely vibrating -OH groups are sufficiently large to prevent formation of the multiple water-hydroxyl ecause of the large distance, the adsorption due to a bridge mechanism, in which each water molecule adsorbson two hydroxyl groups, of little probability [16]. The isolated -OH groups are more strongly attached to the silica surface than the hydrogen-bonded ones. During t outgassing the latter are removed from the surface at temperature of about 800 whichis much lower than the 1200 K found for removal of the isolated hydroxyls [77]. Removal of surface hydroxyl groups results in a much lower surface affinity for water and significant reduction in the water uptake, i.e., a hydrophobic surface is developed 1411. The adsorbed water on the partly hydroxylated surface is held less strongly than water adsorbed on the fully hydroxylated surface 161. O~~tgassing at room temperature of pyrogenic silica is sufficient for removal of the physically adsorbed water, but evacuating at 100°C is required for removal of water adsorbed on the more hydrated surface of precipitated silica (HiSil). ~ x p e r i ~ e n on t s the adsorption kinetics (variation with time of the weight of adsorbed water per gram of silica at constant relative humidity, the amount of water vapor taken up by fumed silica increased with increasing adsorption time [80]. The process of water uptake can be divided into two stages. During the first stage the adsorption proceeded more intensively and adsorbed weight increased very fast. During the second stage the adsorption process was much slower and theadsorbedamounttendedtoreach its saturationweight. The timeneeded to reachthesaturationadsorptionstronglydepended on the RH and was smaller for the smaller one. At large RH, during the second stage the adsorbedweight still increased appreciably, whereasat low R Such variation of the uptaken weight can be explained by the surface properties of fumed silica. At the initial stage of adsorption first water molecules sit “oxygen down” on theisolated hydroxyls [41,77]. At the pointswhere surface hydroxylsare closer to each other, water molecules can form two-dimensional clusters which, as the adsorption proceeds, grow and some three-dimensionality of the clusters can be developed [51,52]. At higher coverages, butstill smaller than one monolayer, water molecules are adsorbed around the molecules already adsorbed on theisolated OH groups to forrn small clusters of liquid water. orm mat ion of the clusters, or nuclei of a liquid phase, starts even before all the isolated -OH groups havebeen occupied by water molecules. Such adsorption behavior of water is typical for the partly hydrophobic surface. It is caused by the relationship between the adsorption energy per water moleculein liquid clusters and the energy per molecule adsorbed on anisolated hydroxyl group. In the first case the energy 10.5 kcal/mol, whereas in the second case it only 6 kcal/mol. Therefore, it energetically more advantageous for watermolecules to forrn liquid clusters than to bond to asingle hydroxyl group 181. ~ u a n t u mmechanical simu ions showed that the stable water clusters consist of five or six molecules [81,82]. olecules in the clusters are more restricted

in their motion than molecules adsorbed on theisolated silanol groups and, therefore, have a smaller entropy. Thus, in the case of the partially dehydroxylated surface and small coverages, the adsorption of water on the single isolated silanol groups is favored by entropy, whereas the formation of liquid clusters is favored by energy stabilization of water in the clusters of liquid [18]. Competition between these two mechanisms determines the adsorbed structure of water on the partly hydrophobic surface of silica. Clusters formed in such a way have different sizes which depend not only on the density of silanol groups but mainly on initial stage, during which water clusters are formed and, consequently, the adsorption weight increases abruptly, is termed “physical adsorption” [41]. The clusters formed supply molecules for the rehydroxylation of the hydrophobic siloxane surface [77].Formation of the new hydroxyl groups takes place owing to the rupture of silicon-oxygen bridging bonds, Si-0, by the reaction of the “residual valences” of the surface silicon atoms with molecular water Si-0-Si

H20

2Si-OH

and by hydroxyl groups or “ionized water” such that Si-0-Si

OH-

Si-0

where an ionized silanol surface complex is formed [41,83]. With time the hydroxylated surface increases along the boundary between silanol and siloxane surfaces. Formation of the first adsorption layer on the siloxane or silanol surfaces is termed “chemisorption.” This layer is formed by the covalent or ionic binding of water molecules with the surface Si atoms [4l]. The rehydroxylationof the hydrophobic surface of fumed silica with vapor is slow process W can be realized only partly.Therehydration is morepronouncedforhigherand is almostsupIt is interesting that formation of liquid clusters is also possible on the surface concentration of surface hydroxyl groups,at least at the initial adsorpielastic neutron scattering measurements of water adsorbed on silica (silanol number 4.6) showed that, at coverages smaller than one monolayer, the adsorbed water existed in two phases with different dynanlics. The ed immobile phase was formed due to inter~inking groups with “first down’’ water molecules l ,52,73 sequent adsorption network formed mobile, molecular clusters, whi doubly hydrogen bonded [M]. That a certain amount of clustering occurs on the hydroxylated silica surface at the early stage of the adsorptionwas shown by experiments on freezing of the adsorbed water. It was observed that even at coverages smaller than three monolayers (this is the thickness of the bound water which remainsunfrozendown to verylow temperatures) part of the adsorbed water does freeze [84]. (Freezing of water adsorbed on silica surfaces is discussed in Section VIA.)

o ~ ~ ~ ~of t Finely i o n Divided It has beendiscussed above that the structure of waternearthe silica surface depends on the degree of hydroxylation of the surface. It varies in nature from a

highly structured form on the fully hydroxylated surfaceto aloose, disordered form on thepartly hydroxylated surface[85]. In the lattercase the adsorbed structure can be approximated by a population of size-distributed nanometer-scale droplets or finely divided water (FDW) droplets. Formation of FDW droplets on the surface of fumed silica and demonstration that a reduced amount of the adsorbed water gives rise to smaller water droplets can be described as follows (25,801. Water molecules adsorb on the surface hydroxyl -OH which are of two different types: large amount of the isolated -OH groups and small amountof the hydrogen-bollded -OH groups. Since surface ups are statistically distributed, their density is different throughout the The fraction of the surface receptive to water is only between 0.15 and 0.25 times its specific surface area 16,861. Condensing water molecules formliquid clusters more easily onthepatcheswith larger density of -OH groupsthan adsorb on thelow-density isolated -OH groups. The clusters formed are not comme~suratewiththesubstrate, i.e., theyaremuchsmaller than the sizeof primary particles. The clusters canhave also limited contactwiththe surface, i.e., they are attached to the surface only with several hydroxyl groups, that determines their relatively larger freedom in motion. the adsorption proceeds, the size of the clusters increases and the adjacent ones merge (to reduce the total surface free energy) due to the contact between them to form larger liquid domains. The size of the domains to a large extent depends on the partial water vapor pressure, Since at large RH concentration of water molecules in the vapor phase is large too, the size of condensing clusters will be larger than. at small RH. This means that merging events will occur more often, which, in turn, leads to larger liquid domains. The isolated -OH groups disposed between the clusters can serve ng the merging process. Since even for the largest adsorbed weight (z94%) the ratioof the volumeof silicaand the adsorbed water to thetotal volumeoccupied by thesample was smallerthan 1/40, these liquid domains have a large free surface [80]. All these allowed us to consider the liquid domains as a population of microdroplets or FDW droplets. Besides RH, size of microdroplets depends also on geometrical configuration of chains of silica particles and concentration of the surface hydroxyl groups. The size of the microdroplets can be smaller, comparable or even larger than the size of primary particles Formation of microdroplets is schematically depicted in Fig. 1. A possible mechanismwhich can promote the formation of microdroplets on the surface of silica can be similar to that experimentally observed forlarge (from one hundred to a few thousand atoms) antimony clusters, nonepitaxially oriented on graphite surfaces (871. It was found that clusters had a surprisingly large surface diffusivity atroomtemperature.The result was ratherunexpected, since most diffusion mechanisms which have been considered for explanation of the surface diffusion of the clusters involve a combinationof single atom diffusion events, and gave a diffusion coefficient much smaller than that observed experimentally[88]. It has been supposed that the drawing mechanism of the observed phenomenon was the mismatch between the lattice parameters of the clusters and the substrate. The interplay between the vibrations (photons) of the substrate and the internal vibration of the clusters can create a random force on the clusters which will execute a

Scheme of the formation of finely divided pure water and solution microdroplets on the surface of fumed silica. rownianmotion in the weak external field. The similar mechanismcould, in principle, exist also in the caseof fumed silica-water system. Indeed,liquid clusters formed on silica surfaces are much smallerthan thesize of the primaryparticles and they can be attached to the surface onlyby several hydroxyl groups. Therefore, the clusters are not strongly locked by the substrate and can vibrate relatively freely. Random forces, caused by the mismatch between the vibration of clusters and vibration of the primary particle can bring about the diffusion of clusters along the surface and thus increase the number of merging events. Diffusion could be promoted,for instance, by mediation of the isolated -OH groupsdisposed between the clusters, At small coverages theelectrical interaction between microdroplets can prevent their merging and formationof larger droplets. Theorigin of this mechanism can be described as follows. Very small water clusters possess larger entropy than thelarge clusters or microdroplets consistingof thousands of molecules. Indeed, dueto very large surface-to-volu~eratio it is of little probability that smallnumber of molecules can form the packed structure resembling that of bulk water, in which there exists a short ordering. Such very small clusters merge to reduce the total surface free energy as was described above. With increasing the number of molecules in acluster its entropy decreasesand arrangementof water molecules from the center of the cluster increasingly resemble the structure of bulk water. But due to the large surface curvature the orientation and interaction of the surface molecules will be different from that in bulk water. Large curvature can increase theamount of the surface molecules with a non-hydrogen-bond (free -OH) projecting into

the vapor phase (reducing in such away surface excess entropy) in comparison with that of bulk water in which the amount of such molecules was found to be about 20% of a full monolayer [89]. Increasing orientational structuring of the surface molecules results in increasing value of the surface potential, which for the case of bulk water has been found to be about 0.16 at 300 K [go]. This potential can bring about electrical interaction between themicro-droplets, resulting in their repulsion. Inotherwords,ata certain sizeof microdroplets their “hydrophobicity” can be large enough to prevent merging of microdroplets to form larger droplets. With increasing coverage, whichwill takeplace at larger RH, microdropletscan merge duetocontact, since growingmicrodropletscansurmount therepulsionforce between them.

There hasbeen Little study of the influenceof acids on the adsorptionof water from the vapor mixtures. In .the previous paragraphs we have discussed the fact that water vapor uptake by fumed silica increases monotonically with increasing partial watervapor pressure. Amechanism of the f o r ~ a t i o nof population of droplets on a partly hydrophobic surface of fumed silica has been proposed. It is reasonable to expect that the presence of moleculesother than watercan changethe water-silica interaction,which,inturn,may result inachange of both the amount of theadsorbedwaterandtheadsorbedstructure itself, As mentioned finesilica particles andcrustaldust particles containing silicon in large amountsare widely presentintheatmosphere wherethey can serve as condensation and ice-forming nuclei. Atmospheric acids HCl, HN03, and adsorbed on these particles canmodify their physicochemicalsurface properties and change in such a way the character of heterogeneous water droplets and ice ~lucleation[80]. This study was also partly motivated by the attempt ssible removal mechanism of pollutant gases from the atmoin this field may also be useful for industrial applications and sses. The adsorption kinetics of the binary H20-HN03 and H20-HCl and ternary C1 vapor mixtures on the surface of fumed silica has been studied for several acid concentrations. The adsorption experiments described elsewhere [25,80,91]. In short, a well-defined dry powder of fumed silica (S 5380, product of Sigma Chemical Co., surface area 255 Irr: 15 m2/g [49], and Aerosil with surface area 256 m2/g produced in theInstitute of ChemicalKinetics andCombustion, ussia) was equilibrated with saturated vapor mixtures obtained in desiccators with solutions different acid concentration. The adsorbed amount was determined by concentration ofadsorbed theacids was calculated from the Adsorption experiments showedthat the adsorbed weight per gram of silica was practical~y the same for pure water vapor and mixtures obtained over solutions with small nitric acid concentrations,0.001, 0.01, 0.1, and .0 wt% The character of the freezing and melting transitions (see below) in the obtained silica-solution systems did not differ from those obtained in the case of a pure water-silica system.

From these Eacts a conclusion has been drawn that at low concentrations of nitric acid in vapor mixtures the amount of the adsorbed acid molecules was so small that they hardly influenced the adsorption process and phase transformations. The character of the adsorption kinetics changed when the concentration of nitric and hydrochloric acids in the solutions was 5 wt% [80]. The total weight (acid water) adsorbed from the vapor mixtures depended on acid concentration and on the type of acid itself. An exampleof variation with timeof the total weight adsorbed from the vapor mixtures obtained over solutions with 15, and 45 wt% HN03 is shown in Fig. A similar character of the adsorption kinetics has also been observed forH2O--HC1vapor mixtures obtained over solutions with 15, and 25 wt% HCl (RH FZ 72, and 44%, respectively), In the figure, time of the p measure~entsare designatedby the sign The first feature which onecan see fro the figure is that, in contrast to the case of pure water vapor, the total adsorbed weight did not decrease monotonically with decreasing RH: atfirst it decreasedand then increased abruptly. Since the amount of the adsorbed acids increased monotonically with acid concentration in mixtures, a nonmonotonic dependence of the total adsorbedweight wasdue to theweight of the uptakenwa at smallest RH 53% (EI~O-HN03mixture) and R H 44% the uptakenweight of water per gramof silica was more than two times larger than thattakenupat RH 75% andRH72%, respectively, This indicates th small REI both H N 0 3 and HCl promoted the water vapor uptake. At all 0.5

0.4 0.35 R0

0.3

g

0.25

b4

g B

0.2

0.1 0.05

0.0

0

100

200

300

400

500

h

Variation with time of the adsorbed binaryH20-HN03 system from vapor mixtures obtained over solutions with 15, 30, and 45 wt% concentration of HN03. For given acid concentrations relative humiditywere 75%, and respectively. The sign indicates the time of the pH measure~en~s.

weight of the uptaken water from the H20-HCl mixture was larger than that taken up from theH20-HN03 mixture, i.e., hydrochloric acid promoted the water vapor uptake in larger amounts than nitric acid [SO]. Calculations showed that the concentration of acids adsorbed onsilica surfaces was much larger than the concentrationof acids in vapor mixtures,i.e., fumed silica accumulated HN03 and HCl and the accumulation was larger for nitric acid [SO]. On the other hand, a series of consecutive pH measure~ents performed on the same binary system adsorbed on Aerosil showed that for the smaller adsor (smaller coverages), concentration of nitric acid was larger than for larger weights (large coverages), i.e., at initial adsorption time molecules of nitric acidadsorbed on thesurface -OH groups faster than watermolecules 1911. With time the incoming water molecules adsorbed ontop theof acid molecules already adsorbed (hydrationeffect) to form liquid clusters and microdroplets and, therefore, the concentration of the adsorbed acid became smaller. It was said above that the effect of acids on the water uptake by fumed silica was different for different relative humidities.At RH 93% the presenceof acids of the uptaken water in comparison with the case of pure 94%) and the reduction was larger for nitric acid. On the cter of the adsorption processitself did not change noticeably in comparison with the case of pure water vapor, i.e., the mechanism of the formation of the microdro~letswas not significantly disturbed by acid molecules. eduction in the uptaken water could be due to the smaller RH (about and due to partof the surfwe hydroxyl groups being occupied by acid molecules which formed the surface complexes through the reactions Si-OH

HCl

SiOHzCI

and

HN03

SiOH2N03

In the case of the small amount of the adsorbed acid the hydration effect, which promoted the water uptake,was insignificant in comparison with disturbed formation of clusters caused by Competition between acid and water molecules to occupy the surface --OH groups. This competition led to a decrease in the water uptake. At larger acid concentration (RH 75%, Fig. 2), the character of the adsorption process changed signi~cantly. In contrast to the case of pure water vapor, when at the initial time the adsorbed weight increased abruptly and the state of saturated adsorption was reached relatively fast, the adsorbed weight from the mixtures increased continuously and the state of saturation adsorption was not reached for more than one month [SO]. Such adsorption kinetics might be caused by an amount of the adsorbed acid that was large enough to causeenhanced disruption of the siloxane surface according to the reactions Si-0-Si

HC1

Si-Cl

Si-OH

and Si-0-Si

HN03

Si--N03

Si-OH

Chemisorbed in such a way, molecules of HCl and HN03 formed new hydroxyl groups, which then could formrelatively large hydrophilic patches. Thus the intensive destruction of the siloxane surfaceby acid molecules led the adsorption structure to become more layered in comparison with the case of pure water. This process manifested itself in that the adsorbed weight increased uniformly over the adsorption time. The reduction in the uptaken waterin the case of RH 75% was dueto the fact that the relative humidity was much smaller than that (93%) in the caseof low acid concentration and the hydration effect, which on the contrary promotes the water uptake, could not compensate it. But what is interesting is that such behavior was not observedin the caseof the third sample (Fig. 2). Here, in spite of very small relative humidity, RH 53% (largest acidconcentration)therewas anabrupt increase in water uptake. The increase has been brought about by the hydration effect which became much larger in comparison with the previous case[80]. The character of variation of the adsorbed weight obtained from ternaryH20H C l - H ~ 0 3vapor mixtures did not change significa~tly fromthe case of binary mixture with the largest acid concentrations. In this case only the adsorbed weight was larger and increased monotonically with increasing acid concentration [80]. Although these simple adsorption experiments cannot serve as an argument for the existence of a removal mechanism for the atmospheric acids caused by silica particles, the results obtained allow us tosuggest that such a mechanism mayexist. The fine aerosol silica particles, which have a long airborne lifetime during the transport through the atmosphere, can adsorb the atmospheric acids, and such a sink may be operative and its existence cannot be ruled out. Further studies are needed to clarify the matter. An indirect con~rmationof this assumption canbe the findings concering scavenging of sulfuric acid by aerosol particles of crustal origin [40]. It was found thatin the atmosphere, aerosolparticles containing large amounts of silicon were coated with a sulfuric acid layer. This layer has been supposed to be formed by oxidation of gaseous on the surface of particles or condensationof finesulfuric acid aerosols. Coating has been observed on particles of both fine and coarse size ranges. The sulfuric-rich particles, accounting for up to 73% of sulfur and 40% of silicon were found in the Arctic winter atmosphere. Oxidation of SO2and condensationof sulfuric acid aerosols on fine particles lead to alarge increase in their size. All these suggest that silica and crustal particles may supply a surface for condensationof gaseous sulfur compounds andsulfuric acid aerosols, with subsequent scavenging from the atmosphere. Also in the high troposphere and stratosphere, submicron solid particles coated in a liquid layer have been observed [7]. Such mixed particles can be formed in aircraft exhausts dueto condensation and oxidation of SO2on the surfaceof crustal particles 171. Condensation of sulfuric, hydrochloric, and nitric acids on thesilica and crustal particles can lead to the formation of relatively large particles composed of a solid core enveloped by a liquid solution shell. Since the coarse particles are more influenced by gravity they can deposit faster from the atmosphere.

It is known that the physicochemical properties and character of the phase transformations in very small particles differ significantly from those of bulk material

71

an

[25,24,92-971. Such particles have a large surface-to-volume ratio because the number ofmolecules (or atoms) at thesurface is anonnegligible part of the total number ofmolecules and, therefore, potential due to surface effects dominates over bulk. The nanometer-scale particles are of special interest because they lie at the transition between bulk material and molecular clusters, in which the influence of the properties of individual molecules playsan importantrole. The same reasonings apply also to the adsorbed films of nanometer thickness. But in this case the role of the substrate can be considerable. Knowledge of properties and understanding the character of the phase transfornlations in the systems with alarge surfaceto-volume ratio are of great importance for fundamentalinterest and for resolving many industrial problems connected with emulsification, colloid stability, etc., and also for understandingbiological processes, since water confined within living a cell also has a large surface-to-volume ratio. The knowledge in this field is also very useful in manyexperimentalsituationssuch asobtainingthinadsorbed films, nanostru~tures, nucleation, synthesis offinely divided materials, aerosols and cloud particle formation in the atmosphere, dust in free space, etc. 1931.

Depending on the method of preparation, very small systems candiffer in their size and amount of free surface. When based on the free surface criterion, the small particles can be of the following different types: Particles suspended in the unlimited vapor phase or deposited on the inert substrate. In this case, all or almost all of the surface of the particles free. Such systems would be ideal for investigation of properties and phase transformations in small systems, since the influence of the substrate is absent or reduced to minimum. But obtaining such systems of very small particles is a difficult problem, especially in the case of liquid microdroplets, due to the Kelvin effect. ~icroemulsions,produced by dispersion of a liquid (usually water and aqueous solutions) in a continuous oil phase containing some conce~trationof surfactant molecules providingthe oil-liquid interface. The size distribution of droplets in microemulsions is usually narrow and determined by the relative concentration of the oil and surfactant [71,98]. In this case the free surface of the microdroplets is zero. The influence of the oil-liquid interface on properties and phasetransitions in thesubmicron-sizedropletsmay be large. ~icroemulsionswith nanometer-scale droplets have not yet been studied. Liquid or solid “particles” obtained by confinement within. the nanometer-scale pores of silica glasses.The free surface of the dispersed materialis very small or even absent and, therefore, the influenceof the matrix, which sometimes is difficult to assess, can be large. Finely divided water droplets obtained by the adsorptionof water vapor on the partly hydrophobic surface of fumed silica. In this case the free surface of the particles can vary over a large range depending on the surface properties of silica and morphology of the chains of primary particles (geometrical config-

uration of the chains). The role of the silica matrix is reduced dueto thelimited contact of droplets with the surface. The study of the properties of small particles has a long history. The effect of size on the meltingand freezing behavior of small particles has been of considerable experimental and theoretical interest since the early 1900s and extensive experimental material has been reported [30-32,99-1041. When the contribution from the surface free energy becomes large and comparable to the bulk energy, the equilibrium melting temperature is much lower than that of bulk material. The first explicit expressionshowing that meltingtemperaturemonotonicallydecreases with particle size appeared in 1909 [99]. Since that time, several theoretical models have been proposed for the explanationof the experimentally observed size-dependent reduction of the melting point [32,92,93,101]. The common result of these theories is that melting temperature decreases linearly or quasilinearly with the radius of particle, rp. The decrease is expressed by the Gibbs-Thomson equation 1021

where is liquid-solid interfacial energy, density of solid particle, AH(J/g) the bulkenthalpy of fusion, and theequilibriumbulkmeltingtemperature. In equilibrium thermodynamicsof all parameters entering Eq.(1) assume well-defined quantities which are independent of size. But as will be discussed below,it does not seem very likely that nanometer-scale particles should possess macroscopic properties. Systematic experimental studies of the melting and freezing behavior of small particles have been performed by optical, dilatometric, calorimetric, and spectroscopic techniques [32,98,100,101]. The main features by which the phase transformations in small systems differ from those of their microscopic limits, are: (l) the melting and freezing temperatures, Tm and Tf,decrease with decreasing particle size;(2) there exists thermal hysteresis between the transition temperatures with TIn T f ; (3) the phase transformations spread over a broad temperature range. Some recent experiments and theoretical simulations have shown that, in addition to these features, there may exist a fourthone, namely,(4) the enthalpy of fusion of the small systems is smaller than that of bulk material and decreases with decreasing size of particles 1,96,97]. Besidesthe anomalies in the phase transitions the cohesive properties can change withsize of particles as well. For instance, when the diameter of the silica particles is less than 100 nm, the particles spontaneously adhere together in loose aggregates. But when the particle diameter is larger than 5000 nm the properties approach those of bulk silica; the cohesive forces become very weak, particles do not attracteach other, and the powderis very “dusty” 1411. In contrast to the metals, where astable population of very small particles can be created by chemical vapor deposition on an inert substrate, it is very difficult or even impossible to create a systemofvery smallwater clusters or droplets suspended in. the unlimited vapor phase or formed on a suitable substrate which would be in stable equilibrium with the vapor phase for a relatively long time. (Nanometer-scale water clusters have been obtained with a very short lifetime, of

7

the order of microseconds, by condensation of water vapor in supersonic flow through a miniature Lava1 nozzle [103,104].) Due to the Kelvin effect the larger particles would grow at the expense of the smaller ones. Even if a population of microdroplets of uniform diameter has been obtained they wouldeasily evaporate due to the very large vaporpressure needed for their existence. It is easier to obtainthe second andthethirdtypes offinely dispersed liquid, namely, by dispersing of water or aqueous solutions in the oils (microemulsions) or by loading of liquids (or adsorption of vapors) into the nanometer-size pores of silica glasses. A number of experiments in ther~odynamics,relaxation, and spectroscopy have been performed on populations of pure water and solution droplets suspended in different oils [105,106]. These studies showed that pure water and aqueous solutions confined to droplets of the order of a micrometer in diameter have the same bulk properties. These results are physically reasonable since droplets with such a size are large enough that the surface effects cannot dominate over bulk. Measurements of the ice homogeneousnucleationtemperatures and nucleation rate have shown that the surrounding medium (disperrsant phase) has not influenced the freezing process [106]. But microemulsions, in which the diameter of water droplets was less than l00 nm, have not yet been studied in the low-temper~~ture region, though these studies would obviously be of considerable interest for resolving the large surface-to-volu~eproblems. Obtaining microemulsions withnanometer-sizedropletsruns up against severe problemsconnected with their stability at low temperatures [105]. The propertiesof water confinedin the nanometer-size pores of silica glasseshave been subject of considerable interest for many years [86,107-1091. The interest has been caused not only byscientific curiosity concerning the anomalous behavior of water on the surfaceof silicabut also by the fact that the study of water confined in very small pores can give insight into the fundamental understand in^ of the properties of water itself and of finely divided materials in general. Besides water, a number of organic and inorganic liquids (ethanol [l lo], nitrobenzene and carbon disulfide l l], helium 1 12,1131, hydrogen [1 14,1151,oxygen [l 15,l 161, indium metal [95], etc.) have been studied in porous silica matrices. The similarity of the experimental data concerning the modified properties and anomalies during the phase transitions in different liquids, not all of which formed hydrogen bonds with the surface of silica,allows oneto suggest that the modificationswere ofa generalnature and have been brought about by the small size of the systems under study. Despite the wealthof experimental and theoretical data concerning the studyof properties and phase t r a n s f o ~ a t i o nin small systems, clearly there still remain many fundal~entalquestions to be answered. The main oneis whether theintensive thermodyna~icproperties, such as surfacetension, enthalpy of fusion, and possibly density, are size-dependent or not. in equilibriu~theremodynamics these parameters are well-defined values which are usually dependent on temperature but, by definition, independent of size. Most attention has been devoted to the study of the curvature effect on surface tension of small liquid droplets. In his well known work, Tolman 1171 has theoretically shown that surface tension of small droplets decreases with decreasing radius. An attempt has been made to extend Tolman’s idea to the case of ice crystals It has been shown that, as long as the Gibbs

surfacemodel is applied tothe crystal facets, thesurfacetension of small ice particles can also be size dependent. Since interfacial free energy of small particles is notoriously difficult to measure directly, especially when one of the phases is solid, it seems that the experimental verification of the curvature effect on surface tension does not appear probable,since the effect becomes apparent atsizes smaller than 100 nm. Therefore, the only tools in the study of this phenomenon are theoretical simulations. In the case of the enthalpy of fusion the situation different. In recent years several experimental and theoretical works have appeared which show that the enthalpy of melting decreaseswith size[25,30,31,96,97]. Thishas been observed with the differential scanning nanocalori~etryfor nanometer-scale tin particles, formed by successive deposition of pure Sn on the SiN substrate [30]. Reductions of the enthalpy of fusion by 46% compared to the bulk values, have been observed for a free spherical cluster consisting of 139 sodium atoms [96,97]. ~omputationalsimulations performed for gold clusters also showed a size-dependent reduction of [31]. Enthalpy of fusion of indium confinedin porous silica glasses has been observed to consist of about one-third of the bulk value 1951. It seems that water is the only substance, the densityof which, was reported from time to time in the literature, may havedifferent value dependingof closeness to the surface. For instance, ~ e a s u r e ~ e nof t swater density showedthat the water structure in the interfacial region immediately adjacent to the surfaces of clay was less dense than that of water in bulk. Density of the surface-held water in montmorilonite was of the order of 3% less than density of bulk water at the same temperature [38]. The differFnce became smaller at larger distances from the surface but was still detectable 50 distant. a reminder of the “polywater” era, which has been recognized to be as a result of contal~inationand dissolution silicates from the glass capillaries, can serve the work reporting that even bulk water can have excess density in capillaries smaller than l 0 pm (1181. The origin of this excess density has been recognized as unknown. More recent experiments on small-angle neutron scattering showed that densityof water adsorbed in porous Vycor changed gradually and was smaller of about 10% than that of bulk water [l 19,1201, Upon freezing, the density change between water and icewas6.6 f0.2% whereas in the bulk case the change is about 9%. The origin of the observed changes in water density near the surface may be due to the fact that the influence the polar hydrophilic surface on water tends to change the packing of water molecules in the first adsorbed layer [16]. Adsorption of the subsequent layers leads to the formation of a highly oriented adsorption film in which orderingpropagatesuptothreemonolayers,where entropy is smaller thanin thecase of bulkwater.Theenhancedstructural arrangement of molecules near the surface of very small pores can be the cause of the observed difference in water density. The question of whether the density change has been brought about only by the influence of the surface or might be dueto very small sizeof waterdomains(thecase of nanometer-scalepores) remains open. To the knowledge of the author there are no studies concerning the measurements of density of very small water droplets having large free surfaces.

From the variety of problems connected with a large surface-to-volume ratio, the one concerning ice-phase formation from supercooled pure water is strongly connected with the water-fumed silica system. A system of microdroplets formed on fumed silica can serve as a suitable “model” systemofsufficient simplicity for investigation of freezing and melting in aqueous systems with a large surface-tovolume ratio. In order tocircunlvent the Kelvineffect and to obtain acollection of microdroplets which would be in relative stable equilibrium with respect to the vapor phase, one should use as a host a material in which the volume available for each individual droplet would sufficiently small [25,29j. For instance, let the radius of a void space containing acluster of radius 10 nm be 90 nm. Then calculations show that for variation of cluster radius of only nm the changein vapor pressure arising from evaporation from (or condensation on to) the cluster is a factor of 3.1 and the change of vapor pressure resulting from the variation of the cluster radius due to the Kelvin effect practically equals zero. For this purpose, its unique physicochemical surface characteristics is an excellent h the void spaces between the chains of silica particles are interconnected, the slow exchange of molecules between the voids could not rapidly equilibrate the abrupt change in pressure causedby evaporation (condensation)if it occurred suddenly in some void. The molecules would be forced to condense (evaesides the problem of the ice-phase formation in the atmosphere, the understanding of freezing and melting behaviorof finely divided pure waterand aqueous solutionsdroplets at subzerotemperaturesforcryobiology, glaciology, isvery important and needs no stressing. Since supercooling of finely dispersed aqueous systems is large, low temperature and freezing impact the physicochel~ical properties of water, such asits acid-base behavior, the hydrogen-bond energy, the microscopic diffusional motion and configuration of molecules in liquid water and ice. Apart from observation of crystallization and melting temperatures,l ~ e a s u r e ~ e n t s of other parameters characterizing thefreezing and melting of individual droplets have not so far been judged feasible even for sufficiently large water droplets, let alone those of sub~icroscopicsize [105]. In terms of therl~odynamics,freezing is a first-order phase transition in which some order parameters undergo a discontinuous change. The classical ice nucleation theory treats the water-ice phase transition in terms of the passage of an ice embryo over the energy barrier separating metastable liquid phase from thestable ice phase. Equilibrium freezing of water is of statistical origin, i.e., it is caused by the density and configuration fluctuations of water molecules which are functions both of cooling rate and time [lz]. Inthe classical ice nucleation theory, description of the formationof the ice phase from supercooled water starts from the nucleation of an ice embryo of dimension (approximat~lyof the orderof l nm, depending on temperature). The theory provides a simplified model for quantitative estimation of critical free energy and critical radius. The usual practice has been to calculate these values using the bulk properties of water and ice. An approximate expression for the critical radius is 121:

orw

m

where is ice-water surface tension, t is temperature in “C, and pi and are the average values of density and enthalpy of fusion over the temperature range between and The critical free energy is calculated according to the formula 1121

where critical radius r; is given by Eq. This is the classical expression for the free energy which must be supplied by the density and thermal ~uctuationsin the supercooled water in order for nucleation to occur. Usually a question arises concerning the surface tension since it appears to the third power in the formula for calculationof free energy. Thisconstitutes so-called c a ~ i Z Z ~ r i t ~ a p p r o ~which i ~ u t states i ~ ~ ? that the surface tension of the microscopic ice embryo is the same asthat of bulk ice and is independent of size. Also enthalpy of fusion and densities of water and ice are assumed to be the same as those of bulk. The enthalpy of fusion AHm and ice density are squared in the expression for the critical free energy and, therefore, a small variationin their value can also influence the ice nucleation rate. Similarto the caseof metals, for water canalso be size dependent [25]. The simplified approach used in classical theory is problematical, since it not likely that clusters of molecules with size *l nm can possess macroscopic properties. Therefore, capillarity approximation has longbeen disputed [28]. In earlier studies the value of AHm for water was measured when water was confined in porous silica glasses, pore walls of which were fullyhydrophilic. In this case thefree surface of liquid was almost absent and, therefore, the role of the silica matrix could be very large. The data obtained fromthe experiments were found to contradicteachother, i.e., theyrevealed areduced increased 11211, or unchanged [l221 value of These couldbe caused by different p ~ y s i c o c ~ e ~ i c a l surface properties of the materials used. Theoretical calculationof the enthalpy of fusion for water using the methods of molecular dynamics runsup against the lack of fundamental understanding of the structure and low-temperature behavior of supercooled water (supercooling increases with decreasing droplet size). Insight into the problem could be gained from the measurenlents of enthalpy of fusion of droplets with sizes ranging from 1 nm to Unfortunately, as mentioned above such measurements are impossible due to (l) the Kelvin effect and (2) the difficulty of measuring the heat involved/consumed duringfreezing and melting of very small separate dropletswhich are too miniscule to be measured using existing techniques. This obstacle can be circumvented by increasing the total mass of the microdroplets under investigation. In other words, a “compacted” collection of droplets must be used,thetotal weightofwhich would besufficient in orderfortheheatevolved/consumedduringthephase transforl~ationsto be detected 125). Besides the thermodynamic properties also crystalline structure of microscopic crystal can be different from that of bulk ice. Usually it is supposed that the crystal structure of an ice embryo is hexagonal (Ih). But recent experiments showed that

the crystalline structure of ice formed during freezing of water confined to very small dimensionsis different. Electrop diffraction analyses showedthat water clusters with sizebetween 63 and 71 A(about 4000-6000 molecules) produced in supersonic flow and subsequently cooled by evaporation froze to form cubic ice at temperatures as low as about200 K [103,123]. These freezing temperatures are far below the -45°C (228 K) limit, which is assumed to be the mini mu^ temperature at which water can exist in the liquid phase [12,105]. Cubic ice is metastable relative to I h and has never been obtained in pure form from bulk liquid. At normal pressures and temperatures up to 153 K, bulk water freezes to stable hexagonal ice [62]. Cubic ice is usually prepared either by vapor deposition on an appropriate substrate attemperatures below 190 K or during warmingof vitreous ice above 120 K. Ice converts rapidly to hexagonal ice at te~peraturesexceeding =l53 K. Experiments showedthat ice can also be formed in the caseof the third typeof finely divided water, i.e., when liquid water is confined in nanometer-scale porous media. For instance, the crystalline structure of frozen water in porous Spherisorb S 2 0 (specific ~ surface area, F32 m2/g, silanol number 4.6 n/rn2, pores size distributed with mean diameter90 A) hasbeen studied with neutrondiffraction [124,125]. The results of the measure~entsshowed that crystallization occurredindependently in separate pores and that both cubic ice and hexagonal ice Ih had been developed. The formation of cubic ice started at tenlperature as high as 242 K and itsamount increasedwithdecreasingtemperature,whereasthe amount of formed atthe warmertemperaturesremainedconstant.The results can be explained as follows: water in the large pores froze heterogeneously at warmer temperatures as I h . \;Vater in smaller pores froze homogeneously to cubic ice at lower temperatures. A very interesting finding was that on warming, the normally metastable (in the temperature range0°C to -80°C and normal pressure) cubicice remained stable at temperatures up to the melting point [125]. The formationof cubic ice has also been observed duringslow warming near the glass transition temperatureof quenched electrolyte solutions (LiCl). Similar to the cases described abqve, the formation and stability of cubic ice also required small crystallites 300 A) 11261. Though the formationof cubic ice from water adsorbed in small pores of silica glasses can be induced by the surface of silica (for instance, the simple silicate glasses invariably crystallize to the metastable cubic silica structure, p-cristobalite, which has the same structure as cubic ice [126]) it seems that the similarity in the structure may only help the nucleation of but the main reason remains the requirement of small crystallites. Therefore, although thermodyna~ically unstable under normal conditions, the stability of the cubic ice can be kinetically correlated to the very small size of the crystallites.

Since the beginning of the 1900s the reduction of the bulk freezing point has been observed duringfreezing of water adsorbed onsilica gel 1271. Since that time a vast number of experimental studies have been reported on the solid-liquid phase transitions in different organicandinorganic liquids (includingwater)confined in

different silica matrices. Usually the study methods were NMR, differential scanning calorimetry IR spectroscopy, and neutron diffraction [l 15,120,124, 125,128-1301. A short reviewof therecentexperimental data obtainedwith NMR and on the phase transitions in water adsorbed on silica surface will begiven. Heterogeneousmechanisms of theice-phaseformation on the partly hydrophobic surface of silica will be discussed well.

1

NMR

Besides the study of the structure and diffusional mobility of water near the silica surface, NMR spectroscopy has widely been used for investigation of freezing and melting behavior of silica-water systems [74,79,91]. NMR is powerful method which allows one to study relatively subtle effects occurring on the molecularlevel, i.e., the small changesin the structure and motionof water molecules causedby the variation of temperature. Themodi~cationin the molecular motion determines the changes in the spin-lattice and times, TIand T2,which in turn determine theintensity and shape of the resonanceline. The resonanceline from an isolated water molecule is very sharp (about 200-300 Hz in width) and this is the case for liquid water. Of course, in the liquid state the interactionsbetween neighboring water molecules are large, but they vary so rapidly that only the average nearest environmentis detected. Therefore, in both cases the shapeof the resonance line is sharp. In the case where water molecules are constrained in the crystalline structure of ice the interactions between water molecules are much stronger and more static, which makes the form of the resonance line very broad (about 20-30 kHz). Since the widths of NMR signals from the protonsof water molecules in the ice phase and surface hydroxyl groups are greaterthan 20 kHz, they are unable to cause signi~cantdistortions in the integral intensities of power spectra obtained from liquid water. An exampleof the temperature-dependent integral intensity of the NMR signals obtained in the spectral range around 1 kHz fromthe unfrozen water adsorbed on the surface of Aerosil is shown in Fig. 3 [91]. The surface area enclosed by the curves has been taken measure of integral intensity (relative units). ing theNMR signal during coolingof the sample one can see what fraction water molecules participates in the freezing process. It is interesting tonotethatno signi~cantdeviations of the integral intensities have been observed during temperature stabilization for times from 10 min to 1 h for given experi~entalconditions, i.e., the adsorbedwater-ice system was in equilibrium. This mayserve as anindication that nucleation of ice in the population of microdroplets formed on partly hydrophilic surfacesof Aerosil was not function of time but mainly function of temperature. It is seen from Fig. 3 that integral intensity lowers with decreasing temperature, since more and more water molecules become locked in the lattice of ice. Gradual lowering of the resonance signal shows that freezing does not occur abruptly, it would be in the case of bulk water. Such freezing behavior can serve a confirmation that the adsorbed structure represents population of microdroplets of

Integral intensity of the NMR signalsfromtheunfrozenmobilewater adsorbedonthesurfaceenicsilica.Thearrowshowsthedirection of ternperature change. (From

different sizes which freeze at different temperatures. At -60°C the signal is very broad and represents almost a straightline: here practically all water nlolecules are ith increasing temperature the intensity of the signal gradually increases, since more and more water molecules pass from ice into the liquid state. ng melting the value of the signal remains smaller than during the ing process for all correspo te~peratures,i.e., thefraction of water molecules locked in the ice phase is larger at melting than at freezing. In other words, the thermal hysteresis between freezing and melting takes place The ter XI (Fig. 4) has been determined from the formula and Bo are values of the integral intensities from the ter molecules at the measured temperature and the equilibrium melting temperature The general picture which arises from the silica-water systems is that not all water participates in free layer of water in the vicinity the silica surface remains unfrozen at very low temperatures, sometimes, even lower than -45°C limit [74]. Unfrozen water was found to correspond to an amount from two to three monolayers In this layer, water possesses high viscosity and exhibits also a structur ent from that of bulk water even at temperature above 0°C Indeed, experilnents showed that water adsorbed on the surface of Spherisorb silica remained always diffusionally mobile in a layer with thickness approximately equal to three monolayers for all coverages larger than threemonolayers Adjoiningwatermoleculesare strongly influenced by the silica surface andformanamorphousstructureto

0.8

R

0.2

..L.

0.0 200 210

240 250 260 270

280 290 300

Temperature (K)

Fraction, of thefrozenpurewaterandbinary H20-HN03 system adsorbed on pyrogenic silica as a function of temperature. Weight of water in the NMR samples correspondingto the given coveragesm 8 mg. (Adapted from Ref.91 accommodate the complex geometry of the surface electrical fields as best as they can. The larger the degree of hydroxylation of the silica surface the more oriented are watermolecules in thevicinity of the surface. It is the strengthof the interaction between polar molecules of water and surface hydroxyl groupswhich determines a healing length away from the surface where the structure of bound water crosses over to the crystalline order of bulk ice. It is believed that only water outside this layer possesses bulkproperties and participates in freezing [74]. The perturbed water is supposed to be bound with an energy of greater than 6.03 kJ/mol, otherwise it would freeze in the form of ice crystals. It worth noting that unfrozen iff fusion ally mobile water at low temperature may have an origin different from that described above. was discussed above, water confined in a very small volume can, besides stable hexagonal ice, freeze to cubic ice. An unfrozen disordered stateof water can also be caused by the stacking faults (when layers of cubic ice are intermixed in hexagonal ice Ih) within the crystalline phase 11251. Withinsuchadisorderednetworkwater molecules can remain diffusionally mobile through the lattice defects. This assumption has been put forward for explanation of the origin of the unfrozen water during a study using neutron diffraction of freezing behavior of water adsorbed on Spherisorb S20W silica 1251.

C~~~s~rements The DSc method is sufficiently sensitive for studying the phase transformationsin condensed matter, since the amount of material required can be quite small-often

less than 1mg [131]. The DSc device consists of two holders. The first one contains a sample to be studied (test sample), and the second holder contains a dummy sample with similar heat capacity and heat transfer characteristics (reference sample). Both holders are supplied with a heater and a thermometer. To keep the temperatures in the holders equal at all stages of the scanning program, heat is differentially supplied or withdrawn. It is this differential rate of supply/withdrawal of energy (dyldt) which is measured. Since the temperature scanning rate ( d T / ~is~ ) constant, the actual quantity measured during the scanning experiment is the differencebetween the specific heat of the test and reference samples. Therefore, scanning experiments, in which the temperature is changed at a constant rate, are often used to determine the freezing and melting temperatures and the enthalpy of fusion (latent heatof fusion) by directly measuring the heat consumed or released during the solid-liquid transitions. Data concering the transitions are obtained from the shape of the curves. DSc ensures monotonic cooling and w a r ~ i n gand, it is supposed, that no heat is missed [l 151. Since the amount of liquid in the fumed silica-water or silica-solution samples can be relatively small and, on the other hand, the adsorbed structure on fumed silica can be considered as a population of finely divided water or solution (in the case of low acid concentration) droplets the DSc method is very usefulfor study of the phase transformationsin aqueous systems with alarge surface-to-volume ratio 1253. The samples for theDSc study can easily be prepared by equilibrating a welldefined weight of dry fumed silica with pure water vapor or vapor mixtures at constant relative humidity in desiccators [25,80,91]. Varying relative humidity (or concentration of acids in the solutions) and the adsorption time, one can obtain different adsorbed weights, i.e., populations of microdroplets withdifferent average size. A model for the fumed silica-water system, in which a reduced weight of the adsorbed water gives rise to a smaller size of microdroplets, has been described above. An example of the freezing exotherms and melting endotherms obtained during cooling and warming (3 K/min) of pure water adsorbedon fumed silica is shown in Fig. 5. The main features of the phase transitionsare: (1) the peaks of the freezing and melting temperatures, and are far below the equilibriul~melting point and decreasewithdecreasingthe amount of theadsorbed weight in the DSc samples; (2) a strongly pronounced thermal hysteresis between Tfand is clearly seen; (3) both the freezing exotherms and melting endotherms are spread over a broad temperature region. In contrary to the smooth melting endotherms, a stepwise l ~ a x i m ~in m the freezing exotherms indicates that in the latter case heat evolved stepwise. This fact, togetherwiththeasymmetricalshape of the DSc curves (at warner temperatures a larger amount of heat isevolved/consLlmed) can serve as confirmation of the assumption that the adsorbed structure on the surface of fumed silica is not a continuous film or a single liquid plug, but rather a population microdroplets withwide size distribution. In the case of a single plug, the DSc curves wouldbe sharp peaks, since even the smallest adsorbedweight .7 pg) is sufficiently large in order to have bulk properties. If the adsorbed structure were a ~ o ~ t i n u ofilm u s then there would be no freezing because even for the largest adsorbed weight (5.8 mg) the thickness of the film would be only four monolayers.

,

.

.

,

,

.

.

~

,

s

,

8

t

0

Temperature

Temperature(“C)

DSG curves obtained during freezing (a) and melting of pure water adsorbed on fumed silica. 1.7, 2.1, and 5.8 mg are weights of water in the samples. (From Ref. 25.)

For the smaller adsorbed weights the film thickness would be only 1.3 and 1.0 monolayer which, as was discussed above, remain unfrozen. Depression of Tfand the stepwisefreezing process, spread over alarge temperature range, have been brought about by the following: (1) the statistical origin of the freezing phenomenon, which states that the probability of an ice embryo formation is smaller in smaller droplets; and (2) distribution in the values of the contact angle forice, which characterizes theability of the silica surface to heterogeneously nucleate ice crystals. It is known from theclassical nucleation theorythat the contact angle strongly influences the nucleation rate.If the contact angle forice is small then the surface favors formation of ice. In the sample with the largest adsorbed weight themajority of waterfroze heterogeneously in the temperature region between -15 and -30°C. Such a wide temperature region canserve as indication of the large distribution in values of the contact angle and size of droplets. On the other hand, low freezing temperatures indicate that fumed silica is not a good ice nucleant (see below). good ice nucleant initiates ice at temperatures slightly below 0°C. Different valuesof contact angle are determined by different configurations of electrical fields created by the surface hydroxylgroups and their interaction with polar water molecules. It is known that ice-like structuring of liquid water increases withdecreasingtemperature [12,98]. At some temperature (called the characteristic temperature) fluctuations in the droplets can produce suchan arrangement of water molecules which, interacting with the surface --OH groups, can give rise to the formation of ice. The probability of heterogeneous freezing is larger for larger droplets, since their contact area with thesilica surface is larger. This is confirmed by the asymmetrical

form of the D S c curves, i.e., at warmer temperatures a larger amount of heat has been evolved. A low-temperature tail in the D S c curve is attributed to freezing of smaller droplets. In the case of smaller adsorbed weights, freezing started below -40°C (Fig. 5). Atthesetemperatureshomogeneous freezing of small ultrapure water droplets usually takes place If, in our case, freeezingof part of the water had occurred homogeneously, then this could be due to the presence of smaller microdropletscompared to the previous case. It is themicrodropletsformedonthe ordered hydrogen-bonded hydrophilic patches that freeze homogeneously, since the HVBL prevents heterogeneous freezing. Since the surface area of hydrophilic patches isvery small (see above), some water could also freeze heterogeneously even at low temperatures. This could occur in the case of very small droplets, the contact area of which with the surfacewas small enough to freeze heterogeneously at warmer temperatures. “Soft” smooth endotherms are due to melting of ice crystals with a wide size distribution. Their melting temperature is usually expressed by the Cibbs-Thomson equation, Eq. It is seen from Fig. 5 that for the smaller adsorbed weights the melting endotherms arevery broad and represent almost a straight line. This could be a cause of the reports that have appeared in the literature stating that at coverages smaller than three monolayers no melting transitions had been observed in water adsorbed from the vapor phase on the surfaceof Spherisorb silica in the temperature range between and K Melting of ice (and solids in general) starts always from the ice-substrateinterface. When the “substrate” is the vapor phase then meltingis referred to as surface melting If the substrate is a foreign solid (in our case a silica surface) then the process is called interfacial melting. The origin of these melting processes is the diffusive nature of the solidliquid interface i.e., melting always begins from the surface and propagates slowly into the solid core. Surface melting, whichis enhanced by the curvature (size) effect occurs in a continuous manner over a broad temperature range. In contrast to metal particles, wherethesurfacemelting is aprecursor of the complete (core) melting, which occurs abruptly at some temperate depending on the size of the particle, in the case of small ice crystals the abrupt core melting is absent due to the low thermal conductivity of ice and large latent heat of fusion. Here the solid-liquid interface gradually propagates into the ice phase, The slow melting process may partlybe caused by the larger weight of silica, the specific heat of which is different from that of water The weight of silica in the D S c samples was about five times larger than that of the adsorbed water. The underlying origin of the large thermal hysteresis between the freezing and melting temperatures is the essential asymmetry between freezing and melting phenomena. In thecase of freezing an energy barrier must be overcome. In the case of melting an energy barrier is absent or it may be very small Therefore, substantial supercooling of water is normal, whereas there is no superheating of ice. Besides the transition temperatures DSC allowsalso measurement of the freezing and melting enthalpies, and AHm. Distinction between AHf and AHm is caused by the strong temperature dependence of the enthalpy of fusion of water According to Kirchhoffs equation, d H / d T C, Ci, where and Ci are

the specific heats of water and ice, respectively. Temperature dependenceof Ciand are given in Ref. 12 [Eqs. (3-12) and (3-16)]. Therefore,theheatsevolved/ consumed during freezing/melting are not the same when the transitions occur at different temperatures. Since AHf AHm, measurements showed (see Table l) that the experimental values of the freezing and melting enthalpies, A@ and AH:, were smaller than those of bulk water, AH: and AH:, calculated for the peak transition temperatures according to

n=O

with A H in IT cal/ t in "C, and 79.7, -0.1200, -8.0481 10-3.2376 -4.2553 1121. In the third and fourth columns of Table l are also given amounts of the uptaken water per gram of silica and the total adsorbed weightsin the samples as functions of relative humidity for the cases of pure water vaporand binary H20-HN03 and H20-HCI mixtures. The adsorbed weights of water for the H ~ O - H ~ 0and 3 H20-HC1 systems are 4.1 and 2.4 mg, respectively. Since the amount of the adsorbed acids is small, the heat evolved/ consumed during the phase transformationsis mainly due to freezing out of water. Theexperimental data collected in the table clearly show thatbothtransition temperatures, and melting and enthalpies, AH; and AH:, monotonically decrease with decreasingweight of water in the DSC samples. In classical therl~odynamics,the enthalpy of fusion is an intensive property, usually dependent on temperature but, by definition, independent of size. In this context the question arises whether the reduction was actual or it has been caused by experimental or some other cause, for instance: (l) part of the adsorbed water was missed (evaporated) during loading of the pans; (2) part of the heat had been missed during the measurements;(3) part of the water remained unfrozenand, therefore, did not contribute any heat to the phase transitions; (4) material obtained during adsorption from the vapor phase does not represent normal water [70,136]. Let us consider them in succession.

Pure Water andAcid Solutions: RH, Adsorbed Weight of H 2 0per Gran1of Silica, Adsorbed Weight in DSC Sample, Concentration of the Adsorbed Acid, Freezing and Melting Temperatures, Experimental Enthalpyof Freezing AH;and Melting AH;, Freezing AH: and Melting AH:l Enthalpies

H20 5.8 0.31 0.229 H,O-HCl 0.137 HZO-HN03 H20 0.101 H20

94 93 93 90 80

0.082

4.2 72.6 2.1 1.7

332.2 288.4 149.5 131.3 -3.2-20.2 -5.2 120.1 136.3 1.4rtO.l -36.8 f 0 . 7 -36.8 74.0 79.3 -13 -42.1 -21.4 72.9 79.2 -43.3 -24.8 53.3 54.9

244.6 244.6 209.3 197.8

328.9 307.9 285.6 278.2

Measurements of the adsorbed weight showed that missing water during loading of the D S c pans was smaller than 3% and it was smaller for smaller adsorbed weights. Therefore, the missing weight cannot strongly disturb the obtained results. ue to the fact that D S c ensures the monotonic cooling and warming, it is supposed that no heat is missed [l 151. Nevertheless, someauthors assume that heat loss during the calorimetric measurements canbe a serious problem [137]. If the heat loss really takes place then in the case of very small systems this problem may really disturb the experimental results. In this work, it supposed that no heat was missed. Unfrozen water could be either in a quasiliquid surfacelayer on the surface of ice crystals or in the HVBL between the silica surface and ice. To the author's knowledge there is no evidence that such a ~uasiliquidlayer exists on ice at such low temperatures. Clancing-angle x-ray scattering experiments performed on the basal and prism facets of bulk ice crystals Ih showed that a disordered surface layer existed only at temperatures warmer than -13.5 ic 2.5"C [138]. Therefore, since the main amount of the adsorbed water froze at temperature lower than =-2O"C, it was assumed that such layer was absent. should be noted that this assumptioncan be consideredasquestionable. Similar to the case of metal particles the formation of the surface-induced disorder of the hydrogen bonds in ice can be enhanced by the very small size of ice crystals [30]. On the other hand,a disordered phase canalso exist due to the stacking faults originated by intermixing cubic and hexagonal ice [125], was discussed above.) The weight of the unfrozen waterin the HVBL has been calculated using the specific surface area and silanol number. If all surface hydroxyl groups constituted one large hydrophilic patch, on which a continuous could be formed, then for the layer with thickness of three molecular diameters the weight of the unfrozen water would be about 46 mg per gram of silica. This weightgave the fraction of theunfrozenwater 0.15, 0.45, and 0.59 for adsorbed amounts of water in the D S c samples of 5.8,2.1, and 1.7 mg, respectively 11251. The fractionof unfrozen water caneasily bedetermined usingNM scopy. Such ~easurementsperformed for the comparable adsorbed mg gave the fractionof the unfrozen water in the pyrogenic silica-water system (similar specific surface area 256 m2jg and similar silanol number) smallerthan 0.05 at -60°C [91J (seeFig. 4). Therefore, the calculatedfractions of the unfrozen water,most likely, are overestimations, especially for smaller adsorbed weights. The overestimation can be due to: (1) not all isolated -OH groups may bear water molecules. Indeed, it was discussed above, that the bonding energy per water molecule adsorbed on the isolated -OH group is almost two times smaller than the adsorption energy per molecule in liquid clusters. For this reason part of the isolated -OH groups may be free from water, especially in the case of small coverage; (2) since freezing of part of the water has ed by the surface, the unfrozen layer between ice and the surface was t even the subtraction of these overestimatedweights gave corrected

values for the enthalpies morethan two times smallerthat those calculated for bulk water at Tfand Tm [25]. 4. The hypothesis that material resulting from the chemical attack of water molecules condensing from the vapor phase could cause the unusual behavior of waternearthe silica surfacehas been put forward in order to explain the anomalous freezing of water on silica surfaces [70,136]. It was assumed that the material consists of two components, which cannotbe interconverted by reversible chemical reaction. One component has been supposed to be some gel-like solute silicic acid gel) and the other one normal water. i n principle, such material could be formed due to dissolution of silica. ~nfortunately,it seems that from the time the hypothesis was proposed, there have been no attempts to determine the nature and properties of the material. The reason may be that the study of this phenomenon involves measurements near the limits of detection. Dissolution of silicahas onlybeen studied in the caseof liquid water [41,139], This mechanism can be described briefly follows. In an excess of water, the dissolution of silica involves hydration catalyzedby OH" ions with the formation of monosilicic acid Si(OH)4,

Solubility of massive anhydrous nonporous amorphous silica at 25°C corresponds to approximately 70 ppm SiOz. iss solution of the hydroxylated aggregates of finelydivided silica particles is larger and consists of 100-130 ppm Si02 [41]. During the dissolution theSi02 molecule does not remain intact, in the case of an aqueous solution of sugar, and does notdissociate to formions in the case of an aqueous NaCl solution. In the case of silica new substance is formed, which is nonionic in neutral solution and does not conduct electric current [41]. Maybe it is this substance which is responsible for the unusual properties of water in the first three monolayers adjacent to highly hydroxylated silica surface. But it seems that there has been no study co~lcerningthe solubility of silica from the vapor phase. Can the material resulting from the dissolution of silica cause the observed anomalies in the freezing and melting be~aviorof water adsorbed on fumed silica? it isvery difficult to answer this question no studies exist on the dissolution of silica from the vapor phase. Nevertheless, some reasonings can be drawn from the mechanism of dissolution of silica in water. Regardless of the type of silica involved, the dissolution process requires the presence of catalyst which can be chemisorbed and which can increase the co-ordination number of the surface silicon atom to more than four [41]. In water such catalyst is the hydroxyl ion, OH" [83]. Since in the case of fumed silica the fraction of the surface receptive to water smaller than 1/4 of the argon surface [16,85,86], the microdroplets formed on the statistically distributed groups cannot have the same contact with the silica surface in the case of fully hydroxylated surface. Therefore, the number of events where OH- ions can reach surfacesilicon atoms and be adsorbed will be smaller too. result the concentration of the monosilicic acid Si(QH)4 will also be smaller than in

the case of fully hydroxylated surface. Therefore, in the case of water adsorbed on fumed silica the amount of the dissolved material cannotbe large enough to significantly disturb the freezing and melting behavior of the adsorbed water.

From these discussions in the framework of a model in which a reduced amount of water adsorbed on fumedsilica gives rise to smaller microdroplets, the observed reduction in the enthalpy of fusion can be considered as size dependent [25]. Of course, the statement would be stronger if the effect could be observed together with the direct measurement of the dropletsize. Unfortunately, direct measurement of very small water droplets is very difficult, or even impossible: the microdroplets will easily evaporate during the measurements. It should be noted in passing that the reductionin the enthalpy of fusion A H in small systems mayalso be responsible for the soft phase transitions, the width of which depends on A H Approximate radii of microdropletshave been calculatedfromtheGibbsThomson equation, Eq. (l), using thecorrected and bulkvalues of AHm [25]. The first sets of radii were about two times larger than those calculated using the bulk enthalpy of fusion. Using both sets of radii, the fractions of silica surface receptive to water have been calculated for the adsorbed weights 2.1 and 1.7 mg. The obtained valueswere 0.1 8 and 0.22 for the first set of radii, and 0.41 and 0.45 for the second ones. It is seen that the first values are in a good agreement with the fractions of surface receptive to waterexperimentallyobtainedfor this type of silica, 0.18-0.25 16,861. This agreement iseasy to explain remembering that at rehydration of the surface ted c and, therefore, microdroplets were formedmainly on theparent At larger RH, size of microdroplets was larger andrehydrationproc probable e that, in turn, resulted in increasil~g the size of droplets. Inthis case capillarity condensation led to increasing the fraction of surface covered with water,which was difficult to calculate 1251. Another interesting moment concerning the experimental enthalpies is that the difference AHEl AB; decreased with decreasing the transition temperatures. This is in contrast to the bulk water when decreasing temperature leads to increasing the AH!. The origin of this phenomenon might be due to the difdifference AH: ferent response of microdroplet and bulk water to supercooling. It is known that with decreasing temperature the structure of water becomes less dense and more and more ice-like. In th case ofvery small droplets formed on the disordered arrangement of the groups the large surface-to-volume ratio and the influence of the surface could additionally change the structure of water, which, for instance, could become less dense in the comparison with bulk water. As result a of this additional decreasein density the specific heat of water could be smaller, i.e., it could be closer to thatof ice. According to Kirchhoffs equationsuch a reductionin the difference between the specific heats of water and ice results in the smaller tel~peraturedependence of the enthalpy of fusion for very small droplets.

The heterogeneous mechanisms, responsible for virtually all of the ice formation in natureinvolve very complexprocesses which have not been satisfactorily

explained [l411 and, probably, cannot be expressed theoretically [27]. The reason is that each substrate (surface) possesses its own individual physicochemicalproperties, which determinehowwatermolecules interact withthesubstrate.This interationplaysadominant role in theformation of the ice phase. Inother words, the mutual configuration of the electrical field resulting from the interaction between the substrate and water molecule should be suitable for the formation of an ice embryo.Therefore, necessary conditionforunderstanding heterogeneous ice formation is knowledge of the surface properties of the substrate. Silica particles can serve as a suitable substrate for the study of heterogeneous ice nucleation, since levelof knowledge of its surfaceproperties is sufficiently high and, on the other hand, it is readily amenable to modification. It was found long that the partially hydrophobic surface of silica obtained by annealing of the precipitatedsilica HiSilcan possess a goodice nucleating ability [17]. Seeding of water droplets in a cloud chamber showed that a large fraction of droplets froze at temperatures warmer than10°C. On the other hand, the partially hydrophobic surface of pyrogenic silica CaboSil was unable to nucleate ice at that temperature, which might seem strange at first sight. Also modified HiSil was a poor ice nucleant when its silanol number was 2.72, i.e., similar to that of CaboSil, 2.19. The good nucleant had fraction of surface receptive to water of approximately half theavailable surface area,i.e., two timeslarger than that of CaboSil. In this case, besides the isolated -OH groups, a considerable part of the surface hydroxyl groups are hydrogen bonded. These results indicate that to be a good ice nucleant silica must possess someoptimumratio between the isolated and hydrogen-bonded -OH groups. Heterogeneous mechanism of the ice formation on the surfaces of good nucleating silica can be described follows. Since the fraction of the surface receptive to water is only 1/2, the concentration of -OH groups is too small to form a continuous oriented film. On the other hand, the concentration is large enough for the formation of hydrophilic patches of different sizes. Both the hydrophilic patches and the isolated hydroxyls are still statistically distributed. Microdroplets formed on the surface will have a larger contact area with the silica surface than in thecase of asmaller silanol number.Watermoleculesinthe first layers adjacent to thesurface willpossess enhancedentropy in comparisonwiththe case of the oriented film formed on a fully hydroxylated surface. This entropy originates from the disordered configuration of electrostatic fields caused by interaction of thedisorderedsurface -OH groupswithpolarwatermolecules.A factor which plays a crucial role in initiating the ice is the presence of the isolated -OH groups. indeed, as was discussed above, water molecules adsorb oxygendown onthe isolated -OH and, therefore, can possess acertainfreedom of motion even in microdroplets. These molecules can also possess rotational symmetry whichwas found to be an important factor for good ice nucleability, since they supply their active H-bonding groups for maximum interaction with oncoming water molecules(12,1421. The probability that, at some temperature, the arrangement of the fields at some point of the adjacent layer gives rise to the formation of an ice embryo is much larger than in the case of a highly oriented film

The concentrationof the surface -OH groups atwhich the fraction of the silica surface receptive to water is =1/2 seems to be an optimum for good a ice nucleating efficiency of silica. At larger concentrations, the of -OH will be formed, the strong and directional fields of which, caused by the oriented arrangement of molecules, prevent the formationof ice.At small concentrationsof -OH, as in the case of fumed silica, the size of the microdroplets is also small. This results in a smaller probability that the configuration of the electrical fields in the adjacent layer give rise to the formation of ice. Heterogeneous freezing of smaller microdroplets can occur at cooler temperatures if a suitable configuration of the fields in the adjacent layer can be created. Otherwise, they will also freeze homogeneously. Very small clusters remain unfrozen even at very low temperatures. Since, as was discussed above, theNMR experiments on freezing ofwater adsorbed on pyrogenic silica showed that the time factor was almost negligible for nucleation of ice [91], temperature (or cooling rate) is the main factor which governs the ice nucleation rate in the system of very small water droplets.

In thefollowing section, stratospheric ozone and the role of in ozone destruction willbebriefly discussed. The role of theatmospheric acids onthephase transformations in water adsorbed on fumed silicawillbegiven in connection with the implication for the stratosphere. The possible role of aerosol silica particles in heterogeneous formation of and photodeco~positionof chloro~uorocarbons in the troposphere will be discussed.

In the atmosphere, ozone, constitutes a minute quantity. If all ozone in the atmospheric column were subjected to the mean pressure and temperature at sea level then its mean thickness would be about 0.3 cm. Approximately 90% of the ozone is present in the stratosphere with thelargest concentration between 12 and 30 km and a maximL~mat 25 km height. In spite of the minute quantity, ozone is one of the most important stratospheric constituents, absorbing solar ultraviolet radiation in the range h 290 nm, whichisvery important for protecting the earth’s biosphere. If there was noozonethentheshort-wavelength ultraviolet radiation would reach the earth’s surface and bring about a concomitant increase in skin cancerin humans, decreasein agricultural production, and disruptionof the oceanic food chain, and even climate 11431. Thus, the stratospheric ozonelayer can be regarded as a global sunscreen protecting the biosphere from possible damage by excess ultraviolet radiation [10,143]. In the stratospherea steady-state concentration of ozone is reached by a series of ~hotocllemicalreactions, involving and molecular oxygen, 02,during which ozone is formed and destroyed. Molecular oxygen absorbs the solar ultraviolet radiation in thewavelengthregion h 200 nm and photodissociatesintotwo free oxygen atoms which subsequently recombine with molecular oxygenin the

presence of another molecule M (generally nitrogen) to produce ozone. O3 molecules, in turn, strongly absorb ultraviolet radiation in the band 210 h 290 nm, which leads to ozone destruction dueto photodissociation 1431. The formation and destruction processes 02+hv 0 O+02+M”+O33” O3 0 2 + 0

yield a net natural balance of ozone in the stratosphere. Now it is supposed that the upset of the natural balance and sharp decrease in concentration of thestratosphericozoneduringthepolarwinterand spring, observed since the early 1970s,is caused by human activities, especially due to the emission into the atmosphere of industrial chlorofluorocarbons and chlorocarbons. These substances havebeen used widely in technological processes, in refrigerants, foams, spray propellants, etc., over several decades and, therefore, their concentration in the atmosphere has significantly increased. Chlorofluorocarbons and chlorocarbons are relatively nontoxic and chemically inactive species in the troposphere. Chlorofluorocarbons released to the atmosphere at the ground level are believed to diffuse (without being destroyed) to the stratosphere. In the stratosphere,chloro~uorocarbons absorb solar radiation of shorter wavelengths and photodecomposition takes place mainly in the “wind~w”between 175 and 220 nm. The result of decomposition is the generation of the free chlorine atoms (Cl) which are effective in destroying ozone (221. Since atmospheric gases, such as nitrogen dioxide(NO2) and methane (CH4), react with C10 and C1 to trap thechlorine in the inert chemical reservoirs of chlorinenitrate(CIONOz) and hydrochloric acid (HCl), ozone depletion in the gas phase is minimal [22,23,143]. Now it is well established that play a central role in the chlorine-catalyzed destruction of ozone in the polar lower stratosphere during winter and spring 1441. Surface-catalyzedchemicalreactions on surfaces of particles promote the conversion of relatively inactive stable gaseous chlorine reservoir forms, CIONOz and HC1, into a photochemically active form (Cl2), enabling the formation of the chlorine radicals, Cl and C10, that destroy ozone very efficiently. In addition, the particles facilitate the conversion of NO2 to the stable species remains mostly in the condensed phase. Reduced concentration of NO2 increases the efficiency of ozone destruction. Liquid sulfuric acid aerosols, whichconstitutethebackground stratospheric aerosols, are believed to the initial material on which PSCs are formed. Different models of the formation mechanisms of depend to a large extent on whether the background stratospheric aerosolsare frozen or liquid [l45 are frozen, i.e., represent sulfuric acid tetrahydrate (SAT), then the solid can be formed by condensation of water and nitric acid vapors on th produce solid nitric acid tryhydrate (NAT) particles. If the sulfuric acid aerosol is liquid then the Type Ib can be formed, i.e., particles represent ternary HNO~-HzSO4 solution droplets 1146-1481. Though it has beenrecognized that

nitric acid playsan important role in the formation and composition of PSCs [l49], there isstill an uncertainty about the phase and composition of the nitric acid containing Type Ia PSCs. More specifically, the particles of this type of clouds are believed to be NAT, nitric acid dehydrate (NAD), oreven an amorphous nitric acid-water mixture Chemical reactivity with respect to chlorine activation of PSC particles depends on their size and physical state which can be solid, liquid, or glassy. The physical state of particles can influence the degree of chlorine activation both directly from dissolved HCl and HOC1 and indirectly via removal of H N 0 3 from the gas phase. Chlorine activation rates increase with decreasing temperature because of the combined increases in particle surface area and the rate coefficients of many reactions Thus understanding of the formation mechanismof PSCs and the knowledgeof composition and the phase state of PSC particles is essential for the analysis of the present ozone depletion and the accurate theoretical prediction of future trends.

It has been described above that adsorbed H N 0 3 and HCl molecules, influencing the interaction between water and surfaces of fumed silica, change the amount of uptaken water vapor and can modify the adsorbed structure. It is reasonable to expect that the character of phase transformations in the adsorbed aqueous solutions will also differ from those in pure water. Below, some initial experimental results concerningthe solid-liquid phasetransitions in solutionsadsorbed on fumed silicawillbe presented. These studies have been motivated by the actual problem concerning the phase state and the formation mechanismof PSCs. Despite a significant amount of effort (a large number of field, laboratory, and modeling studies) aimed at explaining the phase stateand the formation mechanisms of PSCs at temperatures above the ice frost point 188 K depending on the stratospheric conditions l51]), a definitive answer is still outstanding. Solutions of the atmospheric acids €%NO3,HCl, and H2S04 adsorbed on the surface of fumed silica, similar to the case of pure water, have a large surface-tovolume ratio and, therefore, can serve as a simple model for study of the phase transitions taking place in the stratosphere. Varying the adsorption time and acid concentration in the solutions one can obtain populationsof microdroplets or thin films with compositions similar to that occurring in the stratosphere. On the other hand, since aerosol silica particles are often met in the stratosphere(see above) it is possible that they may serve as heterogeneous nuclei for the formation of The main feature that can be drawn from theavailable data obtainedwith D S c is that, similar to the case of pure water, the general character of the phase transitions in small systemssignificantly differs from that of bulk solutions: phase transitions spread over a broad temperature range, freezing and melting temperaturesare much below the bulk equilibrium melting point, and the experimentally observed freezing and melting enthalpies are smaller thanthat of bulk solution and decrease with decreasing adsorbed weight in the D S c sample. On the other hand, the temperatures and are lower than those observed in the case of the adsorbed pure

water, and the depression depends on concentration and type of the adsorbedacids. For instance, for equal coverages of the adsorbed pure water and binary H2QHNQ3 solution, a concentration of the adsorbed nitric acid of 1.5 lowered the transition temperatures about 9 K (see Fig. 4). For comparison, in the M/L, thereduccase of bulksolution with a larger acid concentration, 7.9 of the freezing temperature amounts to only 0.281 K [l 521. epression of in comparison with pure wateris due to thepeculiarities of the freezing process in aqueous solutions. When a solution with small acid concentration freezes, water separates out as pure ice crystals, which leads to an increased concentration of acids at the solid-liquid boundary. The diffusion of water molecules to the advancing surface of ice crystals is smaller than that in supercooled pure water dueto a nonuniform Concentration gradient, which in turn is influenced by the temperature gradientresulting from the dissipation of heat of fusion. In the case of electrolyte solutions there exists so-called freezing potentials, which result from the electrostatic effects of the nonuniform distribution of ions The ability of acids to modify the response of water to supercooling is related to the manner in which ions of the dissociated acid molecules interact with water, Since the ice-like structure is formed by hydrogen bonds,which are of electrostatic origin, ions prevent the formation of ice due to the hydration and size effects (in general, the sizes of ions differ from sizes of water molecules). In the case of the dispersed solution the influence of these factors can be enlarged by increasing surface-tovolume ratio. Another cause of the reduction in Tf is that the adsorbed acids may neutralize the ice-n~lcleatin~ sites on the silica surface (and maybe on the surface in general) depressing in such a way heterogeneous ice nucleation. The neutralization means that the adsorbed acid molecules (dissociated ions) can lower the probability of the formation of ice near the surfaceby disturbing thea r r a n g e ~ e nof t water molecules inthe ice-like structure.Thisassumption is based ontheexperimental result obtained from the adsorption of binary H2Q- NO3 vapormixture on Aerosil [91], and on an analysis of the experimental data obtained on freezing of bulk sulfuric acid solutions [154]. The adsorption experiments showed that concentration of nitric acid in the layers adjacent to the silica surface was larger than in the remote layers. Freezing of pure distilled water used in [l541 forpreparingthe aqueoussolutions of H2SQ4 occurred atabout whereasa sufficiently pure sample of bulk water can be supercooled to a temperature as low as -30°C [105]. Such a "high" freezing temperature indicates that the distilled water contained a significant number of suspended seeds (impurities) which were responsible for its heterogeneousfreezing. Decreasing of the freezing temperature withincreasing acid col~centrationin solutions could be due to neLltralization of the ice-nucleating centers on the impurities caused by the adsorbed ions of sulfuric acid. In the case of droplets of finely dispersed solutions, formed by condensation of vapors on fumed silica, some microdroplets can be free from such impurities or containtheminamountsmuchsmallerthan in thecase of bulk solutions. Therefore,thedepression of theheterogeneous freezing temperature is mainly due to the neut~alizationof ice-nucleating centers on the silica surface. Qn the other hand, at a sufficiently large acid concentration, heterogeneous formation of

hydrates can take place. Under certain experimental conditions (cooling rate,large concentration of solutes) supercooled solutions can. pass, with sufficient cooling, to the amorphous glassy state in which caseboth heterogeneous and homogeneous ice (hydrates) nucleation have been circumvented, i.e., the glass transition can occur 1551 (see below). In Fig. 6 the comparison of the curves obtained from pure water and binary H2O-HN03 and H2O-HCl solutions of low acid concentration is shown. It seen from the figure that, exceptforthe shift of the curves to low temperature region, the total character of the freezing and melting behavior is almost the same, i.e., concentration of the adsorbed acids was small enough in order to significantly change the adsorption structure of water. In the case of the system the freezing curve has astepwise shape similar to the caseof pure in the case of the H20-HN03 system the freezing curve is different. It represents likely overlapping of the two peaks corresponding to freezing out of ice (large peals) and to a eutectic mixture ice hydrate (small peak atlow temperature

T~~~~rature, Comparisonbetweenthe DSc curvesobtainedduringfreezing(upper curves) and melting(lowercurves) purewater and binarysolutions, H20H N 0 3 and H20-HC1, adsorbed on fumed silica. Weight of the adsorbed water in the DSc samples (also in the case solutions) is 5.8, and 2.4 mg, respectively. Concentration of the adsorbed HC1 is 1.4 f 0.1 wt% and €€NO37 f0.7 wt% (see Table

shoulder), The hydrate may possibly be nitric acid tryhydrate ( N ~ T or ) a higher hydrate, for instance, HN03 10H20.Since the peak spreads over the temperature range K, it does not represent freezing of a single liquid plug but rather a system of microdroplets with a narrower size distribution. The large thermal hysteresis, with Tf cooler than and the shape of themeltingendotherms, are almost the same for water and solutions. The main t h e r ~ o d y n a ~features ic of the solid-liquid phase transitions in the adsorbed binary solutions arelisted in Tables 1 and 2. Comparison of the weights of water in the DSc samples (see Table 1) shows that theobserved depressionin the transition temperatures can be caused both by the presence of acids (due to the mechanisms described above) and due to decreasing the weight of water (size of microdroplets) in the samples. In spite of the larger weight of water in the H20-HCl system and about five times smaller concentrationof the adsorbed HCl, the depression of was similar to the case of Hz0-HNO3. This shows that the adsorbed HCI molecules depress thefreezing process more stronglythan nitric acid. Such a tendency remains also for a larger concentration of the adsorbed acid (see Table 2). The data in Table 2 show that for a11 binary solutions the transition temperatures decreased with increasing concentration of the adsorbed acids. (The binary H2O-H2SO~ system on the surface of fumed silica has been obtained by the subsequent injection of gas and water vapor into the desiccator.) Of course, these limited data do not reflect the real dependence of the transition temperatures on concentration of the adsorbed acids in very small systems. In the case of bulk solutions the phase equilibrium diagrams obtained upon warming and the kinetic phase diagram obtained for binary H20-HN03 solution upon cooling have cornplex behavior Also it was found that ashift of the kinetic phase diagram I)

Acid-Water Systems: Adsorbed Weight in the DSc Samples, Concentration Freezing Tf, Crystallization and Melting of the Adsorbed Acids, Glass Transition Tg, Tm Temperatures, Experimental Freezing AH; and Melting ANnl Enthalpies

HZO-HCl

H20-HN03

4.2 2.9 7.8

1.4 f0.1 32f3 54 f 10

2.6 2.3 8.0

7 f0.7 53 rfr 5 70f l1

20 17.2

48 f 5 60 f l0

-36.8 -76.7 -150

-132 -70 -103 -36.8 -80.9

-98.9 -130

-73-99 -63 -100

-120

-75.4 -36.8 -64

-5.2 -40. 1 -112

120.13 27.57

136.26 36.09

-13 49.7

74.02 40.03

79.27 59.57

-58 48.35 -37 -80.6 51.69 30.15 59.02

obtained on freezing of micron-sized solution droplets (suspended in an organic material) had the same trend [159]. Additional measurements are needed to determine whether, in the case of submicron solution droplets, the changes in thefreezing/melting temperatures against acid concentration will have the same behavior. Similar to the case of pure water, in all binary systems there remains a relation AH: In thesampleswith low acidconcentration, in which the weight percent of water is largest, the difference between AH: and is mainly determined by a strong temperature dependenceof the enthalpy of fusion for water. In the samples with larger acid concentration thedifference can be due to a temperature dependence of the enthalpy of fusion for hydrates. Reduction in AH: and AH: for the first two samples (H20-HN03 and H20-HCl systems) can be due to the fact that enthalpy of fusion for water is larger than that of hydrates. On the otherhand,takingintoaccountthat in thesecondsamplesthe weightof the adsorbed water was smaller, the reduction could partly be due to the smaller size of microdroplets which freeze out.

lass Tr~nsitionsin It is seen from Table 2 that in the H20-HCl and 2SO4 systems with the largest acid concentration there was no freezing, but the g ss transitions (vitrification) occurred at temperatures -150°C and -120"C, respectively. In the case of 20-HN03 system, besides the glass transition at 130"C, also freezing took place with a peak at -99°C. The existence of freezing and the glass transition indicates thepresence of liquid domainswith different concentrations of HN03. This situation could be caused by the fact that at different times the diffusion of water and nitric acid molecules in void spacesof fumed silica was not the same. The change couldbe due to thePact that with time the hydrationeffect could prevent the diffusion of water molecules deep into the silica. As a result, the acid concentration was different at the top and bottom of the silica sample. In Fig. 7 anexample of the glass transition is shown for theH20-H2S04 system. The horizontal line in the figure shows that there had been no freezing during cooling (lower curve) up to a temperature -120°C at which the glass transitions occurred. During warming twocrystallization peaks (a large one at -75.4"C and a smaller one at -64°C) and one melting peak at a temperature -36.8"C are seen. Behavior of the glassy state of the H2O-HNO3 and H20-HCl systemsduring warming was similar but the numbers of crystallization and melting peaks were different, which means the formation and subsequent melting of different types of hydrate (Table 2). The natureof the glassy state is the most poorly understoodof the fundamental phases of matter and still remains a matter of controversy 1601. It has been suggested that it represents either first-order, second-order, third-order, or even no phase transition at all, i.e., kinetic freezing 11611. The glass transition has historically been viewed as an anomaly involving aspecific heat of supercooled liquid, cp, which is greater than that of the crystal. If this situation were to continue to a low temperature, the entropy of the supercooled liquid would becomeless than the entropy of the crystal. This is the so-called Kauzmann paradox or an entropy c u t u s ~ r o which ~ ~ e limits the liquid undercooling to the temperature where the liquid and crystal entropies would be equal. In all known cases the glass transition inter-

-140

-00

-60

-40

-20

0

Tempe~~ure,

DSC curves showing the glass transition during cooling (lower curve) a lization and melting peaks during warming. Cooling and warming rates, 3 min. venes and cp drops at a slightly higher temperature than where this catastrophe would occur [l 621. Vitrification is commonly understood as the process by which a liquid loses its ability to flow and becomes brittle during cooling, i.e., its viscosity becomes much larger than that of liquid in an equilibrium state. At the glass transition temperature Tg,thesupercooledstructure of liquid becomes“frozenin” and is retained unchanged at all lower temperatures. The formation of the glassy state is caused by the fact that during cooling, aqueous solutions do not tend to reach the thermodynamicstate oflowest free energy.Completeconversion (crystallization) to another, stabler state is subject to ~ i ~ e t i c [153,155]. The growth of a crystal embryo (ice or hydrates) once initiated is terminated with increasing viscosity. Thisinterrupted freezing processcontinuesduringwarming when the exothermicprocess (crystallization) starts (Fig, 7). During crystallization, molecules migratefromthesmallest crystals to larger ones, i.e., the larger crystals grow at theexpense of thesmallest crystals, which havethehigher interfacial free energy. In such a way the system reduces its total surface area and achieves its e~uilibriumstate at lower energy. The recrystallization process depends mainly on temperature. Usually glass transitions in aqueous solutions have been studied for the bulls case [163,164]. But there is little study on glass transition in highly dispersed aqu-

eous solutions. Whether the temperature of glass transition can be size dependent, similar to freezing and melting temperatures, is not clear from the available data. To elucidate thematter,thequestiondemandsfurther careful consideration. Knowledge of this would be useful for better understanding of the phase state of the particles and background stratospheric sulfuric acid aerosols, which are supposed to be either amorphous, liquid, or solid [165-1671.

size of solution microdroplets formedon the surfaceof fumed silica iscomparable to(H20-H2S04 system) or snlaller (H20HN03 system) than that encountered in the stratosphere. Taking into account the statistical origin of the freezing phenomenon, the observed freezing temperatures can be considered as thelow limit. Larger droplets withsimilar acid concentrations encountered in the stratosphere can have warmer freezing temperatures. A common conclusion that can be drawn from the published theoretical and experimental studies (both on bulk solutions and finely dispersed aerosols) is that NAT particles likely cannot be formed homogeneously above the ice frost point but a heterogeneous freezing mechanism might be possible 145,168,169]. However, attempts to identify the nature of suitable freezing nuclei for NAT gave no positive esults of the DSc measurements shown in Table 2 indicate that samples with an acid concentration of 53 f vvt% HN03 (which is close to the NAT stoichiometry, 54%) froze at a temperature192 K, that is, x 4 K above the ice frost point. Though DSC cannot distinguish between the different hydrates it seems that the solution froze as NAT, since its freezing temperature is close to that found with Fourier transform infrared spectroscopy (FTIR) in the case of freezing of a thin NAT film with asimilar concentration of nitric acid 1701. If in ourcase freezing has been induced by the silica surface then, in the stratosphere, where ultrafine silica particles constitute a considerable fraction of the particulate matter,silica particles can trigger heterogeneous freezing of NAT. The formation of crystalline p-NAT films on silicon substrates at temperatures 3-7 K above the ice frost point is also consistent with these DSC measurements if the oxide film on the substrate is taken into account Evidence that the glass transition in the sample with concentration 70 f 11 wt% HN03 occurred at temperatures as low as --13O"C, i.e., far below the characteristic stratospherictemperatures, indicates thatatmorphous HN03 solutions cannot be possible candidates for PSCs. If the peak at -99°C is due to freezing of nitric acid dihydrates (NAD) then, in principle, NAD may be formed under stratospheric condi~ionsif several factors, such as the concentration of nitric acid, size of the droplets and cooling rate, would favortheir formation. It is supposed that the background sulfuric acid aerosols, in which the concentration of sulfuric acid varies within wide limits, from 40 to 80 wt% of H2SO4, are difficult to freeze homogeneously at the stratospheric temperatures [145,171,172]. The freezing temperature of the sample with 48 f 5% (Table 2) determined from our DSc measurements showsthat larger and moredilute sulfuric acid droplets can freeze heterogeneously on the silica particles at the stratospheric temperatures. Formation of the larger sulfuric acid aerosols may be promoted by silica particles

and silicon-containing crustalparticles, which, as was discussed above, accumulate sulfuric acid on their surface [7,37,40]. Since the chlorine activation clearly requires liquid water as a reactant, it is believed that heterogeneous reactions on/in supercooled liquid droplets (Type Ib PSCs and liquid background aerosols) or on solid particles embedded in a thin quasiliquid overlayer [144,173] proceed faster than on solids [147,148]. This implies that the effective chlorine activation can take place even in the absenceof solid particles. The formation of a quasiliquid HCl solution overlayer around ice and hydrate particles starts from an i~corporation of HC1 vapor into the surfacelayer with subsequent solvation to form aqueous ions. Dissolution of HCI in almost binary H ~ O - ~ N O solutions ~ dropletsincreaseswithdecreasingtemperature 1741. Unfortunately, there are few data available on the concentration of HCI in stratospheric particles. Molina et al. [l441 assumed that in liquid droplets it should be larger than 0.01% at temperatures below 198 K if ambient partial pressure of HCl is1.5 Torr at 100 mbar.They also assumed thatthecol~centration of HC1 in a quasiliquid overlayer consisting of several molecular layers is from to 40% 144.1. Concentration of HCl in the liquid-like layer *33.5% has been estimated in the study by Tabazadeh and Turco [l751. Clearly, knowledge of the response to cooling of the very small H2O-HCl solution droplets, ratherthanbulk solutions, can beuseful for verification of the existence of a quasiliquid overlayeraround the solid PSC particles. Since the freezing temperature -76.7"C of the sample with 32 f 3% HCl was found to be within the limits of the characteristic stratospheric temperatures(see Table 2) it is possible that a quasiliquid overlayeron the surface of ice can exist at this and larger surface concentrations of HCl.

Since solar radiation reaching sea level consists mainly of wavelengths above 320 nm, which is far beyond the absorption threshold of gaseous chloro~uorocarbons, they are not absorbed by these substances and, therefore, photolytic dissociation does not occur in the troposphere [176]. On the other hand, chloro~uorocarbons have very small solubility in water and, therefore, they are not removed by rainfall from the troposphere. The insolubility in water together with chemical stability of CFC13 and CFzClzmolecules preventalso their rapid removal by dissolution in the ocean 1771. All these result in the fact that chloro~uorocarbonshave infinite tropospheric lifetime. In this context, it is reasonable to question whether there is a mechanism other than stratospheric photodecomposition for the removal of chlorofluorocarbons from the atmosphere,and (2) under what conditions the photodecomposition of chloro~uorocarbonswould also occurinthetroposphere. It is known that photodecompositionofmanyorganic and inorganicmoleculesadsorbed on oxidesurfaces at wavelengthsbeyondthegas-phaseabsorptionthreshold is of general occurrence [176]. At the beginning of the chapter it was discussed that in the atmosphere the silicon dioxide surface can be regarded as representative of

dan

many natural dusts and productsof anthropogenic activity. Therefore, it is reasonable to examine silica and silicate surfaces concerning their possible role in the photodecomposition of chlorofluorocarbonsinthetroposphere and scavenging of chlorofluorocarbons from the atmosphere. IJnfortunately, there has been little study of the subject. The available laboratorystudies showed that irradiation of chlorofluorocarbons adsorbed on the surfaceof different types of sand and fused quartz in the presence of C2H6 undergo photodecomposition at wavelengths extending up to approxiof mately 400 nrn, that is, at well beyondthegas-phaseabsorptionthreshold these substances 1761. Adsorption of CF2C12 on various samples of natural sand and synthetic amorphous silica surfaces (Aerosil 200, Kiesegel 60) revealed only a small influence of chemical composition of the sample on the amount adsorbed 178,1791. It has been detected that the formation of mineralization products such as CO2, CO, and HC1 on irradiation both of the CFC13 and CF2C12adsorbed on the surface of silica gel with ultraviolet light from a high-pressure Hg lamp and on exposure of the sample to natural solar radiation at ground level. The photodecomposition rate strongly depended on the wavelength of the incident light. ~nfortunately,their results didnotprovideinformation relating thegas-phase adsorption characteristics and surface-inducedphotodecomposition rates under the tropospheric conditions.

In this work an attempt hasbeen made toshow that the specific interaction between water and fumed silica surfaces allows one to obtain a population of finely divided pure water and solution droplets which can serve as asimple model for studyof the liquid-solid phase transitions taking place in the atmosphere. The finely divided aqueous systems obtainedin such a way are suitable for study withNMR and of the problems concerning alarge surface-to-volume ratio, the resolutionof which is important for both the fundamental and many applied interests. For example, strongly connected with the large surface-to-volume ratio problems are the problems of ice nucleation ill the atmosphere and the formation mechanismof PSCs, which play a crucial role in the stratospheric ozone depletion. There are still large gaps in our knowledge and understandingof the formationof ice from supercooled water but itseems that the correct questions arebeing asked, forinstance, questions concerningcharacter of thephasetransformations in thesystems which are between the bulk matter and small clusters where the influence of the individual molecule is large. In the framework of the model proposed for the fumed silica-water system, anomaliesduringthephasetransformationsin dropletssuchas“soft” phase transitions, depressions of the freezing and meltingtemperatures, and enthalpy of fusion in comparison with those of bulk water are explained by the very small sizeof droplets. Whether the observed reduction in the enthalpy of fusion is only inherent to the specific interaction between water and silica surfaces or is of general character,i.e., caused by the intrinsic property of the systems with a large surface-to-volume ratio, further studies are needed in order to elucidate the

matter. More elaborate experimental techniques and equipment for measuring very small amountsof heat evolved during the phase transformation in small systemsare neededin orderto excludethepossibleheat loss and assignamoreaccurate quantitative relation between the size of droplets and the enthalpy of fusion. Being widely present in the troposphere and stratosphere, isit possible that silica aerosols and crystal particles containing silicon in large amount can play an appreciable role in the different physicochemical processes taking place in the atmosphere, including the formation of Solutions to outstanding problems in understanding of the physical state and the formation mechanism of will come from the interaction of many scientific disciplines, of which those dealing with the problem of heterogeneous ice nucleation are of great importance. The initial, but important, work outlined in the chapter should highlight the need for enhanced activity in the study of systems with a large surface-to-volume ratio to further elucidate their unusual properties.

The author is indebted to the Academy of Finland and Department of Physics of the University of Helsinki for support of this work. He is also grateful to M. Mulmala, A. Laaksonen, A. R. MacKenzie, C. L). O’Dowd, J. C. Petit, and H. Balard for stimulating discussionsof its various aspects. The coverageof the materials discussed in this chapter, especially those concerning the unusual propertiesof the silica-water system, large surface-to-volume ratio problems and the formation mechanism of are far from comprehensive and many relevant papers have not been referred to. Limitations of space and time have prevented the inclusionof others, many of which are equally worthy of inclusion.

Sci. 2943565 (1994). 1. P. M. Dove. Am. r a l sF. , White and 2. P. M.Dove, in Chemical ~ e a t ~ e rSia~t egs o ~ S i l i c a ~ e ~ i n e(A. S. L. Brantley, eds.), Reviews in Mineralogy 31, Mineral Soc. Am., W~shington, DC 1995, pp. 235-290. 3. S. Joussaume. J. Geophys. Res. 95:1909 (1990). 4. S. Caquineau,A. Gaudichet, L. Comes, M.-C. Magonthier, and B. Chatenet. Geophys. Res. Lett. 25:983 (1998). 5. S.-M. Li and J. W. Winchester. Atmos. Environ. 27A32907 (1993). Gotz, E. MCszaros, and G. Vali, ~ t ~ ~ s ~Particles ~ e r i c Nuclei, Academiai 6. Kiado, Budapest, 1991, Chaps. 2,5. 7. F. P. Parungo, C. TT. Nagamoto, P. J. Sheridan, and R. C. Schnell.Atmos. Environ. 241$:937 (1990). W.Winchester, R. C. Schnell,S.”. Fan, Li, B. A.Bodhaine, P. S. 8. Naegele, A. D. A. Hansen, and H. Rosen. Atmos. Environ. 192167 (1985). 9. C.-C. Chen, L.-C. Hung, andH.-K. Hsu. J. Colloid Interface Sei. 157:465 (1993). P. Peixoto and A. Oort, Physics of Climate, American Institute of Physics, 10. New York, 1992, Chap.

11. S. E. Schollaert and J. T. Merrill. Geophys. Res. Lett. 25:3529 (1998). 12. H. R. Pruppacher and J. D. Klett, ~icrophysicsqf Cloudsand Prec~itation, luwer, Dordrecht, 1997, Chaps. 3,7,9. E. Hagen, R. J. Anderson, and J. L. Kassner Jr. J. Atmos. Sci. 38: 1236 (1981). 13. 14. A.Laaksonen, J. Hienola, M. Kulrnala, and F. Arnold.Geophys.Res.Lett. 243009 (1997). 15. U. Chen, S. M. Kreidenweis, L. M. McInnes, D. C. Rogers, and P. J. DeMott. Geophys. Res. Lett. 25: 1391 (1998). 16. D. R. Bassett, E. A. Boucher, and A. C. Zettlemoyer. J. Colloid Interface Sci. 27:649 (1968). 17. D. R. Bassett, E. A. Boucher, and A. C. Zettlemoyer. J. Colloid Interface Sci. 34436 (19'70). 18. K. Klier and A. C. Zettlemoyer. J. Colloid Interface Sci. 58:216 (1977). 19. E. Boucher, in ~ u c l e a ? i (A. o ~ C. Zettlemoyer, ed.) Marcel Dekker, Inc., New York, 1969, pp. 566-568. 20. M. Kulmala, A. Laaksonen, P. Korhonen, Vesala, T. Ahonen, and J. C. Barrett. J. Geophys. Res. 98:22949 (1993). 21 Kulmala, A. Laaksonen, R. J. Charlson, and P. Korhonen. Nature 388:336 (1997). 22. 0. B. Toon and R. P. Turco. Scient. Am. June:40 (1991). 23. P. Hamill and 0. B. Toon. Physics Today 44:34 (1991). 24. M. A. Tolbert. Science 272:1597 (1996). 25. Bogdan, M. and N. Avramenko. Phys. Rev. Lett. 81:1042 (1998). 26. A. Bogdan. J. Chem. Phys. 1061921 (1997). c s Academic Press, New York Dufour and R. Defay, T ~ e r ~ o d y n a ~ ~Clouds, 27. and London, 1963. 28. P. V. Physics, Clarendon Press, Oxford, 1974. 29. R. Defay and Prigogine, Surface Tension and A ~ ~ ~ Q r Long~ans, ~ t i o ~ , 1966. 30. L, Lai, J. Y. Guo, V. Petrova, C. Ramanath, and L. H. Allen. Phys. Rev.Lett. 77:99 (1996). 31. F. Ercolessi, W. Andreoni, and E. Tosatti. Phys. Rev. Lett. 66:911 (1991). 32. Buffat and J.-P. Borel. Phys. Rev. A. 13:2287 (1976). 33. W. C. Hinds, Aerosol ~ e c h n o l o gJohn ~ , Wiley Sons, New York, 1982, Chap. 1. 34. S. A. E. Johansson and J. L. Campbell, A ~ec~nique.for Ele~en Analysis, John Wiley Sons, Chichester, 1988, Chap. 12. 35. C. H. Twohy and B. W. Gandrud. Geophys. Res. Lett. 25:1359 (1998). 36. U. M. B i e r ~ a n n ~ Presper, T. T. Koop, J. MoDinger, J. Crutzen, and Th.Peter. Geophys. Res. Lett. 23:1693 (1996). 37. P. J. Sheridan. Atmos. Environ. 23:2371 (1989). 38. NI. Anderson. J. Colloid Interface Sci. 25:174 (1967). 39. Li and J. W. Winchester. J. Geophys. Res. 95:13897 (1990). 40. Li and J. W. Winchester. J. Geophys. Res. 95:1797 (1990). 41. R.K. Iler, The C ~ e ~ i s t Silica, ~ y John Wiley Sons, New York, 1978, Chaps. 5.6.

42. M. Neville, R. J. Quann, B. S. Haynes, and A. F. Sarofim,in Eig~zteenth Combustion Institute, Symposium international^ on Combustion, The University of Waterloo, Canada, 1981, p. 1267. Taylor, in ~ i g ~ t e e n t h S y ~ n p o s i u ~ ~ I n t e r n a t i o ~ a l ~ 43. R. C. Flagan and D. Combustion, The Combustion Institute, University of Waterloo, Canada,1981, p. 1227. A. M.McDonnell, ed.), John Wiley Sons, 44. D. W. Hughes, in Cosmic Dust New York, Chap. 1978. 45. T. Kane and C. Gardner. Science 259:1297 (1993). M.Hunten, R. P. Turco, and 0. B. Toon. Atmos. Sci. 37:1342 (1980). 46. 47. P. Hamill, A. Tabazadeh, S. Kinne, 0. B. Toon, and R. P. Turco. Geophys. Res. Lett. 23:753 (1996). 48. H. Barthel, L. Rosch,and Weis, in~rganosiliconChemistry II(N. Auner and J. Weis, eds.), VCH, Weinheim, 1995, pp. 761-778. 49 Sigma Chemical Company, P. Box 14508, Saint Louis, Missouri 63178, USA. 50. H. Barthel. Colloids Surfaces A: Physicochem. Engin. Aspects 101:217 (1995). 51. P. A. Sermon. J. Chem. Soc. Faraday Trans. I 76:885 (1980). 52. P. Hall, A. Pidduck, and C. J. Wright. J. Colloid Interface Sci. 79:339 (1981). 53. P. G. Hall, R. T. Williams, and R. C. T. Slade. J. Chem. Soc. Faraday Trans. I 813347 (1985). 54. T. Takei, M.Ataku, M. Fuji, T. Watanabe, and M.Chikazawa, in Silica 98, International Conference, Mulhouse, France, September, 1998, p. 561. 55. H. Balard, private communication. 56. M. Zaborski, A. Vidal, G. Ligner, H. Balard,E. Papirer, and A. Langmuir 5:447 (1989). 57. C. A. Angell, P. A. Cheeseman, and Tarnaddon. Science 218:885 (1982). 58. P. H. Poole, M. Hemmati, and C. A. Angell. Phys. Rev. Lett. 79:2281 (1997). 59. H. Kanno, R. J. Speedy, and C. A. Angell. Science (1975). 60. F. Prielmeier, E.W. Lang, R. Speedy, and H.-D. Liidernann.Phys.Rev. Lett. 1128 (1987). D. Bernal and R. H. Fowler. Chem. Phys. 1:515 (1933). 61. 62. N. H. Fletcher, The Chemical Physics of Ice, CambridgeUniversityPress, London, 1970, Chap. 2. 63. C. A. Angell and H. Kanno. Science l 121 (1976). 64. R. A. Robinson and R. H. Stokes, Electrolyte Solzltio~s,Butterworths, London, 1959, Chap. 1. 65. P. H. Poole, F. Sciortino, T. Grande, H.E. Stanley, and C. A. Angell. Phys. Rev. Lett. 73: 1632 994). 66. R. L. Mozzi and B. E. Warren. Appl. Crystallogr. 2164 (1969). 67. Eithenberg and W. Kauzmann, The Structure and Properties of Water, Oxford University Press, New York, 1969. 68. F.Franks (ed.), Winter: a C o m ~ r e h e ~ ' ~Treatise, ive Plenum Press, New York and London, 1981,Vols. 69. M. L. Hair, Infra-red Spectroscopy in Surface Chemistry, Arnold, London, 1967. 70. Clifford, in Water: a C o ~ p r e ~ e n s i vTreatise e (F. Franks, ed.), Plenum Press, New York and London, 1975, Vol. 5, Chap. 2.

71. W.Derbyshire,in Water:a Treatise (F. Franks, ed.),Plenum Press, New York and London, 1982, Vol. 7, Chap. 4. 72. K. Overloop and L. Van Gerven. J. Mag. Res. Ser. 101:179 (1993). 73 K. Overloop and L. Van Gerven. J. Mag. Res. Ser. 101:147 (1993). 74. R.T. Pearson and W. Derbyshire. J. Colloid Interface Sci. 46:232 (1974). 75. F. S. Baker and K. S. W. Sing. J. Colloid Interface Sci. 55505 (1976). 76. W. Hertl and M. L. Hair. Nature 223:1150 (1969). 77. P. Barrac~ough and P. G. Hall. J. Chem. Soc.Faraday Trans.1 74: 1360 (1978). 78. A. V, Kiselev, in The Structure and Properties OfPorous ater rials (D, H. Everett and F. S. Stone, eds.), Butterworths, London, 1958, p. 195. 79. K. J. Packer,in Progressin Magnetic NuclearResonance Spectros~o~y (J. W. Ernsly et al., eds.), Pergarnon Press, London/New York, 1967, Chap. 3. 80. Bogdan and M. Kulmala. J. Colloid Interface Sci. 191:95 (1997). 81. J. Del Bene and J. Pople. J. Chem. Phys. 52:4858 (1970). 82. H. Kistenmacher, H. Popkie, and E. Clernenti. J. Chem. Phys. 58:1689 (1973). 83. P. M. Dove. J. Geophys. Res. 100:22349 (1995). 84. J. Clark, P. G. Hall, J. Pidduck, and C. J. Wright. J. Chern. Soc. Faraday Trans. 81:2067 (1985). 85. A. C. Zettlernoyer, F. J. Micale, and K. Klier, in Water: a C o ~ ~ r e ~ e n sTreatise ive (F. Franks, ed.), Plenum Press, New York and London, 1975, 5, Chap. 5. 86. M. N.Plooster and S. N. Gitlin. J. Phys. Chem. 75:3322 (1971). 87. L. Bardotti, P. Jensen,A. Hoareau, M. Treilleux, and B. Cabaud. Phys.Rev. Lett. 744694 (1995). 88. P. Deltour, J.-L. Barrat, and P. Jensen, Phys. Rev. Lett. 78:4597 (1997). 89. Q. Du, R. Superfine, E. Freysz, and Y. R. Shen. Phys. Rev. Lett. 70:2313 (1993). 90. M. Matsurnoto and Y . Kataoka. J. Chern. Phys. 88:3233 (1988). 91. A.Bogdan, M. Kulmala, B. Gorbunov, and A. Kruppa.J. Colloid Interface Sci. 177:79 (l 996). 92. R. R. Vanfleet and J. M. Mochel. Surf. Sci. 341:40 (1995). 93. M. Wautelet. Phys. D. Appl. Phys. 24343 (1991). 94. B. E. Wyslouzil, J. L. Cheung, G. Wilemski, and R. Strey. Phys. Rev.Lett. 79:431 997). 95. K. M. Unruh, E. Huber, and C. Huber. Phys. Rev. B 48:9021 (1993). 96. M.Schmidt, R. Kusche, W. Kronrnuller, B. von Issendorff, and H. Haberland. Phys. Rev. Lett. 79:99(1997). 97. G. Bertsch. Science 277: 1619 (1997). 98. H. Rasrnussen and A. P. MacKenzie. J. Chern. Phys. 59:5003 (1973). 99. P. Pawlow. Phys. Chem. 65:545 (1909). 100. M.Wautelet. Eur. J. Phys. 16283 (1995). 101. P. R. Couchman and W. A. Jesser. Nature 269:481 (1977). 102. L. Jackson and G. B. McKenna. J. Chem. Phys. 93:9002 (1990). 103. J. Huang and L. S. Bartell. J. Phys, Chern. 99:3924 (1995). 104. L. S. Bartell and J. Huang. J. Phys. Chem. 98:7455 (1994). 105. C. Angell, in Water: a C o ~ ~ r e ~ e n sTreatise ive (F. Franks, ed.), Plenum Press, New York and London, 1982, Vol. 7, Chap. 1. 106. G. R.Wood and G. Walton. J. Appl.Phys.41:3027(1970).

107. 108. 109. 110. 111. 112. 113. 114. 115.

116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130. 131. 132. 133. 134. 135. 136. 137. 138. 139. 140.

A. A. Antoniou. J. Phys. Chem. 682754 (1964). G. Litvan and R. McIntosh. Can. J. Chem. 41:3095 (1963). G. C. Livan. Adv. Colloid Interface Sci. 9:253 (1978). M. B. Ritter, D. D. Awschalom, and M. W.Shafer.Phys.Rev.Lett.61:966 (1988). J. Warnock, D.D. Awschalom, and M. W. Shafer.Phys.Rev.B 34:475 (l 986). D. F. Brewer, J. Rejendra, N. Sharma, A. L. Thornson, and Jin Xin. Physica B (1990). D. F. Brewer, J. Rejendra, N. Sharma, A. L. Thomson, and Jin Xin. Physica B 165&166:577 (1990). D. F. Brewer, Rejendra, N. Sharma, and A.L.Thomson.PhysicaB 165&166:569 (1990). E.Molz, A. P. Y. Wong, M. H. W. Chan and J. R. Bearnish.Phys.Rev.B. 485741 (1993). D. D. Awschalom and J. Warnock. Phys. Rev. B 35:6779 (1987). R. C. Tolman. J. Chem. Phys. 17:333 (1949). D. Hare and C. M. Sorensen. Chem. Phys. 87:4840 (1987). J.-C. Li, D. K. Ross, and M. J. Benham. J. Appl. Cryst. 24794 (1991). J.-C. Li, D. K. Ross, and R. K. Heenan. Phys. Rev. B. 48:6716 (1993). A. Patrick and W. A. Kemper. J. Phys. Chem. 42:369 (1938). Rennie and J. Clifford. J. Chem. Soc. Faraday Trans. 73:680 (1977). G. D. Stein and J. A. Armstrong. J. Chenl. Phys. 58:1999 (1973). D. C. Steytler, J. C. Dore, and C. J. Wright. Phys. Chem. 87:2458 (1983). M. Dunn, J. C. Dore, and P. Chieux. J. Cryst. Growth 92:233 (1988). A. Elarby-Aoulzerat, F. Jal, C. Ferradou, J. Dupuy, P. Chieux, and A. Wright. J. Phys. Chem. 87:4170 (1983). H. W. Foote and B. Saxton. J. Am. Chem. Soc. 39:1103 (1917). R. Mu and M. Malhora. Phys. Rev B 444296 (1991). J. Warnock, D. D. Awschalom and M. W.Shafer.Phys.Rev.Lett.57:1753 (1986). J. L. Tell and H. J. Maris. Phys. Rev. B 285122 (1983). F. Spaepen and C. V. Thompson. Appl. Surf. Sci. (1989). L. A. Wilen, J. S. Wettlaufer, M. Elbaum, and M. Schick. Phys. Rev. B 52:12426 (1995). C . Dash,Haiying Fu, and J. S. Wettlaufer.Rep.Prog.Phys. (1995). J. H. Bilgram. Phys. Rep. 153:l (1987). J. C. Petit, private com~unication. H. Everett, J. M. Haynes, and P. J. McElroy. Nature 226:1033 (1970). S. L. Lai, G. Ramanath, and L. H. Allen. Appl. Phys. Lett. 67:1229 (1995). A. Lied, H. Dosch, and J. H. Bilgrarn. Phys. Rev. Lett. 72:3554 (1994). P. M. Dove and J. D. Rirnstidt, in Silica: G~oc~~~istry, ~ ~ p l i c u t i o n(P. s J. Heaney, C. Prewitt, and G. Gibbs,eds.), Mineralogical Society America, 1994, Ch. 8. P. Labastie and R. L. Whetten. Phys. Rev. Lett. 65:1567 (1990).

141. R. J. Anderson, R. C. Miller, J. L. and D. E. Hagen. J. Atmos, Sci. (1980). 142. N.Fukuta. J. Atmos. Sci. 23191 (1966). 143. B. Finlayson-Pitts and J. N.Pitts, ~ t ~ o s ~ ~Chemistry, e r i c John Wiley lk Sons, New York, 1986, Chap. 1. 144. J. Molina, R. Zhang, P. J. Wooldridge, J. R. McMahon, J. E. Kim, H, Chang, and K. D. Beyer. Science 261:1418 (1993). 145. R. MacKenzie, Laaksonen, E. Batris, and M. Kulmala. J. Geophys. Res. 103:10875 (1998). 146. K. Carslaw and Th. Peter. Geophys. Res. Lett. 24:1743 (1997). 147. K. S. Carslaw, Th. Peter, and R. Muller. Geophys. Res. Lett. 24:1747 (1997). 148. A. R. Ravishankara and D. R. Hanson. Geophys. Res. 101:3885 (1996). 149. D. R. Hanson. Geophys. Res. Lett. 17:421 (1990). 150. B. S. Berland, D. R. Haynes, K. L. Foster, M. Tolbert, S. M. George, and 0. B. Toon. Phys. Chem. 98:4358 (1994). 151. K. S. Carslaw, B. P. Luo, S. L.Clegg, Th. Peter, P. Brimblecombe,andP. J. Crutzen. Geophys. Res. Lett. 212479 (1994). 152. andb book C h e ~ i s t r yand Physics, 67th ed., CRC Press, Inc.Boca Raton, Florida, 1987, p. D-239. 153. F. Franks, in Water: C o ~ ~ r e ~ e n sTreatise, ive Vol. 7 (F. Franks, ed.), Plenum Press, New York and London, 1982, Chap. 3. 154. T. Ohtake. Tellus 4513:138 (1993). 155. A. P. MacKenzie. Phil. Trans. R. Soc. Lond. B 278:167 (1977). 156. C. Gable, H. F. Betz, and S. H. Maron. J. Amer. Chem. 72:1445 (1950). 157. F. Kuster and R. 2;. Kremann. Anorg. Chem. 41:l (1904). 158. K. Ji, F. Mahi: and C. Petit. in Atmospheric Ozone, Proceedings the XVIII Quadrennial Ozone Symposium(R. D. Bojkov and G. Visconti, eds.), L’Aquila, Italy, 12-21 Sept. 1996, p. 923. 159. K. Ji, C. Petit, P. Negrier, and Y. Haget. Geophys. Res. Lett. 23:98 1 (1996). 160. S. D. Bembenek and B. B. Laird. Phys. Rev. Lett. 74:936 (1995). 161. H. Fecht, Z. Fu, and W. L. Johnson. Phys. Rev. Lett. 64:1753 (1990). 162. N. Birge and S. R. Nagel. Phys. Rev. Lett. 542674 (1985). 163. C. A. Angell and E. J. Sare. Chem. Phys. 49:4713 (1968). 164. C. Angell and E. J. Sare. J. Chern. Phys. 52:1058 (1970). 165. A. Tabazadeh, B. Toon, and P. Hamill. Geophys. Res. Lett. 22:1725 (1995). 166. R. Zhang, M.-T. Leu, and M. J. Molina. Geophys. Res. Lett. 23:1669 (1996). 167. S. K. Meilinger, T. Koop, B. P. Luo, T. Huthwelker, K. S. Carslaw, U. Krieger, P. J. Crutzen, and Th. Peter. Geophys. Res. Lett. 22:3031 (1995). 168. Bertram and J. J. J. Geopllys. Res. 103:3553, 13261 (1998). 169. T. Koop, B. Luo, U. Bierrnann, P. J. Crutzen, and Th. Peter. Phys. Chem. 101:11 17 (1997). 170. A. Middlebrook, B. G. Koehler, L. S. McNeill, and M. Tolbert. Geophys. Res. Lett. 19:2417 (1992). 171. S. E. Anthony, R. T. Tisdale, R. S. Disselkarnp, M. Tolbert, and J. C. Wilson. Geophys. Res. Lett. 22:1105 (1995).

172. M. L. Clapp, R. F. Niedziela, L. J. Richwine, T. Dransfield, R. E. Miller, and D. R. Worsnop. Geophys. Res. 1028899 (1997). 173. P. Abbatt, K. D. Beyer, F. Fucaloro, J. R.McMahon, P. J. Wooldridge, J. Molina. J. Geophys. Res. 97: 15819 (1992). R. Zhang, and 174. A. Tabazadeh, R. P. Turco, and M. Z. Jacobson. Geophys.Res.99:12897 (1994). 175, A. Tabazadeh and R. Turco. J. Geophys. Res. 12727 (1993). (1977). 176. P.Ausloos, R. E,Rebbert, and L. Glasgow. J. Res. Nat. Bur. Stand. 177. M. Molina and F. Rowland. Nature 249:810 (1974). 178. Gab, Schrnitzer,H. Thamrn, and F. Korte. Angew. Chern. Int. Ed. Engl. I7366 (1978). Gab, J. Schmitzer, H. W. Tharnm, H. Parlar, and F. Korte. Nature 270:331 179. 977).

This Page Intentionally Left Blank

Acceptor number of molecular probes for IGC, 213 Acid base properties of silica, 21 1--214 Adsorbed aqueous solutions freezing behavior of, 730-734 Adsorption and disorder and confinement, 139-141 and disorder and curvature, 141-143 Adsorption energy distribution of, 214-21 6, 220-224, 229 estimated by computer simulation, 234 Adsorption energy distribution functions principle for determination of, 214-2 16 of silylated silica, 230 of various silica samples, 220-224 Adsorption isotherms by IGC, 214 of neutral and charged polymers, 483-488 Alkaline-earth cation adsorption on silica, 309-3 I role of pH on adsorption, 3 l 1-3 13 Amorphous silica crystallographic surface models for, 299

[Amorphous silica] surface structure of, 299 Application of immobilized biocatalysts, 546-552 of immobilized enzymes, 543-546 of modified silica as catalyst support, I00 of modified silica for separation of carbohydrates, 109 of modified silica for separation of fullerenes,107 of modified silica in liquid chromatography, 102-1 14 of modified silica for separation of proteins, l06 for separation of polyaromatics, 106 Biologically active compounds immobilized on silica, 556-557 Catalysis paradigm of heterogeneous, 370-372 modified silica for, 57, 100 C~lorofluorocarbons and silica in stratosphere, 737-738 Chromatography kinetic aspects of, 573-577

7

~Chromatography] [Computer simulation] optimum resolution of, 578-579 of silica surface, 233 properties of columns for, 585-587 of specific interaction, 321-323, 328 properties of silica particles for, of covered silica, 228-231 579-585 separation processes of, 572-573 Dissociation constants with silanized silica, 591-593 of silanols in organic-water mixtures, t~ermodynamicsaspects of, 568-571 359-362 Colloidal silica Dissolution preparation of aqueous suspensions of, enhanced by alkaline ions, 290 475 Donor number Colloidal silica systems of molecules, 21 3, 347 equilibrium properties of, 332-334 DSC neutron scattering on, 336-339 measurments on adsorbed water, osmotic pressure ~easurementson, 7 19-726 334-336 Colloids Electrical double layer DVLO theory for, 472-474 of silica-water interface, 280 flocculation kinetics of polymerconcept of, 323 covered, 475 Electrokinetic properties steric stabilization of polymer-covered, in almost dry organic solvents, 474 345-350 Computer simulation in humid organic solvents, 350-355 of acetone molecular dynamics in measurement by electrophoresis, cylindrical pores, 265-27 1 324-328 of adsorbed molecules, 142 in water rich organic solvents, of adsorption and condensation 355-356 processes,145 Electron microscopy of CO2 addition to ion bombarded scanning, 580 silica, 386-387 of silica aggregates, 153 in comparison with IGC results, of silica fume, 152 234-237 of silica gel, 130, 131, 154 of cyclohexane in cylindrical pores, of silica glass, 135, 139 25 1-257 Enzymes of fullerene in cylindrical pores, 257 activity and stability of immobilized of fullerene diffusion, 257 enzymes, 540-542 of li~u~d-plastic transition of adsorption and chemisorption cyclohexane in silica cavities, equilibrium, 552-556 27 1-274 analytical application of immobilized of molecular dynamics dependence on enzymes, 543-546 pore size, 259-260 catalytic activity of immobilized of molecular dynamics of molecules in enzymes, 546-552 pores, 247-25 1 direct attachment on aminoof molecules in closed pore, 245 organosilicas, 526 of molecules in model pore, 244-245 immobilization on activated aminoof molecules in open B pore, 246 organosilicas, 528

[Enzymes] influence of geometrical characteristics of organosilicates on kinetics of immobilization of, 537-539 influence of pH on immobilization of, 539 physicochemical properties immobilized enzymes, 537-543 ESR of adsorbed polymer, 478 EXAFS for heavy metal cations sorption on silica, 414-4 16 Flocculation as a function of polymer concentration, 455 Flotation of quartz, 457-460 Fullerene computer simulation of diffusion in cylindrical pores of, 257 grafting on silica of, 95 separation, 1l l Fumed silica effect of acids on water uptake by, 706-709 surface properties of, 695-697 Glass transition of adsorbed acid-water systems, 733-736 Heavy metal cations adsorption on other oxides, 436-437 experimental (macroscopic) methods to study adsorption 404-412 experimental (spectroscopic) methods to study adsorption of, 412-416 factors influencing adsorption 422-428 influence of electrolytes on adsorption of, 405-41 sorption isotherms of, 417-421 sorption on silica, 399-404 spectroscopies for study 412-41 6

(Heavy metal cations] surface complexation models for, 428-436 Historical review classical period, 1 mesostructured silica, 2-4 perspectives, 4 porous silica, 2 silica gel chemistry, 2 Hydrated ions Competition for adsorption sites, 306-308 Hydrogen bonding evidenced by DRIFT, 49-50 H bond energy, 78, 86-87 with water, 83-86 Hydroxide ion adsorption, 305-306 Hydroxyl ion mechanisms of adsorption 305 Hydroxylated silica surface bound water on, 700-701 Hydroxylation of pyrogenic silica, 40-43 in relation with the preparation method, 40-45 of silica gel and precipitated silica, 43-45 Image analysis, 134-1 39 Infra-Red spectroscopy of adsorbed polymer, 477 of adsorbed water, 67-68 after adsorption various compounds, 75-77 after deuterium exchange, 65 after Pt deposition, 100-101 after silane treatment, 92 of chlorosilane modified silica,90-92 DRIFT, 50 of heat treated silica, 19 of ion-bombarded silica, 383-385 low-wave number spectra of silica, 15-18 of loaded silica, 58 of pyrogenic silica, 65-66

I

[Infra-Red spectroscopy] of silanol groups, 10-15, 64-68 of silica in vacuum, 11-1 5 of siloxane groups, 18-20 Inverse Cas Chromatography for acid-base surface properties measurements, 21 1--2 14, 220, 227 advantages and limitations of, 206-207 application to various silicas, 216-23 1 and computer modeling, 231--238 determination of nanorugosity by, 208 determination of surface energy by, 207, 210 in finite concentration conditions, 214-216 in infinite dilution conditions, 207-214 principle of, 206 Ion exchange properties, 46-48 Ion-bombarded silica addition Ar, 381-383 reactions with, 387-396 computer simulatioll of 386-387 and expected adducts, 382-383 and expected CzH4 adducts, 387 imparted damage of, 377 IR examination of 383-385 IR examination of 389-390 oxygen-bridge vacancy of, 378-380 radiation defects of, 380-381 silicon-link vacancy of, 378 theoretical study of 385-386 theoretical study of C2H4,391-396 Ionization constants determination of, 320-32 Large surface-to-volume ratio general description problems due to, 710-713 in ice nucleation, 713-716 MCM-41, 665-685 (see Mesoporous silica) Mesoporous silica dense silane monolayer formation on, 682-683

[Mesoporous silica] direct silanization of, 679-68 1 historical review of, 2-4 hydration of calcined, 68 1-682 inorganic functionalization by cocondensation, 669--671 inorganic functionalization by postcalcination metalization, 672-674 inorganic functionalization by surfatant displacement, 67 1-672 and interfacial chemistry, 666-669 organic functionalization by cocondensation, 674-676 organic functionalization by surfactant extraction, 676-678 porosity of, 193-1 96 self-assernbled silane monolayer on, 678-683 state of calcined surface, 678-679 Molecular dynamics simulation of molecules in cylindrical silica pores, 25 1-271 of molecules in model pores, 244-251 Nanoscale texture and surface roughness, 150-1 54 NMR 'H spectroscopy of silica, 20-26 spectroscopy of silica, 29-30 29Si spectroscopy silica, 26-29 of adsorbed polymer, 477 of adsorbed water, 718 using CRAMPS, 24-26 Organosilicas activation of amino-organosilicas, 526-533 activation of silica carrying silicon hydride groups, 535-536 activation of vinylsilicas, 533-534 methods for surface activation of, 524537 Pathogenicity due to inhaled particles, 646-648

[Pathogenicity~ health effects of silica, 646 Phase transition of cyclohexane in silica cavities, 27 1-274 Point of zero charge after metal ions exchange, 46-49 determination of, 316-3 18 Polydimethylsilo~ane adsorption of bimodal polymer, 625 adsorption kinetics of, 6 10-61 2 adsorption of pure polymer, 624 chain adsorption law, 623 chain anchoring on silica, 604-606 end grafting on silica, 61 6-6 17 exchange in adsorbed layer, 612-61 6 formation of surface layers, 598-600 gelation in presence of silica, 635-641 initial state of adsorption of, 626 kinetics of chain adsorption, 630-634 network formation with silica, 627-630 parameters controlling surface excess Of,600-604 thickness of adsorbed layers in good solvent, 606-6 10 Polyelectrolyte adsorption of, 468-472 adsorption measurement of, 476-480 general features of adsorbed layers of, 463464 theoretical basics of adsorption of, 468472 Polymer adsorption of, 460-468 adsorption in relation with polymer conformation, 456 adsorption isotherms of neutral and charged, 483-488 adsorption ~~easurement of, 476-480 adsorption mechanism of, 454456 displacement and exchange of adsorbed, 502-504 fraction of adhered segments of, 488-494 general features of adsorbed layers of, 463-464

[Polymer] theoretical basics of adsorption of, 464-468 thickness of adsorbed layer of, 478-479, 494-500 Polymer-Surfactant mixtures adsorption on silica of, 456-460 Porosity of heated silica, 196-200 of mesoporous silica, 193-196 of mixed silica samples, 192-193 and molecular dynamics simulation, 243-275 of molecular sieves, 193-196 of nuclear membranes, 182-1 86 of silica gels with bonded alkyl grafts, 186-1 92 of silica glass, 136 of silica mixes, 192-193 from temperature-programmed desorption, 169-1 82 of various silicas, 182-1 92 Porous silica activated by ion bombardment, 381 cation adsorption on, 309-3 13 Proton association constants experimental determination of, 304-305 theoretical determination of, 302-304 Quartz ball milling of, 447 flotation of, 457-460 wetting, in presence of surfactant of, 446 zeta potential of, 444 Reactive silica by crushing, 374 from dispersed silica, 372-373 by pyrolysis of methoxylated silica, 373 by radiation damaging, 376 by straining of silica network, 375 Roughness experimental methods and evidence for, 154-163

(Roughness] and nanoscale texture, 150-1 54 and scale invariance, 146-1 50 Silane treatment of mesoporous silica, 588-59 1, 679-68 1 Silanol groups activity of, 45 content, 72-74 exchange properties of, 46-49 of heat treated silica, 12,16, 71 Hydrogen bond energy, 86-87 by IR spectroscopy, 10 by Raman spectroscopy, 19 reactivity of, 52-54 in relation with surface area, 70 types of, 40 Silica dispersion aggregation kinetics of, 510 rheology of, 482 stability characteristic parameters for, 504-5 10 stability-~occulationmeasurements of, 480-482 stabilization of, 5 10-5 2 Silica dissolution depoly~erizationmechanism of, 301 enhancement by alkaline-earth groups, 290 Silica gel chemistry history 2 Silica in atmosphere of antropogenic origin, 694-695 of extraterrestrial origin, 695 natural origin, 694 Silica morphology, 695 Silica surface structure and hydration strength of ions, 308 and water adsorption, 301 Silica toxicity mechanisms 648-655 modulation by adsorption of inorganics, 659-661

[Silica toxicity] modulation by adsorption of organics, 657-659 role of surfactant in, 653-657 Silica water chemistry, 442444 Silica-water interface chemistry of, 442446 dissolution studies, 288-292 electrical double layer in, 280 electrical double layer of, 280-284 generalities on, 697-698 model 278-284 properties of water at, 284-285 structure of, 279-280 thickness structured water at, 285-288 Small-angle neutron scattering on colloidal silica systems, 336-339 Small-angle X-ray scattering on porous silica glasses, 157-161 Sorption of heavy metal cations, 399-404 Stratosphere and silica particles, 736-737 Surface charge adsorbed polymer, 479 at silica-water interface, 281-284 density versus pH, 320 determination from titration data, 318-320 effect of deuterated water on, 364-365 effective charge concept of, 331-332 in organic water mixtures, 356-359 origin 343-345 Surface chemistry of silica adsorbellts, 63-72, 584-585 Surface energy of heat treated silica, 224 nonspecific component of, 207-208, 216-2 18, 224-226 of silica grafted with alkyl chains, 224225 of silylated silica, 225-226 specific componentof, 2 l 1,219-220

[Surface energy] variation with surface coverage of silica, 227-228 of various silica samples, 217 Surface functionality of silica, 36-38 Surface ionization models of, 313-316 by adsorption, 97-100 Surface modification with alcohols, 55, 224 with boron trichloride, 54, 55 chemical, 87-93,587-591 with chlorosilanes, 88-93 for chromatographic separations, 54-59,579-594 with deposited C layers, 93-95 with diols, 224 by end grafting of polydimethylsiloxane, 616-6 17 by grafting of fullerene, 95-97 with silane, 55, 88, 99, 225-230, 589-591 of silica for HPLC, 587-591 of silica with mercapto groups, 534 Surface morphology basic notions of, 122 curvature and confinement in relation with, 122-128 curvature and connectivity in relation with, 128-1 34 by image analysis, 134-1 39 mathematics of, 123-1 34 morpholo~yindices measured by XGC, 218 Surface roughness and nanoscale texture, 150-1 54 and scale invariance and hierarchy, 143-1 50 Surface rugosity by IGC, 208-210,218-219 Surface titration for surface charge determination, 318-320

Topology index of molecular probes for IGC, 209, 219 Types of surface groups and activity of, 45-59, 584 Vibrational spectroscopy genera1 features and surface silanol groups, 10-1 5 low-wave number spectra, 15-1 8 and siloxane sites, 18-20 Water effect of acids on adsorption of, 706-709 droplets formation on silica surface, 703-706 freezing on silica surface, 717-726 ice formation on partly hydrophobic silica surface, 726-728 interaction with silica surface, 699-706 IR of adsorbed, 67-68 models of behavior at silica surface, 289-290 properties at interface, 284 similarity with liquid silica, 698-699 thickness of structured water at interface, 285-288 Water adsorption interaction with silica, 699-706 isosteric heat of, 82 isotherms, 8 1-84 Zeta potential of polymer covered silica particles, 478,483, 501-502 of quartz, 444 of silica in organic medium, 362 of silica in presence of Pb, 41 1412 of silica in presence of surfactant, 453454