Group 13 Chemistry III. Industrial Applications

105 Structure and Bonding Managing Editor: D.M.P. Mingos

Editorial Board: A.J. Bard-I.G. Dance. P. Day. J.A. Ibers- T. Kunitake T.J. Meyer-H.W. Roesky. J.-P. Sauvage A. Simon-F. Wudl

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Group 13 Chemistry III Industrial Applications Volume Editors: H.W. Roesky and D.A. Atwood

With contributions by D.A. Atwood, P.H.M. Budzelaar, D.J. Linton, A.R. Hutchison, D.M. Schubert, G. Talarico, W. Uhl, A.E.H. Wheatley, Y. Zhang

Springer

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Editorial Board Prof. Allen I. Bard Department of Chemistry and Biochemistry University of Texas 24th Street and Speedway Austin, Texas 78712, USA E-rnaih [email protected]

Prof. fan G. Dance Department of Inorganic and Nuclear Chemistry School of Chemistry University of New South Wales Sydney, NSW 2052, Australia E-mail: [email protected]

Prof. Peter Day, FRS Director and Fullerian Professor of Chemistry The Royal Institution of Great Britain 21 Albemarle Street London WIX 4BS, UK E-mail: [email protected],uk

Prof. James A. Ibers Department of Chemistry North Western University 2145 Sheridan Road Evanston, Illinois 60208-3113, USA E-maih [email protected]

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Preface

The present issue of Structure and Bonding is dedicated to applied group 13 chemistry, particularly for the elements boron and aluminum, and to a lesser degree gallium and indium. Although boron is a trace element (0.01 g kg -1) in the earth's crust, it has been concentrated in a few locations by geochemical processes and is relatively easy to mine as borax. Aluminum, on the other hand, is the most abundant metal in the earth's crust (82 g kg -1) and dispersed widely throughout the globe. Thus, boron and aluminum are readily available and their associated products or compounds are usually inexpensive and thereby easy to commercialize. The chapters were chosen to encompass both applied and fundamental aspects of their subiects. The first chapter 'Borates in Industrial Use' provides a complete, and perhaps, quintessential, coverage of compounds containing boron-oxygen bonds. In the chapter Schubert explains the close relationship between the basic properties of the boron compounds and their associated uses. The remaining four chapters focus, to some degree, on aluminum. Since a great deal of literature exists in this area, these chapters are more focused on areas of emerging utility, and contain a great deal of fundamental information. Uhl's contribution in Chapter 2 provides basic synthesis and structural information for aluminum and gallium hydrazides. These types of compounds are being explored as potential molecular precursors to metal nitrides such as the important blue-green laser material gallium nitride. In the third chapter Wheatley and Linton describe new developments in the remarkable chemistry of aluminum and oxygen. While compounds featuring an aluminum-oxygen bond find widespread utility (as catalyst supports and co-catalysts, for example) there is still a great deal that is not understood in how these compounds form. This chapter provides the necessary basic information upon which further applications may be developed. Aluminum reagents are used in a wide range of catalytic and polymerization processes. Karl Ziegler's 1950s discovery of the now legendary 'Aufbau Reaktion' is one dramatic example of aluminum used in this way. In this and many other reactions the ability of the aluminum atom to undergo insertions or hydride transfer reactions is of key importance. Budzelaar and Talarico, in Chapter 4, describe how both of these processes occur, and in so doing provide a great deal of guidance for the use of alkyl aluminum reagents for much broader applications.

VIII

Preface

Chapter 5 surveys the remarkable structural diversity associated with compounds containing five-coordinate group 13 elements. There is a growing awareness that many of these compounds can be used as Lewis acidic reagents or catalysts. Their coordination numbers are more fixed in comparison to 'traditional' compounds with four-coordinate elements and they are far less air sensitive. Thus, they should be more amenable to applications. Overall, these five chapters should provide a general reader with a broad understanding of the uses, both current and potential, associated with the majority of the group 13 elements. The chapters were written to include basic background information for the edification of students and industrial scientists, as well. November 2002

H.W. Roesky D.A. Atwood

Contents

Borates in Industrial Use D.M. Schubert . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Aluminum and Gallium Hydrazides W. Uhl . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

41

The Synthesis and Structural Properties of Aluminium Oxide, Hydroxide and Organooxide Compounds D.J. Linton, A.E.H. Wheafley . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

67

Insertion and ~-Hydrogen Transfer at Aluminium P.H.M. Budzelaar, G. Talarico . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

141

Compounds Containing Five-Coordinate Group 13 Elements D.A. Atwood, A.R. Hutchison, Y. Zhang . . . . . . . . . . . . . . . . . . . . . . .

167

Author Index Volumes 101-105 . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

203

Subject Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

205

Contents of Volume 103 Group 13 Chemistry I Fundamental New Developments V o l u m e Editors: H.W. Roesky, D.A. A t w o o d Structure and Bonding in Boron Containing Macrocycles and Cages - Comparison to Related Structures with Other Elements Including Organic Molecules H. H6pfl Multiple Bonding Between Heavier Group 13 Elements P.P. Power The R2M+ Group 13 Organometallic Fragment Chelated by P-Centered Ligands L. Mahalakshmi, D. Stalke Synthesis, Structure and Reactivity of Group 13/15 Compounds Containing the Heavier Elements of Group 15, Sb and Bi S. Schulz

Contents of Volume 104 Group 13 Chemistry II Biological Aspects of Aluminum V o l u m e Editors: H.W. Roesky, D.A. A t w o o d Acute Aluminum Intoxication K. Berend, G.B. van der Voet, F.A. de Wolff A New Effect of Aluminum on Iron Metabolism in Mammalian Cells S. Oshiro The Complexity of Aluminum-DNA Interactions: Relevance to Alzheimer's and Other Neurological Diseases S. Anitha, K.S.J. Rao Aluminum: Interaction with Nucleotides and Nucleotidases and Analytical Aspects of Its Determination M.R.C. Schetinger, V.M. Morsch, D. Bohrer Aluminofluoride Complexes in the Etiology of Alzheimer's Disease A. Struneck~i, J. Pato~ka Fluoroaluminate Chemistry B. Conley, D.A. Atwood

Borates in Industrial Use David M. Schubert U.S. Borax Inc., 26877 Tourney Road, Valencia, California 91355, USA e-mail: [email protected]

Boron compounds find extensive use in a wide range of industrial applications, nearly all involving boron-oxygen compounds. Although quite diverse, these end uses depend on the same fundamental aspects of the structure and bonding patterns of boron. The most important industrial uses of boron compounds are discussed along with recent developments in the understanding of fundamental chemistry of crystalline and vitreous borates that underlie these applications. The formation of ester linkages to boron leads to industrial uses and provides a basis for the biological interactions of boron. Full recognition of the essential role that boron plays in biological systems has only come about in the last few years. The use of boron in the manufacture of glass and other vitreous products accounts for more than one half of all boron use. Perborates, agriculture, wood preservation, and fire retardancy are also important application areas. New applications are being developed for the use of borates in the pulp and paper and ceramics industries, among others. Keywords: Boron, Borate, Glass, Ester, Bioessentiality

1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Aqueous Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Crystalline Borates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Vitreous Borates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Organic Ester Formation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Biological Aspects of Boron . . . . . . . . . . . . . . . . . . . . . . . . . . .

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6.1 6.2 6.2.1 6.2.2 6.3

Natural Products . Bioessentiality . . . Plants . . . . . . . . . Animals . . . . . . . Toxicity . . . . . . .

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Industrial Boron Products . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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7.1 Refined Borates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.1.1 Borax Pentahydrate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.1.2 Borax Decahydrate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Structure and Bonding, Vol. 105
Springer-Verlag Berlin Heidelberg 2003

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D. M. Schubert

7.1.3 7.1.4 7.1.5 7.1.6 7.1.7 7.2 7.2.1 7.2.2 7.2.3 7.2.4 7.2.5 7.2.6

Boric Acid . . . . . . . . . . . . . Anhydrous Borax . . . . . . . . Boric Oxide . . . . . . . . . . . . ‘Disodium Octaborate’ . . . . . Zinc Borates . . . . . . . . . . . . Beneficiated Mineral Borates Colemanite . . . . . . . . . . . . . Tincal . . . . . . . . . . . . . . . . . Kernite . . . . . . . . . . . . . . . . Ulexite . . . . . . . . . . . . . . . . Hydroboracite . . . . . . . . . . . Metamorphic Borates . . . . .

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Major Industrial Uses of Borates . . . . . . . . . . . . . . . . . . . . . . . .

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8.1 8.1.1 8.1.2 8.1.3 8.2 8.3 8.3.1 8.4 8.5 8.5.1 8.5.2 8.6

Vitreous Applications . . . . . . . . . Fiberglass . . . . . . . . . . . . . . . . . . Specialty Glass . . . . . . . . . . . . . . Ceramic Glazes/Porcelain Enamels Agricultural Use . . . . . . . . . . . . . Biostatic Use . . . . . . . . . . . . . . . . Wood Preservation . . . . . . . . . . . Fire Retardants . . . . . . . . . . . . . . Perborates . . . . . . . . . . . . . . . . . Detergency – Consumer Bleaching Detergency – Laundry Builder . . . Emerging Applications of Borates

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1 Introduction Boron finds use in a wide range of industrial applications, the vast majority of which involve boron-oxygen compounds. This class exhibits a rich structural chemistry, featuring both crystalline and non-crystalline solids of industrial importance, and complex solution chemistry. The full extent of the role of boron as an essential element in fundamental life processes has only become clear during the past decade. The emerging biological roles of boron are relevant to both the commercial and environmental aspects of boron science. Although boron-oxygen chemistry is diverse and complex, it is ultimately based on the electrophilic character of trivalent boron. This review explores the chemistry of boron as it relates to its large-scale technological roles. Although boron has a relatively low natural abundance (ca. 0.001% of the earth’s crust) it is widely dispersed in the environment, occurring in rocks, soils, and natural waters at low but significant levels [1]. Owing to boron’s unique

Borates in Industrial Use

3

properties, it is more susceptible to fractionation by geological processes and can become concentrated to a greater extent than other elements of similar natural abundance. With the exception of a few rare fluoride minerals, boron always occurs on the earth in combination with oxygen, either in crystalline borate minerals or in water, primarily in the form of naturally occurring boric acid, B(OH)3. Boron is found in sea water at an average level of 4.6 ppm. This level is too low to permit economical extraction of industrial quantities. Thus, commercial supplies of boron are derived from mineral deposits, which are found as alkali or alkaline earth metal borates and borosilicates, with sodium and calcium borates being the most important. Although borate minerals are numerous, large commercially exploitable borate ore deposits are rare. However, when such deposits occur they can be vast. Owing to the relatively high water solubilities of borate minerals, these deposits are always found in arid regions. Most borates are produced in California and Turkey, with the rest coming from Russia, China, and the Andean regions of South America [2]. An estimated 1.25 · 106 metric tons of B2O3 is currently consumed annually in the form of the various borate products. Fig. 1 shows the approximate proportions of boron consumed by various large industries. It is notable that the use of borates in the manufacture of various types of vitreous materials, particularly fiberglass, ceramic glazes, and specialty borosilicate glasses, accounts for more than half of all boron used by industry. The uses of borates in peroxygen bleaching systems for household laundry detergents, micronutrient fertilizers, cellulose insulation (fire retardant), and cleaning products are also important. Following these applications there are hundreds of other minor uses of borates, a small selection of which are listed in Table 1. For the most part, these other applications utilize boron oxides, but a small fraction of uses involves non-oxide materials such as boron hydrides and engineering ceramics. Although these non-oxide materials have received great attention in academic circles in recent decades, they are yet to gain much industrial significance in the overall scheme of the borate industry.

Fig. 1. Estimated total borate use by major industrial applications in B2O3 equivalents for the year 2001

4

D. M. Schubert

Table 1. Examples of miscellaneous uses of borates

Adhesives Automotive coolants Biocides Boron hydrides Brake fluids Buffers Cement (set retardant) Cleaning products Corrosion inhibitors Cosmetics/lotions

Capacitors Fire retardants Fuel additives Iron and steel production Leather tanning Lubricants Metallic glasses Metal refining Non-oxide ceramics Nuclear industry

Nylon production (catalysis) Pharmaceuticals Photography Magnets Metal working fluids Solders Textiles Waxes/polishes Water treatment Wire drawing

Boron has the electron configuration [He]2s22p1, suggesting that it might form univalent compounds. Boron instead utilizes three valence electrons to combine with other elements to form trigonal planar compounds of the type BR3 having covalent bonds, where R is hydroxy, alkoxy, aryloxy, alkyl, halogen, etc. One of the defining features of boron is the very small size the B3+ ˚ ; four coordinate boron, 0.11 A ˚ ), cation (three-coordinate boron, 0.01 A leading to a high ionic potential and a particular affinity for electronegative elements. Boron does not form ionic compounds under any ordinary conditions. The element forms particularly strong covalent bonds with oxygen and forms stronger bonds only with fluorine [3]. Since boron possesses one more valence orbital (s, px, py, pz) than valence electrons, trivalent boron compounds have a pronounced tendency to form complexes with electron donors. This includes reversible reactions with anionic Lewis bases to form tetrahedral anions in which boron has a formal negative charge. Much has been made of this so-called ‘‘electron deficiency’’ of boron, as illustrated by the extensive chemistry of the boron hydrides, including the polyhedral boranes and their numerous derivatives. These compounds often contain multicenter bonds or more extensive electron delocalization [4]. In contrast, boron-oxygen compounds do not normally form multicenter bonds [5]. Nevertheless, the electrophilicity of boron plays a pervasive role throughout boron-oxygen chemistry, leading to a complex interplay of trigonal planar BO3 and tetrahedral BO4 structural units. This electrophilicity, coupled with a related tendency to readily and reversibly form esters and anhydrides through condensation with compounds bearing hydroxy groups, including B-OH compounds, accounts for virtually all large-scale industrial uses of boron, as well as its essential biological roles. An important aspect of boron is the restricted range of bond angles around both boron and oxygen atoms bound to it. The extensive range and structural complexity of borate compounds is comparable to that of the silicates. However, significant differences exist that define the distinctive industrial roles of these two classes of compounds. Moreover, the confluence of borate and silicate chemistries leads to the technologically important borosilicate glasses, as well as to the relatively abundant borosilicate minerals. This also holds true to some extent in comparison with aluminates. The strength of the BAO bond is similar to that of the SiAO bond. However, molecular orbital calculations

Borates in Industrial Use

5

indicate that optimal BAOAB and BAOASi bond angles (137 and 125, respectively) are small in comparison to optimal SiAOASi and AlAOAAl bond angles (143 and 151, respectively). These calculations correlate well with the range of bond angles observed in mineral and synthetic compounds. Also, the barriers to linearity of these linkages are much higher for the linkages involving boron, with the optimal angles occurring in a deep energy well. This leads to a strong tendency toward cyclization [6, 7]. These observations, coupled with the interplay of both BO3 and BO4 structure units, are important to understanding how the structural chemistry of borates differs from the silicates and aluminates, as well as how borates behave in vitreous systems.

2 Aqueous Chemistry Boric acid dissolves endothermically in water, its solubility increasing greatly with increasing temperature. Its solutions are mildly acidic, which is a result of the electron-acceptor character of boron and not a tendency toward proton donation. Boric acid reacts with water, as given by Eq. (1), to give an equilibrium concentration of the [B(OH)4]) anion. The equilibrium constant ð1Þ for this reaction (Ka ¼ 5.80 · 10)10 mol L)1) is small enough that the proportion of boric acid in near neutral pH dilute solutions is > 99%. The relative concentration of the tetrahydroxyborate anion increases with increasing pH and becomes the dominant species above roughly pH 9. In dilute solutions boric acid is monomeric, but polymeric species become significant at concentrations above about 0.1 molar. Formation of these species is a direct result of boron’s tendency to form complexes with electron-donor species including oxygen attached to boron itself. The most important polyborate species observed in solution are the triborate anions [B3O3(OH)4]) and [B3O3(OH)5]2), the pentaborate anion [B5O6(OH)4]), and the tetraborate anion [B4O5(OH)4]2). The population distributions of these species found at 25 C in 0.4 molar boric acid equivalent solution as a function of pH is shown in Fig. 2 [8]. Concentrations of these species may vary considerably with temperature and overall boron concentration, and other minor polyborate species not represented are also likely to exist in solution. Fig. 2 should be regarded as a snapshot of a complex system taken under one set of conditions. The polyborate oxoanions can be regarded as acid anhydrides, formed by elimination of water following nucleophilic attack of a B-OH oxygen on a trigonal boron center. The polyborate anions contain one or more tetrahedral boron centers and have a negative charge resulting from acceptance of the OH) ion that is shared by boron atoms within the polyborate. Polyborate

6

D. M. Schubert

Fig. 2. Population distributions of borate species as a fraction of total boron for a 0.4 mol L)1 boric acid equivalent solution at 25 C

anions do not occur in significant concentrations at low pH values, the ) concentration B(OH) 4 ion becoming the major species when the OH approaches the level where charge sharing between boron atoms is no longer required. Thus, the occurrence of polyborate species is only important at relatively high concentrations in the intermediate pH range, with B(OH)3 and B(OH) 4 remaining the dominant species at low and high pH values, respectively. A plot of the BO4/BO3 molar ratio for this system as a function of pH produces a smooth curve having an inflection point at about pH 9. This curve corresponds to increasing basicity of the borate anion and is useful in determining the solution conditions under which various solid state borates, both mineral and synthetic, are formed.

3 Crystalline Borates Crystalline mineral and synthetic borates exhibit considerable structural variety and complexity. More than 200 borate minerals and hundreds more synthetic borates are known. To date more than 100 borate minerals and many synthetic borates have been structurally characterized. These compounds contain boron oxide anionic components built from trigonal planar [BO3]3) and tetrahedral [BO4]5) groups as primary structural units. Despite the apparent simplicity of these fundamental units, borate structures are often quite complex. These units may exist as isolated anions, link together by sharing oxygen to form rings and cages, or further polymerize into chains, sheets, or extended networks. Furthermore, borate anions interact with cations in complex ways and often display extensive hydrogen-bonding integrating their structures.

Borates in Industrial Use

7

Borates can be thought of as containing boron oxide or boron oxide hydroxide structures that are tightly integrated anionic arrays with excess charge balanced by interstitial cations. Polymeric borate structural units contain well-defined repeating structural units called fundamental building blocks (FBBs). The occurrence of an extensive range of FBBs has given rise to various classification schemes [9–13]. These schemes classify FBBs in terms of the number of boron atoms, the number of 3- and 4-coordinate boron atoms in their repeating units, and the mode of condensation between FBBs to give isolated, modified isolated, chain, modified chain, sheet, modified sheet, and three-dimensional network structures. The most recent refinement of such schemes was proposed by Burns to classify borate minerals [13]. This scheme has since been used as a basis to classify the numerous known structures of borate minerals [14]. The classification nomenclature of Burns is given with the anions shown below. There are two broad categories of borates referred to as hydrated and anhydrous. Hydrated borates contain B-OH groups and may also contain interstitial H2O in their structures. Hence, the fully hydrated forms of the fundamental borate units above are B(OH)3 and [B(OH)4]), as found in aqueous solution. The majority of borate minerals and synthetic borates of commercial interest are hydrated. Compositions of borates can be resolved into the form a MxO Æ b B2O3 Æ c H2O (x ¼ 1 for dications M2+; x ¼ 2 for monoanions M+) in much the same way that carbohydrates can be represented by the general formula [C(H2O)]x. Such formulas are useful since the ratio b/a, referred to as Q, relates to the BO3/BO4 ratio in the compound. The value of Q can be regarded as a measure of ‘‘alkalinity’’ of the borate, with lower Q corresponding not only to lower B2O3 content of compositions but also to higher pH of the solution from which the borate crystallized. The value c in the resolved oxide formula often relates to the degree of condensation of the borate. A number of borates are known that contain isolated [B(OH)4]) units, including the mineral frovolite [15] {Ca[B(OH)4]2 ¼ CaO Æ B2O3 Æ 4 H2O}, shown in Fig. 3a. The condensation of two [B(OH)4]) units produces the [B2O(OH)6]2) anion found in minerals such as pinnoite [16] {Mg[B2O(OH)6] ¼ MgO Æ B2O3 Æ 3 H2O}, shown in Fig. 3b. Further condensation leads to infinite chains of oxygen-linked 4-coordinate boron found in the mineral vimsite [17] {Ca[BO(OH)2]2 ¼ CaO Æ B2O3 Æ 2 H2O} shown in Fig. 3c. These minerals have similar resolved oxide formulas (CaO Æ B2O3 Æ c H2O, with c ¼ 4, 3, and 2) that vary only by the degree of hydration. However, the relationship between c and the extent of anion condensation may be complicated by the presence of H2O versus B-OH groups in the compound. For example, due to the presence of interstitial H2O, the mineral pentahydroborite [18] {Ca[B2O(OH)6] Æ 2 H2O ¼ CaO Æ B2O3 Æ 5 H2O} has c ¼ 5 and contains the same anion as pinnoite (Fig. 3b). In practice, the two forms of water can usually be distinguished by TGA experiments since the dehydration reaction given by Eq. (2) generally occurs at significantly higher onset temperatures compared with loss of interstitial H2O. Instead of forming ð2Þ

8

D. M. Schubert

Fig. 3. Schematic structures of the calcium borate minerals (a) frovolite, CaO Æ B2O3 Æ 4 H2O,

(b) pinnoite, MgO Æ B2O3 Æ 3 H2O, and (c) vimsite, CaO Æ B2O3 Æ 2 H2O (n ¼ infinity)

simple linear chains, B(OH)3 and/or [B(OH)4]) units often condense to form 6-membered rings of the types shown in Fig. 4. This tendency to cyclize is a characteristic feature of borate chemistry. The neutral species shown in Fig. 4a, known as orthorhombic metaboric acid or trihydroxyboroxine [19], is formed by partial dehydration of boric acid (monoclinic and cubic metaboric acid also occur). If boric acid is considered as equivalent to B2O3 Æ 3 H2O, then the metaboric acids are B2O3 Æ H2O. The triborate mono- and dianions [B3O3(OH)4]) (Fig. 4b) and [B3O3(OH)5]2) (Fig. 4c) occur in substantial concentrations in borate solution (see Fig. 2) and are found in a number of mineral and synthetic crystalline borates. These include the minerals ameghinite [20], Na[B3O3(OH)4], inderite [21], Mg[B3O3(OH)5] Æ 5 H2O, and the synthetic zinc borate Zn[B3O3(OH)5] Æ H2O (2 ZnO Æ 3 B2O3 Æ 7 H2O) [22]. The [B3O3(OH)6]3) anion, formally derived from the condensation of three [B(OH)4]) is less common but is found in the calcium borate minerals nifontovite [23], Ca3[B3O3(OH)6]2 Æ 2 H2O, and olshanskyite [24], Ca2[B3O3(OH)6]OH Æ 3 H2O. Two triborate moieties may link together by sharing a common tetrahedral boron atom to form the pentaborate anions shown in Fig. 5. The pentaborate monoanion shown in Fig. 5a is observed in solution (see Fig. 2) and occurs in synthetic and mineral borates. Examples are the synthetic commercial ammonium pentaborate, NH4[B5O6(OH)4] Æ 2 H2O [ ¼ (NH4)2O Æ 5 B2O3 Æ 4 H2O], and the mineral sborgite Na[B5O6(OH)4] Æ 3 H2O [25], which both contain isolated pentaborate monoanions. The pentaborate trianion (Fig. 5b) is found in the industrially important mineral ulexite [26], NaCa[B5O6(OH)6] Æ 5 H2O (see discussion below).

Fig. 4. Orthorhombic metaboric acid (a) and triborate mono-, di-, and trianions (b–d)

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Fig. 5. Pentaborate mono- and trianions

Triborate moieties may also link together through shared tetrahedral boron atoms to form bi- and tricyclic systems, most notably the [B4O5(OH)4]2) and [B6O7(OH)6]2) anions shown in Fig. 6. These anions can be viewed as the result of fusing two and three triborate dianions (Fig. 4c), respectively. The [B4O5(OH)4]2) anion is found in the industrially important borax decahydrate, Na2[B4O5(OH)4] Æ 8 H2O (Na2O Æ 2 B2O3 Æ 10 H2O, discussed below) as well as in a number of other borate minerals such as hungchaoite [27], Mg(H2O)5[B4O5(OH)4] Æ 2 H2O. The [B6O7(OH)6]2) anion has less industrial importance but is found in a number of minerals including aksaite [28], Mg[B6O7(OH)6] Æ 2 H2O. The overall charge on most borate anions is equal to the number of tetrahedral boron atoms it contains. The [B6O7(OH)6]2) anion is a notable exception, due to the presence of an unusual three-coordinate oxygen that carries a formal positive charge. Most of the isolated borate anions shown above may polymerize in a variety of ways to form extended chains, sheets, and networks. Polymerization occurs by either sharing an exocyclic oxygen atom or by sharing an intracyclic tetrahedral boron atom. These are exemplified by the structures of the important industrial minerals colemanite [29], Ca[B3O4(OH)3] Æ H2O (2 CaO Æ 3 B2O3 Æ 5 H2O, Fig. 7a), and kernite [30], Na2[B4O6(OH)2] Æ 3 H2O (Na2O Æ 2 B2O3 Æ 4 H2O, Fig. 7b).

Fig. 6. The (a) [B4O5(OH)4]2) and (b) [B6O6(OH)6]2) anions

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D. M. Schubert

Fig. 7. Schematic structures of (a) colemanite, Ca[B3O4(OH)3] Æ H2O, and (b) kernite, Na2[B4O6(OH)2] Æ 3 H2O

A great many polymeric borates exist, some having quite complex structures. These may have polyborate building blocks that link together into chains, sheets, or three-dimensional frameworks. An example of a borate having a sheet structure is the mineral nobleite [31], Ca[B6O9(OH)3] Æ 3 H2O, shown in Fig. 8, which contains a hexaborate FBB that links together by sharing oxygen. This calcium borate can be prepared synthetically [32], and processes have been developed for the manufacture of synthetic nobleite for use in various commercial applications. The linkage of BO3 and BO4 groups via oxygen-sharing can, in theory, produce a great many polyborate structures. Despite the variety and complexity of known borate structures, a much smaller subset of all topologically possible structural permutations is actually observed. This is because certain structures are energetically more favored by virtue of having preferred bond angles and lengths. A survey of synthetic and mineral borate structures reveals an abundance of B3O3 rings. It can be noted that each of the borate structures shown in Figs. 4–8, including all the common polyborate anions known to exist in aqueous solution, contain the B3O3 rings. Statistical analysis of known borate mineral structures supports the predominance of B3O3 ring structures, and molecular orbital calculations show this to be a particularly stable arrangement [33]. This is particularly true for rings made

Fig. 8. Structure of the mineral nobleite, Ca[B6O9(OH)2] Æ 3 H2O

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up of one trigonal BO3 and two tetrahedral [BO4]) groups. Hence the structural unit shown in Fig. 9 is the dominant structural motif of borate chemistry. The structures of borates represented above are greatly simplified because the important interactions of the anionic boron oxide components with their cationic counterparts and other interstitial species, as well as with neighboring boron oxide components, were largely ignored. To achieve a more complete understanding of solid state borates we must consider these interactions in detail. One approach is through the use of the binary structural representation developed to study mineral structures [34–36]. This model divides borate structures into two components: the borate structural unit and the interstitial complex. The structural unit is usually an anionic boron oxide or boron oxide hydroxide array that is held together by relatively strong covalent bonds. The interstitial complex consists of metal ions, or other cations, and all other interstitial species, such as H2O, OH), or Cl). The overall charge of the interstitial complex is the sum of the charges of the species it contains and matches the overall charge of the structural unit to maintain electroneutrality. The contents of the interstitial complex, which may or may not bond together, participate in relatively weak cation-anion or hydrogen bonding interactions with the structural unit, with B-O-B and B-OH oxygen in the coordination sphere of metal ions. In addition, all hydrated borates display extensive hydrogen bonding, having H-bonds emanating from their B-OH groups that link it to adjacent structural units and/or interstitial components, often water. Occasionally one finds intra-structural unit H-bonds. Hence crystalline borates are highly integrated systems where both cations and H-bonds play important roles in controlling overall structure. Although bonding interactions between structural units and interstitial complexes and neighboring structural units are relatively weak, they are typically numerous enough to comprise a major factor in determining structure and stability. Since the latter interactions are generally the weakest ones in the system, they are important in determining the bulk physical properties of crystalline borate materials. Borates can be viewed as complex salts in which the Lewis basicity of the structural unit must match the Lewis acidity of the interstitial complex to produce a stable structure. The relationship between basicity of the borate structural unit and its BO4/BO3 ratio was mentioned above, indicating that each specific structural unit has an associated basicity. It was recently shown that the Lewis basicity of borate structural units correlates with the average coordination number of the oxygen atoms they contain, counting H-bonds

Fig. 9. The dominant structural motif of crystalline borates

12

D. M. Schubert

and bonds to cations in the interstitial complex [37]. However, a plot of this correlation reveals a range of oxygen coordination numbers for each structural unit, suggesting that the structural unit may adjust its basicity to some extent to match the Lewis acidity of the interstitial complex. Once the Lewis acidity of the interstitial complex moves beyond the range in which adjustment can be made through change in coordination number of structural unit oxygens, the system becomes unstable and a different structural unit must occur. Methods for the numerical estimation of Lewis basicity structural units and Lewis acidity of interstitial complexes have been proposed [37]. Most mineral and synthetic industrial borates are hydrated, and a majority contains interstitial H2O, i.e., water that is part of the interstitial complex in addition to HO) groups in the structural unit. The relationship between the overall water content of a borate, as apparent from the resolved oxide formula, and the degree of polymerization of borate structural units was discussed above. Water in the interstitial complex also plays a further critical role in determining borate structure. The various roles that water can play here have been described [38]. If water is not bonded to an interstitial cation it is usually involved in an H-bonding network, being a donor of two H-bonds and an acceptor of two H-bonds, and thus contains 4-coordinate oxygen. Since water in this role both accepts and donates H-bonds, it does not change the net strength of chemical bonds but simply interconnects the overall structure. Water that is bound to a metal cation performs a different role. Water may coordinate to one or more metal ions in the interstitial complex and simultaneously participate in H-bonding with the structural unit. Although water may do this in a variety of modes, in many cases water serves to divide one relatively strong metal-oxygen bond into two weaker hydrogen bonds. Water in this role serves to moderate the Lewis acidity of the interstitial complex and has been termed transformer water [37, 38]. Methods of estimating the Lewis acidity of the interstitial complex take into account the number of transformer and non-transformer water molecules present in the structure [37]. Application of these methods leads to the conclusion that the presence of transformer water is necessary to the stabilities of many borate structures. For example, one can predict that in order to obtain overlap between the Lewis acidities of borate structural units and Lewis basicities of interstitial complexes, interstitial Mg2+ will bond to transformer water whereas Na+ will not. This prediction agrees well with observations of mineral structures. Although other arguments involving the relative hydration energies of these cations can also explain the occurrence of hydrated and non-hydrated metal ions in borate structures, the role of interstitial water in borate structures is clearly important. The binary representation is applicable to various other oxide materials. However, an important distinction can be made between borates and other main group element oxide systems, such as aluminates and silicates. In the latter systems cations predominantly reside at sites created by the demands of rigid anionic oxide frameworks. Although some degree of structural control may be obtained by varying cations or by use of template synthesis, the oxide frameworks of these systems tend to be relatively inflexible in comparison with

Borates in Industrial Use

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extended borate structural units. As a result of the greater flexibility of boron oxide structures, allowing them to conform with greater freedom to the steric and electronic demands of cations and H-bonds, the interstitial complex plays a more important role in directing the spatial arrangement of the structural unit in borates than is found for other main group element oxide systems. Many borates are known that contain non-metal cations. Examples include the ammonium borates [39], such as the commercial products ammonium pentaborate, NH4[B5O6(OH)4] Æ 2 H2O, and ammonium tetraborate, (NH4)2[B4O5(OH)4] Æ 2 H2O, both of which contain isolated borate anions shown above (Figs. 5a and 6a, respectively) [40]. Isolated borate anions containing more than six boron atoms are rare. However, the mineral ammonioborite {(NH4)3[B15O20(OH)8] Æ 4 H2O ¼ (NH4)2O Æ 5 B2O3 Æ 2 2/3 H2O), which may be prepared synthetically by prolonged reflux of a saturated solution of ammonium pentaborate, contains the unusual isolated anion [B15O20(OH)8]3), shown in Fig. 10 [41]. This anion is not known to occur in any metal borate, its formation is apparently directed exclusively by the ammonium cations which can donate H-bonds to interstitial water and the borate structural unit. Direction of borate structures by specific cations is also illustrated by the formation of the unusual nonaborate anion in the presence of the imidazolium, [C3H7N2]+, and guanidinium, [C(NH2)3]+, cations [42]. The reaction of imidazole with three molar equivalents of boric acid in aqueous solution results in the spontaneous formation of the imidazolium salt of the [B9O12(OH)6]3) anion, shown in Fig. 11, associated with three [C3H7N2]+

Fig. 10. Structure of the [B15O20(OH)8]3) anion found in the mineral ammonioborite,

(NH4)2O Æ 5B2O3 Æ 2 2/3 H2O

Fig. 11. The nonaborate anion, [B9O12(OH)6]3) (6D3h:<2Dh>–––<2Dh>)

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D. M. Schubert

cations and no interstitial water. An analogous nonaborate compound forms with the guanidinium cation. In each case, cation hydrogen atoms participate in H-bonding with oxygens of the nonaborate anion. These are the only known examples of an isolated nonaborate anion. Related nonaborate FBBs are found in the polymeric minerals preobrazhenskite [43], Mg3[B11O15(OH)9], and NaCa2[B9O14(OH)14] Æ 2 H2O [44]. Despite having analogous structures, the imidazolium and guanidinium nonaborates exhibit different chemistries in contact with water. Whereas the imidazolium compound [C3H5N2]3[B9O12(OH)6] can be recrystallized from water at room temperature, its guanidinium analogue [C(NH2)3]3[B9O12(OH)6] exhibits a temperature-dependent equilibrium, Eq. (3), in aqueous solution with guanidinium tetraborate [C(NH2)3]2[B4O5(OH)4], where the nonaborate is stable in hot water and reversibly converts to the tetraborate when cooled to room temperature. 2½CðNH2 Þ3 3 ½B9 O12 ðOHÞ6  ƒ! ƒ 3½CðNH2 Þ3 2 ½B4 O5 ðOHÞ4  þ 6BðOHÞ3

ð3Þ

Another example of cation control over borate structure is found in the recently elucidated structure of the industrially important synthetic zinc borate Zn[B3O4(OH)3] [2 ZnO Æ 3 B3O3 Æ 3 H2O], shown in Fig. 12 [45]. This compound has a polytriborate structure reminiscent of that found in colemanite (Fig. 7a) [29]. Yet, due to the different coordination requirements of the Ca2+ and Zn2+ cations, the anionic polyborate chains in these two compounds have different spatial arrangements. In the case of colemanite, parallel triborate chains are linked together into extended sheets by coordination with interstitial Ca2+ (and H2O). These sheets are bound together primarily by relatively weak interstitial complex-structural unit interactions, resulting in the characteristic cleavage planes exhibited by the mineral colemanite. In the zinc borate Zn[B3O4(OH)3], on the other hand, zinc is coordinated twice by oxygens (B-O-B and B-OH) of one polytriborate chain and once coordinated by oxygen (B-OH) of each of two other chains. Thus, zinc links three separate polytriborate chains in a more three-dimensional arrangement. As a result, this form of zinc borate is a much harder material than colemanite. These examples illustrate not only how cations direct the structures of isolated borate anions and the spatial arrangements of extended borate polymers, but also how they determine chemical and physical properties of

Fig. 12. The

structure of (2 ZnO Æ 3 B2O3 Æ 3 H2O)

the

synthetic

industrial

zinc

borate

Zn[B3O4(OH)3]

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borate materials. A better understanding of how various organic and inorganic cations control borate structures may lead to improvements in the directed synthesis of borate material having desirable properties.

4 Vitreous Borates Glass and related vitreous products have enormous technological importance, and the use of borates in the manufacture of vitreous products, particularly various kinds of borosilicate glasses, ceramic glazes, and porcelain enamels, is the most important industrial application area of boron, accounting for more than half of the world’s boron consumption. Although intensive physical and chemical investigations on boron-containing glasses have been done, a detailed and comprehensive understanding of these complex systems remains elusive. Nevertheless, much progress has been made over the past decade toward developing a fundamental understanding of the structure and bonding of boron in glasses. Boric oxide is often referred to as a flux in glass manufacture. However B2O3 has limited solubility in silica and does not act as a flux in the usual sense. In the production of silica-based glasses, modifiers such as Na2O, K2O, CaO, and MgO are used to reduce the melting temperature and viscosity of molten silica to a practical range for the manufacture of glass articles. These modifiers serve to break the rigid silicate network by cleaving SiAOASi bonds and form nonbridging Si-O) groups that carry a negative charge. Boric oxide, on the other hand, acts as a network former and strengthener rather than a network breaker. It is an excellent flux for metal oxides and assists the incorporation of these into the glass matrix. However, this does not explain the unique and important role that B2O3 plays in glasses. Although this role is not yet completely understood, the following discussion sheds some light on what is currently known. The detailed structure of vitreous B2O3 has been the subject of considerable debate over the past several decades. It is established that all boron in this substance is present in trigonal BO3 groups and that all oxygen is bridging between two borons. Also, a significant fraction of the BO3 groups are in sixmembered B3O3 boroxol rings (Fig. 13). What remains controversial is the actual fraction of boron in these boroxol rings. A large body of experimental data, including X-ray and neutron diffraction, and NMR, NQR, and Raman spectroscopic data, seems to indicate that a high mole fraction of boron is

Fig. 13. Boroxol rings occurring in glass structures

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D. M. Schubert

present in boroxol rings, perhaps 0.6–0.7 [46]. However, theoretical calculations predict that a much smaller fraction of boron should be involved in boroxol rings [47]. The most recent NMR data lend support to the contention that 72 mole percent of the boron in B2O3 is present in boroxol ring structures [48]. The structures of binary alkali metal borate glasses are also controversial. Of particular interest is the behavior of vitreous B2O3 upon addition of alkali oxide network modifiers (e.g., Na2O, K2O), leading to the so-called boron anomaly [48]. This phenomenon relates to the relatively abrupt change in physical properties, such as density and thermal expansion coefficient, that occurs as the alkali oxide content reaches 15–20 molar percent metal oxide. As network modifiers are added the trigonal BO3 groups are converted to BO4 groups having attached non-bridging oxygens that carry a negative charge. This appears to occur at BO3 groups in boroxol rings as well as at isolated BO3 groups. At higher modifier levels BO4 groups are replaced by BO3 groups having one or more non-bridging oxygens (Scheme 1). In the past, the boron anomaly had generally been attributed to a rapid decrease in the relative proportion of BO4 groups at 15–20 molar percent alkali oxide. However, more recent NMR and NQR studies show that the concentration of BO4 groups continues to increase up to about 40 molar percent alkali oxide and then drops off to near zero by about 70 molar percent. This means that the boron anomaly does not correlate with a sudden disappearance of BO4 groups, as previously supposed. It is more likely that the rapid change in physical properties at 15–20 molar percent alkali oxide is related to the appearance in the glass network of a significant amount of more bulky BO4 groups and larger polyborate superstructures. This leads to more compact regions in the glass network having greater connectivity and rigidity. The formation of large superstructural polyborate units reduces the internal degrees of freedom of the system since these units have a limited range of bond and torsion angle variability. This is consistent with a model proposed by Krogh-Moe, which contends that glasses should contain structural grouping similar to those found in crystalline compounds of the same glass-forming system [50]. The nature of the superstructural units found in borate glass networks is informed by the extensive chemistry of crystalline borates, which provides insights into deciphering the structures of borate glasses [51]. In fact, experimental data

Scheme 1. Structural changes in boron centers with increasing levels of Na2O (or other alkali metal oxide) modifier

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indicate that a number of the common polyborate anions found in crystalline borates are present in borate glass anhydrous forms. The difference is that most commercially important glasses are formed at higher temperature (>1000 C) in which the components are present in high energy states. Upon cooling the viscosity increases to the point at which relaxation can no longer occur, the fictive temperature, and thus glass at room temperature still resembles this higher energy state. This also means that cooling rates may have a strong bearing on glass structure. The family of alkali borosilicate glasses has considerable technological importance, being the basis of many commercial glasses produced on large scale. In such glasses, BO3, BO4, and larger borate groupings contain oxygens that are bonded to silicon. One interesting finding from NMR experiments is that at low alkali content nearly all the alkali oxide is associated with borate structures in the glass network, converting BO3 groups to BO4 groups, rather than being divided between borate and silicate. At higher alkali levels the alkali oxide begins to divide between borate and silicate. The structural diversity of borates provides a basis to approach an understanding of the vitrification phenomenon. Although far from well understood, it is clear that the ability of boron to form tightly integrated superstructural units having restricted degrees of freedom underlies the unique and useful properties of borate-containing glasses.

5 Organic Ester Formation Synthetic borate esters produced by reaction of boric acid with alcohols find use as components of many types of industrial fluids, including lubricants, hydraulic fluids, metal working fluids, etc. Borate esters produced from polyols find use in applications ranging from toys to starch adhesives for cardboard boxes. A small-scale but increasingly important use of borate esters is as intermediates in a range of organic reactions important to the production of specialty chemicals and pharmaceuticals. Borate ester formation is also the fundamental process underlying the biological role of boron, as discussed below. Although the mechanism of ester formation has been the subject of some controversy, this process is related to polyborate formation in that both cases likely involve nucleophilic attack of an HO) oxygen on B(OH)3 followed by elimination of water from a tetrahedral intermediate, Eq. (4). BðOHÞ3 þ ROH ƒ! ƒ ðHOÞ2 BOR þ H2 O

ð4Þ

The rate of the esterification reaction is very fast, with the BAO bonds breaking in this process, not the CAO bond. The equilibrium of Eq. (4) does not lie far to the right, since water competes as a nucleophile in the reverse reaction, and rapid equilibration occurs unless water is removed to drive the reaction to the right. Hence most simple borate esters are readily hydrolyzed, displaying pseudo-first order hydrolysis kinetics. In most cases esterification

18

D. M. Schubert

proceeds to give a large proportion of fully esterified orthoborate, B(OR)3, or metaborate, B3O3(OR)3, products only if water is removed by chemical or physical methods, Eqs. (5) and (6). BðOHÞ3 þ 3ROH ! BðORÞ3 þ 3H2 O

ð5Þ

ð6Þ

Reaction of boric acid with 1,2- and 1,3-diols in aqueous media results in a drop in pH, a consequence of the formation of cyclic diesters having enhanced kinetic stability (Fig. 14). These also have a general tendency toward hydrolysis, but exist in an equilibrium with boric and diol that is shifted further to the right. Although most borate esters are quite susceptible to hydrolysis, some exhibit remarkable hydrolytic stability. These include esters of hindered phenols, as shown in Fig. 15a [52] and the ester of tripropanolamine, shown in Fig. 15b. Hindered phenolic borates, as a class, have remarkable hydrolytic stabilities with half-lives ranging up to two years (e.g., at 25 C in 90% aqueous dioxane solvent). The hydrolysis half-life of the ester shown in Fig. 15a, where R is isopropyl, has been measured at 379 days, which makes it millions of times more stable than triisopropyl borate. Tri-n-propanolamine borate is reported to have indefinite stability under the same conditions [53]. These compounds illustrate two different modes of protection of boron from nucleophilic attack by water. In the case of the hindered phenol, the t-butyl groups lie above and below the trigonal boron center, and hydrolytic stability results at least partly from steric protection with assistance from bulky R groups such as isopropyl. The borate ester of

Fig. 14. Cyclic borate esters derived from 1,2- and 1,3-dihydroxy compounds (R = alkyl,

aryl)

Fig. 15. Hydrolytically stable borate esters of hinder phenols (R = alkyl) and tri-n-

propanolamine

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tri-n-propanolamine, which contains tetrahedral boron, presents an interesting case in which a transannular interaction between boron and the lone pair of electrons on nitrogen renders boron unsusceptible to nucleophilic attack by water. Other borate ester chelates containing tetrahedral boron, shown in Fig. 16a, have been prepared as precursors to solvent-supported borate cations (Fig. 16b) that show catalytic activity toward olefin polymerization [54]. An important class of borate esters is the anionic spirocyclic esters, shown in Eq. (7). The reaction of two molar equivalents of 1,2- or 1,3-diol with boric acid to form such spirocyclic esters is accompanied by dramatic decrease in pH, a consequence of the kinetic stability of these esters. This reaction ð7Þ provides the basis of methods for the analysis of borates. For example, reaction with mannitol or sorbitol allows titration of boric acid as a strong acid, the consequence of an increase in acid equilibrium constant by a factor >104. Spirocyclic ester formation allows borates to function as cross-linking agents for industrial products, such as starch adhesives. This reaction also underlies the essential role of boron in biochemical processes. Since anionic ester formation proceeds with release of a proton, the reaction is also favored by the presence of proton acceptors or high pH. The observation that ester formation is favored at high pH has led some researchers to conclude that the reaction involves [B(OH)4]) as reactant rather than B(OH)3 [55]. However, careful analysis of equilibrium expressions for this process leads to the conclusion that a tetrahedral reactant is not required, and other evidence supporting a tetrahedral reactant is unconvincing [56]. The borate cross-linking reaction is often demonstrated by the addition of borate to a dilute solution of polyvinyl alcohol, resulting in formation of the so-called ‘‘slime’’ gel, a semi-solid fluid containing more than 95% water by weight and that exhibits non-Newtonian flow properties. While greatly appreciated by grade school children, this experiment also illustrates the dramatic effect that borates can exert on polyhydroxylic compounds, including those commonly found in biological systems. NMR studies show that this is a highly dynamic system in which the borate linkages form and break with a frequency order of 100 Hz. Similar dynamic processes involving boron that take place within living cells are essential to life.

Fig. 16. Borate ester chelate complex (a) serves as precursor to catalytically active borate ester cation (b). R = alkyl, X = H, Cl, Br

20

D. M. Schubert

6 Biological Aspects of Boron 6.1 Natural Products

Boron-containing natural products are not as rare as once supposed. Traditional methods for the isolation and purification of organic natural product compounds involve procedures that typically result in the hydrolytic removal and loss of boron. Thus only in recent years have significant numbers of boron-containing natural products become recognized. Among the first boron-containing natural products characterized were macrolide antibiotics, including aplasmomycin [57], boromycin [58], and tartralon B [59], shown in Fig. 17. Another called borophycin is also known [60]. These ostensibly similar molecules derive from quite disparate biological sources, aplasmomycin and boromycin from Streptomycetes, tartrolon from Myxobacteria, and borophycin from blue-green algae (Cyanobacteria). Studies of the metabolic roles of boron in higher plants have now revealed numerous other boron-containing natural products. 6.2 Bioessentiality

Life evolved in the presence of a ubiquitous low level of boron, and thus it is not surprising that the fundamental processes of life should make use of this omnipresent element. Much progress has been made in the area of boron bioessentiality in the past decade, proving unequivocally that boron plays an essential role in the cycle of life. Boron is required for the growth of plants, diatoms, and marine algal flagellates [61, 62]. More recently boron was shown to be essential for the proper development of at least some higher animals, and mounting evidence now indicates that boron is important to the health of humans. The essentiality of boron in animals has taken a long time to be

Fig. 17. Boron containing macrolide antibiotics

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revealed, largely because boron is so widely distributed in the natural environment and is so difficult to remove from foodstuffs that deficiency disorders have been hard to conclusively demonstrate. While the specific mechanistic details are only slowly emerging, the role of boron is now understood to be both subtle and pervasive, having a vital influence on a number of fundamental metabolic processes in both plants and animals. 6.2.1 Plants It has long been recognized that boron is required by higher plants [61, 62], and recent research indicates the involvement of boron in three main aspects of plant physiology: cell wall structure, membrane function, and reproduction. In vascular plants, boron in solution moves in the transpiration stream from the roots and accumulates in the stems and leaves. Once in the leaves, the translocation of boron is limited and requires a phloem transport mechanism. The nature of this mechanism was only recently elucidated with the isolation of a number of borate polyol compounds from various plants [63–65]. These include sorbitol-borate ester complexes isolated from the floral nectar of peaches and mannitol-borate ester complexes from the phloem sap of celery. The implication is that the movement of boron in plants depends on boratepolyol ester formation with the particular sugar polyol compounds used as transport molecules in specific plants. It is now clearly established that boron plays a fundamental role in plant cell wall structure [66]. The cell walls of higher plants determine size and shape during growth and provide strength and resiliency to plant parts. The mechanical properties of cell walls can be controlled by cross-linking between major cell wall components, including cellulosic polymers and matrix polymers such as hemicellulosic and pectic polysaccharides. It is relevant that up to 90% of boron in plants is localized in the cell wall. Many years of study of plant nutrition reveals a strong relationship between boron and cell wallrelated phenomena. For example, boron deficiency in plants is associated with abnormalities in growth and cell wall organization, with diagnostic symptoms including brittle leaves and cracked stems. Borate cross-linking between hydroxy groups of cell wall carbohydrates, and possibly glycoproteins, appears to be important in the maintenance of proper plant cell wall structure and function. Rhamnogalacturonan II (RGII) is an extremely complex pectic polysaccharide found in plant cell walls. The presence of RGII in all plants suggests it is important to either structure or growth, and even small changes in this molecule result in growth defects. It was recently shown that the function of RGII depends on its ability to dimerize through ester cross-linking by boron [66]. Such RGII dimers (RGII-B) have been isolated from the plant cell [67–69]. Other hydroxy compounds in plants may also to be involved in borate cross-linking, including hydroxyproline-rich glycoproteins and proline-rich proteins such as extensin. Boron cross-linking of other plant cell wall carbohydrates, and perhaps glycoproteins and proteins, may play a vital role in organizing, assembling, and providing structural elasticity in plants.

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Many studies have shown that plants require an ample and continuous supply of boron for flowering, pollen tube growth, pollen germination, and fruiting. In fact, most plants require much more boron for reproductive growth than for normal vegetative growth [62]. The physiological reasons for higher boron demand during reproductive growth are not well understood. It is likely that cell wall composition determines the requirement for boron, with reproduction cell walls having special compositions. For example, graminaeceous plants (grasses) have low cell wall pectin levels and also have relatively low boron requirements for maintenance of vegetative growth, yet these plants need as much boron as other plants during the reproductive growth stage. Boron deficiency causes symptoms ranging from sterility to flower malformation in a wide variety of plant types. Borate ester complexes are proposed to be involved in pollen tube extension and germination processes. Pollen contains many glycoproteins, the oligosaccharide components of which contain significant levels of mannose and fucose, both of which form ester complexes with borate. Studies have shown that the presence of boron strongly effects the phase change patterns of membrane lipids [70]. Also, complete inhibition of pollen tube germination above 21 C in the absence of sufficient boron has been observed. This may relate to the importance of boron in crops that reproduce during the warm season, such as maize. Lipid thermostability data suggest that boron is involved in maintaining proper membrane structure and function, and pollen growth indicates an involvement of boron in cell wall structure [62]. Functions of cell membranes include isolation and delimitation of cells, regulation of ion and metabolite transport, reception and production of extracellular signals, and catalysis of enzymatic reactions. Membranes are composed of lipids, carbohydrates, and proteins including a number of important enzymes. These components contain chemical species possessing cis-dihydroxy functionalities, including phosphoinosinates, glycoproteins, and glycolipids, and therefore have the ability to interact with boron. The addition of boron to boron-deficient plants induces rapid changes in membrane function that cannot be explained by the role of boron in cell walls described above. Although found in membrane fractions at much lower levels than in cell walls, boron is known to influence ion uptake. For example, uptake of phosphorus and potassium by the roots of various plant species, such as maize and sunflowers, is inhibited by boron deficiency. Uptake of these ions returns to normal within one hour after boron is restored [62]. Boron also appears to be involved in redox metabolism in cell membranes. Boron deficiency was shown to inhibit membrane H+-ATPase isolated from plant roots, and H+-ATPase-associated proton secretion is decreased in borondeficient cell cultures [71]. Other studies show an effect of boron on membrane electron transport reactions and the stimulation of plasma reduced nicotinamide adenine dinucleotide (NADH) oxidase upon addition of boron to cell cultures [72, 73]. NADH oxidase in plasma membrane is believed to play a role in the reduction of ascorbate free radical to ascorbate [74]. One theory proposes that, by stimulating NADH oxidase to keep ascorbate reduced at the cell wall-membrane interface, the presence of boron is important in

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maintaining a reducing environment in the plant apoplast, supporting ion uptake [75]. Notably, both ascorbate and NADH oxidase activity are associated with plant growth processes. These data strongly suggest that boron plays a role in membrane structure and/or function; however the details of this role have to be elucidated. Studies have shown leakage of metabolites, including potassium, sucrose, phenolics, and amino acids, from the cells of plants deficient in boron [76]. One theory contends that boron complexes polyhydroxy membrane components and thereby maintains plasma membrane structure to keep either enzymes or channels in functional conformations. This could involve ester complex formation and also H-bonding with glycolipid or glycoprotein membrane components. Other studies observed that certain enzymes that are normally bound in cell membranes can be released and activated under boron-deficient conditions [77]. These enzymes include ribonuclease, glucose-6-phosphate dehydrogenase, phenylalanine ammonia lyase, b-glucosidase, and polyphenol oxidase. The release of these enzymes can significantly alter plant metabolism by decreasing RNA levels and increasing the production of phenolic compounds which can build up to toxic levels in plant cells. The use of borates as enzyme stabilizers and inhibitors lends evidence to their ability to directly bind and influence the activity of enzymes. Many reactions are catalyzed by enzymes through a lowering of the activation energy brought about by binding to an active site in the configuration of the transition state. Serine hydrolase enzymes are inhibited by borates and borate derivatives, which react to form tetrahedral complexes (Fig. 18) in which the borate forms an ester linkage with a serine hydroxy group, hydrogen bonds with the imidazole group of an adjacent histidine, and is further stabilized by hydrogen bonding with serine and other amino groups [78]. 6.2.2 Animals Recent research shows that boron is necessary at least in the early stages of animal life and probably in later life as well. This includes studies showing that a deficiency of boron adversely affects the development of frog and fish [79, 80] embryos. Other studies suggest that low boron levels hinder reproduction

Fig. 18. Borate complex in serine hydrolase inhibition

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of rats and mice [81]. Studies with humans indicate that a sufficient dietary intake of boron is necessary to maintain health [82]. Dietary deficiency of boron in the African clawed frog (Xenopus laevis) causes necrosis of eggs and abnormalities in developing eggs that are suggestive of disrupted cell membrane function or structure. More than 80% of boron deficient embryos died before 96 h of development, whereas more than 75% of embryos supplemented with boron survived at 96 h. Cultivating frogs in a medium containing only 3 lmol L)1 B resulted in severe developmental abnormalities and difficulties in metamorphosis to adult frogs [79]. It is now clear that earlier studies indicating that very low levels of boron are toxic to fish embryos were actually documenting the effects of boron deficiency [83]. More recent studies on fish show that embryonic growth is stimulated by boron in a dose-dependent manner [80]. While very high boron levels are toxic, insufficient boron levels (<9 lmol L)1 B) were shown to result in adverse effects on rainbow trout and zebrafish development. Studies on humans strongly support a need for boron in the diet to maintain proper health. Although the details are yet to become clear, evidence suggests that boron may play roles in the optimal utilization of other vitamins and minerals, bone metabolism, and brain function, among others. The human requirement for boron is estimated to be in 1–2 mg/day range. Because of the widespread occurrence of boron in plants and the environment, this level is found in a normal healthy diet that includes vegetables. However, some studies suggest that heath benefits could result from the intake of somewhat higher boron levels [82]. 6.3 Toxicity

Most inorganic salts display some level of toxicity. Since boron occurs naturally throughout the environment and is essential for plant life, it is a natural part of normal foodstuffs and a component in a healthy human diet. Common inorganic borates display low acute mammalian toxicities. These borates include boric acid and borax, other sodium borates as well as the ammonium, potassium, and zinc borates, all of which dissociate to boric acid under physiological conditions. With regard to human exposure, these compounds are not skin irritants or skin sensitizers and have no mutagenic or carcinogenic effects. Toxicological studies employing very high chronic exposure levels have shown developmental toxicity and male reproductive system effects in several mammalian species, with very similar toxicological effects across species. These data are important to consider in regard to the proper uses of borate products. However, the toxic threshold levels are considerably higher than humans are likely to be exposed to in ordinary handling and use [84, 85]. Thus, borates are regarded as safe ingredients for use in consumer products, including topical ointments, lotions, soaps, etc. They are not generally allowed for use as intentional food additives.

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7 Industrial Boron Products All boron products consumed on a large scale by industry today are boronoxygen compounds. These are used at vastly greater tonnage levels than any non-oxide compounds of boron. Borate products in the marketplace are of two main types, refined borates and mineral borates. The refined borates are, of course, manufactured from mineral borates but are either converted into other borate compounds or refined to the point of being substantially pure chemical compounds. Mineral borate products are beneficiated ores, upgraded by methods that might include crushing, screening, washing, magnetic separation, and blending. Mineral borate products contain impurities comprised mostly of clays and other constituent minerals that tend to be associated with clays. The most important borate products currently used on a large industrial scale are described below. 7.1 Refined Borates

7.1.1 Borax Pentahydrate Borax pentahydrate is by far the most important refined borate in use today. The term borax is used in reference to two different crystalline hydrated sodium salts of the [B4O5(OH)4]2) anion (Fig. 6a), which are Na2[B4O5(OH)4] Æ 8 H2O and Na2[B4O5(OH)4] Æ 3 H2O, shown in Fig. 19. The former compound is also known as the mineral borax, also called tincal, and the latter as the mineral tincalconite. As articles of commerce, these are often designated by their resolved oxide formulas, Na2O Æ 2 B2O3 Æ 10 H2O and Na2O Æ 2 B2O3 Æ 5 H2O (also written as Na2B4O7 Æ 10 H2O and Na2B4O7 Æ 5 H2O). Use of these formulas leads to the various names applied to these compounds in commerce, the former being known as borax decahydrate or borax 10-Mol, the latter as borax pentahydrate or borax 5-Mol. These are white crystalline granular or powder materials. Borax pentahydrate is by far the more important of the two hydrates, with annual worldwide production in 2001 in excess of one million metric tons. Its lower water content makes it more economical to ship, and it has better handling properties than borax decahydrate, including a lower tendency toward caking.

Fig. 19. Borax structure, where x = 8 for borax decahydrate and x = 2.67 for borax pentahydrate

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Most borax is produced from tincal (native borax) ore in a process that amounts to a large scale continuous recrystallization. Borax pentahydrate crystallizes in the Na2O-B2O3-H2O system at appropriate concentrations above 67 C. It contains two crystallographically unique 6-coordinate Na+ positions and three unique water positions. A more recent refinement of its structure shows that one of these water positions is not fully occupied, giving an actual composition of Na2[B4O5(OH)4] Æ 2.67 H2O [86]. 7.1.2 Borax Decahydrate Although borax decahydrate, Na2[B4O5(OH)4] Æ 8 H2O, is the traditional form of borax, it has less industrial importance today than borax pentahydrate for the reasons mentioned in the preceding section. It is favored, however, for some applications such as cleaning products, where the mild abrasivity and dissolving rate of its crystals are valued. Crystallizing below 67 C in the Na2O-B2O3-H2O system at appropriate concentrations, the structure of borax decahydrate contains, in addition to the [B4O5(OH)4]2) anion, two crystallographically unique Na+ positions, each being octahedrally coordinated by H2O. These octahedra share edges to form chains that cross-link the polyborate anions to form parallel sheets that are further integrated linked by a network of H-bonds [87]. 7.1.3 Boric Acid The most versatile of all borates as a synthetic reagent, orthoboric acid, B(OH)3, is the second most important refined borate product in terms of industrial tonnage. It is typically manufactured by reacting a borate ore, particularly kernite or colemanite, with sulfuric acid, Eqs. (8) and (9). Na2 ½B4 O5 ðOHÞ2  3H2 O þ H2 SO4 þ 2H2 O ! 4BðOHÞ3 þ Na2 SO4 kernite Ca½B3 O4 ðOHÞ3  H2 O þ H2 SO4 þ H2 O ! 3BðOHÞ3 þ CaSO4 colemanite

ð8Þ

ð9Þ

Boric acid has a relatively low thermal stability and begins to lose some water when heated to as low as 100 C, well below its melting point of 170.9 C. Metaboric acid, HBO2, which exists in three different crystalline modifications, is formed upon controlled dehydation of boric acid. The structure of orthorhombic metaboric acid is shown in Fig. 4a. Whereas boric acid is equivalent to the trihydrate of boric oxide, B2O3 Æ 3 H2O, metaboric acid is the monohydrate, B2O3 Æ H2O. In theory, metaboric acid should have enhanced value relative to boric acid in vitreous applications since its has better melting characteristics and lower dehydration energy requirements compared with orthoboric acid. Its lower water content also reduces shipping cost per unit of

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Fig. 20. The structure of boric acid

B2O3. Production processes have been developed for metaboric acid, but to date it has not become commercially important and is not currently produced in any substantial quantities. Boric acid is a white solid that crystallizes from water as waxy platelets. The crystal structure of boric acid consists of triangular B(OH)3 molecules ˚ apart, (Fig. 20) interconnected by hydrogen bonding into planar layers 3.18 A [88]. This structure accounts for the slippery feel of boric acid and the cleavage planes observed in boric acid crystals. Boric acid provides lubricity in products ranging from cosmetics to pencil leads. Although the tribological properties of boric acid have been known and utilized for many years, its lubricating properties were rediscovered in recent years by the materials science community [89]. 7.1.4 Anhydrous Borax Anhydrous borax, Na2B4O7, can exist in several crystalline forms, but the commercial product is always an amorphous material produced by fusing hydrated borax in a furnace at about 1000 C. The resulting molten borax is drawn out of the furnace as a continuous ribbon and cooled between chillrollers to produce a solid glass. This glass that is crushed, screened, and packaged as granular products. Anhydrous borax is usually used in the manufacture of vitreous products where the release of water during the melting process is undesirable. 7.1.5 Boric Oxide Boric oxide, B2O3, is also referred to as anhydrous boric acid. The article of commerce is an amorphous material produced by fusing boric acid in a furnace. The structure of vitreous B2O3 is discussed above. Using a process

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similar to the production of anhydrous borax, molten B2O3 is chilled to a glass that is crushed, screened, and packaged as granular products. Boric oxide begins to soften around 400 C and becomes a pourable liquid above about 500 C. Reaction of B2O3 with water is exothermic and gives boric acid. Vitreous boric oxide is only mildly hydrogroscopic at room temperature, and the commercial product typically contains about 1% water, primarily in the form of a surface layer of boric acid. Although boric oxide is extremely difficult to crystallize, several crystalline phases are known. However, these have no particular industrial importance. 7.1.6 ‘Disodium Octaborate’ Disodium octaborate is not a discrete compound but rather a largely amphorous material approximating the composition Na2O Æ 4 B2O3 Æ 4 H2O. It is a white powder produced by spray drying a concentrated sodium borate solution of the appropriate composition. The advantage of this composition is its high water solubility, which is approximately twice that of borax or boric acid. Due to polyborate formation, borate salts exhibit irregular solubility patterns with sodium borates reaching maximum solubility at a B2O3/Na2O mole ratio near 4. Disodium octaborate dissolves in water to give solutions having pH values near neutral (pH 7–8), which are used in applications requiring high borate concentrations, such as wood preservation. 7.1.7 Zinc Borates The industrial importance of the zinc borates has increased substantially during the past decade. Although produced in small volumes compared to the other industrial borates discussed here, annual world production of zinc borates is now about ten thousand metric tons. At least seven unique crystalline forms of hydrated zinc borate are known, each of which can be prepared selectively by reaction of zinc oxide with boric acid in aqueous media under controlled conditions. The zinc borates Zn[B3O3(OH)5] Æ H2O (2 ZnO Æ 3 B2O3 Æ 7 H2O) and Zn[B3O4(OH)3] (2 ZnO Æ 3 B2O3 Æ 3 H2O) were mentioned above. The former compound has limited usefulness because it has a relatively low dehydration onset temperature of about 100 C. The most important commercial zinc borate, Zn[B3O4(OH)3], has an unusually high dehydration onset temperature of about 390 C, making it suitable for use in applications requiring relatively high processing temperatures. The composition of this zinc borate is now known to be 2 ZnO Æ 3 B2O3 Æ 3 H2O, but it is usually referred to as an article of commerce by the formula 2 ZnO Æ 3 B2O3 Æ 3.5 H2O (or ZB 2335) since its structure, shown in Fig. 12 above, was only recently revealed by a single-crystal X-ray study [45]. Other commercial zinc borates include one of composition 4 ZnO Æ B2O3 Æ H2O (dehydration onset ca. 415 C) [90] and another of approximate composition ZnO Æ B2O3 Æ  2 H2O (dehydration onset ca. 200 C). The last two zinc borates are discrete crystalline compounds that

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have not been structurally characterized. The exceptional high dehydration temperature of 4 ZnO Æ B2O3 Æ H2O allows for its use in applications requiring even higher processing temperatures. The principle uses of the zinc borate Zn[B3O4(OH)3] are as a polymer additive and preservative for wood composites, such as oriented strand board (OSB). As a polymer additive it functions as a fire retardant synergist and modifier of electrical and optical properties. Its function as a fire retardant additive is discussed further below. A substantial amount of Zn[B3O4(OH)3] is used to improve the tracking index, which is an important performance criterion for polymers, such as polyamides (nylon) and polybutyl teraphthalates (PBT), used in electrical applications. Many borate compounds exhibit strong, broad infrared absorption characteristics, and this property leads to the use of zinc borates to modify the optical properties of polymers. For example, zinc borates can be used to produce thermally insulating polymer products and to facilitate the laser marking of plastic parts. In contrast to zinc oxide, zinc borates have generally relatively low refractive indices, allowing the formulation of substantially transparent or translucent plastic products. This property is particularly striking for the zinc borate 4 ZnO Æ B2O3 Æ H2O [90], which contains approximately 80% ZnO by weight and yet has a low refractive index, permitting it to substitute for ZnO in applications where opacity is undesirable. 7.2 Beneficiated Mineral Borates

7.2.1 Colemanite The structure of the mineral colemanite {Ca[B3O4(OH)3] Æ H2O ¼ 2 CaO Æ 3 B2O3 Æ 5 H2O} is shown in Fig. 7a above. This is one of the most important industrial borate minerals, serving both as a raw material for the manufacture of boric acid and as a borate product available in various beneficiated grades. These mineral concentrate products contain from 27–42% B2O3 along with clay and refractory mineral impurities. Higher grades of commercially available colemanite concentrates contain about 42% B2O3, and are ground and blended to improve homogeneity. The principle use of beneficiated colemanite is the manufacture of textile fiberglass, discussed below. 7.2.2 Tincal Tincal, the mineral borax, is primarily used in the production of refined borax but also finds use as a mineral concentrate in some industrial applications where impurities are acceptable. Commercial concentrates, mostly produced in Turkey, typically contain about 32% B2O3.

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7.2.3 Kernite Kernite (Na2[B4O5(OH)2] Æ 3 H2O = Na2O Æ 2 B2O3 Æ 4 H2O) is also an important industrial mineral that is mainly used to produce refined borates, boric acid and borax. The polymeric borate structure of the kernite is shown in Fig. 7b above. Upon hydration, kernite converts to the monomeric borate tincal, which dissolves more readily. Thus, kernite may be hydrated to facilitate processing into borax. Although primarily a raw material for the manufacture of refined borates, concentrates of this mineral have been used as industrial products. 7.2.4 Ulexite Ulexite is a sodium calcium borate mineral of composition Na2O Æ 2 CaO Æ 5 B2O3 Æ 16 H2O (43.0% B2O3) and structural formula NaCa[B5O6(OH)6] Æ 5 H2O. Higher grades of this mineral are supplied as concentrates containing about 38% B2O3, mainly for use in the manufacture of fiberglass. Lower grades are used as boron micronutrient fertilizers. The structure of ulexite [26] is based on the isolated pentaborate trianion shown in Fig. 21. This polyborate anion is associated with one crystallographically unique 6-coordinate Na+ cation and one unique 8-coordinate Ca2+ cation. The cations form chains by sharing coordinated oxygen, and these chains link the polyanions together into parallel sheets. These sheets are associated by coordination with interstitial Ca2+ cations and hydrogen-bonding. 7.2.5 Hydroboracite The mineral hydroboracite is a calcium magnesium borate of composition CaO Æ MgO Æ 3 B2O3 Æ 6 H2O. Concentrates of this mineral typically containing about 44% B2O3 are primarily used in the manufacture of ceramic glaze compositions. This mineral, of structural formula CaMg[B6O8(OH)6] Æ 3 H2O, contains infinite polyborate chains based on the same ( FBB) type as found in

Fig. 21. Structure of the mineral ulexite

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colemanite [91] (Fig. 7a). These polyborate chains are linked by coordination with Ca and Mg into a three-dimensional framework that is further integrated by hydrogen bonding involving B-OH and H2O hydrogens. Thus hydroboracite is a harder material than colemanite, which has a layered structure and exhibits distinct cleavage planes. 7.2.6 Metamorphic Borates Although not actual borate products, metamorphic borates deserve mention because these minerals dominate the world’s known borate reserves [2]. Mostly borosilicates, these are called metamorphic borates because they were formed in secondary geological processes from primary alkali or alkaline earth metal borate minerals by incursion of silicate minerals. The most abundant of these are calcium borosilicates. Borosilicate minerals are not commonly used as commercial sources of refined borates because they are typically lower in B2O3 grade and more difficult to process than primary borates. Some borosilicate minerals as such datolite, 2 CaO Æ B2O3 Æ 2 SiO2 Æ H2O, and danburite, CaO Æ B2O3 Æ 2 SiO2, are used in the production of boric acid, most notably at Bor in Russia. Although most commercial borates are now produced from primary borate deposits, the borosilicate minerals may increase in importance in the future.

8 Major Industrial Uses of Borates 8.1 Vitreous Applications

The manufacture of glass and other vitreous products accounts for more than one half of all borate use worldwide. The history of borate use in glass manufacture has been reviewed [92]. 8.1.1 Fiberglass The manufacture of fiberglass is one of the largest industrial uses of boron. There are two primary types of fiberglass: insulation and textile fiberglass, which have different compositions, use different manufacturing methods, and consume different borate products. However, in each cases boric oxide serves both to improve quality and performance characteristics of the finished glass as well to facilitate the manufacturing process. Insulation fiberglass is glass wool used for thermal and acoustic insulation of homes, commercial buildings, industrial equipment, and automobiles. It has a sodium aluminoborosilicate-based composition, often with additional modifiers, such as calcium, magnesium, and potassium, which vary by

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producer and specific use. Since this is a sodium-containing glass composition, borax can be used to supply both Na2O and B2O3 to the glass batch. Insulation fiberglass typically contains 3–7% B2O3. Important performance characteristics of insulation fiberglass are R-Factor (an insulation value related infrared absorption), corrosion resistance, and resiliency. The presence of B2O3 in the glass composition contributes positively to each of these characteristics while also reducing fiberizing temperature and improving other manufacturing parameters. Textile fiberglass is a family of continuous filament glass products used for reinforcement applications in plastic composites, circuit boards, and pipe and tank wrappings [93]. In contrast to the sodium-rich insulation fiberglass compositions, textile fiberglass is a calcium aluminoborosilicate glass, which generally contains limited amounts of sodium. For this reason, either boric acid or colemanite are used as the B2O3 source in the manufacture of textile fiberglass. Textile fiberglass typically contains 5–10% B2O3. Corrosion and electrical resistance are two important performance characteristics of textile fiberglass. Over time the fiberglass industry has arrived at commercial glass compositions that simultaneously optimize both performance and production parameters. The simultaneous optimization of both durability and practical fiberization temperatures (the latter being typically 40–120 C above the liquidus temperature) has been the driving force behind the use of B2O3 in commercial fiberglass compositions. In the production of most commercial silicate glasses, Na2O, K2O, CaO, and MgO are important network modifiers used to adjust the melt viscosity, while Al2O3 is often used to improve glass durability. However, substitution of SiO2 by Al2O3 typically increases the liquidus temperature, as does CaO. On the other hand, both Na2O and B2O3 reduce the liquidus temperature upon SiO2 substitution, but Na2O has a strongly adverse influence on durability. In this balancing act it turns out that B2O3 has the unique property of reducing the workable fiberizing temperature without having a strongly negative effect on corrosion resistance of glass fibers. For these reasons it is difficult to manufacture fiberglass without the use of B2O3 [94]. Viscosity is one of the most important parameters in the glass-forming process. Tight control over viscosity must be maintained through precise temperature control. The presence of B2O3 in the glass batch results in a flattening of the temperature viscosity curve, allowing greater operating latitude in temperature control. This ultimately leads to improved productivity in the glass-making operation. 8.1.2 Specialty Glass Many different types of specialty borosilicate glasses are manufactured. Notable among these are heat-resistant glass, pharmaceutical glass, optical glass, and sealing glass. Heat-resistant glass, typically containing 12–15% B2O3, has a low thermal expansion coefficient for uses such as cookware and laboratory glassware. Pharmaceutical glass has a similar composition and is used for

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medicinal vials and medical appliances. In optical glasses B2O3 is an important component allowing formulation of specific optical dispersion requirements. Sealing glasses contain 8–30% B2O3 and are designed to have the appropriate thermal expansion and other engineering characteristics to form metal to glass seals, as needed in lighting fixtures such as automotive headlamps. 8.1.3 Ceramic Glazes/Porcelain Enamels Ceramic glazes are specialized glass formulations that are applied to the surfaces of ceramic bodies such as wall and floor tiles and other ceramic wares. Porcelain enamels are also special glass compositions used to coat metal surfaces such as appliances, cookware, sinks, and bathtubs. The compositions of these vitreous materials are designed for specific manufacturing and performance requirements. In most cases the borate is introduced in a frit glass, which is a borosilicate glass, containing perhaps 11–13% B2O3, that is ground to a powder for use in glaze and enamel formulations. 8.2 Agricultural Use

Deficiency of boron is a more prevalent worldwide agricultural problem than for any other agricultural micronutrient [62]. Adequate boron supplies are critical for maintaining high crop yields as well as good crop quality. For this reason it is common agricultural practice to apply boron-containing fertilizers in areas where soil boron deficiency is found. Boron deficiency is particularly prevalent in light-textured soils in which water-soluble borates are gradually leached down the soil profile and become unavailable to plants. Heavier, more loamy soils tend to retain more boron because they contain an abundance of compounds, such as humic acids, that can complex boron. Certain crop types have higher boron requirements and benefit most from supplementation. These include soybeans, cotton, peanuts, oil palm, apples, and almonds. In addition to borax, which can be applied directly as a fertilizer, special borate products are manufactured for fertilizer use. These include granulated beads for blending with other common fertilizer products and soluble borates for application as liquid foliar sprays. Boron fertilizers are typically applied at a rate equivalent to about one to two pounds elemental boron per acre, depending on the specific crop and soil conditions. As with most fertilizers, excessive application can have adverse results. Borate fertilizers are phyotoxic at high levels and are sometimes used as non-selective herbicides. 8.3 Biostatic Use

Although boron is essential for life, high concentrations of borates can inhibit the growth of harmful insects, bacteria, and fungi. The ability of borates to act

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as biocides or biostats without presenting a serious human health or environmental risk provides the basis for increasingly important areas of industrial use. 8.3.1 Wood Preservation Moderate concentrations of borates can inhibit the growth of wood-destroying fungi and repel attacks by wood-boring insects, such as termites, carpenter ants, and beetle larvae. The low mammalian toxicity of borates combined with their broad activity against organisms detrimental to wood makes them attractive in this application. As with other wood preservatives, borates are generally applied to whole lumber by pressure treatment using aqueous solutions. Higher solubility borates, particularly disodium octaborate, are typically used for this purpose. Because of their positive toxicity and performance characteristics borates are increasingly used as a safer alternative to traditional wood preservatives, such as the widely used Copper-Chrome-Arsenate (CCA), now being phased out of commercial use. In contrast to CCA, borates generally have the disadvantage of being more leachable from wood, and thus they lack persistence in outdoor exposure situations. Although effective in prolonging the outdoor functional lifetime of wood, borates are not usually applied to whole lumber intended for outdoor service. Instead, borates are generally applied to interior framing lumber. Much current research is directed toward developing methods of chemically ‘‘fixing’’ borates in wood to render them leach resistant [95]. The production of wood composite products, such as oriented strand board (OSB), particle board, wood-plastic extruded composites, etc., has increased dramatically in recent years. Although these products tend to have excellent engineering characteristics, they are often quite susceptible to wood destroying fungi. The mode of manufacture of these products provides an opportunity to incorporate sparingly-soluble borate wood preservatives. The zinc borate Zn[B3O4(OH)3] is particularly effective as a preservative for wood composite products, such as siding and sheathing used in home construction. The use of zinc borate in these products has increased significantly in recent years. 8.4 Fire Retardants

Borates are widely used in fire retardant applications. For example, cellulose insulation products used in homes and cotton batting used in mattresses and other furnishings are typically treated with boric acid to inhibit smoldering combustion. Borates are also used as fire retardants or fire retardant synergists in plastics, rubber products, and paints, where specialized borates such as zinc borate may be used. Fire retardants are generally considered to operate by one of two modes, referred to as vapor phase and solid phase mechanisms. Vapor phase fire

Borates in Industrial Use

35

retardants, usually a halogenated organic compound together with a synergist such as antimony oxide, volatilize during pyrolysis of the substrate and act as free radical scavengers to extinguish the flame. These additives have an important role to play but also have disadvantages. For example, although vapor phase fire retardants can be quite effective in controlling flaming combustion, they generally increase smoke evolution and may do little to control smoldering combustion. While vapor phase fire retardants may render materials less ignitable and/or more extinguishable, the intense radiant heat experienced in real fire situations can nevertheless force materials containing these additives to burn, resulting in the production of large amounts of smoke, often with deadly consequences. Solid phase fire retardants generally act by forming a char barrier at the surface of the substrate that inhibits the transmission of heat and oxygen. Carbonaceous materials may be diverted from fuel to incorporation into vitreous char, a process that can reduce smoke evolution. The process of char generation can ultimately extinguish a fire or at least significantly slow flame spread. The formation of a strong glassy char can also stabilize the burning substrate and prevent the formation of flaming drips and flying embers that can propagate fire. Borates, through their ability to act as glass network formers, can act as excellent char formers and drip suppressants in fire retardant applications. In many cases this involves processing into polymeric materials, leading to specific requirements for thermal stability and particle size. Most common borate materials, however, exhibit relatively low dehydration temperatures and may be unsuitable for use in many polymer systems. Zinc borates are often used because they have unusually high dehydration onset temperatures and can be produced as small particle size powders. 8.5 Perborates

The reaction of borate salts with hydrogen peroxide in alkaline aqueous solution results in formation of peroxoborates. This reaction is rapid and occurs by nucleophilic attack of the perhydroxyl anion on boric acid, Eq. (10) [96]. ½BðOHÞ4  þH2 O2 ¼BðOHÞ3 þ½HO2  þH2 O¼½ðHOÞ3 BOOH þH2 O

ð10Þ

The only significant peroxoborate species present in dilute solution is the [(HO)3BOOH]) anion (Fig. 22). Peroxoborate salts, commonly referred to as perborates, can be selectively crystallized from solution as various alkali and

Fig. 22. The [(HO)3BOOH]) anion

36

D. M. Schubert

Fig. 23. The structure of sodium perborate, x = 0 for PBS-1 (NaBO3 Æ H2O); x = 6 for PBS-4

(NaBO3 Æ 6 H2O)

alkaline earth metal salts. The most important of these are the sodium salts. Two sodium peroxoborates are available as articles of commerce, Na2[B2O4(OH)4] and Na2[B2O4(OH)4] Æ 6 H2O. These crystalline granular or powder materials are usually referred to in the literature, as well as the commercial realm, by their equivalent formulas NaBO3 Æ H2O (PBS-1) and NaBO3 Æ 4 H2O (PBS-4), respectively. Both of these borates contain the cyclic structure shown in Fig. 23. Peroxoborate salts can be used as oxidizing agents in a wide range of industrial and laboratory oxidations. Immediately upon dissolution in water, the peroxoborate structure shown in Fig. 23 hydrolyzes to the [(HO)3BOOH]) anion (Fig. 22), which further decomposes according to the reverse of Eq. (10). Since H2O2 predominates in solution, the perborate salts provide a means to supply active oxygen in a safe and convenient form. The use of sodium perborate as a versatile oxidizing agent for organic synthesis has been reviewed [97]. Beyond simply acting as a latent source of H2O2 the peroxoborate species appear to enhance reactivity in many cases. This may be a result of the ability of the [(HO)3BOOH]) anion to act as a donor of perhydroxyl anion, HO2), at lower pH than normally found with H2O2 alone. Hydrogen peroxide has a pKa value of about 11.6 and thus requires high pH levels to be effective in nucleophilic oxidations. In addition, tricoordinate boron that is present at moderate pH may act as a hydroxyl ion acceptor to assist reaction, Eq. (11), ½ðHOÞ3 B-O-O-H] þ E ! E-O-O-H þ B(OH)3 ! E¼O þ ½BðOHÞ4  ð11Þ where E is an electrophilic substrate. Oxidative processes have great industrial importance. The least expensive oxidizing agents are air, chlorine, and nitric acid. In practical applications, each of these may have disadvantages, such as a lack of selectivity, a need for special catalysts, or environmental problems. The next least costly oxidizing agent is hydrogen peroxide. This is a relatively weak oxidizer that may require activation for specific oxidations. Peroxoborates are widely used in connection with the supply of peroxide for oxidative applications, and can provide certain advantages. 8.5.1 Detergency – Consumer Bleaching The most important use of peroxoborates is as an active oxygen carrier in consumer laundry detergent formulations [98]. Domestic laundry products

Borates in Industrial Use

37

contain synthetic detergents and enzymes for removal of oily and proteinaceous stains. In addition, a bleach component is often added for removal of socalled ‘‘bleachable’’ stains. These stains are typically highly colored materials that bind strongly to fabric fibers and cannot be solubilized by detergents alone. Typical bleachable stains are materials present in tea, red wine, and fruit juices, such as theaflavins and anthocyanins. The bleach components are oxidizing agents that act to chemically degrade these staining materials and decolorize them. Important oxidizing agents in this context are the hypochlorite ion, OCl), and hydrogen peroxide, H2O2. Hypochlorite salts can be added to laundry formulations, but they are relatively unselective for stains and may damage fabrics. Hydrogen peroxide is a milder and more selective oxidizing agent. It is an effective bleaching agent above about 60 C, and below this temperature requires the presence of activators. Hydrogen peroxide is an unstable liquid that cannot easily be formulated into household laundry powders. Addition of sodium perborates, PBS-1 or PBS-4, to these formulations is now widely practiced in the detergent industry for the manufacture of heavy-duty laundry detergents. Perborates in these products are sometimes referred to as ‘‘color safe bleach’’. In most cases an activator is also added to allow effective bleaching at lower wash temperatures. The most common activator used in consumer products for this purpose is tetraacetylethylenediamine (TAED). This compound reacts with peroxide to form peracetic acid, which is a more effective oxidizer. 8.5.2 Detergency – Laundry Builder In addition to the use of peroxoborates in detergent bleach formulations, borates are also widely used in non-oxidizing detergents and other cleaning products. Appropriate alkalinity is important to the cleaning process. Borax and boric acid are used to adjust pH or provide a source of alkalinity in cleaning products. Many cleaning products contain borax as a builder, which provides alkalinity and pH buffering, electrostatic stabilization of suspended soil particles, reduction of interfacial tension between water and oily soils, and sequestration of calcium. The same benefits are found in peroxoboratecontaining detergents since the borate by-products provide the functions of a laundry builder [98, 99]. 8.6 Emerging Applications of Borates

Many new applications of borates continue to be developed. These include, among others, new uses in construction products, pulp and paper, and ceramics industries. Borates can be used as a carrier of alkalinity in the kraft pulping process, resulting in reduced raw materials and energy savings [100]. Although borates are traditional components of ceramic glazes, recent research shows that addition of borates to ceramic bodies can reduce firing

38

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temperature and required furnace residence time, increase firing temperature range, and increase green strength [101, 102]. Borates have unique and versatile properties which are being applied to the development of many other new applications are being developed.

9 References 1. Argust P (1998) Biological Trace Element Research 66: 131 2. Garrett DE (1998) Borates, Handbook of Deposits, Processing, Properties, and Use, Academic Press, San Diego 3. Brotherton RJ (1994) In: King RB (ed), Encyclopedia of Inorganic Chemistry. John Wiley & Sons, Chichester, 1: 374 4. Schubert DM (1993) Boron Compounds – Boron Hydrides, Heteroboranes and Their Metalla Derivatives. In: Kirk-Othmer Encyclopedia of Chemical Technology, 4th Edn, John Wiley & Sons, New York and 5th Edn (2002) published on-line 5. A notable exception is the icosahedral ‘‘aromatic’’ cluster [B12(OH)12]2) and related compounds which contain twelve equivalent BAO bonds, see Maderna A, Knobler CB, Hawthorne MF (2001) Angew Chem Int Ed 40: 9 6. Geisinger KI, Gibbs GV, Navrotsky A (1985) Phys Chem Minerals 11: 266 7. Navrotsky A (1996) In: Grew ES, Anovitz LM (eds), Reviews in Mineralogy. The Mineralogical Society of America, Washington DC 33: 165 8. Ingri N (1963) Sven Kem Tidskr 75: 199 9. Christ CL (1960) Am Mineralogist 45: 334 10. Edwards JO, Ross VF (1960) J Inorg Nucl Chem 15: 329 11. Tennyson C (1963) Fortschr Mineral 41: 64 12. Heller G (1970) Fortschr Chem Forsch 15: 206 13. Burns PC (1995) Can Mineralogist 33: 1167 14. Grice JD, Burns PC, Hawthorne FC (1999) Can Mineralogist 37: 731 15. Simonov MA, Kazanskaya EV, Egorov-Tismenko YK, Zhelezin EP, Belov NV (1976/77) Dokl Akad Nauk SSSR 91: 230 16. Paton F, McDonald SGG (1957) Acta Crystallogr 10: 653 17. Simonov MA, Egorov-Tismenko YK, Belov NV (1976) Soviet Phys-Cryst 21: 332 18. Kazanskaya EV, Chemodina TN, Egorov-Tismenko, Yu K, Simonov MA, Belov NV (1977) Sov Phys Crystallogr 22: 35 19. Coulson CA (1964) Acta Crystallogr 17: 1086 20. Dal Negro A, Martin-Pozas JM, Ungaretti L (1975) Am Mineralogist 60: 879 21. Corraza E (1976) Acta Cryst B 32: 1329 22. Ozols YK, Tetere IV, Ievins AF (1973) Akad Nauk Latv SSR, Ser Kim 7: 3 23. Simonov MA, Egonov-Tismenko YK, Kazanskaya EV, Belokoneva EL, Belov NV (1978/79) Sov Phys-Dokl 23: 159 24. Callegari A, Mazzo F, Tadlini C (2001) Can Mineralogist 39: 139 25. Merlino S, Santori F (1972) Acta Cryst B 28: 3559 26. Ghose S, Wan C, Clark JR (1978) Am Mineralogist 63: 160 27. Wan C, Ghose S (1977) Am Mineralogist 62: 1135; Erd RC, McAllister JF, Eberlein GD (1979) Am Mineralogist 64: 369 28. Hanic F, Lindqvist O, Nyborg J, Zedler J (1971) Coll Czech Chem Commun 36: 3678; Dal Negro A, Ungarretti L, Sabelli C (1971) Am Mineralogist 56: 1553 29. Burns PC, Hawthorne FC (1993) Can Mineralogist 31: 297 30. Cooper WF, Larsen FK, Coppens P, Giese RF (1973) Am Mineralogist 58: 21 31. Erd RC, McAllister JF, Vlisidis AC (1961) Am Mineralogist 46: 560 32. Schubert DM (1997) US Patent 5,688,481 33. Burns PC, Grice JD, Hawthorne FC (1995) Can Mineralogist 33: 1131

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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. 59. 60. 61. 62. 63. 64. 65. 66. 66. 67. 68. 69. 70. 71. 72. 73. 74.

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Hawthorne FC (1985) Am Mineral 70: 455 Hawthorne FC (1986) Can Mineral 24: 625 Hawthorne FC (1990) Can Mineral 192: 1 Schindler M, Hawthrone FC (2001) Can Mineralogist 39: 1225 Hawthorne FC (1992) Z Kristallogr 201: 183 Bowden GH (1980) Boron Oxygen Compounds. Mellor’s Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol V. Longman Group, London Ammonium tetraborate is known as an article of commerce by the name ‘‘ammonium biborate’’ Merlino S, Sartori F (1971) Science 171: 377 Schubert DM, Visi MZ, Knobler CB (1999) Inorg Chem 39: 2250 Rumanova IM, Razmanova ZP, Belov NV (1972) Sov Phys Dokl 16: 518 Yamnova NA, Egorov-Tismenko YK, Pushkarovskii DY, Malinko SV, Dorokova GI (1993) Krystallogrfiya 38: 71 Schubert DM, Alam F, Visi M, Knobler CB (2003) Chemistry of Materials, in press Bray PJ (1997) In: Wright AC, Feller SA, Hannon AC (eds), Borate Glasses, Crystals & Melts. Society of Glass Technology, Sheffield Teter M (1997) In: Wright AC, Feller SA, Hannon AC (eds), Borate Glasses, Crystals & Melts. Society of Glass Technology, Sheffield Kroeker S, Stebins JF (2001) Inorg Chem 40: 6239 Uhlmann DR, Shaw RR (1969) J Non-Cryst Solids 1: 347 Krogh-Moe J (1965) Phys Chem Glasses 6: 46 Chryssikos GD, Kamitsos EI (1997) In: Wright AC, Feller SA, Hannon AC (eds), Borate Glasses, Crystals & Melts. Society of Glass Technology, Sheffield Hunter DL, Steinberg H (1960) US Patent 1,203,698 Groszos SJ, Day NE (1960) US Patent 2,942,021 Wei PW, Atwood DA (1998) Inorg Chem 37: 4934 Knoeck J, Taylor JK (1969) Anal Chem 41: 1730 Acree TE (1973) Adv Chem Ser 117: 208 Okami Y, Okazaki T, Kitahara T, Umezawa H (1976) J Antibiot 29: 1019; Nakamura H, Iitaka Y, Kitahara T, Okazaki T, Okami Y (1977) J Antibiot 30: 714 Hutter R, Keller-Schierlein W, Knusel F, Prelog V, Rodgers GC Jr, Suter P, Vogel G, Voser W, Zahner H (1967) Helv Chim Acta 50: 1533; Dunitz JD, Hawley DM, Miklos D, White DNJ, Berlin Y, Marcesic R, Prelog V (1971) Helv Chim Acta 54: 1709 Hemsheidt T, Puglisi MP, Larsen LK, Patterson GML, Moore RE, Rois JL, Clardy J (1994) J Org Chem 59: 3467 Herbert I, Schummer D, Gerth K, Hofle G, Reichenbach H (1995) J Antibiotics 48: 26 Warington K (1923) Ann Bot 37: 627 Blevins DG, Lukaszewski KM (1998) Annu Rev Plant Physiol Mol Biol 49: 481 Brown PH, Hu H (1996) Annals of Botany 77: 497 Hu H, Brown, PH (1997) Plant Physiol 113: 649 Hu H, Penn SG, Lebrilla CG, Brown PH (1997) Plant Physiol 113: 649 O’Neill MA, Eberhard S, Albersheim P, Darvill AG (2001) Science 284: 846 Kobayashi M, Matoh T, Azuma (1996) J Plant Physiol 110: 1017 O’Neill MA, Warrenfeltz D, Kates K, Pellerin P, Doco T, et al. (1996) J Biol Chem 271: 22923 Ishi T, Matsunaga T (1996) Carbohydr Res 284: 1 Kaneko S, Ishi T, Matsunaga T (1997) Phytochemistry 44: 243 Jackson JF (1991) In: Randall DD, Blevins DG, Miles CD (eds), Current Topics in Plant Biochemistry and Physiology. Columbia Univ Mo press, 10: 221 Lawrence K, Bhalla DG, Misra PC (1995) J Plant Physiol 146: 1143 Barr R, Bottger M, Crane FL (1993) Biochem Mol Biol Int 31: 31 Barr R, Crane FL (1991) In: Randall DD, Blevins DG, Miles CD (eds), Current Topics in Plant Biochemistry and Physiology. Columbia Univ Mo press, 10: 290 Morre DJ, Navas P, Penel C, Castillo FJ (1986) Protoplasma 133: 195

40 75. 76. 77 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96. 97. 98. 99. 100. 101. 102.

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Brown JC (1979) Plant Nutr 1: 39 Cakmak I, Kurz H, Marschner H (1995) Physiol Plant 95: 11 Shkolnik MY (1984) Trace Elements in Plants, New York, Elsevier Levy HA, Lisenky GC (1978) Acta Crystallogr B34: 3502 Fort DJ, Propst TL, Stover EL, Strong PL, Murray FJ (1998) Biological Trace Element Research 66: 237 Rowe RI, Bouzan C, Nabili S, Eckhert CD (1998) Biological Trace Element Research 66: 261 Chapin RE, Ku WW, Kenney MA, McCoy H (1998) Biological Trace Element Research 66: 395 Sutherland B, Strong P, King JC (1998) Biological Trace Element Research 66: 193 Black JA, Barnum JB, Birge WJ (1993) Chemosphere 26;1382 Richold M (1998) Biological Trace Element Research 66: 121 Hubbard S (1998) Biological Trace Element Research 66: 343 Powell DR, Gaines DF, Zerella PJ, Smith RA (1991) Acta Crystallogr C47: 2279 Levy AH, Lisensky GC (1978) Acta Crystallogr B34: 3502 Zachariesen WH (1954) Acta Crystallogr 7: 305 Bayer AJ, Erdemir A (1991) Adv Mater Process 140: 40 Schubert DM (1995) US Pat 5,472,644; Schubert DM (1994) US Patent 5,342,553 Sabelli C, Stoppioni A (1978) Can Mineralogist 16: 75 Smith RA (1997) In: Wright AC, Feller SA, Hannon AC (eds), Borate Glasses, Crystals & Melts. Society of Glass Technology, Sheffield Loewenstein KL (1993) The Manufacturing Technology of Continuous Glass Fibers, 3rd edn., Elsevier, Amsterdam Harding FL, Bauer JF, Russel HH, Xu, X (1997) In: Wright AC, Feller SA, Hannon AC (eds), Borate Glasses, Crystals & Melts. Society of Glass Technology, Sheffield See for example, Schubert DM, Manning MJ (1997) US Patent 5,612,094 Pizer R, Tihal C (1987) Inorg Chem 13: 117 McKillop A, Sanderson WR (1995) Tetrahedron 51: 6145 Greenhill-Hooper MJ. Tenside Surfactants Deterg 33: 366 Greenhill-Hooper MJ (2000) World Surfactants Congr 5: 664 Tran H, Bair CM, McBroom RB, Strang W, Morgan B (1999) Proceeding of the 1999 TAPPI Engineering Conference 1: 163 Noirot MD (1999) Amer Ceram Soc Bull, August Cook SG (2002) Ceramic Engineering and Science Proceedings 23: 47

Aluminum and Gallium Hydrazides Werner Uhl Fachbereich Chemie der Philipps-Universita¨t Marburg, Hans-Meerwein-Straße, 35032 Marburg, Germany e-mail: [email protected]

Aluminum and gallium hydrazides have found considerable interest in current research because they are potentially useful as single source precursors for the epitaxial growth of semi-conducting AlN or GaN films. Furthermore, they may exhibit singular structural properties owing to the capability of the hydrazido group to act as a bidentate ligand. The synthesis of those hydrazides succeeds by several methods, such as the alkane and hydrogen elimination, which requires the treatment of alanes or gallanes with hydrazine derivatives containing at least one N-H function. Furthermore, the formation of salts by the employment of lithium hydrazides and chloroaluminum or -gallium compounds or the hydroalumination of 2,3-diazabutadiene derivatives are facile synthetic methods. These hydrazides adopt a great variety of structures with four-, five-, and six-membered heterocycles or with polycyclic frameworks and cages. With the exception of sterically insufficiently shielded derivatives of unsubstituted hydrazine N2H4, their thermal stability is high enough to allow a secure handling. Only at elevated temperature does thermolysis occur, which in some cases yielded the nitrides AlN and GaN. Keywords: Aluminum, Gallium, Hydrazine

1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

42

2

Hydrazine Adducts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

43

3

Aluminum Hydrazides Derived from Free Hydrazine N2H4 . . . . . .

45

4

Reactions with Hydrazine Derivatives . . . . . . . . . . . . . . . . . . . . . .

46

4.1 Alkane Elimination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2 Hydrogen Elimination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.3 Salt Elimination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

46 50 55

5

Formation of Hydrazides by Hydroalumination . . . . . . . . . . . . . .

57

6

Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

60

6.1 Synthetic Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.2 Structural Motifs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.3 Structural Parameters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

61 61 65

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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Structure and Bonding, Vol. 105
Springer-Verlag Berlin Heidelberg 2003

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1 Introduction The synthesis of a few organoelement aluminum and gallium hydrazides was reported for the first time about 40 years ago. However, their characterization remained insufficient, in particular with respect to the prediction of conclusive molecular structures [1–5]. They have found renewed interest in current research, which essentially may be caused by two reasons: (i) The coordination chemistry of aluminum and gallium amides has been investigated intensively, and a fascinating variety of different compounds and structures has been reported. Owing to their capability to act as bidentate ligands hydrazido derivatives may open the access to an even larger variability of molecular structures with the formation of novel rings and cages. (ii) Furthermore, aluminum and gallium hydrazides may find some technical applications in chemical vapor deposition processes as starting compounds for the formation of the important inorganic materials aluminum nitride (AlN) and gallium nitride (GaN). Both compounds are wide-gap III–V semiconductors and may be employed, for instance, in high-temperature devices or for the high-quality surface passivation of other III–V semiconductors such as GaAs [6]. Owing to its bandgap of 3.4 eV GaN is a promising material for near-ultraviolet optoelectronic devices. In conventional OMCVD processes these compounds are formed by the reaction of trialkylelement derivatives with ammonia. These reactions were usually performed at very high temperatures (>900 C), which may result in some thermal stress and nitrogen deficiency in the deposited layers. Lower temperatures are also required for the epitaxial growth of the nitrides on the surface of other semiconductors such as GaAs, in order to prevent their decomposition. In the recent literature, several reports appeared which describe the deposition of AlN and GaN on GaAs or other surfaces by the thermolysis of trimethylaluminum or trimethylgallium in the presence of hydrazine or dimethylhydrazine under milder reaction conditions [7–12]. GaN was formed at temperatures above 450 C. However, thorough investigations did not reveal a significant difference to the corresponding reaction with ammonia. The rate-determining step is the surface-catalyzed decomposition of the organometallic component in all cases. Only the excess of ammonia must be larger than that of hydrazine [11, 12]. AlN was deposited at very low temperature (220 C) [11]. Also the deposition of GaAs and GaN epilayers has been reported with triethylarsine and diethylarsine as As sources and dimethylhydrazine as the nitrogen source [13]. Thus, with respect to the potential application of hydrazine derivatives for the formation of AlN and GaN films by OMCVD processes it may be of particular interest to synthesize simple aluminum and gallium hydrazides as organoelement single source precursors [14]. However, this report is not directed to that possible technical application of those compounds, but it is focused merely on their syntheses and quite interesting structural features.

Aluminum and Gallium Hydrazides

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2 Hydrazine Adducts The intermediate formation of adducts of hydrazine derivatives with alkylaluminum or -gallium compounds containing coordinatively unsaturated aluminum or gallium atoms was postulated in many reactions [1, 3, 4, 7, 10]. However, owing to their high reactivity these adducts often gave a secondary reaction by the transfer of a proton from hydrazine to a carbanionic alkyl group, which resulted in the release of an alkane molecule. Pure adducts were isolated and completely characterized in few cases only. Treatment of the dialkylaluminum chlorides R2AlCl [R = CMe3, CH2-CMe3, CH(SiMe3)2] with bis(trimethylsilyl)hydrazine afforded stable adducts [1 to 3, Eq. (1)] in high yields, two of which were characterized by crystal structure determinations (1, 3). By dismutation tris(trimethylsilyl)hydrazine and trimethylsilylhydrazine were formed, the last one is unstable in solution at room temperature and stabilized here by coordination to the aluminum atoms of the chlorides. The central atoms of these compounds have a distorted tetrahedral coordination sphere with long Al-N distances (>201 pm) which are characteristic of AlAN ‘‘dative’’ bonds. Hydrogen bonds between the chlorine atoms and N-H protons of the b-nitrogen atoms may be derived from the conformation of the molecules, the elongated AlACl bond lengths (>218.5 pm) and short intramolecular H


N contact distances of 260 to 274 pm. A similar adduct, Me3C-AlCl2 Æ H2N-N(SiMe3)2, containing the intact starting compound 1,1-bis(trimethylsilyl)hydrazine was isolated in trace amounts as a by-product of the synthesis of 1 [15]. Tert-butylhydrazine as a ligand was observed in the compound (Me3C)2AlCl Æ NH2-N(H)CMe3 [16], in which the sterically less shielded nitrogen atom is attached to the aluminum atom and which has structural features quite similar to those of 1 and 3 including a possible NH


Cl interaction. All these stable compounds have at least one electronegative chlorine atom bonded to aluminum.

ð1Þ

44

W. Uhl

Several complexes of hydrazine derivatives with trialkylaluminum or gallium compounds were reported in literature, however, to the best of my knowledge, they were not characterized by crystal structure determinations. An adduct of tetramethylhydrazine and trimethylaluminum was characterized by 1H-NMR spectroscopy and molar mass determination [1, 3]. Gallium compounds of the type R3Ga Æ NH2-NR¢R¢¢ [4 to 7, Eq. (2)] were obtained almost quantitatively by mixing of the components at room temperature [17]. They were characterized by elemental analyses, NMR, and IR spectroscopy. In all cases, the sterically less shielded NH2 nitrogen atom was suggested to coordinate to the gallium atoms. While tert-butylhydrazine H2N-N(H)CMe3 forms an adduct (5) that is stable at room temperature [17], the analogous phenylhydrazine derivative gave a rapid secondary reaction under similar conditions (see below) [18].

ð2Þ

Another type of adducts [8, Eq. (3)] was formed by the reaction of di(tertbutyl)aluminum chloride with dilithium bis(trimethylsilyl)hydrazide in low yields below 30% [19]. The structure of 8 consists of a distorted heterocubane with four vertices occupied by nitrogen atoms, two of which are connected by an intact NAN bond across one face of the cube. The cation positions are occupied by two aluminum and two lithium atoms, of which the last ones bridge the NAN bond. Part of the hydrazide molecules was cleaved, and the aluminum atoms are bonded to one tert-butyl group only. On the basis of the NMR spectroscopic characterization many unknown by-products were formed in the course of that reaction, and no information is available concerning the reaction mechanism. Compound 8 may be described as an adduct of dilithium bis(trimethylsilyl)hydrazide to a dimeric iminoalane containing a four-membered Al2N2 heterocycle. Further

Aluminum and Gallium Hydrazides

45

derivatives were obtained in which the lithium atoms are coordinated by ether molecules or by di(tert-butyl)chloroaluminum moieties [19]. However, these reactions of aluminum chlorides with dilithium bis(trimethylsilyl)hydrazide proceed in a rather non-specific way and some aluminum amides were formed by complete cleavage of the NAN bonds under different conditions [19].

ð3Þ

3 Aluminum Hydrazides Derived from Free Hydrazine N2H4 In few cases only free hydrazine has been employed in the synthesis of aluminum or gallium hydrazides, which is mainly due to the extremely hazardous properties of the products isolated [2, 5, 20]. Trimethylaluminum or aluminum hydride were reported to yield solids which are dangerously explosive upon touch. These products are only poorly characterized and all proposed formulas are rather speculative. An adduct of H2N-NH2 with triisopropoxyaluminum was isolated and characterized by elemental analysis [4]. Recently, stable alkylaluminum hydrazides were described which had two bulky tert-butyl groups attached to their aluminum atoms [21]. Di(tertbutyl) aluminum hydrazide 9 was obtained in 86% yield by the reaction of di(tert-butyl)aluminum hydride with hydrazine at low temperature ()100 C) and slow warming to room temperature [Eq. (4)]. Compound 9 is a dimer in the solid state and has a six-membered Al2N4 heterocycle in a twist-conformation possessing two intact NAN bonds. Further treatment of 9 with two equivalents of the hydride gave a singular tricyclic, aluminum-rich compound [10, Eq. (4)] in which both NAN bonds of 9 are bridged by two

46

W. Uhl

di(tert-butyl)aluminum groups [21]. All Al-N distances of 10 are in a very narrow range (194 to 197 pm). Both compounds 9 and 10 are quite stable and decompose slowly only above 110 and 160 C, respectively, to yield a gray powder. No sensitivity to mechanical shock was observed, and even larger samples of both compounds (gram scale) can be thermolyzed without indications of a spontaneous, explosive decomposition. Attempts at sublimation remained unsuccessful.

2 (Me3C)2AlH + 2 N2H4

H

H

- 2 H2

H

N

N CMe3

Me3C Al Me3C

+ 2 (Me3C)2AlH

Al N

H

H

- 2 H2

CMe3

N H

ð4Þ

9

Me3C H Me3C

CMe3 H

Al N

CMe3

N Al

Al N

Me3C

N Al

H

CMe3 H

CMe3

Me3C 10

4 Reactions with Hydrazine Derivatives 4.1 Alkane Elimination

By far most of the reactions have been done with substituted hydrazines owing to their higher thermal stability. The most simple method for an

Aluminum and Gallium Hydrazides

47

effective synthesis of aluminum and gallium hydrazides is the reaction of trialkylelement derivatives with hydrazines that contain at least one protic hydrogen atom attached to nitrogen so that alkanes can be eliminated. Accordingly, trimethylaluminum or the corresponding gallium compound reacts with 1,1-dimethylhydrazine [22] or 1,1-diphenylhydrazine to yield the hydrazides 11 to 14 [23] [Eq. (5)]. Compound 11 was obtained at room temperature already, while the synthesis of the remaining compounds required hot solutions (100 C in heptane for 12, boiling dichloromethane for 13, boiling toluene for 14). Three compounds were characterized by crystal structure determinations (11, 13, and 14). They have four-membered centrosymmetric E2N2 heterocycles with two exocyclic NAN bonds. All molecules exhibit a trans-configuration with the NAN bonds on different sides of the heterocycles in the solid state. In solution, cis/trans isomerization occurred with free energies of activation between 10 and 15 kcal/mol. Compound 11 was sublimed in vacuum without decomposition. Such dimeric structures containing four-membered heterocycles with exocyclic NAN bonds were proposed in an early publication on aluminum hydrazides [1].

ð5Þ

Alkane elimination was also observed upon heating of the adducts R3Ga Æ NH2-NR¢R¢¢ [4 to 7, Eq. (2)] in boiling toluene, which gave the hydrazides 15 to 18 in excellent yields [Eq. (6)] [17]. One compound (17) was characterized by a crystal structure determination. Similar gallium hydrazides were obtained directly at room temperature without the isolation of stable adducts when the phenylhydrazine derivatives H2N-N(H)C6H5 [18] and H2N-N(C6H5)2 [17] were employed. The products (15 to 18) form dimers

48

W. Uhl

with Ga2N2 heterocycles and exo NAN bonds in a trans arrangement. Similar to the compounds described before, a temperature-dependent rearrangement was observed in solution with an equilibrium between cis and trans forms and an average NMR spectrum at elevated temperatures. Mixing of two different dimeric hydrazides gave a new compound, which was identified as the mixed species. A ring-opening mechanism was suggested for all these exchange processes [17, 22].

ð6Þ

Prolonged heating of the dimeric derivatives [Me2GaNHNHR]2 (R = CMe3 (18) [17] or C6H5 [18]) in boiling toluene for 36 hours afforded singular cage compounds [19 and 20, Eq. (7)] by release of further molecules of methane. The structures consist of two hexagonal rings in boat conformation bound together by four GaAN bonds. The gallium atoms and four nitrogen atoms have a coordination number of four in a distorted tetrahedral environment. These nitrogen atoms are in a bridging position between two gallium atoms, while the remaining four nitrogens are attached to one gallium only. These compounds may be described as dimers of the digallium dihydrazides [Me-Ga{N(H)NR}2GaMe] (R = CMe3, C6H5) possessing a six-membered Ga2N4 heterocycle and two endocyclic NAN bonds of two hydrazido dianions. Compounds 19 and 20 were formulated as intermediates on the way from the normally obtained Ga2N4 dimers (e.g., 11 to 18) to the inorganic material GaN, and further heating to 700 C did indeed give hexagonal GaN along with unidentified volatile organic products [18].

Aluminum and Gallium Hydrazides

49

ð7Þ

A dialuminum hydrazide with two dimethylaluminum groups coordinated to one nitrogen atom was formulated as an intermediate in the reaction of 1,1dimethylhydrazine with two equivalents of trimethylalane which proceeds by evolution of methane [24]. This compound was not isolated and characterized, but upon treatment with warm acetonitrile a secondary reaction occurred which afforded a remarkable heterocyclic product possessing a structure similar to that of the porphine ring system [21, Eq. (8)] [24]. Compound 21 was isolated in 40% yield. It contains eight aluminum atoms in a macrocycle and four almost planar AlCN3 heterocycles including the intact NAN bond of the hydrazine. The crystal structure determination revealed a strongly folded molecule with the five-membered rings alternately above and below the average molecular plane. However, in solution a fast exchange process was observed with a splitting of the 1H NMR resonances at 265 K. The activation energy was estimated to 64 kJ/mol.

50

W. Uhl

4 CH3CN

8 Me3Al + 4 H2N-NMe2

- 8 CH4

Me2 Al

Me2 N N

Me2Al

N C

C

N

Me2 Al

Me

Me

Me

Me

NMe2 N

Me2Al

ð8Þ AlMe2

N

C

Me2N

C N

Al Me2

N AlMe2

N Al Me2

N Me2

21

4.2 Hydrogen Elimination

Hydrogen evolution is a similar powerful method for the preparation of hydrazides, although it was employed to a lesser extent than the release of alkanes. Two reactions of this type were described above in conjunction with compounds derived from free hydrazine [9 and 10, Eq. (4)] [21]. Treatment of an excess of 1,1-dimethylhydrazine with the trihydrido compound H3Ga Æ NMe3 at room temperature afforded the dimeric dihydridogallium hydrazide 22 in an almost quantitative yield [Eq. (9)] [25]. The reaction was remarkably slow and required a period of two days. Compound 22 has a centrosymmetric structure with a four-membered Ga2N2 heterocycle and two exocyclic NAN bonds. It melted at 62 C and remained stable as a liquid up to 90 C. Decomposition was observed at higher temperatures and resulted in the formation of elemental gallium. However, another product [23, Eq. (9)] was obtained in 63% yield when 22 was heated under reflux in 1,1-dimethylhydrazine (70 C) as a solvent for four days [25]. By the further deprotonation of hydrazine molecules and evolution of hydrogen a very interesting new cage compound was formed in which two Ga2N2 heterocycles are bridged by two dianionic hydrazido ligands Me2N-N2). Their negatively charged nitrogen atoms are coordinatively unsaturated and bridge two gallium atoms (see for comparison compound 24 below). They adopt a trigonal planar geometry with the sum of the angles being 360 , which may be caused by the high ionic character of the GaAN bonds or by the delocalization of the lone-pairs to the

51

Aluminum and Gallium Hydrazides

gallium atoms. All hydrogen atoms attached to gallium were consumed, and the remaining terminal positions at the gallium atoms of 23 were occupied by -NH-NMe2 groups. The four-membered rings are slightly tilted, and their exocyclic NAN bonds adopt a trans position similar to the monocyclic compound 22. NAH


N hydrogen bonding was discussed between NMe2 and bridging NH groups. 4 H3Ga . NMe3 + 4 H2N-NMe2

Me

Me N

H

N

H 2

- 4 H2, - 4 NMe3

H

H2N-NMe2

Ga

Ga H

- 8 H2 H

N H

N

Me

Me 22

ð9Þ Me

Me N

H N

N

Me

Me N

H

Ga H

Me

N

N

Me

Me

N

Me Me

H N

Me

H

N

Ga

Me N

Me N

N

N

Me

Ga

N H

Me

Ga Me

Me N Me

N

N

N H

N Me

H

N Me

Me

23

A similar reaction with the formation of elemental hydrogen was observed when dimethylaminoalane (Me2NAlH2)3 was treated with 1,1-dimethylhydrazine [Eq. (10)] [26]. All dimethylamino substituents were replaced by terminal hydrazido groups (24; 66% yield). Dianionic hydrazido ligands analogous to compound 23 were formed. They occupy bridging positions

52

W. Uhl

between two and three aluminum atoms with trigonal pyramidal and distorted tetrahedral geometries at the inner nitrogen atoms, respectively. A ladder-type structure with three fused four-membered Al2N2 heterocycles, unique in aluminum or gallium hydrazide chemistry, had formed in which two Al2N2 chains are connected by four AlAN bonds. As a quite remarkable structural feature the side-on coordination of NAN bonds via the lone pairs of the nitrogen atoms was observed with relatively long Al-N distances in the chain (209.1 pm) and shorter ones (195.7 pm) to the NMe2 nitrogen atoms in b-position [see drawing in Eq. (10)]. Thus, two aluminum atoms possess a coordination number of five in a highly distorted trigonal bipyramidal coordination sphere.

4/3 (Me2NAlH2)3 + 10 Me2N-NH2

Me Me Me

H

N N

Me

Me N N

Me

- 4 Me2NH, - 9 H2

Me

H

N

Al

Me N

N

N

H

Al

N

Me

ð10Þ

N Me

H

Me N

N

Me

Al

N N

N H N Me

Al

Me

Me

Me Me

N

N N H

Me

Me

N Me

24

In Eq. (10) the simultaneous release of an amine and hydrogen led to the formation of an aluminum hydrazide possessing aluminum atoms exclusively coordinated by hydrazido ligands. The simultaneous liberation of alkane and hydrogen was observed in the reaction of bis(trimethylsilyl)hydrazine with di(tert-butyl)aluminum hydride [Eq. (11)] [27]. Compound 25 was isolated, which, to the best of my knowledge, is the only monoalkyl dihydrazide known in the literature. The central atom is attached to one carbon atom of a tertbutyl group and to the two negatively charged nitrogen atoms of two hydrazido ligands. Similar as in 24 the aluminum atom has a coordination number of five by the side-on coordination of the NAN bonds via their lonepairs to give two three-membered AlN2 heterocycles. The Al-N distances differ strongly (181.4 to 207.2 pm), the larger ones correspond to the ‘‘dative’’ bonds between aluminum and the N(H)SiMe3 groups. The hydrazido ligands are not in a plane, nevertheless, the coordination geometry resembles more that of a

53

Aluminum and Gallium Hydrazides

square pyramid than that of a trigonal bipyramid. The nitrogen atoms attached to hydrogen have a strongly distorted tetrahedral configuration, and the environment of the amido nitrogen atom is not planar (sum of the angles 347.4). Compound 25 was isolated in 60% yield; it decomposes above 130 C. Sublimation in vacuum succeeded at 100 C but was accompanied by considerable decomposition [27].

(Me3C)2AlH + 2 N2H2(SiMe3)2

SiMe3 N Me3Si

H

N

ð11Þ Al

CMe3

N H

+ H2 + HCMe3

SiMe3

N SiMe3 25

Another type of compounds resulted when lithium alanate LiAlH4 was employed instead of the neutral alanes or gallanes considered so far. The reaction with 1,1-bis(trimethylsilyl)-2-phenylhydrazine, which contains only one protic hydrogen atom, afforded the lithium trihydridohydrazidoaluminate 26 by the release of one equivalent of hydrogen [28]. This compound is dimeric in the solid state possessing an eight-membered heterocycle in which two lithium and two aluminum atoms are bridged by four hydride ions. The hydrazido ligands are in terminal positions. Probably owing to some hyperconjugative interactions, its nitrogen atoms have an almost ideally trigonal planar surrounding with the sum of the bond angles of 359.1 for the nitrogen atoms attached to aluminum and 358.3 for the remaining nitrogen atoms. A six-membered Al2N4 heterocycle [27, Eq. (12)] was obtained by the reaction of LiAlH4 with the diprotic hydrazine derivative Me3C(H)N-N(H) SiMe3 [28]. Both protons were replaced by aluminum atoms, and a dialanate was isolated with two lithium counter ions. One cation is located above the ring and bridges those two opposite nitrogen atoms which are attached to tert-butyl groups. Owing to a less strong hyperconjugative interaction these atoms may have a higher negative charge than the silyl-substituted ones.

54

W. Uhl

OEt2

Et2O

C6H5

SiMe3

Li H

H Me3Si

H

Al N

N

Al

N

Me3Si

N

H

C6H5

H

Li

Et2O

SiMe3 H

OEt2 26

+ 4 Et2O Me3Si + 2

C6 H 5 N

- 2 H2

N H

Me3Si

2 LiAlH4

ð12Þ

+ 4 THF Me3C

SiMe3 N

+ 2

- 4 H2

N H

H

(THF) Li Me3C

H N

Me3Si

Al

N

H N

H

Al

Li(THF)3 SiMe3

N CMe3

H 27

Furthermore, this lithium ion is coordinated by two hydrogen atoms of different AlH2 moieties. The second lithium atom bonds to a terminal hydrogen atom and to three THF molecules.

Aluminum and Gallium Hydrazides

55

4.3 Salt Elimination

Salt elimination is a further powerful and widely employed method for the synthesis of aluminum and gallium hydrazides. Monomeric compounds, which were stabilized by adduct formation with THF or diethyl ether, were obtained by the reaction of lithium hydrazides LiN2R3 (R = SiMe3, C6H5) with dimethylaluminum chloride [Eq. (13)] [28]. While the aluminum atoms have a tetrahedral coordination sphere, the geometry at the nitrogen atoms of all products (28 to 30) is almost ideally trigonal planar with sums of the angles near 360. This may be caused by hyperconjugative interactions with the phenyl or trimethylsilyl substituents.

ð13Þ

In the absence of a donor solvent the coordinative saturation of the aluminum atoms can only be achieved by dimerization. Thus, the reaction of lithium trimethylsilyl-tert-butylhydrazide with dimethylaluminum chloride afforded the dimeric hydrazide 31 [Eq. (13)], which has a six-membered

56

W. Uhl

heterocycle in a chair conformation [28]. The deprotonation occurred at those nitrogen atoms which were attached to trimethylsilyl groups, because of the better stabilization of the negative charge by hyperconjugation at this position. The N(H)CMe3 groups remain intact and coordinate to the aluminum atoms by a ‘‘dative’’ bond. The different bonding situation of the nitrogen atoms causes different Al-N distances of which the shorter ones are directed to the negatively charged nitrogen atoms (187.2 compared to 201.6 pm). Similarly, gallium hydrazides were obtained by the treatment of in situ generated lithium dimethylhydrazide or diphenylhydrazide with dimethyl- and diethylgallium chloride, respectively [29]. The dimethyl compound was also obtained in other routes and has been discussed before [Eqs. (5) and (6)]. It was employed as a single source precursor in a detailed study on the deposition of GaN on GaAs or sapphire substrates, but these films (generated at 580 C) were found to be polycrystalline possessing a poor surface morphology [14]. The diethylgallium derivative [Et2Ga-N(H)-N(C6H5)2]2 (32) forms a dimer in the solid state with a Ga2N2 heterocycle and two exocyclic NAN bonds [29]. Another approach to the synthesis of aluminum hydrazides by salt elimination is the treatment of the adduct 1 with n-butyllithium [Eq. (14)] [30]. The product (33, 69% yield) is a dimer in the solid state and possesses a six-membered Al2N4 heterocycle in a chair conformation with two endocyclic NAN bonds. There are some indications that probably owing to some steric stress in the molecules partial monomerization occurs in benzene solution. As described before [Eq. (13)], the deprotonation of the hydrazine derivative occurred at that nitrogen atom which is attached to the trimethylsilyl group. Each aluminum atom is coordinated by a negatively charged nitrogen atom and a NH2 group with different Al-N distances of 186.7 and 200.4 pm, respectively. Compound 33 sublimes at 120 to 140 C in vacuum without decomposition. However, a rearrangement by the shift of a hydrogen atom was observed, and a new product [34, Eq. (14)] was isolated in almost quantitative yield [30]. Compound 34 has a unique structure and contains a five-membered heterocycle with one endocyclic and one exocyclic NAN bond. An N-H group is in a bridging position between both aluminum atoms. As was clearly shown by NMR spectroscopy this structure is retained in solution. A similar compound (35, Scheme 1) was obtained directly at room temperature when the adduct (Me3C)2AlCl Æ NH2-N(H)C6H5 was treated with the alkyllithium derivative [30]. Quantum-chemical calculations with model compounds revealed that the five-membered heterocycle is the energetically most favorable one followed by the six-membered ring in a chair conformation. The least likely structure contains the four-membered ring with two exocyclic NAN bonds [30]. Interestingly, such a structure (36) was obtained when the adduct of tertbutylhydrazine (Me3C)2AlCl Æ NH2-N(H)CMe3 was employed [16]. Thus, three different structures were observed in the series of the di(tert-butyl)aluminum hydrazine derivatives depending on the kind of the substituents attached to nitrogen. The reason for that different behavior may be a finely balanced situation between electrostatic attraction (Ald+-Nd)) and repulsion (transannular Nd)


Nd)and Ald+


Ald+) on the one hand and steric stress on the other. The transannular repulsion favors the formation of the larger heterocycles with

57

Aluminum and Gallium Hydrazides

a larger distance between atoms of the same charge, while the attraction may favor the four-membered ring. The formation of 36 may be caused by steric reasons because in both other forms at least one tert-butyl group of the hydrazido ligand must approach the tert-butyl groups attached to aluminum. 36 adopts a trans configuration in the solid state, unknown resonances in the 1H NMR spectra may be assigned to the cis isomer. Probably owing to the unfavorable structure of 36, slow decomposition was observed in solution at room temperature [16]. H

H

H

N

Me3C

N SiMe3

Al

2 Me3C

+ 2 BuLi - 2 BuH, -2 LiCl

Cl 1

H

Me3Si N

Me3C Al Me3C

CMe3 Al

N H

H

N

CMe3

N

ð14Þ SiMe3

H 33

H

Me3Si N

Me3C Al Me3C

H

N

CMe3 Al CMe3

N H N Me3Si

H 34

5 Formation of Hydrazides by Hydroalumination This method has been employed for the synthesis of aluminum hydrazides in only few, but quite interesting cases. The reaction of azobenzene with an arylaluminum dihydride [aryl = tri(tert-butyl)phenyl] was reported to afford a

58

W. Uhl

Scheme 1.

five-membered Al2N3 heterocycle [37, 22% yield, Eq. (15)] with an endocyclic NAN bond and a bridging amido group [31]. Thus, compound 37 may be described as a 1,2-diphenylhydrazido dianion attached to two aluminum atoms in the 1,2-position. The inner ring is not planar but adopts an envelope conformation with one of the hydrazido nitrogen atoms above the plane of the remaining atoms and a folding angle of 26. 3 [H2Al-C6H2(CMe3)3]2 + 4 H5C6-N=N-C6H5

ð15Þ C6H5 N 2 (Me3C)3C6H2

Al

Al N

C6H2(CMe3)3

+ 6 H2

N

C6H5

C6H5 37

+ [(Me3C)3H2C6AlNC6H5]2

59

Aluminum and Gallium Hydrazides

Unique compounds were formed upon treatment of 1,1,4,4-tetramethyl-2, 3-diazabutadiene with the alane amine adduct AlH3 Æ NMe2Et and upon complete hydroalumination of its two C@N double bonds [Eq. (16)] [32]. A cage compound (38) was isolated in a low yield (14%) by cautious work-up of the reaction mixture at low temperature. Compound 38 has a complicated structure formed by three 1,2-di(isopropyl)hydrazido dianions and four aluminum atoms, two of which are attached to two hydrogen atoms, the other two are bonded to only one hydrogen atom. As a particular interesting structural motif 38 has an NAN bond side-on cordinated to an aluminum atom through its lone pairs of electrons. Similar structures have been discussed before with the compounds 24 and 25. Compound 38 is almost insoluble in pentane. It decomposes in benzene and other polar solvents so that its characterization in solution is quite poor. On sublimation at 30 to 70 C in vacuum a heterocubane-type Al4N4 molecule was formed by the complete cleavage of all NAN bonds [32]. Azopropane was detected as a byproduct of that decomposition reaction, which may give some insight into the mechanism of the thermolysis of hydrazides to finally deposit AlN. CH3 H3C C 4 [AlH3(NMe2Et)] + 3

N

N

- 4 NMe2Et C CH3

H3C

i Pr

H

ð16Þ

H H

H

N

Al

i Pr

Al H

i Pr

N Al

N N

N Al

i Pr

i Pr

N

i Pr

H 38

Another main product [39, Eq. (17)] was formed and isolated in 62% yield upon heating the crude product of the reaction of diazabutadiene with the alane adduct in a closed vessel so that gas evolution was prevented [32]. Although 39 is an isomer of 38 with three hydrazide dianions and four aluminum atoms the direct rearrangement was not observed by heating pure samples of 38. Possibly, an excess of the alane adduct or the presence of the free amine is required for the reaction to proceed. A crystal structure determination of 39 revealed a highly symmetric structure with a heterocycle of three di(isopropyl)hydrazide dianions bridged by three AlH2 groups. This

60

W. Uhl

heterocycle includes the fourth aluminum atom by the side-on coordination of all three NAN bonds. That inner atom has a coordination number of six in a strongly flattened octahedral coordination geometry. The constitution remains intact in solution and two resonances were found in the 27Al-NMR spectrum at )9 and +113 ppm in the characteristic ranges of aluminum atoms with coordination numbers of six and four, respectively. Me Me

C N

4 [AlH3(NMe2Et)] + 3

N

- 4 NMe2Et C

Me

Me

H

H

Me

H

Me Al

C Me

H C

N

Me H

ð17Þ

Me

N

Me C

N

C

N

H

Al

Me H

Me

Al

Al N

H Me

N

H C

C Me

H

Me

Me

H

H 39

6 Discussion Aluminum and gallium hydrazides were described in the literature for the first time about forty years ago, however, the products reported were insufficiently characterized, and no conclusive suggestions concerning the structures of these compounds were possible. That situation did not change until the middle of the last decade, when the first structurally authenticated hydrazides of aluminum and gallium were reported. The renewed interest in hydrazine chemistry was stimulated by the observation that hydrazine derivatives are useful starting materials for the deposition of semiconducting materials such as AlN or GaN. Hydrazide chemistry with the heavier elements of group 13 has developed now to a well established field with about 40 completely characterized compounds and with a growing understanding of reactivity and structure.

Aluminum and Gallium Hydrazides

61

6.1 Synthetic Methods

The synthesis of aluminum and gallium hydrazides succeeds by several methods, which were used as a rough ordering scheme in the first part of this article: – Reaction of alkylaluminum and -gallium compounds with hydrazines bearing at least one protic N-H function by the release of alkanes. – Similarly, the employment of hydrides of these elements by the release of elemental hydrogen. – Treatment of alkylelement halides with lithium hydrazides and salt elimination. – Hydroalumination of 2,3-diazabutadiene derivatives. All these methods have been employed in many reactions and work with great success yielding the hydrazides in often high yields. The first three procedures are standard methods, while the hydroalumination is restricted to a few examples with a narrow range of application up to now. All hydrazido derivatives described here are rather stable at room temperature, they can be handled under inert gas atmosphere even in gram amounts, and hazardous properties have not been reported. However, derivatives of free hydrazine N2H4 with small substituents attached to aluminum must be handled with the greatest care, because they are highly explosive. Only the introduction of bulky tert-butyl groups prevented spontaneous decomposition and yielded compounds (9 and 10, see above) which are not shock sensitive and decompose slowly only at elevated temperatures. 6.2 Structural Motifs

Monomeric dialkylaluminum hydrazides with terminal hydrazido groups (A, Scheme 2) were observed only in those cases in which donor ligands such as ether molecules (28 to 30) or hydride ions (26) saturate the aluminum atoms to yield coordination numbers of four. In the absence of Lewis bases dimers are formed of the monoaluminum and monogallium monohydrazido derivatives, a trimer has never been detected by crystal structure determinations. Three different types of structures were reported: – Four-membered heterocycles (B). They have the negatively charged nitrogen atoms of the hydrazido ligands in a bridging position between two metal atoms. Both NAN bonds are exocyclic and adopt a trans position in the solid state. In solution, an equilibrium between cis/trans isomers was detected by temperature dependent NMR spectroscopy in many cases. – Five-membered heterocycles (C). That structural motif was determined in few cases only (34 and 35). It has one endo- and one exocyclic NAN bond, and only one of the negatively charged nitrogen atoms is in a bridging position. The rings adopt a twist conformation. The structure seems to be preserved in solution, and no dynamic processes were determined by NMR spectroscopy.

62

W. Uhl

Scheme 2.

– Six-membered heterocycles (D). Both NAN bonds are included in the central ring of the molecules, which usually adopt a chair conformation. Each aluminum atom is coordinated by a negatively charged ‘‘amido’’ nitrogen atom and by an ‘‘amino’’ NR2 group of different hydrazido ligands. Structures C and D have never been observed for gallium hydrazides, which exclusively form four-membered heterocycles. In contrast, all kinds of structures were observed with aluminum. Quantum-chemical calculations showed that for aluminum the five-membered heterocycle is the most

Aluminum and Gallium Hydrazides

63

favorable one, while the four-membered heterocycle is most unfavorable. The finally observed structure is probably determined by a balanced situation of electrostatic attraction between atoms of opposite charge and a transannular repulsion between atoms of the same charge (Al or N, respectively) as well as steric requirements. Owing to the relatively high electronegativity of gallium, charge separation in gallium compounds may not be as important as in aluminum compounds so that repulsion caused by the very narrow inneratomic distances in the four-membered rings may play a minor role only. That structure seems to be realized in aluminum chemistry only in those cases in which steric interactions prevent the coordination of the two-fold alkyl- and aryl-substituted amino nitrogen atoms to the metal atom (13, 36). The formation of six-membered rings (9, 31, 33) may be kinetically favored, the rearrangement to give the five-membered heterocycle was achieved in one case [Eq. (14)] at elevated temperature. A monoalkylaluminum dihydrazide was detected only with compound 25 (E). It is a monomer in the solid state possessing a unique structure with two AlN2 triangles and a coordination number of five at the aluminum atom. Homoleptic trihydrazido compounds are unknown. However, compound 24 may be derived from a trihydrazido compound in which one hydrazido group is two-fold deprotonated and gives a dianionic N-NR2 ligand. In the same sense, compound 23 could be described as a derivative of a trihydrazidogallium compound. Beside these simple structures three-dimensional cages (8, 19, 20, 23, 38) or more two-dimensional polycycles (10, 21, 24, 39) were described, up to now a general scheme does not exist for their classification. They show, however, the ability of the hydrazido groups to form interesting structures by their varied coordination behavior. The N-N groups can adopt several types of coordination modes, those which were determined by crystal structure determinations are summarized in Scheme 3. Neutral hydrazine molecules adopt exclusively terminal positions with one of their nitrogen atoms coordinated to one aluminum or gallium atom (F). A bridging function has never been observed. Probably owing to the short distance between both nitrogen atoms the two-fold coordination by organoelement compounds is sterically unfavorable. Electrostatic repulsion may also contribute to the non-existence of compounds with bridging hydrazines. A greater structural variety (G to J) was observed with hydrazide monoanions which can adopt one of four different structures. Monomeric compounds with terminal hydrazido ligands require the coordinative saturation of the central atoms by a donor. The structural motif H is the most widespread one of the hydrazides at all. It occurs in the four- and five-membered heterocycles. Structure I is part of the five- and six-membered rings, while J has been observed in a single case only (dihydrazide 25). The higher the negative charge the higher is the number of structures, and six different types were observed for the hydrazinediide anions with up to four bridged metal atoms. The bridging of two metal atoms can occur by a geminal (K) or a 1,2-coordination mode (M). Similar are the structures L and N, in which a third metal atom is side-on coordinated to yield three-membered AlN2 heterocycles. The exclusively terminal coordination of three or four metal

Scheme 3.

64 W. Uhl

Aluminum and Gallium Hydrazides

65

atoms was observed in O and P. The last motif was detected in the cage compound 38 only, which has three different coordination modes of hydrazides realized in one molecule (N, O, P). Scheme 3 impressively shows the fascinating structural variety of hydrazido compounds. 6.3 Structural Parameters

The most common structural motif of aluminum and gallium hydrazides is the centrosymmetric four-membered heterocycle with two exocyclic NAN bonds. The N-N distances of all these compounds fall in a very narrow range of about 145.5 pm on average with the most significant deviation (143.7 pm) detected for the diphenylhydrazido gallium derivative 14. The E2N2 heterocycles are planar with almost ideally equidistant E-N separations. As expected, the AlAN bonds of these compounds (196 to 199 pm) are shorter than the GaAN bonds (199 to 207 pm). Generally, the EAN bond lengths depend on the coordination number and the charge of the coordinating nitrogen atoms. Those discussed before have a coordination number of four with a monoanionic nitrogen atom. Further examples of that type were observed, for instance, in the tricyclic compound 10 (Al-N = 195 to 197 pm), in the cage compounds 19 and 20 (Ga-N = 199 to 201 pm), or 23 (Ga-N = 199 to 203 pm). A coordination number of three at nitrogen reduces the E-N distances of hydrazides considerably to about 185 pm. Those structures were observed in compounds containing six- (31, 33) or five-membered heterocycles (34, 35), in adducts with donor atoms attached to aluminum (26, 28 to 30), or in some cages or polycyclic derivatives (19, 20, 23, 24) containing terminal hydrazido groups. While the Ga-N distances to the two-fold negatively charged bridging nitrogen atoms of the Me2N-N2) groups in 23 are quite similar to that standard value, the corresponding Al-N distances in 24 are elongated to more than 195 pm. Similarly long distances were detected for 9 bearing the unsubstituted hydrazido ligand, which, however, may be caused by some disorder of the molecules. Neutral hydrazines attached to aluminum atoms in some adducts (1 to 3) gave Al-N distances of 199 to 203 pm. Similar values were observed in those heterocycles in which a NR2 group is attached to aluminum (six- and five-membered heterocycles 31, 33 to 35). Both known five-membered rings with one endocyclic and one exocyclic NAN bond (34, 35) show strongly differing AlAN bond lengths to the bridging nitrogen atom of the exocyclic N-N group (193 and 201 pm). The shorter one was observed to that aluminum atom which is further coordinated by the NR2 group. The dihydrazido compound 25 has extremely different Al-N distances to the pentacoordinated aluminum atom of 181.1 (very short) and 207.2 pm (very long) for the tri- and tetracoordinated nitrogen atoms, respectively. The N-N distances depend on the charge of the hydrazido groups, the steric stress in the molecules, and to a smaller extent on the electronic properties of the substituents. Trimethylsilyl groups usually give longer N-N distances than alkyl or aryl groups. The N-N distances span a range of 140 to 159 pm. Neutral hydrazines in adducts (1 to 3) and monoanionic hydrazides have values below

66

W. Uhl

150 pm, while hydrazides bearing a negative charge at each nitrogen atom (8, 10, 27, 38, 39) have longer bonds up to 159 pm, which may be caused by an electrostatic repulsion. Lower values were observed in the gallium cages 19 and 20 (149 pm) and in compound 37 (144.3 pm). Also, the monoanionic hydrazido ligands of the monoalkyl dihydrazido compound 25 deviate from this general scheme with a long NAN bond of 152.1 pm. As was shown here, aluminum and gallium hydrazides have been developed to a nice class of well characterized compounds in the last decade. Beside their potential applicability in vapor deposition they have found great interest because of their fascinating structural variability. The capability of the hydrazido group to bridge up to four metal atoms makes it a very promising starting material for the synthesis of further derivatives, in particular of cages and polycyclic compounds with up to now unknown structures and properties.

7 References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32.

Fetter NR, Bartocha B (1961) Can J Chem 39: 2001 Paterson WG, Onyszchuk M (1961) Can J Chem 39: 2324 Fetter NR, Bartocha B, Brinckman FE Jr, Moore DW (1963) Can J Chem 41: 1359 Bains MS, Bradley DC (1962) Can J Chem 40: 1350 Bock H (1962) Z Naturforsch 17b: 429 Pearton SJ, Ren F (2000) Adv Mater 12: 1571 Gaskill DK, Bottka N, Lin MC (1986) J Crystal Growth 77: 418 Okumura H, Misawa S, Yoshida S (1991) Appl Phys Lett 59: 1058 Miyoshi S, Onabe K, Ohkouchi N, Yaguchi H, Ito R, Fukatsu S, Shiraki Y (1992) J Crystal Growth 124: 439 Lee RT, Stringfellow GB (1999) J Electron Mater 28: 963 Mizuta M, Fujieda S, Matsumoto Y, Kawamura T (1986) Jpn J Appl Phys 25: L945 Fujieda S, Mizuta M, Matsumoto Y (1987) Jpn J Appl Phys 26: 2067 Okumura H, Yoshida S, Misawa S, Sakuma E (1992) J Crystal Growth 120: 114 Lakhotia V, Neumayer DA, Cowley AH, Jones RA, Ekerdt JG (1995) Chem Mater 7: 546 Uhl W, Molter J, Saak W (1999) Z Anorg Allg Chem 625: 321 Uhl W, Molter J, Neumu¨ller B, Saak W (2000) Z Anorg Allg Chem 626: 2284 Peters DW, Bourret ED, Power MP, Arnold J (1999) J Organomet Chem 582: 108 Peters DW, Power MP, Bourret ED, Arnold J (1998) Chem Commun 753 Uhl W, Molter J, Koch R (1999) Eur J Inorg Chem 2021 Janik JF, Duesler EN, Paine RT (1993) Chem Ber 126: 2649 Uhl W, Molter J, Neumu¨ller B (2001) Inorg Chem 40: 2011 Kim Y, Kim JH, Park JE, Song H, Park JT (1997) J Organomet Chem 545–546: 99 Cho D, Park JE, Bae B-J, Lee K, Kim B, Park JT (1999) J Organomet Chem 592: 162 Gibson VC, Redshaw C, White AJP, Williams DJ (1999) Angew Chem Int Ed 38: 961 Luo B, Gladfelter WL (2000) Chem Commun 825 Silverman JS, Abernethy CD, Jones RA, Cowley AH (1999) Chem Commun 1645 Uhl W, Molter J, Neumu¨ller B (2000) Organometallics 19: 4422 No¨th H, Seifert T (1998) Eur J Inorg Chem 1931 Neumayer DA, Cowley AH, Decken A, Jones RA, Lakhotia V, Ekerdt JG (1995) Inorg Chem 34: 4698 Uhl W, Molter J, Koch R (2000) Eur J Inorg Chem 2255 Wehmschulte RJ, Power PP (1996) Inorg Chem 35: 2717 Uhl W, Molter J, Neumu¨ller B (2001) Chem Eur J 7: 1510

The Synthesis and Structural Properties of Aluminium Oxide, Hydroxide and Organooxide Compounds David J. Linton, Andrew E.H. Wheatley Department of Chemistry, University of Cambridge, Lensfield Road, Cambridge, CB2 1EW, UK e-mail: [email protected]

By virtue of the oxophilicity of aluminium, many compounds exist in which at least one of the metal’s formal valencies is occupied by oxygen or, alternatively, where the coordination state of the metal is raised to 4, 5 or 6 by the donation of electron density from oxygen. This review presents aluminium oxides first and thereafter aluminium hydroxides and organooxides. Although the discussion concentrates on the solid-state structural properties of such systems, solution structural, theoretical and reactivity studies are also presented. Keywords: Aluminium, Hydroxide, Oxide, Organooxide, Solid-state, Solution

1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

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2

Aluminium Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

69

2.1 2.2 2.2.1 2.2.2

Solution, Theoretical and Reactivity Studies . . . . Solid-State Structural Studies . . . . . . . . . . . . . . . Oxygen Mono-Bridged Complexes . . . . . . . . . . . Single and Edge-Fused Aluminium-Oxygen Rings

. . . .

69 70 70 73

3

Aluminium Hydroxides and Organooxides . . . . . . . . . . . . . . . .

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3.1 3.2 3.2.1 3.2.2 3.2.3 3.2.4 3.2.5

Solution, Theoretical and Reactivity Studies Solid-State Structural Studies . . . . . . . . . . . Monomeric Complexes . . . . . . . . . . . . . . . Mono-Bridged Complexes . . . . . . . . . . . . . Oxygen Mono-Bridged Complexes . . . . . . . Di-Bridged Complexes . . . . . . . . . . . . . . . . Oxygen Di-Bridged Complexes . . . . . . . . . .

4

Concluding Remarks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131

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References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132

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Structure and Bonding, Vol. 105
Springer-Verlag Berlin Heidelberg 2003

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D. J. Linton Æ A. E. H. Wheatley

List of Abbreviations acac binol Bz Cp CRAMPS DME DPT EXSY HETCOR HMPA MBMP Mes M. O. neol-H2 py pyO pz salen-H2 Me2-salen-H2 Et2-salen-H2 2,4-(t-Bu)2-salen-H2 THF TMP TriMEDA

acetyl acetonate 1,1¢-binaphthyl-2,2¢-dioxide benzyl cyclopentadienyl combined rotation and multiple pulse spectroscopy dimethoxyethane 1,3-diphenyltriazenide excitation spectroscopy heteronuclear correlation hexamethylphosphoramide 2,2¢-methylenebis(6-tert-butyl-4-methylphenoxide) mesityl molecular orbital 2,2-dimethylpropane-1,3-diol pyridine pyridine N-oxide pyrazolyl N,N¢-ethylenebis[salicylideneimine] N,N¢-ethylenebis[salicylidene(methyl)imine] N,N¢-ethylenebis[salicylidene(ethyl)imine] N,N¢-ethylenebis[2,4-bis(tert-butyl)salicylideneimine] tetrahydrofuran 2,2,6,6-tetramethylpiperidide N,N,N¢-trimethylethylenediamine

1 Introduction The manifold applications of aluminium-containing systems have led to extensive research in the field [1]. As a result, the chemistry of inorganic and heteropolymetallic systems incorporating aluminium is widely recognised and the subjects have recently formed the bases of several reviews [2, 3]. However, the synthetic utility of organoaluminium reagents – species which have been employed in regio- and stereoselective transformations both stoichiometrically and, in some instances, catalytically – has led to a great deal of interest in that field and to the preparation and use of a wide variety of such materials [4]. Of course, organometallic compounds of aluminium exhibit Lewis acidity by virtue of the vacant orbital on the metal centre. This means that they are ideal precursors to so-called ‘’ate complexes’, where the Lewis acid is coordinatively saturated by virtue of the predilection of aluminium for conjoining with the anionic components of early Main Group metal salts. This mode of behaviour has led to the development of a large field – the subject of a recent review [5]. The present article will, however, cover only homometallic systems. Nevertheless, the affinity of aluminium for electron density has a profound effect on

The Synthesis and Structural Properties of Aluminium Oxide

69

the structures and, consequently, on the reactivity of such species. The same property of aluminium which gives rise to ’ate complexes, combined with the thermodynamic favourability of aluminium-oxygen bond formation, means not only that alanes (AlR3, R ¼ organic group) will normally fail to persist in an oxygen [6] or moisture [7, 8] rich environment or in the presence of other oxygen-containing reagents but also that mixed aluminium-oxygen fragments demonstrate extensive aggregation behaviour. In the past, the field of aluminium organooxide chemistry has been the subject of intensive study [9] and this review aims both to bring the story upto-date and to combine it with the fields of aluminium hydroxide and oxide chemistry. However, more generally, inorganooxide species (those which do not contain CAO bonds – most commonly polymeric siloxide or phosphonate species [10]) have been excluded. The article covers the period 1990 to mid2001 and is divided into two main sections, the first of these dealing with oxide species, with the structurally similar fields of hydroxides and organooxides covered thereafter. Each section is sub-divided firstly into solution, theoretical and reactivity studies (including structure elucidation by solution NMR spectroscopy) and, secondly, into solid-state structural studies. The latter of these is further categorised firstly by aggregation state and thereafter according to the experimental technique employed and will incorporate investigations for the most part by single-crystal X-ray diffraction but, in some cases, by powder diffraction or solid-state NMR spectroscopy.

2 Aluminium Oxides 2.1 Solution, Theoretical and Reactivity Studies

This review will be dominated by advances in the structural chemistry of aluminium oxide, hydroxide and organooxide systems. While the reactions of such species will be considered, we will not be discussing the catalytic applications of Group 13 oxides in general and methylalumoxane (MAO) in particular. Following on from early work by Sinn and Kaminsky [11], ZieglerNatta polymerisation has recently been the subject of several review articles [12–17] and the reader is particularly directed to that by Chen and Marks for a detailed discussion of MAO [18]. Of course, studies into the synthesis and properties of MAO are very much ongoing [19] and attention is drawn to one article which presents a 31P NMR spectroscopic route for the determination of trialkylaluminium (AlR3) content in alumoxanes. This method utilises the treatment of the aluminium species with an excess of PPh3; rapid exchange of the phosphine and R3Al Æ PPh3 occurs, with the observable 31P signal representing a weighted average of the two [20]. As already pointed out, recent advances in the synthesis and structural characterisation of discrete molecular alumoxanes have lately been the subject of review [8]. While developments in the crystallisation and structural

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D. J. Linton Æ A. E. H. Wheatley

characterisation of such species have been dependent on the introduction of bulky Al-bonded alkyl or aryl groups (in place of methyl substituents), 1H NMR spectroscopy has proved to be a valuable tool in the development of our understanding of the solution structural properties of alkylalumoxanes. In this context and leading on from formative studies [21], the hydrolytic behaviour of AlMes3 has recently been probed. The addition of water at )60 C has been found to afford (H2O Æ AlMes3) Æ nTHF with mesityl elimination at )10 C yielding the dimer (THF)2 Æ [Mes2Al(l-OH)]2 [22]. Tris(aluminium) complexes incorporating 2-methyl- and 2-phenyl-substituted azaindoles have been reported (Section 2.2.1) to contain a mutually bridging, trigonal planar oxide with three out of four azaindole N-centres spanning pairs of metal atoms and the fourth such ligand bonding to one Alcentre via its pyrrole nitrogen [23]. However, variable-temperature 1H and EXSY NMR spectroscopy reveals that these complexes are fluxional in solution. Hence, whereas both (MeAl)3(l3-O)(2-R-7-C7H4N2)4 (R ¼ Me, Ph) exhibit asymmetric structures in the solid state such that the four R-groups are in distinct environments, 1H NMR spectroscopy suggests that these groups are identical in CDCl3 solution at room temperature. This implies rapid dynamic behaviour with static structures revealed only at )60 C along with the appearance of other isomers of these complexes. For R ¼ Ph, EXSY reveals the interconversion of major with minor isomers via the intramolecular migration of coordinated azaindole ligands [23]. The exothermic reaction of [(t-Bu)Al(l3-O)]6 (Section 2.2.1) [24] with AlMe3 in benzene at ambient temperature has recently been monitored by 1H, 13C and 1 13 H, C-HETCOR NMR spectroscopy and mass spectrometry [25, 26] and found to yield a mixture of two isomers of stoichiometry Al7(l3-O)6(t-Bu)6Me3. These species both appear to comprise a hexagonal Al6(l3-O)6(t-Bu)5Me core, one Al-O-edge of which has undergone the insertion of AlMe2(t-Bu) such that the proximity of the core-Me group and the cleaved Al-O interaction represents the only difference between the two structures. Spectroscopic studies reveal that methyl-exchange is possible between the deuterated zirconocene Cp2Zr(CD3)2 and the complexed AlMe2(t-Bu) moieties but not the core methyl groups. Furthermore, comparative investigations into the activities of [(t-Bu)Al(l3-O)]6 and Al7(l3-O)6(t-Bu)6Me3 for the co-catalytic polymerisation {with the fluorenone complex [Me2C(Cp)(C13H8)]ZrBz2 [27]} of 1,5-hexadiene reveal that the latter species is significantly more efficacious, with further improvements to be had by introducing trimethylaluminium to [(t-Bu)Al(l3-O)]6 in six-fold excess for a 1 : 1 hexamer:zirconium ratio [25]. 2.2 Solid-State Structural Studies

2.2.1 Oxygen Mono-Bridged Complexes The most simple species which contain single oxide bridges between aluminium centres are monomers of the type (R2Al)2O. However, the

The Synthesis and Structural Properties of Aluminium Oxide

71

recrystallisation of complexes containing such a motif is achieved only if the R-groups are sufficiently bulky [8]. Thus, an investigation into the structural properties of (R2Al)2X systems [R ¼ (Me3Si)2C(H), X ¼ chalcogen] has yielded complexes whose solid-state structures are significantly bent at the Group 16 centre for X ¼ S [28] and Te [29] but which for X ¼ O reveal a linear Al-O-Al ˚ ], moiety [30]. Concomitantly, the Al-O bond is extremely short [1.6877(4) A suggesting the presence of p-character in metal-chalcogen bonding [31]. However, the external stabilisation of such a motif has been found to alter the angle at oxygen (and the extent to which bonding interactions reveal p-character). Accordingly, the mono-solvate [py Æ (t-Bu)2Al]2(l-O) retains ˚ , Al-O-Al ¼ 180.0(1)] [24] while its linear Al-O-Al bridge [Al-O ¼ 1.710(1) A the N,N,N¢-trimethylethylenediamine stabilised bis(dialkylaluminium)oxide [TriMEDA Æ (t-Bu)2Al]2(l-O) is generated by the preference of [Me2N(CH2)2 NMe]Al(t-Bu)2 [32] for hydrolytic cleavage of an Al-N rather than an Al-C ˚ , Al-O-Al ¼ 173.0(4)] [33]. bond and is only slightly bent [Al-O ¼ 1.690(7) A Moreover, the treatment of {[(Me3Si)2C(H)]2Al}2(l-O) with trimethylamine oxide has been found to give both asymmetrically mono- and symmetrically bis-coordinated complexes [34]. In this context, {Me3NO Æ [(Me3Si)2C(H)]2Al} ˚ (to the {[(Me3Si)2C(H)]2Al}(l-O) affords Al-(l-O) bond lengths of 1.687(4) A coordinatively unsaturated metal ion, suggesting some p-interaction by virtue ˚ (to its saturated analogue) and an Al-O-Al of its shortness) and 1.753(3) A angle of 162.3(2) with the asymmetry inherent in this structure being preserved in hydrocarbon solution up to 90 C. The symmetrically bissolvated {Me3NO Æ [(Me3Si)2C(H)]2Al}2(l-O) reveals Al-(l-O) bond lengths of ˚ and 1.739(2) A ˚ (no p-interaction) and a 161.4(2) Al-O-Al angle. 1.732(2) A ˚ [34]. Dative Al-ONMe3 bonds are in the range 1.846(3)-1.865(2) A The employment of more complex aluminium-bonded organic residues has been shown to facilitate intramolecular support of the Group 13 metals forming the single Al2O bridge of the type discussed above. A wide variety of salen-type ligands exist. Their coordination chemistry with respect to Group 13 metal centres has been widely studied – as this review will show – and recently reviewed [35, 36]. The first use of 2,4-(t-Bu)2-salen-H2 in conjunction with Group 13 elements was achieved by heating [2,4-(t-Bu)2-salen]AlN(H)(t-Bu) to initially form [2,4-(t-Bu)2-salen]AlOH. This compound undergoes facile condensation to form {[2,4-(t-Bu)2-salen]Al}2O (Fig. 1) which, in turn, can be hydrolysed back to the hydroxide precursor. In {[2,4-(t-Bu)2-salen]Al}2O each of the two aluminium-bonded bis(phenoxide) groups stabilise the metal centre to which they are formally coordinated by using both imine N-centres to afford a contiguous array of 5-membered C2N2Al and 6-membered C3NOAl rings [37]. In this way, each metal centre is rendered five-coordinate – a bonding environment which appears to retain its integrity in solution [38, 39]. Further, the bond angle at the bridging oxygen centre is [at 159.5(5)] greater than that of 152.0(3) which has been reported in the analogous complex [(salen)Al]2O [40]. In a similar vein, tris-aluminium complexes have been reported in which related 2-methyl- and 2-phenyl-substituted azaindoles use their two N-centres to bridge two Group 13 metal centres. In either case the oxide itself adopts a trigonal planar geometry (angles sum to 360.0 and 359.7, respectively) with

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D. J. Linton Æ A. E. H. Wheatley

t-Bu

N N

t-Bu t-Bu

O

Al O t-Bu O

t-Bu t-Bu

t-Bu O N

Al

O

t-Bu

N Fig. 1. {[2,4-(t-Bu)2-salen]Al}2O [37]

two azaindole ligands both bridging one Al    Al divide, a third bridging another such gap and the fourth and final ligand adopting a monocoordinative mode to just one aluminium ion via its pyrrole N-centre (Fig. 2) [23]. In consequence, although NMR spectroscopy (Section 2.1) suggested that ligand migration was responsible for the rapid fluxionality of these species in solution, in the solid state they both reveal structures in which each Al-centre is in a distinct chemical environment. A further tris(aluminium) azaindole complex containing a mutually-bridging oxide has been noted. Hence, investigations have been undertaken into certain organoaluminium compounds which incorporate azaindole ligands and which exhibit blue fluorescence due to p ‹ p* transitions in these ligands. The 2 : 1 : 1 reaction of AlMe3 with 7-azaindole and bis(trifluoromethyl)isopropanol has led to the isolation and structural characterisation of (C7H5N2)4[Al(OR)]2Al(Me)(l3-O) [R ¼ C(H)(CF3)2]. This contains only terminal bis(trifluoromethyl)isopropoxide moieties, with the Al(l3-O) core being supported by the bridging action of the aromatic ligands. In this instance two azaindole moieties span the two MeAl    AlOR divides with the remaining two aromatics bridging Me Me N N

N

Me

Al N Me Al

N

O Al N

N

Me N

Me Me Fig. 2. A 2-methylazaindole complex of aluminium [23]

73

The Synthesis and Structural Properties of Aluminium Oxide

ROAl    AlOR [41, 42]. While the latter azaindole ligands are disordered, the former are not and reveal that the deprotonated N-centre interacts more ˚ versus 2.092(9) A ˚ ] with aluminium than does the neutral strongly [1.868(9) A (pyridyl) N-atom. Higher association state species have been noted to incorporate mono-oxide bridges between Group 13 metal centres. Indeed, the treatment of ArAlH2 [Ar ¼ 2,4,6-(t-Bu)3-C6H2] with freshly sublimed (Me2SiO)3 has yielded (ArAlO)4 wherein the association of four arylaluminium oxide fragments ˚ , mean affords a simple 8-membered (AlO)4 ring core (mean Al-O ¼ 1.689 A Al-O-Al ¼ 150.9) [43]. The crystallographic data for a tetra(aluminium) complex which includes an [Al(l-O2S)2Al(l-O)]2 core has been deposited on the Cambridge Crystallographic Database (CCDC 101483) along with details of several structures which are discussed later in this review [44]. The superstructure of {[(TMP)AlOS(@O)(CF3)O]2(l-O)}2 results from the interaction, via two oxide centres, of a pair of metallocycles which each incorporate two aluminium atoms and two sulfate-based SO2 moieties (Fig. 3). Lastly, the positions of hydrogen atoms have been located in a series of lowsurface-area aluminium oxide-hydroxides. Hence, the long known solid-state structures of diaspore (a-AlOOH) and boehmite (c-AlOOH) have been recently probed using 1H CRAMPS [45] along with X-ray and neutron diffraction techniques [46]. For diaspore results point to the presence of Al3(l3-OH) groups with 6-fold and 4-fold coordination at aluminium and oxygen, respectively, while Al2(l-OH) moieties with 6-coordinate aluminium and 3-coordinate oxygen are revealed in boehmite [46]. 2.2.2 Single and Edge-Fused Aluminium-Oxygen Rings A single example has been reported of a bis(aluminium) oxide-alkoxide based on an Al(l-O)2Al motif where one oxygen centre represents the oxide while the other one is part of the organooxide moiety [47]. [(i-Pr)OH]3 Æ ClAl(l-O) [l-O (i-Pr)]AlCl2 reveals a 4-membered metallocycle with the stericallyavailable bridging oxide centre also coordinating to a molecule of AlCl3. TMP

F3 C

O

O

Al O O

S

S

CF3 O

TMP

Al

O O

O O S F3C Al TMP

O

TMP

Fig. 3. {[(TMP)AlOS(@O)(CF3)O]2(l-O)}2 [44]

CF3

S Al

O

O

O

O

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D. J. Linton Æ A. E. H. Wheatley

Meanwhile, isolated (AlO)2 rings incorporating only oxide centres have been reported in the context of tetraaluminium systems such as [(t-Bu)2Al]2[l-OAl(t-Bu)2]2 [24] and {Me[(2-C5H4N)2N]Al}2(l-OAlMe2)2 [48]. More recently, bis(trifluoromethyl)isopropoxide ligands have been noted to bond terminally to aluminium in the complicated oxide-alkoxide tetraaluminium species which results from the 3 : 2 : 1 reaction of 7-azaindole with AlMe3 and hexafluoro-2-propanol [42]. This species reveals a cyclic (AlO)2 core incorporating the two oxide centres, with the remaining two metal ions each bonding to one oxide centre and acting as links to the bis(trifluoromethyl)isopropoxide groups. Extensive electronic support for this arrangement is afforded both by the neutral and deprotonated nitrogen centres of six azaindole groups which each interact with one O(l-Al)O metal centre and one AlO centre. The result is that the core (AlO)2 metallocycle resides at the centre of an array of six edge-fused AlOAlNCN rings. More frequently, edge-fused (AlO)n rings have been noted in both symmetric and asymmetric contexts for mixed aluminium oxide-hyroxides/ organooxides. Regarding the former (symmetric) type, the simplest such examples of these systems is the trimetallic, bicyclic pyridine solvate (py)3 Æ ClAl(l3-O)[(l-OEt)AlCl2]2 [49]. Higher-order (tetra-, penta- and hexametallic) motifs have also been reported. For example, t-BuAl(l3-O)(l-OH) [l-OAl(t-Bu)2][Al(t-Bu)2]2 contains three types of Group 13 centre. These bond to a mutually bridging oxide centre. The two heterocyclic arrays are closed by l-OH and l-OAl(t-Bu)2 fragments (Fig. 4) [50]. The more complex Al4O4 ladder core of dimeric [Et2Al(l3-O)(l-OAlEtAr)AlAr]2 [Ar ¼ 2,4,6(t-Bu)3C6H2] reveals two l3-oxide bridges which complete the four-membered heterocycle in the ladder core. In this instance both supporting heterocycles are closed by l-OAlEtAr – cf. {[2,4,6-(t-Bu)3-C6H2]AlO}4 (above) [43]. Bridging carboxylate groups are noted to exhibit a variety of bonding modes in the complicated system (AlO)2[O2CN(i-Pr)2]8 [51]. The core of this cluster incorporates two l3-oxide centres into an (AlO)2 metallocycle to which two further (outer) metal centres are bonded. Supporting this core is the organic periphery wherein three carboxylate-types are revealed: two mono-bridge core and outer aluminium ions while four double-bridge such centres and the final two chelate outer aluminium ions only [51] to afford a motif closely related to that noted elsewhere for an azaindole-containing oxide-organooxide [42]. A single penta(aluminium) cluster based on four edge-fused four-membered rings has been reported. In the oxide-hydroxide t-BuAl(l3-O)2(l3-OH)2 t-Bu t-Bu Al O

t-Bu t-Bu

Al

Al O

O Al t-Bu

t-Bu

Fig. 4. t-BuAl(l3-O)(l-OH)[l-OAl(t-Bu)2][Al(t-Bu)2]2 [50]

t-Bu

75

The Synthesis and Structural Properties of Aluminium Oxide

t-Bu HO Al

t-Bu

Al O HO O

OH Al

t-Bu

t-Bu Al O

O

Al

O

t-Bu

OH

O

Ph

Ph

Fig. 5. The oxide-hydroxide t-BuAl(l3-O)2(l3-OH)2[Al(t-Bu)]4(l-OH)2(l-O2CPh)2 [52]

[Al(t-Bu)]4(l-OH)2(l-O2CPh)2 a single tert-butylaluminium moiety participates in each Al(l-O)(l-OH)Al metallocycle while the remaining metal centres are supported by bridging hydroxy and benzoate groups (Fig. 5) [52]. Lastly, just as one example of an aluminium oxide-alkoxide complex has already been discussed [47] so the dimeric oxide-alkoxide {t-BuAl[l-OC(H)MeCH2CO2][(l3-O)Al(t-Bu)]2}2 reveals an unusually extended Al6O6 ladder in which each metal centre bears a tert-butyl group. The ladder ends incorporate the alkoxide centres of each of two – OC(H)MeCH2CO2– ligands, with the adoption by the ladder of a boat-shaped configuration resulting from the coordinative abilities of the two carboxylate moieties (Fig. 6) [53]. The cyclisation of a hexa(aluminium) oxide system has also been reported, with [(t-Bu)AlO]6 incorporating six edge-fused (AlO)2 metallocycles and with cyclisation concomitantly yielding two trinuclear (AlO)3 rings [24]. The asymmetric structures which incorporate Al-O edge-fused rings of different size all contain (AlO)n metallocycles in which n ¼ 2, 3 or, in an isolated instance, 8. The retention of at least one four-membered heterocycle (n ¼ 2) relates these systems to the symmetric polycycles discussed above in all but one aluminium oxide. Two (AlO)2 rings are incorporated in the closely related structures of {[l-O(H)Al(t-Bu)][l3-OAl(t-Bu)][l3-OAl(t-Bu)O2CCCl3]}2 (Fig. 7) [54] and {[l-OAl(t-Bu)][l3-OAl(t-Bu)][l3-OAl(t-Bu) Æ NH2(CH2)3Me]}2 [55], wherein two (AlO)2 rings are held together by two bridging Al(t-Bu)R or Al(t-Bu) Æ L Me Me

t-Bu

Al

O O

O O

O Al t-Bu Al

t-Bu

O

O Al

Al

O

O Al O

t-Bu

t-Bu t-Bu

Fig. 6. Dimeric oxide-alkoxide {t-BuAl[l-OC(H)MeCH2CO2][(l3-O)Al(t-Bu)]2}2 [53]

76

D. J. Linton Æ A. E. H. Wheatley

t-Bu Al O

O Cl3C

O

Al

t-Bu Al O Al

t-Bu HO

O

OH

t-Bu

Al

O

O

CCl3

O

Al

t-Bu t-Bu

Fig. 7. {[l-O(H)Al(t-Bu)][l3-OAl(t-Bu)][l3-OAl(t-Bu)O2CCCl3]}2 [54]

fragments [R ¼ OC(O)CCl3, L ¼ NH2(CH2)3Me] and two bridging oxide or hydroxide groups, respectively. Moreover, in the former structure both O-centres in each carboxylate moiety are involved in the formation of peripheral (AlO)3 rings. Three 4-membered rings are incorporated in the structures of [(t-Bu)AlO]3[(t-Bu)2AlOH](OH) and [(t-Bu)AlO]4[(t-Bu)2AlOH]2 [50]. [(t-Bu)AlO]3[(t-Bu)2AlOH](OH) is obtained by fractional recrystallisation of the result of the stoichiometric hydrolysis of Al(t-Bu)3 by hydrated aluminium sulfate; its structure reveals two Al-O edge-fused rings in a ladder-type core folded about the common bond, terminal atoms of which are bridged by a hydroxyl group and an [Al(t-Bu)2]O moiety to afford 4-membered metallocycles (Fig. 8). [(t-Bu)AlO]4[(t-Bu)2AlOH]2, obtained during the slow thermolysis of [(t-Bu)2Al(l-OH)]3 to [(t-Bu)Al(l3-O)]6 and thus representing a possible precursor to the hexameric alumoxane, demonstrates the (t-Bu)2AlOH-bridging of terminal centres in a [(t-Bu)AlO]4 open pseudocubic arrangement [50]. Complex arrays of five 4-membered (AlO)2 rings based on a central oxide dianion and bridging alkoxide ligands have been noted in both F3CCH2OH Æ (AlOCH2CF3)3[Al(OCH2CF3)2](l-OCH2CF3)4(l4-O) (Fig. 9) – wherein each metal atom is 5-coordinate and exhibits a distorted trigonal bipyramidal environment [56] – and i-PrOH Æ (AlCl)3[Al(i-Pr)Cl][l-O(i-Pr)]5(l4-O) [57, 58]. A mixture of 4- and 6-membered heterocycles form the superstructure of the octanuclear alkylaluminium oxide complex [(t-Bu)AlO]8 [50]. This species t-Bu

O t-Bu Al

Al

t-Bu

O

Al

O

Al

O O

Al

t-Bu Fig. 8. [(t-Bu)AlO]3[(t-Bu)2AlOH](OH) [50]

t-Bu

t-Bu

t-Bu

The Synthesis and Structural Properties of Aluminium Oxide

77

Fig. 9. (AlOR)3[Al(OR)2](l-OR)4(l4-O) (R ¼ CH2CF3) [56]

has been isolated as a minor by-product during the synthesis of Ga(t-Bu)3 using Al(t-Bu)3 and the cyclic trimer [(t-Bu)2Ga(l-OH)]3 [59]. The structure is best viewed as composed of a stack of two 6-membered, cyclic [(t-Bu)AlO]3 fragments (with alternating short and long bonds), two inter-trimer AlAO bonds having cleaved in order that two adjacent Al- and O-centres in each trimer can instead attach to a cyclic [(t-Bu)AlO]2 arrangement – a motif similar to one previously reported in imidoalane structural chemistry [60, 61]. A novel ammonia-solvated oxide-isopropoxide complex has been isolated in which seven contiguous (AlO)2 rings are seen [62]. The structure is based on a zig-zag (AlO)4 ladder core, with the Al-centres on either side of the ladder each bridged by two further oxides. Thereafter, each of the six 2-coordinate oxides in this arrangement are stablised by a peripheral Al[O(i-Pr)]3 group (Fig. 10). An unusual and complicated array of heterocycles results from the inclusion of hydroxide moieties into a similar tert-butylaluminium oxide system to [(t-Bu)AlO]8 [50]. In this context, the product which results from the hydrolysis of [(t-Bu)Al(l3-O)]6 [24] – [(t-Bu)Al]6(l3-O)4(l3-OH)4 – is best viewed as comprising an octahedron of aluminium centres each face of which is l3-capped by either an oxide or a hydroxide group (Fig. 11) [63]. This species was the first to exhibit penta-coordinate Al-centres in an alumoxane context. Moreover, the polyhedral architecture incorporated an interstitial void which, it was suggested, might facilitate the formation of inclusion complexes. The predilection for (AlO)n (n ¼ 2) metallocycles does not hold for the tetracyclic array of n ¼ 3 rings displayed by the mixed oxide-hydroxide

Fig. 10. A novel ammonia-solvated oxide-isopropoxide (R ¼ i-Pr, L ¼ NH3) [62]

78

D. J. Linton Æ A. E. H. Wheatley

[(Me3Si)3CAl]4(l-OH)4(l-O)2 [64]. This complex results from the reaction of THF Æ Me2AlC(SiMe3)3 with H2O (degassed in THF; 2 eq.) and reveals an inor˚ ) of a type noted both for ganic adamantanoid structure (mean Al-O ¼ 1.798 A the gallium analogue [64] and also for charged cadmium phosphides [65]. Lastly, early work on polyoxycation clusters of aluminium led to tentative claims of the synthesis of Al24O72 [66]. More recently, it has been established, by 27Al MAS NMR spectroscopy, that the sulfate salt of this species has a ˚ , agrees well with that (of 7.2 A ˚ ) calculated for a dimer of radius which, at 7 A Al12 clusters [67]. Thermal treatment of the complex reveals not only 4- and 6-coordinate aluminium centres, but also the formation of a metastable, 5-coordinate metal site.

3 Aluminium Hydroxides and Organooxides 3.1 Solution, Theoretical and Reactivity Studies

A group of simple chiral complexes of the type Me2AlOR {R ¼ (S)-(–)-1methyl-2-pyrrolidinemethanolate, (2S,3S)-(+)-2-amino-3-methyl-1-pentanolate, (1R,2S)-(–)-a-[1-(methylamino)ethyl]benzyl alkoxide, cinchoninate, quinidinate, quininate}, all of which contain potentially chelating N-centred fragments within the R-group, have been derivatised. Whereas the solid-state structure of dimethylaluminium (2S,3R)-(+)-4-(dimethylamino)-1,2-diphenyl3-methyl-2-butoxide (Section 3.2.1) establishes that the nitrogen centre does, indeed, render the metal centre 4-coordinate, dynamic dissociation of this interaction in the aforementioned complexes is detected by NMR spectroscopy with Arrhenius activation energies in the range 12.5(6)-25.0(1.2) kcal mol)1 [68]. The mixed-metal Li(l-O)Al-based complex C6H8O  (THF)2  LiAl[(R)-binol]2 has been a source of great interest recently by virtue of its ability to catalyse [69, 70] the conjugate addition of organolithium species to a,b-unsaturated t-Bu t-Bu Al O Al O t-Bu

O O

Al O Al O

Al

O

Al

t-Bu

O

t-Bu t-Bu Fig. 11. [(t-Bu)Al]6(l3-O)4(l3-OH)4 with disordered hydroxy H-centres omitted [63]

The Synthesis and Structural Properties of Aluminium Oxide

79

ketones [71, 72]. It is, of course, well known that organolithium compounds normally add 1,2 across such groups [73, 74]. However, more recent reports suggest that the presence of the sterically demanding Lewis acids aluminium tris(2,6-diphenylphenoxide) [75–80] or methylaluminium bis(2,6-di-tertbutyl-4-methylphenoxide) [81–83] – the simple, monomeric solid-state structure of which is reported below [84, 85] – also incurs 1,4-addition. Lately, in situ-generated monomeric lithium aluminate complexes have been implicated in this process [85], though straightforward complexes between unsaturated molecules (nitriles and ketones) have also been reported [86]. In the context of the latter, 1H NMR spectroscopic studies have allowed the calculation of equilibrium constants for the relative concentrations of free and coordinated aluminium compound over a limited temperature range in toluene and, subsequently, have made it possible to calculate DH and DS for dissociation of the coordinating Lewis base [86]. The treatment of ethylaluminium bis(2,6-di-tert-butyl-4-methylphenoxide) [¼ EtAl(OR)2] with benzophenone in a hydrocarbon environment led to the synthesis of Ph2C@O Æ Ph2C(H)OAl(OR)2 [87]. The Group 13 tris(2,6-diphenylphenoxide) has also been employed as a selective activator of saturated carbonyls towards a-alkylation by organolithium species [88–91]. Lastly, as part of an investigation into the radical anion complexes of tris(1,3-diphenyltriazenido) aluminium, both Al(OR)2(DPT) and Al(OR)(DPT)2 (R ¼ 2,6-di-tert-butyl-4methylphenoxide) have been found to reveal irreversible reduction properties [92]. 27 Al NMR spectroscopy has recently been employed to study the various aluminium hydroxides formed by the hydrolysis of AlCl3 in benzene solution [93]. Results point to formation of the rapidly exchanging ions AlCl4), AlCl3OH) and also the (AlO)2-based dimer [Cl2Al(l-OH)]2. Along with (t-Bu)2Al(2-OCH2-C5H4N) and (t-Bu)2Al(8-OCH2-C9H6N), (t-Bu)2Al(2-O-C5H4N), (t-Bu)2Al(8-O-C9H6N) and (i-Bu)2Al(8-O-C9H6N) have all been synthesised. The last three of these have been found to be a dimer, monomer and dimer, respectively, in the solid state (Sections 3.2.1 and 3.1.2) [94]. Moreover, for the last of these compounds it has been established that the solid-state motif (an O-bridged dimer with the isobutyl groups on each metal centre adopting an anti-conformation) is retained at ambient temperature in hydrocarbon solution. This is revealed by the 1H NMR spectroscopic observation of inequivalent i-Bu groups in spite of the presence of only one set of signals for the aromatic ligands. In a similar vein, previous reports of anisochronus metal-bonded methylene groups in the isostructural dimer of Et2Al(8-O-C9H6N) [95] are replicated here. Such results point to hindered rotation of the alkyl fragments about the metal-carbon bond. For (i-Bu)2Al(8-O-C9H6N), coalescence of the alkyl signals is noted at high temperature [94]. However, at 68.5(4) kJ mol)1, the value of DG (large for bond rotation) suggests the existence of agostic interactions in solution. Two studies into simple oxygen-bridged dimers have been published [96, 97]. The solid-state structures of a wide variety of straightforward complexes incorporating 4- and 5-coordinate metal centres are mentioned in Section 3.2.5 and Table 4. Moreover, the dissociative behaviour of an

80

D. J. Linton Æ A. E. H. Wheatley

internally coordinating ester group has been monitored in solution, with variable-temperature 1H NMR spectroscopy revealing that {(i-Bu)2Al[l-OC(H)MeC(OEt)¼O]}2 exhibits rapid intra-dimer ligand exchange above 92 C [96]. More recently, it has been reported that the solid-state stabilisation (or lack thereof) of the metal in {R2Al[l-O(CH2)nX]}2 (R ¼ H, Me, i-Bu, t-Bu; n ¼ 2, 3; X ¼ OMe, SMe, NH2, NMe2) is generally dependent on the size of R, while in solution it has been established (by 13C NMR spectroscopy) that dimers such as these exhibit equilibria between species in which the aluminium centres are 4- and 5-coordinate. This equilibrium is affected by a variety of factors: the steric bulk of the metal-bonded R-group, the choice of bridging heteroatom, basicity of the neutral donor group and chelate ring size. Hence, variabletemperature NMR spectroscopy reveals that the bond dissociation energies for interactions between metal and neutral Lewis base centres in these systems are very small (2.3–13.2 kJ mol)1) and that thioethers show a greater ability to stabilise the metal centres than do ethers, with amines being the worst such coordinating agents. As might be expected, ab initio studies point to steric influences as the major contributor to these observations [97]. In a similar vein, 1H, 13C, 17O and 27Al NMR spectroscopies have been utilised to demonstrate the existence of uniquely dimeric species in toluene for a variety of complexes of the type [Me2Al(l-OR)]n [R ¼ i-Pr, i-Bu, s-Bu, t-Bu, CH2(t-Bu); n ¼ 2] with complexes incorporating R ¼ Me, Et, n-Pr, n-Bu, (CH2)2(i-Pr), n-C5H11, n-C6H13, n-C8H17, n-C10H21, n-C12H25 instead favouring a dimer-trimer equilibrium (n ¼ 2, 3) [98]. Moreover based on 1H NMR spectra, equilibrium constants for the {Me2Al[l-O(n-Pr)]}3 to {Me2Al [l-O(n-Pr)]}2 conversion have been calculated over the temperature range 25–140 C with DH and DS subsequently determined to be 95(2) and 260(10) kJ mol)1, respectively (that is 47.5 kJ mol)1 for the conversion of two trimers into three dimers) while DH and DS [55(3) and )120(20) kJ mol)1, respectively] have been computed [98]. It has been reported that the 1 : 2 treatment of the diphenol 2,2¢-(ethyne-1, 2-diyl)bis[6-(1,1-dimethylethyl)-4-methylphenol] with Al(i-Bu)3 gives a species in which two OAl(i-Bu)2 groups associate intramolecularly to yield an (AlO)2 ring (Section 3.2.5 and Table 4) [99]. This species has been probed in solution with a view to discovering whether it is capable of reverting to a potentially bidentate Lewis acid. 13C NMR spectroscopy reveals two Me signals (at room temperature) which coalesce at 55 C – pointing to a DG value of 17.0(2) kcal mol)1 for i-Bu group exchange. However, although they are consistent with formation of the bidentate Lewis acid in solution, observations have been taken to suggest the cleavage of one Al-O bond followed by rapid rotation of the uncoordinated Al(i-Bu)2 group about its remaining metaloxygen bond [100]. In a similar vein, 2,2¢-(1,3-butadiyne-1,4-diyl)bis[6-(1,1dimethylethyl)-4-methylphenol] reacts with two eqivalents of Al(i-Bu)3 to yield the corresponding bis(diisobutylaluminium) aryloxide. This has been crystallographically characterised as an oligomeric DME-solvate in which each metal ion interacts with a single donor centre in the supplied Lewis base (Section 3.2.1). Dissolution in CD2Cl2 allowed the investigation of this complex in solution, with low-temperature ()100 C) 1H NMR spectroscopy suggesting

The Synthesis and Structural Properties of Aluminium Oxide

81

that the predominant species is that in which each O-centre of a single molecule of DME is bonded to a different metal ion within the monomer, such that in solution the bis(diisobutylaluminium) phenoxide does act as a bidentate Lewis acid [100]. The tetraanionic [H2O Æ Al3(l-OH)(l-OR)3]4) [R ¼ C(CH2CO2)2CO2] moiety has been achieved by the combination of (H2O)9 Æ Al(NO3)3 with citric acid [101], with the synthesis establishing the existence of polynuclear complexes of a type previously postulated [102, 103]. 27Al NMR spectroscopic studies have established that the tris(aluminium) cluster (Section 3.2.5) is not fluxional in D2O but instead retains three distinct metal environments over a wide pH range (3.0)9.0) with acid hydrolysis occurring below pH 3.0 and the resultant spectrum revealing the presence of [(H2O)m Æ (RO)Al3]+ and/or [(H2O)n Æ (ROH) Æ Al3]+ and also [(H2O)6 Æ Al]3+ [101]. The usefulness of aluminium tris(isopropoxide) in the Meerwein-PonndorfVerley reduction of cyclic ketones is well established [104]. However, it has more recently been noted that grafting aluminium alkoxide moieties onto siliceous mesoporous MCM-41 by one-step impregnation in hexane [105] enhances their catalytic activity with respect to such processes while facilitating catalyst separation [106]. Homogeneous catalytic processes have also been the subject of study. The neutral aluminium chloride, ClAl [(CH2)2NSiMe3]2NSiMe3, and the chiral, tetracoordinated aluminium cations {ClAl[(CH2)2NR]2NR¢}+ (R ¼ R¢ ¼ SiMe3; R ¼ SiMe3, R¢ ¼ Me; R ¼ i-Pr, R¢ ¼ Me; the counter-ion for each being AlCl4)) have recently been investigated in the context of ring-opening propylene oxide oligomerisation – a reaction known to proceed via an aluminium alkoxide propagating species generated by the nucleophilic attack of displaced Cl) on an aluminium bonded propylene oxide molecule [107]. Results have indicated the formation of low molecular weight polymers (Mn ¼ 318–2158) with narrow molecular weight distributions (Mw/Mn ¼ 1.14–1.25). The fully silylated cation revealed the highest catalytic activity [35% conversion over 2 days at room temperature (Mn ¼ 1418, Mw/Mn ¼ 1.18) which increased to 58% at 80 C (Mn ¼ 2158, Mw/ Mn ¼ 1.24)]. 13C NMR spectroscopy pointed to the products being head-to-tail polymers rich in meso diad and isotactic triad sequences [108]. Remaining with polymerisation catalysts, the Schiff bases 2-[C(H)@NR]-3,5-(t-Bu)2HOC6H2 (R ¼ CH2CH2NMe2, 2-PhO-C6H4, 2-CH2-C5H4N, 8-C9H6N) react with AlMe3 to give complexes – the first two of which have been structurally characterised (see below). These then react with B(C6F5)3 in dichloromethane or toluene to yield the corresponding cationic {2-[C(H)@NR]-3,5-(t-Bu)2Me2AlOC6H2}+ derivatives [109]. For R ¼ CH2CH2NMe2 and 2-PhO-C6H4, catalytic activity for ethylene polymerisation is noted. In the former case a productivity of 50 g mol)1 h)1 bar)1 (0.25 mmol catalyst in 200 ml toluene at 5 bar ethylene pressure and 50 C for 1 hour) was recorded, resulting in polyethylene of Mw ¼ 172,000 and Mn ¼ 2400. For R ¼ 2-PhO-C6H4, productivity improved to 110 g mol)1 h)1 bar)1, Mw ¼ 218,000 and Mn ¼ 5200 [109]. The catalytic activity of the isopropoxy-bridged dimer {Al(MBMP)[l-O(iPr)]}2 has been probed [110]. Hence, in toluene at 50 C it induces the polymerisation of e-caprolactone to yield a high molecular weight product

82

D. J. Linton Æ A. E. H. Wheatley

with a narrow molecular weight distribution (Mw/Mn < 1.5). Lastly, the homopolymerisation of cyclohexene oxide and its copolymerisation with carbon dioxide have been effected in the presence of catalytic ClAl(OR)OR¢, ClAl(OR)2, ClAl(OR¢)2 [R ¼ c-C6H11, R¢ ¼ C(O)C(H)¼C(H)C(O)OC8H17] or a 1 : 1 mixture of the two symmetrical species [111] with 27Al NMR spectroscopy pointing to an equilibrium between 5- and 6-coordinate metal centres which shifts towards the lower coordination state during initiation of the polymerisation process. 3.2 Solid-State Structural Studies

3.2.1 Monomeric Complexes The ability of aluminium to adopt a wide variety of coordination numbers has led to the isolation of a large number of monomeric Al(III) hydroxide and organooxide complexes, many of which incorporate polyfunctional ligands, with the Group 13 metal centre in the coordination states 3–6. Several straightforward, monomeric aluminium alkoxides [44, 68, 112–115] have been noted. More complicated is the mono(aluminium) complex Al(C7H5N2)(C7H6N2)[OCH(CF3)2]2 [116]. Its solid-state structure reveals a distorted tetrahedral metal centre [bond angles in the range 105.0(2)115.5(2)] with an Al-O bond length comparable to that already noted for polynuclear aluminium complexes incorporating azaindole ligands [42]. 1H NMR spectroscopy reveals that at room temperature the monometallic species Al(C7H5N2)(C7H6N2)[OCH(CF3)2]2 exists as two structural isomers while at the same temperature 27Al NMR spectroscopy points to an essentially tetrahedral Al(III) centre in solution. This view is corroborated by agreement between 27Al solution and MAS NMR spectroscopy, which strongly suggests retention of the solid-state structure in solution. The observation of a small nuclear quadrupole coupling constant (1.30 MHz with an asymmetry factor of 1.0) also indicates a limited change in electric field at the metal nucleus. A bis(alkoxide) complex is afforded by the reaction of 1,3,5-triamino-1,3, 5-trideoxy-cis-inositol {[HOC(H)¼C(H)NH2]3} [117] with AlCl3 [118]. The ion-separated tribromide, {Al[OC(H)¼C(H)NH3]3}3+ 3Br) [Al-O ¼ 1.894(2)– ˚ ], contains a cationic component which incorporates two neutral, 1.912(2) A tripodal O-centred ligands and so shows a hexa(oxygen) coordination sphere for the metal centre. In a similar vein, mixed O,O¢,N-tridentate behaviour has been noted for the monoanionic 1 : 2 polycarboxylate complex {Al[(O2CCH2)2NMe]2}) (for which the separate counter-ion is NH4+) [119]. A large number of aryloxide complexes which incorporate lipophilic substituents on a benzenoid moiety have been characterised; those which contain singly deprotonated aryloxides are summarised in Table 1. At the simplest level they contain one or two such ligands in conjunction with bulky but otherwise unfunctionalised organoaluminium fragments [44] or else they are homoleptic bis- or tris(aryloxide) systems [86, 120]. However, the

O

O

N

t-Bu

t-Bu

Ph

Ph

t-Bu

O

N

O

i-Pr

t-Bu

O

i-Pr

R1

t-Bu

t-Bu

Me

t-Bu

t-Bu

¼ R1

¼ R1

Et

Me

TMP

R2

Me

¼ R1

¼ R2

Cl

¼ R2

R3

R2, R3 = other ligands, C.N. ¼ coordination number of the metal

–

O@C(NMe2)H

–

–

–

L

3

4

4

4

3

C.N.

[84, 85]

[76]

[144]

[144]

[44]

Ref.

Table 1. Monomeric AlR1R2R3 complexes incorporating simple, monodeprotonated aryloxide ligands (R1) and, where appropriate, Lewis base (L).

The Synthesis and Structural Properties of Aluminium Oxide

83

O

O

O

O

O

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

R1

Me

Me

Me

Me

Table 1. (Contd.)

¼ R1

¼ R1

H

¼ R1

¼ R1

R2

H

¼ R1

¼ R2

H

¼ R1

R3

OEt2

–

NMe3

H2N(t-Bu)

t-Bu

L

O

4

3

4

4

4

C.N.

[86, 133]

[86, 120]

[86]

[86]

[86]

Ref.

84 D. J. Linton Æ A. E. H. Wheatley

N

N

O

O

i-Pr

i-Pr

i-Pr

i-Pr

Me2N

O

N

O

t-Bu

N

O

t-Bu

t-Bu

t-Bu

¼ R1

Me

Me

Me

t-Bu

¼ R1

H

¼ R2

¼ R2

¼ R2

)

O(CH2)4N+Me2Et

THF

–

–

–

5

5

5

5

4

[112]

[112]

[109]

[109]

[94]

The Synthesis and Structural Properties of Aluminium Oxide

85

O

O

O

O

Me

Me

t-Bu

t-Bu

t-Bu

t-Bu

i-Pr

i-Pr

R1

Me

Me

t-Bu

Table 1. (Contd.)

¼ R1

¼ R1

Cl

¼ R1

R2

Me

Me

¼ R2

Cp

R3

N

py

OEt2

–

L

Me

Me

4

4

4

3

C.N.

[126]

[126]

[125]

[121]

Ref.

86 D. J. Linton Æ A. E. H. Wheatley

O

O

O

O

O

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

i-Pr

i-Pr

Me

Me

Me

Me

Cl

¼ R1

¼ R1

Me

¼ R1

¼ R2

Me

Me

¼ R2

¼ R1

OEt2

O@C(Ph)OMe

O@C(t-Bu)H

O@CPh2

py

4

4

4

4

4

[129]

[128]

[128]

[128]

[126]

The Synthesis and Structural Properties of Aluminium Oxide

87

O

O

O

O

O

t-Bu

t-Bu

t-Bu

t-Bu

F

F

t-Bu

t-Bu

t-Bu

t-Bu

R1

F

F

Me

Me

F

Me

Me

Table 1. (Contd.)

Et

Me

Me

Me

Cl

R2

¼ R2

¼ R2

¼ R2

¼ R2

¼ R2

R3

Me

Me

H2N(t-Bu)

Me

N

Me

N

H 2C

O

Cl)

L

4

4

4

4

4

C.N.

[131]

[131]

[131]

[130]

[129]

Ref.

88 D. J. Linton Æ A. E. H. Wheatley

O

O

O

O

O

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

Me

Me

¼ R1

Me

Me

¼ R1

Et

i-Pr

Cl

¼ R2

Me

¼ R2

N+

N+

–

H2N(t-Bu)

H2N(t-Bu)

-O

-O

3

4

4

4

4

[134]

[132]

[132]

[131]

[131]

The Synthesis and Structural Properties of Aluminium Oxide

89

Ph

Ph

t-Bu

O

Me

O

Me

O

Me

O

O

t-Bu

R1

O

N

p-C6H4Me

O

N

n-Bu

O

N

Et

Me

Table 1. (Contd.)

¼ R1

¼ R1

¼ R1

¼ R1

¼ R1

R2

¼ R1

¼ R1

¼ R1

¼ R1

Me

R3

–

–

–

O Me

N

N

O@CPh2

L

Me O

6

6

6

4

4

C.N.

[140]

[139]

[138]

[137]

[135]

Ref.

90 D. J. Linton Æ A. E. H. Wheatley

O

O

O

O

O

O

O

OMe

Me

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

Me

Me

Me

t-Bu

Me

¼ R1

Me

¼ R1

O

O

t-Bu

O t-Bu Me

O

Me

Me

¼ R2

Me

Me

Et

Me

–

O

Me

4-Me-py

Me

–

–

O

5

3

4

4

4

[145]

[143]

[142]

[142]

[141]

The Synthesis and Structural Properties of Aluminium Oxide

91

O

O

O

N

N

O

O

Me

HO

N

O

N

O

N

O

R1

Br

Table 1. (Contd.)

¼ R1

¼ R1

¼ R1

¼ R1

¼ R1

R2

¼ R1

¼ R1

¼ R1

¼ R1

¼ R1

R3

–

–

–

–

–

L

6

6

6

6

6

C.N.

[149, 150, 151]

[148]

[147]

[146]

[146]

Ref.

92 D. J. Linton Æ A. E. H. Wheatley

N

t-Bu

O

O

t-Bu

Me2N

O

Me2N

O

H2N

t-Bu

t-Bu

Me

N

t-Bu

O

O

t-Bu

¼ R1

¼ R1

t-Bu

t-Bu

–

¼ R1

–

–

–

HMPA

6

5

6

[155]

[153]

[152]

The Synthesis and Structural Properties of Aluminium Oxide

93

94

D. J. Linton Æ A. E. H. Wheatley

propensity of aluminium for accepting electron density has meant that the single known example of a bis(aryloxide) complex in which the Group 13 metal centre is Cp-coordinated [121] reveals a solid-state structure in which ˚ ]. the metal centre is g5-bonded to the Cp ligand [Al-ring centroid ¼ 1.920(9) A This bonding mode, one of several which aluminium can adopt towards a cyclopentadienyl ring [122] had, at the time, been noted in only one other aluminium-containing complex [123]. Furthermore, the overall structure is similar to that reported for methylaluminium bis(2,6-di-tert-butyl-4-methylphenoxide) [84, 85] notwithstanding a lengthening of the Al-O bonds [from a ˚ to one of 1.736(2) A ˚ ] and a narrowing of the O-Al-O bond mean of 1.686 A angle [from 111.9(1) to 101.6(1)] in the Cp complex [123]. The Lewis acidity of aluminium is satisfied by a polydentate triazenide ligand in (PhN)2N Æ Al[2,6-(t-Bu)2-4-Me-OC6H2]2 [124]. Indeed, the coordination number of aluminium is not infrequently raised by virtue of the introduction of a neutral Lewis base (Section 3.1). Hence, single aryloxide ligands are noted in Et2O Æ Cl2Al[2,4,6-(t-Bu)3-OC6H2] [125] and in L Æ RR¢Al[2,6-(t-Bu)2-4-Me-OC6H2] {L ¼ py, R ¼ R¢ ¼ Me [126]; L ¼ NH3, R ¼ R¢ ¼ Me [127]; L ¼ Ph2CO, R ¼ R¢ ¼ Me [128]; L ¼ Et2O, R ¼ R¢ ¼ Cl [129]; L ¼ Me(CH2¼)C(MeO)CO, R ¼ R¢ ¼ Me [130]; L ¼ 2,6-(Me)2-C5H3N, R ¼ R¢ ¼ Me [131]; L ¼ pyO, t-BuNH2, R ¼ R¢ ¼ Me, Et [131, 132]; L ¼ (t-Bu)NH2, R ¼ Me, R¢ ¼ Cl [132]} while a perfluorinated phenoxide ligand is incorporated into the monomer HC(CH2CH2)3N Æ Me2AlOC6F5 [131]. In one case a remarkable co-crystalline system has been found to contain both a trimethylamine-complexed dihydridoaluminium aryloxide monomer and a monohydridoaluminium aryloxide dimer (Section 3.2.4) in the same crystal lattice [86]. Just as two 2,6-dialkylphenoxide moieties exist in the ether- and amine-complexed aluminium hydrides L Æ HAl[2,6-(t-Bu)2-OC6H3]2 {L ¼ Et2O [86,133], (t-Bu)NH2 [86]} and THF Æ HAl[2,6-(i-Pr)2-OC6H3]2 [112] and also in uncomplexed (i-Pr)Al[2,6-(t-Bu)2-OC6H3]2 [134], so 2,4,6-tri(alkyl)phenoxides have been noted in both unsolvated [84, 85] and solvated {by Ph2CO [135], H(t-Bu)CO [128], MeO(Ph)CO [128], (t-Bu)(H)C(CH2)2C¼O [86], 1,7,7trimethylbicyclo[2.2.1]heptan-2-one-3-onate [136], pyO [131], 3,5-(Me)2C5H3N [126], Cl) [for which 2,6-(Me)2-C5H3NH+ is the counter-ion] [129]} methylaluminium systems. Lastly, the acceptor behaviour of aluminium tris(aryloxide) complexes has been noted in the context of synthetic chemistry [75–79]. Accordingly, tris(aryloxide) containing tetra-coordinate Group 13 metal salts have been seen in the presence of N- [112, 126] and O- [76] donors, with 1,4-dimethylpiperazine-2,5-dione acting as a bridge between two tris(aryloxide)aluminium monomers in one instance [137]. Several examples exist in which aryloxide ligands are incorporated in conjunction with other functionalised groups, and these compounds reveal more complicated structural motifs. Use of the 1-(alkyl)-2-methyl-3-hydroxy4-pyridinate ligand (alkyl ¼ ethyl [138], n-butyl [139]) has yielded octahedral aluminium (and, for n-butyl, gallium) complexes on treatment with the relevant MMe3 (M ¼ Al, Ga) reagent. These species show O,O¢-chelation of the metal and are analogous to the compound which results from the employment of 1-(p-tolyl)-2-methyl-3-hydroxy-4-pyridinate [140]. Internal coordination by

The Synthesis and Structural Properties of Aluminium Oxide

95

the carbonyl side-arms of the O-C(t-Bu)MeCH2C(¼O)(t-Bu) and 2-[MeC(O)]4-Me-6-(t-Bu)-OC6H2 (¼ OR) ligands in MeAlOR(OR¢) [OR¢ ¼ 2,6-(t-Bu)2-4Me-OC6H2] causes the metal centre to be tetra-coordinate [141–143] and, in a similar vein, the 2-[(t-Bu)N@C(H)]-4,6-(t-Bu)2-OC6H2 ligand has recently been incorporated into 1 : 1 complexes with AlMeCl+, AlMe2+ and MEt2+ (M ¼ Al, Ga, In) moieties with crystal structures having been reported for AlMeCl+ and AlEt2+ fragments [144]. A 5-coordinate metal centre is noted in MeAl(OR)2 {OR ¼ 2-[C(O)OMe]-OC6H4} [145], while 2-oxazole and 3-oxide groups have rendered the metal centres 6-coordinate in aluminium tris(phenoxide) complexes [146–147]. Remaining with tris(aryloxides), several reports exist which incorporate pyridyloxide components. Whilst the use of 2,3dihydroxypyridine can afford a complex in which the aluminium centre is coordinated only to oxygen [148] it is more common for the pyridyl nitrogen, as it does in the mono(aryloxide) [94], to support the aluminium centre – as is reported for Al(O-7-C9H6N)3 [149–151]. In the presence of strongly coordinating Lewis bases, only two substituted phenoxide ligands bond to aluminium. The reaction of AlCl3 with 2-aminophenol in the presence of HMPA yields an ion-separated chloride for which the cationic component is the extensively hydrogen-bonded Al-containing species {(HMPA)2 Æ Al [2-(H2N)-OC6H4]2}+ [152] and in which the metal centre is rendered 6-coordinate (distorted octahedral) by virtue of trans coordination by the two HMPA molecules. Aluminium is noted to adopt a slightly distorted trigonal bipyramidal bonding environment in the presence of three equivalents of the potenially tridentate ligand 2,6-di-(N,N-dimethylaminomethyl)-4methylphenoxide. It is, in fact, found that variable coordinative behaviour can be displayed by this organic residue with both mono- and bidentate modes being adopted [153]. Examples have also been reported of complexes in which a phenoxide side-arm is, itself, polyfunctional. The potentially O,N,N¢-tridentate Schiff base 3,5-di-(tert-butyl)-(2-pyridylmethyl)salicylidenamine and its dimethylaminoethyl congener react with AlMe3 to yield complexes in which the metal centres describe distorted trigonal pyramidal geometries with relatively short AlN(imine) interactions [109]. As noted above, these species react with B(C6F5)3 to give the corresponding cationic [3,5-(t-Bu)2-OC6H2-C(H)¼NRAlMe2]+ (R ¼ 2-CH2-C5H4N, CH2CH2NMe2) derivatives – the first of which has been shown to act as a catalyst for ethylene polymerisation. In a similar vein, potentially tridentate polyfunctional side-arms on phenoxide ligands have also enabled the formation of 2 : 1 complexes containing hexa-coordinate aluminium centres. Thus, {[3,5-(t-Bu)2-OC6H2-C(H)¼NR]2Al}+ Cl) (R ¼ 2-H2N-C6H4) incorporates a distorted octahedral metal centre surrounded by meridionally-organised ligands with the neutral amine nitrogen atoms bonding ˚ ) to aluminium than do the imine N-centres less strongly (mean Al-N ¼ 2.066 A ˚ ). The amine functions are also involved in extensive (mean Al-N ¼ 1.967 A hydrogen bonding with the chloride counter-ions such that dimerisation results [154]. In some cases, however, such 2 : 1 systems have also revealed a tendency for polydeprotonation as well and, as such, have afforded products in which a single ligand type exhibits several oxidation states [155].

96

D. J. Linton Æ A. E. H. Wheatley

Just as 3,5-di-(tert-butyl)-(2-pyridylmethyl)salicylidenamine and its dimethylaminoethyl analogue have been noted to react with AlMe3 to give monomeric distorted trigonal pyramidal species [109], so many examples have been reported of aluminium complexes which incorporate salen-type ligands (Tables 2 and 3). These compounds fall broadly into two categories depending on whether or not the Al-bonded moiety carries a formal positive charge. Of the recorded examples of mononuclear, neutral salen-type complexes of aluminium, three reveal aluminium-alkyl interactions (alkyl ¼ Me [124, 156, 157], i-Pr [157]), while the remainder show Al-Cl [158], Al-OSiPh3 [157] and, in the case of mixed organooxide complexes, Al-OMe [159], Al-OEt [159] or Al-(2,4,6-Me3-OC6H2) [40] bonds (Table 2). In each of these examples, the Group 13 metal centre resides in a 5-coordinate bonding environment. A series of dinuclear compounds incorporating various salen-type ligands have been noted lately in which the backbones of the salen-type ligands have been modified such that the product complexes contain two metal centres at various (long) displacements with respect to one another [160]. Hence, while the ligand N,N¢-bis[2-hydroxy-3,5-di-(tert-butyl)-benzyl]-1,2-diaminobenzene has been noted to N,O-chelate two AlRR¢+ fragments (R ¼ R¢ ¼ Me; R ¼ Me, R¢ ¼ Cl), unfurling of the more flexible (CH2)n (n ¼ 4, 6) backbones of N,N¢-bis[2-hydroxy-3,5-di-(tert-butyl)-benzyl]-1,4-diaminobutane and -1,6diaminohexane has facilitated the observation of significantly more massive inter-metal displacements within the corresponding dinuclear complexes [160]. In a similar vein to the latter of these compounds, a single example of an analogous neutral, dinuclear system has been reported wherein an additional amine centre is incorporated in the salen-type ligand (see below) [161]. It has generally been found that isolation from a strongly coordinating Lewis base medium is required if the coordination sphere of the metal is to be completed not by a monoanionic ligand but by a neutral donor. The result is an ionseparated complex in which the Al-centred residue is formally monocationic. Several such complexes, incorporating both protic [107, 156, 162, 163] and aprotic [152, 161, 164] donors, have been characterised and these are recorded in Table 3. In all but one case, the metal centre is rendered six-coordinate by virtue of interactions with four donor sites on the salen-type ligand with one Lewis base molecule on each side of the chelating ligand plane. Notably, the use of amino[bis(N-propylenesalicylidenimine)] allows the observation of distinctly different complexes. While a monomeric, THF-solvate is revealed by the encapsulation of aluminium [161], a dinuclear species has been reported to deposit from hydrocarbon media in which the secondary amine is redundant and each phenoxide centre is formally bonded to an AlMe2+ moiety. The possibility of generating a 1 : 1 complex of aluminium by tethering together three aryloxide moieties has been explored. Synthesis [165] of the tripodal, polydentate ligand N{2-[(CH2)2N(H)CH2]-Ar}3 (Ar ¼ 4-chlorophenol) allows subsequent treatment with one equivalent of (H2O)9 Æ Al(ClO4)3 in the presence of excess (H2O)3 Æ NaOAc to afford a tris(phenoxide) complex [166]. More recently, the formation of Al-N(H) interactions has been precluded in the tris(2-hydroxybenzoyl)-2-aminoethylamine and tris(2-hydroxy-3-methoxybenzoyl)-2-aminoethylamine trianion complexes

97

The Synthesis and Structural Properties of Aluminium Oxide

Table 2. Neutral, mononuclear complexes of aluminium with dianionic salen-type (R1)

groups. R2 ¼ other ligand, L ¼ Lewis base, C.N. ¼ coordination number of the metal R1

N

N

O

O

t-Bu

N

N

O

O

t-Bu

t-Bu

N

O

O

t-Bu

t-Bu

N

N

O

O

Me

–

5

[124,156]

Me, OSiPh3

–

5

[157]

OSiPh3

–

5

[157]

i-Pr

–

5

[157]

OSiPh3

–

5

[157]

Cl

–

5

[158]

t-Bu

t-Bu t-Bu

N

N

O

O

t-Bu t-Bu

N

N

O

O

t-Bu

Ref.

Me

t-Bu

Me

C.N.

t-Bu

t-Bu

t-Bu

L

t-Bu

N

Me

t-Bu

R2

Me t-Bu

98

D. J. Linton Æ A. E. H. Wheatley

Table 2. (Contd.)

R1

t-Bu

N

N

O

O

t-Bu

t-Bu

R2

L

C.N.

Ref.

t-Bu

OMe

MeOH

6

[159]

t-Bu

OEt

–

6

[159]

–

THF

6

[161]

t-Bu

N

N

O

O

t-Bu

t-Bu

N H N

N

O

O

Al[OAr-2-C(O)N(H)(CH2)2]3NH+ Cl) (Ar ¼ C6H4, 3-MeO-C6H3) where the coordination number of the metal is raised to six by the action of both car˚, bonyl groups and deprotonated oxygen centres [mean Al-OAr ¼ 1.83 A ˚ ] [167]. Potentially hexadentate triaminocyclohexmean Al-O(¼C) ¼ 1.93A ane-based ligands form complexes of two distinct types. The neutral, cyclohexyl-based ligand [OC(H)C(H)N(H)Me2]3 bonds to aluminium only through its deprotonated O-centres and so acts as a purely tridentate, tripodal group in Al{[OC(H)C(H)N(H)Me2]3}2 [168]. However, in a similar vein to the hexadentate ligands discussed above [166] a triaminocyclohexane group tethers the three aryloxide ligands in Al[4-R-OC6H3-2CH2N(H)]3C6H9 (R ¼ H [169], NO2 [169, 170]) wherein the metal reveals bonds to all six amine N- and phenoxide O-centres. A closely related bonding arrangement, in which both oxygen and nitrogen centres interact with the metal atom, is noted in the distorted octahedral complex Al[OC(O)CH2N(CH2)2]3 [171, 172]. Moreover, the tripodal ligand [(O2CCH2)2(O)CCO2]4) bonds to aluminium via three of its four anionic oxygen centres in the 1 : 2 complex (Al{[(O2CCH2)2(O)CCO2]}2)) NH4+ [173]. Lastly, in the context of tethered polyaryloxide ligands, three examples of calix[4]arene-encapsulated aluminium centres have been noted. Hence coordination of an AlR2+ moiety by the 5,11,17,23-tetra-(tert-butyl)26,28-dimethoxycalix[4]arene-25,27-diolate dianion has been noted for R ¼ H [174], Cl [174], Me [175]. For the last of these, adducts with AlMe3 have also been characterised and these are discussed below. A single example of a doubly-metallated calix[4]arene ligand has also been reported

99

The Synthesis and Structural Properties of Aluminium Oxide

Table 3. Monocationic, mononuclear complexes featuring salen-type groups (R1). L ¼ Lewis base, C.N. ¼ coordination number of metal

R1

Counter-ion

t-Bu

N

N

O

O

t-Bu Me

C.N. Ref.

Cl), p-Me-C6H4) HOMe 6 SO3), BPh4)

[107, 156, 163, 163]

BPh4)

HOMe 6

[109, 162]

Cl)

H2O

[109, 162, 163]

Cl), F3CSO3), I), AlCl4)

HMPA, 6 THF

[152]

BPh4)

THF

[164]

t-Bu Me

N

N

O

O

t-Bu

N

N

O

O

t-Bu

t-Bu

t-Bu

L

N

N

O

O

6

t-Bu

N

N

O

O

t-Bu

t-Bu

6

t-Bu t-Bu

in which two aryloxide groups act as terminal ligands while two act as inter-metal bridges, yielding an (AlO)2 ring core [176]. In the context of bidentate ligands, and in a similar vein to the polyfunctional aryloxides discussed above [141, 142, 146, 147], the chelation of a metal centre by deprotonated b-diketones is a recurrent feature of monomeric organooxide complexes of aluminium. The employment of such ligands results in the observation of simple hexa-coordinate complexes which incorporate three [RC(O)C(H)C(O)R¢]) moieties (R ¼ R¢ ¼ Me [177, 178], CF3 [179], Ph [180, 181]; R ¼ t-Bu, R¢ ¼ CF3 [182]). More recently, the cocrystallisation of M(acac)3 (M ¼ Al, Cr) has allowed the crystallographic study of metal disorder in a series of solid solutions of stoichiometry Al1-xCrx(acac)3 (x ¼ 0.02–0.91) [183, 184]. Chelation of the metal centre similar to that reported in monomeric b-diketonates has also been noted in the presence of

100

D. J. Linton Æ A. E. H. Wheatley

2-methylmalonate ligands, though here only two such organic residues are incorporated in the monomer [185], while just one (N-arylcarbamoyl)phenylamine ligand exists in monomeric [(Me3Si)2(H)C]2Al[(O)CNPh]N(Ph) [(O)CN(H)Ph] [186]. Moreover, the N-nitroso-N-phenylhydroxylamine- and N-cyclohexyl-N-(hydroxo)diazeniumoxide-based complexes Al[ON(Ph) N(¼O)]3 [187] and Al[ON¼N(O)C6H11]3 [188] reveal only O,O¢-donor behaviour. Moreover, this bidentate motif is retained if the ligand incorporates N-donors instead. Thus, a single, chelating 2-(p-tolylamino)-4-(p-tolylimino)pentene ligand binds to the metal via both nitrogen centres in F3CS(¼O)2O Æ MeAl[N(C6H4Me)C(Me)]2CH [189]. The solution properties of the single reported example of a higher order bis(diisopropylaluminium phenoxide) have already been discussed [100]. Whereas treatment of the bis(diisobutylaluminium) phenoxide product of the 1 : 2 reaction between 2,2¢-(1,3-butadiyne-1,4-diyl)bis[6-(1,1-dimethylethyl)4-methylphenol] and Al(i-Bu)3 in DME yields a 1 : 1 complex, X-ray crystallography reveals that the species forms oligomers rather than discrete molecules in the solid state. This arises because instead of each of the two metal centres in the monomeric unit bonding to either donor atom of a DME molecule [that is, instead of the bis(aluminium) complex behaving as a bidentate and chelating Lewis acid], each metal atom interacts with just one donor centre in each of two Lewis bases [100]. 3.2.2 Mono-Bridged Complexes Both aluminium oxide (Section 2.2) and, as the next section will make clear, aluminium organooxide complexes are dominated by the fundamental Al(l-O)nAl building block. However, in the presence of polyfunctional ligands containing more than one type of electron-rich heteroatom, the connection of Group 13 metal centres via a polyatomic bridge is also viable. Thus, the treatment of trichloroacetamide with AlMe3 yielded [Cl3CC(O)N(H)AlMe2]2 with dimerisation resulting from the ability of the N-deprotonated carboxylic amide group to straddle aluminium atoms using both O- and N-centres [190]. Similarly, employment of the more sterically congested carboxylic amide ArC(O)NH2 [Ar ¼ 2,4,6-tris(trifluoromethyl)phenyl] has enabled the formation of a 12-membered metallocycle at the core of the trimer [ArC(OAlMe2)NH]3 [190]. More extensive core metallocycles have been noted if polyfunctional substrates are used. Hence, whereas the dimer [(C9H6N2)2(AlMe2)2]2 has been generated by the double deprotonation of 8-quinolylamine using AlMe3 [191], the addition of trace MeOH to the reaction mixture (or else the introduction of O2 to a solution of the dimer) has been noted to afford a product – MeAl(C9H6N2)2(AlMe2)[Al(Me2)]2OMe – which is resistant to further oxidation and which, in the solid state, reveals a structure based on an array of contiguous rings [191]. This aggregate reveals a core Al(Me)NAl(Me2)OAl(Me2)N 6-membered metallocycle in which an AlMe moiety is coordinated by both N-centres of two quinoline-based dianions such that the deprotonated amine centres are incorporated into the core heterocycle. These two N-centres

101

The Synthesis and Structural Properties of Aluminium Oxide

Et O

Al

Et

Et O

Et

Al

O

O

Al Et

O Et

O

Al

Et Et

Fig. 12. Dimeric tetra(aluminium) complex {2-[Et2AlOC(O)]-C6H4OAlEt2}2 [193]

also interact with the two remaining (methoxy-bridged) AlMe2 units. The fourth Group 13 metal atom straddles the core metallocycle by bridging between the two deprotonated N-centres. In this way the cluster core reveals an (AlN)2 ring as well as a second 6-membered heterocycle: Al(Me2)NAl(Me2)OAl(Me2)N [191]. Recently, the 2 : 1 reaction of Al(t-Bu)3 with anthranilic acid has afforded a 1 : 1 adduct between the unreacted aluminium-containing substrate and the deprotonated salt (t-Bu)2AlOC(O)-2-H2N-C6H4, wherein the Al(t-Bu)2+ moiety bridges between a carboxylate O-centre and the neutral amine donor atom, while the coordinated molecule of tris(tert-butyl)aluminium bonds only to the remaining carboxylate donor site such that the complex reveals an AlOCOAl motif in the solid state [192]. Meanwhile, the first example of an organoaluminium incorporating a hydroxy carboxylic acid ligand was obtained when the 1 : 2 reaction of salicylic acid with AlEt3 resulted in the isolation and characterisation of the dimeric tetra(aluminium) complex {2-[Et2AlOC(O)]C6H4OAlEt2}2 (Fig. 12) [193]. Here one metal centre bridges aryloxidecarboxylate groups to yield a 6-membered C3O2Al chelate ring within each monomeric building-block while the second aluminium centre bridges between the carboxylate functions in each half of the dimer, affording a flexible 12-membered (OAlOCOAl)2 heterocycle with respect to which the bidentate carboxylate groups adopt syn-anti conformation. Such large heterocycle synthesis has also enabled the formation of the similar dimeric polycarboxylate {2-[Me3AlOC(O)]-C6H4C(OAlEt2)O}2 with its intra-monomer 7-membered C4O2Al chelates and 16-membered inter-monomer metallocycle [194]. However, both of these examples were predated by the cyclic core of equivalent size observed in the tetrameric 2-oxobenzoxazole derivative [Me2AlNC(O)-OC6H4]4 (Fig. 13) [195]. This species, in which the O- and N-centres are 2- and 3-coordinate resepctively, was synthesised (as part of an investigation into potential precursors to new Ziegler-Natta catalysts) by the 1 : 1 reaction of AlMe3 with 2-mercaptobenzoxazole. Its structure reveals three types of aluminium centre in the solid state. While one Group 13 metal atom bridges between the N-centres of two organic residues, a second links the 2-oxo groups of two benzoxazole ligands. The remaining two Al-centres interact with the nitrogen atom of one ligand and the 2-oxo side-arm of a

102

D. J. Linton Æ A. E. H. Wheatley

Me Me

Al O

O N

Me Me N O

O

Al N Me O

Al

Me O

O

N

O Al

Me

Me Fig. 13. 2-oxobenzoxazole complex [Me2AlNC(O)-OC6H4]4 [195]

second ligand. The result of this coordinative behaviour is that the benzoxazole moieties are aligned in two nearly parallel pairs. However, the ˚ ) precludes any distance between aromatic systems in each such pair (>4.2 A graphitic p-stacking interactions [195]. Lastly, a 12-membered metallocyclic core lies at the heart of the (extensively metallated) penta(aluminium) salt of p-tert-butylcalix[6]arene [196]. This species incorporates three of the five Group 13 metal centres in the heterocyclic core, with each linking the phenoxide oxygen centres on neighbouring aryloxy groups ) the result being to leave just one O    O divide unbridged. Notably, the macrocycle adopts an unusual shape in which two opposite aromatic rings are nearly perpendicular to the remainder of the calixarene, causing the six oxide centres to be almost coplanar. The remaining two Al-centres bond to aryloxy groups on opposite sides of the calixarene superstructure (Fig. 14). 3.2.3 Oxygen Mono-Bridged Complexes Several examples of mono-bridged dialuminium oxide complexes were treated above. Species of the type Al(l-OR)Al have also been noted for R ¼ H or an organic residue. At the simplest level such species are dimeric or, for different ligand sets on each Al centre, dinuclear in the solid state. Such straightforward behaviour requires the employment of at least one polyfunctional organic residue capable of chelating, and therefore raising the coordination state, of at least one of the Group 13 metal centres. An example of this behaviour is furnished by the structures of Me2Al[l-OC(OMe)¼C(H)NMe2]AlMe2Cl [197], Me2Al[l-OCH2C(H)NH2CH2Ph]AlMe3 [198], Me2Al[l-O(CH2)2NMe2]Al(t-Bu)3 (Fig. 15) [199] and Cl2Al[l-OC(H)(CH2)4C(H)OSiMe3]AlCl3 [200]. In these species the internal stabilisation of one Al-centre by either an amine [197–199] or a siloxy [200] function precludes association beyond that resulting from

103

The Synthesis and Structural Properties of Aluminium Oxide

Me

THF Al

Me

O

t-Bu t-Bu O

t-Bu

Me

Al

O Al

Me

O

O

t-Bu

Me

Al

Me

t-Bu

t-Bu O Me

Al

Me

Me Fig. 14. The penta(aluminum) salt of p-tert-butylcalix[6]arene [196]

t-Bu t-Bu Al

N

O

Me Me

Al

t-Bu Me

Me

Fig. 15. Me2Al[l-O(CH2)2NMe2]Al(t-Bu)3 [199]

alkoxide stabilisation of the otherwise trivalent alkylaluminium [198, 199], mixed alkylaluminium halide [197] or aluminium halide [200] moiety. One unusual example of this motif is particularly noteworthy. A carbonyl bridge is seen in [2-(Me2Al)-3-(t-Bu)-5-Br-C6H2]2CO wherein formation of an Al2O fragment is reported in tandem with the generation of two edge-fused C3O2Al rings which share the bridging carbonyl function [201]. Whereas a mononuclear alkoxide complex of aluminium which incorporates azaindole ligands has been noted [116], so azaindoles have been reported to bridge Al-centres in oxides [23, 41, 42]. This is also true of an oxidealkoxide which incorporates the bis(trifluoromethyl)isopropoxide ligand [42]. In the case of dinuclear aggregate {[(F3C)2C(H)]O}2Al[l-OC(H)(CF3)2]-

104

D. J. Linton Æ A. E. H. Wheatley

[l-C7H5N2]2AlMe each N-centre of the azaindole ligand bonds to a different metal centre such that mono-bridging of the metals by one alkoxide fragment is supported [42]. In the case of trimeric [(t-Bu)2AlOH]3 mono-bridging results from the ability of the aggregate to satisfy the Lewis acidity of aluminium by cyclising to give a 6-membered metallocycle [24]. Such a motif is replicated for the alkoxyaluminium halide (Cl2AlOMe)3 [202]. Akin to that of {[2,4,6-(t-Bu)3-C6H2]AlO}4 [43] (Section 2.2), an (AlO)4based core is revealed by the product of facile reaction between a diethyl ethersolvated, polycyclic tetra(aluminium) oligosiloxane and pyridine [203]. In a similar vein to the esoteric structures of several lithium oligosiloxane aluminates which form part of the same study, the solid-state structure of {Al[O(SiPh2O)2](l-OH Æ py)}4 (Fig. 16) incorporates an 8-membered heterocyclic core which is composed of four hydroxy-bridged aluminium centres and which is supplemented by four annelated Al2Si2(l-O)3(l-OH) rings [203]. The Al-OH bond lengths are of particular note in the context of hydrogen bonding between the bridging hydroxide groups and the coordinating ˚ they are intermediate between those noted in pyridine molecules: at 1.769(1) A the diethyl ether-solvated precursor and the triethylamine coordinated analogue – an observation consistent with the pKa values of the respective Lewis bases. A similar arrangement is noted for the dianionic congener in ({Al[O(SiPh2O)2]}4(l-O)2(l-OH)2)2) 2(HNEt3)+ [204]. Lastly, the previously-mentioned study into the low-surface-area aluminium oxide-hydroxides diaspore and boehmite by 1H CRAMPS (Section 2.2.1) [46] has also been extended to the aluminium hydroxides gibbsite and bayerite, revealing three distinct resonances in a 3 : 2 : 1 ratio in these systems and suggesting the presence of Al2(l-OH) groups with 6-coordinate metal centres.

Ph

Ph

Ph

Si

Si Ph O

O

O

Ph

O

Al

Si Ph

Ph

py·HO Al O Ph OH·py Si

Si Ph

O Ph Al OH·py

O O Ph Si Ph

Fig. 16. {Al[O(SiPh2O)2](l-OH Æ py)}4 [203]

O Ph Si

Al

O

Ph

Si

OH·py

O

Ph

O

Ph

The Synthesis and Structural Properties of Aluminium Oxide

105

3.2.4 Di-Bridged Complexes As noted in Section 2.2, the oxophilicity of aluminium means that, as far as its oxides are concerned, the fundamental structural building block is the (AlO)2 ring. However, while this principle extends to aluminium hydroxides and organooxides (see below), examples have also been reported in which complexes of this type reveal bridging by other relatively hard, electron-rich heteroatoms. The monomeric component of the unusual co-crystalline monomer-dimer system [Me3N Æ Al(H)2OAr] Æ [Me3N Æ Al(l-H)(H)OAr]2 [Ar ¼ 2,6-di-(tertbutyl)-4-methyl-phenoxide] has already been discussed [86]. The dimeric part takes the form of a dihydridoaluminium aryloxide which is based on an (AlH)2 ring core that utilises the bridging capabilities of aluminium-bonded ˚ (in the monomer); Al-H ¼ 1.515 A ˚ , Al-(l-H) hydride [Al-H ¼ 1.731–1.884 A ˚ ¼ 1.630 A (in the dimer) (no standard deviations given)]. The ability of hydrogen to form 4-membered metallocycles is also demonstrated [in tandem with the formation of Al(l-O)2Al fragments – see below] in the tetranuclear complex {(l-H)(H)Al[l-OC(H)Me(t-Bu)]2AlH2}2 [205]. Here, both terminal ˚ ] and bridging Al(H)(l-H) [Al-H ¼ 1.49(5) A ˚ AlH2 [Al-H ¼ 1.38(7)-1.57(5) A ˚ Al-(l-H) ¼ 1.60(4)-1.85(4) A] moieties are noted, the latter being employed in the formation of the Al(l-H)2Al dimer core by virtue of the ability of two {Me2Al[l-OC(H)Me(t-Bu)]}2 dimers to associate (cf. {Me2Al[l-OCMe(t-Bu)2]}2) [205]. Notably, each dimeric component of the doubly hydridebridged aggregate contains asymmetrically-substituted C-centres with the same chirality: S,S in one dimer and R,R in the other. A bridging hydride moiety has also been seen in a more extensive bis(aluminium) heterocycle with a carboxylate bridge wherein each O-centre interacts with a different Group 13 metal atom (cf. mono-bridging carboxylates [193, 194]). Hence, a 6-membered CO2Al2H metallocycle defines the solid-state structure of {[(Me3Si)2C(H)]2Al}2(l-O2CPh)(l-H) [206]. A similar bridging mode to that noted for the carboxylate ligand has been reported for a serendipitous by-product of sequential carbodiimide insertion and thermal rearrangement obtained in the presence of trace oxygen. Hence, [(i-Pr)N]2C ordinarily combines with AlMe2Cl to yield Me2AlN(i-Pr)C(Cl)N(i-Pr). However, the by-product MeClAlN(i-Pr)C(Me)N(i-Pr) is also isolable from the crude reaction mixture, as is the species {(Me2Al)2{l-[N(i-Pr)]2CMe}(l-OMe) if oxygen is not scrupulously excluded from the system [207]. Variable behaviour of the potenially tridentate 2,6-di-(N,N-dimethylaminomethyl)-4-methylphenoxide ligand has already been noted to result in the isolation of a simple mono(aluminium) monomer whose ligands adopt both mono- and bidentate coordination modes (see above) [153]. However, more complex bis- and tris(aluminium) systems have also been revealed in the presence of this ligand. Thus, whereas the simple aluminium tris(aryloxide) monomer results from the 1 : 3 treatment of AlMe3 with the parent phenoxide, the adducts (Me3Al)n Æ Me2Al[2,6-(Me2NCH2)2-4-Me-OC6H2] (n ¼ 1, 2) are yielded by the analogous 2 : 1 and 4 : 1 reactions, respectively. These two

106

D. J. Linton Æ A. E. H. Wheatley

co-complexes both reveal terminally N- and O-bonded trimethylaluminium fragments with bridging of the demethylated alkylaluminium moiety between the deprotonated O-centre of the organic ligand and one N,N-dimethylaminomethyl side-arm yielding 6-membered C3NOAl heterocycles in both instances [153]. Mixed N,O-dibridging of two Group 13 metal centres has been noted in the structure of the co-complex between the internally coordinated Al(III) salt of (2-hydroxybenzyl)ethylenediamine and Al(NMe2)3 [208]. However, although a straightforward 4-membered ring has been observed in the N,O-dibridged complex (H2Al)2[l-N(t-Bu)SiMe2H][l-O(t-Bu)] [209], the formation of spirocyclic arrays has also been noted for such mixed systems. Hence, the reaction of AlMe3 with 2-(HOCH2)-C6H4NH2 has resulted in the isolation of {(Me2Al)2[l-N(H)-C6H4-CH2(l-O)]}2AlMe (Fig. 17) in which the central AlMe2+ group is chelated by both organic residues, resulting in two types of Al-O and Al-N interaction [210]. Whereas one complex which contains an Al(l-N)2Al fragment has already been discussed by virtue of its inclusion of an Al(l-X)Al (X ¼ O, N) motif [191], a uniquely N-bridged, 4-membered ring has also been noted in the species isolated from the derivatisation of pre-formed bis(aluminium) complexes containing salen-type ligands [211]. The product of reaction between N,N¢-bis(2-hydroxybenzyl)-1,2-diaminobenzene and three equivalents of AlEt3 reveals a contiguous array of aluminium-containing rings by virtue of the cis-coordination mode adopted by a single salen-type chelate. Three distinct metal environments are observed, with a core AlEt2+ moiety being O,O¢,N,N¢-bonded, an AlEt2+ fragment being N,N¢-bonded [to give an Al(l-N)2Al motif] and another such group being O,O¢-bonded [yielding an Al(l-O)2Al ring]. The consequence is that both Al-N and Al-O bonds fall ˚ , mean EtAl-N ¼ 2.001 A ˚ , mean into two categories (mean Et2Al-N ¼ 1.994 A ˚ , mean EtAl-O ¼ 1.889 A ˚ ). Meanwhile, the employment of Et2Al-O ¼ 1.844 A Al(i-Bu)3 with N,N¢-bis(2-hydroxybenzyl)-1,2-diamino(4,5-dimethyl)benzene yields a species in which the diaminobenzene-bridging dialkylaluminium fragment is excluded, with dimerisation of the remaining bis(aluminium) complex occurring via (AlN)2 ring formation (Fig. 18) [211]. Recently, reactions of dimeric [Me2Al(l-NEt2)]2 with different tripodal ligands have been reported. Along with an Al(l-O)2Al ring, two Al(l-N)2Al motifs are noted in the core of the crystal structure of the tetra(aluminium) Me N Me

Al

N

O O Al

Al

Me Fig. 17. {Me2Al[l-N(H)-C6H4-CH2(l-O)]}2AlMe [210]

Me

Me

107

The Synthesis and Structural Properties of Aluminium Oxide

Me Me

i-Bu N

Al

O N

Al

O

N

i-Bu

O

Al

i-Bu O

Al

N

i-Bu Me Me Fig. 18. The product of reaction between Al(i-Bu)3 with N,N¢-bis(2-hydroxybenzyl)-1,2diamino-(4,5-dimethyl)benzene [211]

Me Me Al F3C

O

N

S

t-Bu

CF3

S

N

CF3

t-Bu

O

CF3

Al Me Me

Fig. 19. (NSOAl)2-based [ArNS(t-Bu)OAlMe2]2 [Ar ¼ 3,5-bis(trifluoromethyl)phenyl] [213]

salt of MeC(CH2NH2)2CH2OH [212]. A complex array of 4- and 6-membered heterocycles is afforded through bridging of AlMe2 moieties between the two ˚ ] or between pairs of the four N-centres (mean O-centres [Al-O ¼ 1.936(4) A ˚ Al-N ¼ 1.938 A) in two of the multi-functional organic residues, and also through tripodal chelation of a central, trigonal prismatic metal atom ˚ , mean Al-N ¼ 2.018 A ˚ ]. [Al-O ¼ 2.019(4) A While the isolation of ClMeAl[2,6-(t-Bu)-OC6H3] from tert-butylamine has already been reported to result in the simple, monomeric mono-solvate [132], the incorporation of Group 17 elements as inter-aluminium bridges has been revealed in the (AlCl)2-based dimers [ArO(Me)Al(l-Cl)]2 {Ar ¼ 2,6-(t-Bu)2-C6H3 [132], 2,6-(t-Bu)2-4-Me-C6H2 [129]}.

108

D. J. Linton Æ A. E. H. Wheatley

More extensive metallocyclic arrays have been generated using organic residues capable of constituting polyatomic bridges between aluminium centres. One such residue – the delocalised N…C(R)…O moiety (R ¼ organic group) – has been achieved either by O- or N-deprotonation. Thus, 2-hydroxypyridine undergoes facile reaction with Al(t-Bu)3 to give the dimeric, 8-membered heterocycle [2-(t-Bu)2AlO-C5H4N]2 wherein association results from the coordinative properties of nitrogen [94]. A similar motif has also been noted for [ArNS(t-Bu)OAlMe2]2 [Ar ¼ 3,5-bis(trifluoromethyl)phenyl]. This complex is afforded by the employment of the ArN(H)S(¼O)(t-Bu) ˚] racemate in tandem with AlMe3 and it incorporates both Al-O [1.836(2) A ˚ and Al-N [1.932(2) A] interactions in a core (NSOAl)2 heterocycle (Fig. 19) [213]. The sulfur centres in this dimer (and its gallium analogue) exhibit R,S configurations, in contrast with the uniquely R,R or S,S configurations noted if the terminal methyl groups are replaced by more sterically demanding CH2SiMe3 moieties [213]. Similar behaviour has been revealed by a study into the reactivity of primary carboxylic amides towards alanes. In this context, [Cl3CC(O)N(H)AlMe2]2 reveals a dimeric structure in the solid state by virtue of the ability of the N-deprotonated carboxylic amide group to bridge Al-centres [190]. 3.2.5 Oxygen Di-Bridged Complexes It has been widely noted that the (AlO)2 4-membered ring motif which pertains for aluminium oxide systems also represents the basic building-block for hydroxide, alkoxide and aryloxide systems and this section will discuss examples of these in which only oxygen acts as a bridge between Group 13 metal centres. Such bridging activity extends to two ligand types: those in which single O-centres bridge between metal atoms (which results in 4-membered ring cores) and those where the bridging fragments are more extensive (e.g., carboxylate) and each oxygen donor bonds to just one aluminium centre (yielding larger heterocycles). The majority of examples pertain to the former motif. While conforming to the straightforward principles of Al2(l-OR)2R4 structures, they fall into two categories. In the first instance, many dimeric systems are centrosymmetric and/or reveal Al-centres that are bonded to two different terminal organic residues (Fig. 20a and Table 4) though, alternatively and less commonly, each Group 13 metal atom may bond to a different pair of terminal organic residues (Fig. 20b). The observation of a bonding arrangement of the type shown in Fig. 20a in (AlO)2-centred dimers is a recurrent feature of aluminium hydroxides and organooxides of simple, electron rich ligands [44, 99, 110, 125, 134, 204, 205, 214–231]. Those more complex dimers which have been noted take two main forms: both polydentate terminal and bridging ligands allow multiple interactions with the Group 13 metal centres. A handful of bis(hydroxy)-bridged dimeric compounds have been reported in which aluminium centres bond to two different terminal organic residues. While complexes with mixed terminal groups – that is Al2(l-OH)2R2R¢2 [64,

The Synthesis and Structural Properties of Aluminium Oxide

109

Fig. 20. The two basic classes of simple R2Al(l-OR)2AlR2 dimer

232] – result from the presence of monofunctional ligands, species of the type Al2(l-OH)2R4 and Al2(l-OH)2(OR)4 have also been seen. The former class is represented when R ¼ 2,4,6-trimethylphenyl [22] whereas the latter type has been found to result if OR represents 2-(2-oxyphenyl)-2-benzoxazoline (with the organic anion chelating aluminium with its deprotonated phenoxide centre as well as its benzoxazoline nitrogen atom [233]) or the N,O,O¢-tridentate bis(carboxylate) [(O2CCH2)2NH]2) [234]. The inclusion of electron-rich side-arms on organooxide bridging groups often results in the chelation of aluminium and the increase of its coordination number from four to five or, less commonly, six. Pendant primary [96, 198, 235, 236], secondary [235, 236] and tertiary [94, 96, 197, 233, 237, 238] amines have all been incorporated into the bridging groups and (with the exception of {Al[l-O(CH2)3NMe2](t-Bu)2}2 [215]) reveal N-donor behaviour towards the Lewis acid metal centres. In a similar vein, coordination by oxygen renders each metal centre 5-coordinate in several dimeric species. Etherate oxygen centres stabilise aluminium in {Al[l-O(CH2)2OR]R¢R¢¢}2 [R ¼ Ph, R¢ ¼ R¢¢ ¼ Cl, Br; R ¼ Ph, R¢ ¼ Cl, R¢¢ ¼ Et; R ¼ (CH2)2O(CH2)2O(CH2)2Br, R¢ ¼ R¢¢ ¼ Br; R ¼ R¢ ¼ Me, R¢¢ ¼ I; R ¼ Me, R¢ ¼ R¢¢ ¼ CH2SiMe3; R ¼ R¢ ¼ R¢¢ ¼ Me; R ¼ Me, R¢ ¼ R¢¢ ¼ t-Bu; R ¼ R¢ ¼ (CH2)2OMe, R¢¢ ¼ 2,6-(t-Bu)2-4Me-OC6H2] [96, 97, 216, 237, 239, 240] but while similar stabilisation is retained in {Al[l-O(CH2)3OMe]Me2}2 it is absent in {Al[l-O(CH2)3OMe](tBu)2}2 [96]. Furthermore, oxygen-donor hydroxide [242, 243] and carbonyl groups [96, 244–247] have also been noted. However, the inhibition of metal centre chelation has led to the observation of tri-coordinate aluminium for R ¼ O(CH2)2OPh, R¢ ¼ R¢¢ ¼ i-Pr [239], Ph [243] while both modes of behaviour are apparent (with only bridging ligands chelating metal centres) in dimeric {Al[l-O(CH2)2OMe](OAr)O(CH2)2OMe}2 [OAr ¼ 2,4-di-(tertbutyl)-4-methylphenoxide] [237]. Similarly, the behaviour of the 2-methoxyphenoxide fragment in [Al(l-OCH2O-2-OMe-C6H4)R2]2 is rather variable. Whereas for R ¼ Et, coordination of the methoxy groups is reported [248], for R ¼ CH2(i-Pr) both tetra- [239] and (internally stabilised) penta-coordinate [248] metal centres have been observed. For the 2-MeS analogue of bridging 2-methoxyphenoxide, Al-S interactions have been seen [248]. Aluminiumsulfur interactions persist in {Al[l-O(CH2)nSMe]R2}2 (n ¼ 2, R ¼ i-Bu) but are absent for n ¼ 2, 3, R ¼ t-Bu [96]. The chelation of aluminium has been recorded in the presence of formally dianionic bridging units. In two examples the metal centres are bridged by the carbanion-containing organooxide moiety [O(CH2)3CH2]2) with the presence

O

O

O

O

Me

Me

Me

N

O

O

Me

O

O

O

O

Me

Me

CH2(i-Pr)

Et

Me

CH2(i-Pr)

Me

OH

Me

OEt

O

Me

R2

OEt

OH

R1

Me

¼ R2

¼ R2

¼ R2

–

–

–

–

–

C(SiMe3)3

¼ R2

–

–

L

Si(SiMe3)3

¼ R2

R3

5

5

5

5

4

4

4

C.N.

[96]

[96]

[96]

[94]

[64]

[44]

[22]

Ref.

Table 4. Dinuclear complexes with an (AlO)2 core in which each metal centre bonds to two types of terminal ligand (R1–3 are explained in Fig. 20a). L ¼ Lewis base, C.N. ¼ coordination number of the metal

110 D. J. Linton Æ A. E. H. Wheatley

O

O

Me

O

O

O

MeO

O

O

O

O

O

Me

Me

– –

¼ R2 ¼ R2

i-Bu

t-Bu

SMe

–

–

–

–

–

–

¼ R2

¼ R2

¼ R2

¼ R2

Et

¼ R2

t-Bu

Me

t-Bu

i-Bu

t-Bu

Cl

SMe

OMe

SMe

O

Me

O

4

5

4

5

4

5

5

5

[97]

[97]

[97]

[97]

[97]

[97]

[97]

[96]

The Synthesis and Structural Properties of Aluminium Oxide

111

t-Bu

O

O

t-Bu

H2N

O

Me2N

O

Me2N

O

Me2N

O

Me2N

O

R1

Me

Me

Table 4. (Contd.)

i-Bu

t-Bu

H

Me

i-Bu

t-Bu

R2

¼ R2

¼ R2

¼ R2

¼ R2

¼ R2

¼ R2

R3

–

–

–

–

–

–

L

4

5

5

5

5

4

C.N.

[99]

[97]

[97]

[97]

[97]

[97]

Ref.

112 D. J. Linton Æ A. E. H. Wheatley

O

O

OEt

t-Bu

t-Bu

t-Bu

t-Bu

t-Bu

O

O

t-Bu

t-Bu

O(i-Pr)

O(i-Pr)

OMe

t-Bu

Me

Me

Me

t-Bu

O

N O

N

t-Bu

t-Bu

¼ R2

Me

–

–

–

–

–

–

4

4

4

6

[125]

[110]

[110]

[107]

The Synthesis and Structural Properties of Aluminium Oxide

113

O

Me

Me

Me

H

OC(t-Bu)2Me

SiMe3

Et

Me

OSiPh3

Ph

Ph

OMe

¼ R1

CH2(i-Pr)

R2

O(t-Bu)

H2N

O

H2N

O

Me2N

O

O

Me

R1

Table 4. (Contd.)

– –

¼ R2

–

–

–

–

–

L

¼ R2

¼ R2

¼ R2

¼ R2

¼ R1

¼ R2

R3

4

4

5

5

5

4

4

C.N.

[205]

[204]

[198]

[198]

[197]

[134]

[134]

Ref.

114 D. J. Linton Æ A. E. H. Wheatley

i-Pr

O

O

O

O

O

O

O

O

Ph

Ph

Ph

Ph

O(i-Pr)

O

O

i-Pr

NMe2

Cl

Cl

Br

Cl

t-Bu

t-Bu

Me

Et

¼ R2

¼ R2

–

–

–

–

–

¼ R2

¼ R2

–

–

¼ R2

¼ R2

5

5

5

5

4

4

4

[216]

[216]

[216]

[216]

[215]

[215]

[214]

The Synthesis and Structural Properties of Aluminium Oxide

115

O

O

Me

Me

TMP

Me

Me

OEt

CH2

t-Bu

Me

Cl

Me

t-Bu

t-Bu

B

Me

i-Pr

R2

O(i-Pr)

O

O

O

MeO

O

MeO

O

R1

Table 4. (Contd.)

¼ R2

Cl

C5Me5

¼ R2

¼ R2

¼ R2

¼ R2

¼ R2

R3

–

–

–

–

–

–

–

–

L

5

4

4

4

4

4

4

4

C.N.

[222]

[221]

[220]

[219]

[218]

[217]

[216, 227]

[216]

Ref.

116 D. J. Linton Æ A. E. H. Wheatley

O

Me

i-Pr

O

i-Pr

O

MeO

Me

Me

Me

Me

Me

CH2(i-Pr)

Me

Me

t-Bu

OPh

O

H

Me

O(t-Bu)

F

F

H

F

F

¼ R1

O(t-Bu)

O

F

O(t-Bu)

–

¼ R2

¼ R2

¼ R2

¼ R2

–

–

–

–

–

¼ R2

¼ R2

–

–

–

O(t-Bu)

¼ R2

¼ R1

4

4

4

5

4

4

4

4

4

[228]

[228]

[228]

[227]

[226]

[225]

[225]

[224]

[223]

The Synthesis and Structural Properties of Aluminium Oxide

117

O

Me

CH2

OH

O(t-Bu)

OEt

O

Me

O

O(i-Pr)

R1

Me

Me

Table 4. (Contd.)

N

N

i-Pr

i-Pr

i-Pr

i-Pr

¼ R1

N

N

O

Me

H

R2

B

Et

Et

i-Pr

i-Pr

OMe

t-Bu

H

¼ R1

–

O

t-Bu

¼ R2

¼ R2

R3

Me

–

–

–

–

–

–

L

4

4

5

5

4

4

C.N.

[232]

[231]

[231]

[230]

[230]

[229]

Ref.

118 D. J. Linton Æ A. E. H. Wheatley

Me2N

O

OH

OH

OH

Me

O

O

N

O

HN

Me

O

O

i-Pr

i-Pr

i-Pr

i-Pr

O

i-Pr

i-Pr

¼ R2

–

¼ R2

Cl

–

H2O

–

–

5

5

6

4

[235]

[234]

[233]

[232]

The Synthesis and Structural Properties of Aluminium Oxide

119

O

O

Me2N

O

O

HN

O

O

HN

H2N

O

Me2N

R1

Me

Et

Ph

Me

Table 4. (Contd.)

H

Me

Me

Me

Me

R2

O

t-Bu

t-Bu

¼ R1

¼ R2

¼ R2

¼ R2

¼ R2

R3

Me

–

–

–

–

–

–

L

5

5

5

5

5

5

C.N.

[237]

[237]

[236]

[236]

[236]

[235]

Ref.

120 D. J. Linton Æ A. E. H. Wheatley

O

O

Ph Ph

Ph

Me

HO

Me

Me

O

Me

HO

O

O

O

O

O

O

O

O

Me2N

O

O

Br

Me

Cl

Cl

CH2SiMe3

I

Br

Cl

¼ R2

¼ R2

¼ R2

¼ R2

Me

¼ R2

Et

–

–

–

–

–

–

–

4

5

5

5

5

5

5

[243]

[243]

[242]

[241]

[240]

[239]

[238]

The Synthesis and Structural Properties of Aluminium Oxide

121

O

O

O

O

O

O

S

S

Me

Me

C

Me

Me

O

O

Me

OEt

O

O

Me

O

R1

O

Cl

OEt

Table 4. (Contd.)

3

Me

CH2(i-Pr)

F

Et

Et

Me

Me

R2

¼ R2

¼ R2

Me

¼ R2

¼ R2

¼ R2

¼ R2

R3

–

–

–

–

–

–

–

L

5

5

5

5

5

5

5

C.N.

[248]

[248]

[247]

[246]

[245]

[245]

[244, 245]

Ref.

122 D. J. Linton Æ A. E. H. Wheatley

O

O

O

O

H2C

O

H2C

O

O

O

O

O

Me

Me

–

Cl

–

–

Et

CH2(i-Pr)

Me

–

i-Pr

i-Pr

C(SiMe3)3

¼ R2

¼ R2

i-Pr

THF

HOEt

–

–

–

–

6

5

4

5

5

5

[254]

[253]

[250]

[249]

[248]

[248]

The Synthesis and Structural Properties of Aluminium Oxide

123

124

D. J. Linton Æ A. E. H. Wheatley

of the negatively charged carbon centre inducing the metal to participate in a formal Al-C bond [249, 250]. This motif is more commonly displayed when aluminium is able to interact with just one polyfunctional terminal ligand which is capable of both multiple deprotonation and metal ion chelation. In this context, two salen-H2 moieties have each undergone double-deprotonation during the formation of the simple dimer [(salen)AlOMe]2 [251] wherein each metal centre is rendered 6-coordinate by the bridging action of the methoxy groups and the tetradentate nature of the doubly-deprotonated salen ligand. Such polydentate ligand behaviour is mirrored by both the Me2-salen-H2- [107] and the N,N¢-bis(2-hydroxybenzyl)-1,2-diaminoethane-based analogues [252]. Bis(oxygen) chelation of aluminium by a terminal 2,2¢-methylene-bis[4-methyl6-(tert-butyl)-phenoxide] ligand has been reported in the isopropoxy-bridged dimer {Al(MBMP)[l-O(i-Pr)]}2 [110]. Furthermore, the raction of two equivalents of Et2B(l-pz)2AlEt2 with molecular oxygen has yielded the dimeric complex [EtAl(l-OEt)(pz)2BEt2]2. This species reveals an (AlO)2 core peripheral to which each Group 13 metal centre is chelated by one bidentate bis(pyrazolyl)borate ligand such as to afford two 6-membered AlN4B metallocycles [231]. The inclusion of external solvent in a complex of the type discussed here (that is, incorporating chelating, terminal ligands) has been reported only once with aluminium being hexa-coordinate in the hydroxy-bridged dimer {H2O Æ Al(l-OH)[(O2CCH2)2NH]2}2 (see above) [234]. The process of chelating the metal centre in a dimeric complex using either a terminal or a bridging ligand, since it derives from the Lewis acidity of aluminium, normally precludes external solvation. However, for the latter ligand-type, in a few cases the participation of either protic (e.g. HOEt [253]) or aprotic (e.g. THF [254]) solvent molecules in the complex has been reported. In both of these instances, internal coordination is facilitated by the presence of relatively flat alkoxyfuran [253] and (R)-binol [69, 70] bridging groups, though in the first of these instances it is noteworthy that each metal centre carries a formal positive charge (countered by a lattice chloride ion). Of course, exceptions to the structure-types illustrated in Fig. 20 exist. Thus, although [(t-Bu)O]2Al[l-O(t-Bu)]2Al(OR)2 pertains to the structure-type represented in Fig. 20b [231], (t-Bu)O[(t-Bu)OO]Al[l-O(t-Bu)]2Al(OR)2 [R ¼ 2-MeO(O)C-OC6H4] reveals two identical terminal ligands on one metal centre and two different terminal ligands on its counterpart [255]. Oxygenbridging behaviour is adopted by the two 2-oxy-3,5-di-(tert-butyl)-benzyl(tertbutyl)amino ligands in Me2Al{l-O[(t-Bu)2-C6H2-CH2N(t-Bu)]}2Al wherein the oxide centres bridge aluminium atoms while the deprotonated N-centres both (terminally) bond to the same Group 13 metal [256]. Monomeric calix[4]arene complexes of aluminium have already been discussed [173–175]. In one instance, however even such sterically congested systems may associate, with tetradeprotonation of the ligand resulting in a –2 charge on the dimeric unit (balanced by ion-separated trimethylammonium ions) [257]. The residence of a single positive charge on the Group 13 metal-containing cluster also leads to deviations from previously discussed structure types. Thus, the monocation {Me2Al[l-OCPh2-C4H7NH]2AlMe}+ [counterbalanced by (AlMeCl3))] contains both tetra- and penta-coordinate aluminium sites, the latter being chelated by

125

The Synthesis and Structural Properties of Aluminium Oxide

a,a-diphenyl-2-pyrrolidinylmethoxy groups [237]. Whereas Me2-salen has been noted to behave only as a terminal ligand in a straightforward bis(aluminium) complex [162], doubly-deprotonated Et2-salen-H2 has been found to act both as an O- and N-coordinating ligand – bridging the two Alcentres in the cationic component of {Me2Al[l-(Et2-salen)]AlMe}+ (AlMeCl3)), and also chelating the AlMe2+ fragment in this moiety [258]. An aluminium tetrachloride anion counters the charge on the tris(acetylacetonate) monocation [(AlCl)2R3]+ (R ¼ pentane-2,4-dione). The ability of b-diketone ligands to chelate aluminium has already been discussed in the context of monomeric complexes and, while this behaviour is retained in higher order organoaluminium species, the presence of two metal centres allows for both terminal and bridging bidentate ligand behavour. Hence, in this instance one metal ion is chelated by a single acetylacetonate ligand whilst the remaining two such groups both bridge between this and the second metal centre and also chelate the latter Group 13 atom [259]. 18-crown-6 encapsulates the dicationic, bis(hydroxy)-bridged fragment [MeAl(l-OH)2AlMe]2+ (for which there exist charge-balancing aluminium tetrachloride ions) [260]. Moving away from (AlO)2-based species, more extensive ring-cores have been reported for bis(aluminium) complexes which incorporate polyfunctional organic acids. Hence, the 1,3-diol neol-H2 undergoes mono-deprotonation (as evidenced by 1H NMR spectroscopy) in the presence of Al(t-Bu)3 to yield (t-Bu)2Al(neol-H) with dimerisation yielding an 8-membered metallocycle [261]. Three equivalents of a complicated polydentate bis-hydroxypyridinone ligand (H2L, Fig. 21) have been noted to bridge two aluminium atoms to yield a triple helicate of M2L3 stoichiometry which contains chiral (L or D) metal centres [262] and which excludes terminal ligands [263]. This behaviour has been reported in the homochiral helicate (DD,LL)-Al2L3, the stucture of which contrasts with the heterochiral mesocate structure of (DL)Ga2L3. The inter-metal distances differ significantly depending on the ˚ , Ga    Ga ¼ 9.74 A ˚ ]. Moreover, only the structure-type [Al    Al ¼ 7.13 A aluminium helicate contains a significant cavity and, consistent with this, only this species reveals an encapsulated molecule of H2O ) with host-guest interactions implicated for the relatively close approach of the two metal centres as compared with that recorded for the gallium system [263]. In spite of their more complex structures, higher order aluminium organooxides are still based on straightforward (AlO)2 ring motifs of the kind discussed above (Section 2.2). In one instance a tetra(aluminium) system Me

Me

NH

HN N Me

Me N O O Fig. 21.

OH

O HO

O

126

D. J. Linton Æ A. E. H. Wheatley

has been obtained which features two independent (AlO)2 rings. Hence, polyfunctional 2,7-di-(tert-butyl)-9H-fluoren-1,8-diol [¼ (HO)2R] can be doubly deprotonated using AlEt3 to afford (Et2Al)2[(O)2R]2(AlEt2)2 in which the two dianionic conjugate bases each contribute one oxide centre to one 4-membered (AlO)2 ring and the remaining oxide centre to a second such heterocycle [264]. More generally, rather than revealing structures analogous to the edge-fused Al-O metallocyclic arrays common for aluminium oxides and mixed oxide-hydroxides/organooxides (Section 2.2) several tris(aluminium) systems reveal spirocyclic units fused at a central aluminium centre (cf. {Me2Al[l-N(H)-C6H4-CH2(l-O)]}2AlMe [210]). These have been seen with both monofunctional and polyfunctional (chelating) ligands. Amongst the former are hydroxide bridges, such as are noted (in tandem with a bridging carboxylate ligand) in the remarkable material (RO2Al)2(l-O2CMe) (l-OH) {[RO2Al(Æ OH2)(l-OH)2]2Al(l-OH)}2 (RO2 ¼ 2-oxyphenylsalicylideniminate) (Figs. 22a/b). This species has been isolated from the 1 : 1 : 1 reaction of (H2O)6 Æ AlCl3, NaOH and N-salicylidene-o-aminophenol as part of an investigation into the ability of Schiff bases to coordinate to aluminium in a similar way to that which occurs in biological systems and it represents a new type of aluminium clathrate [265]. The {[RO2Al( Æ OH2)(l-OH)2]2Al(l-OH)}2 component in this co-crystalline system reveals a novel spirocyclic pattern whereby three Al-centres participate in RO2-intercepted bis(hyroxy)-bridged chains with the middle metal ions associating via the formation of a core Al(l-OH)2Al motif (Fig. 22b) [265]. Remaining with monofunctional bridging ligands, the isopropoxide RAl{[l-O(i-Pr)]2AlR2}2 (R ¼ H [205], Cl [47]) – the hydride being able to polymerise through the formation of multiple Al(l-H)2Al rings [205] (cf. tetranuclear {(l-H)HAl[l-OC(H)Me(t-Bu)]2AlH2}2 [205]) – and the cyclohexyloxide ROAl[(l-OR)2Al(OR)2]2 (R ¼ C6H11) have been reported [266]. Concerning polyfunctional bridging ligands, diolate moieties have been utilised. These chelate AlH2+, AlCl2+, AlMe2+ and [AlC(H)(SiMe3)2]2+ fragments, respectively, in the spirocycles HAl[(neol)Al(t-Bu)2]2 [261], ClAl[(neol) Al(t-Bu)2]2 [261], MeAl{[(l-O)CMe2CH]2AlMe2}2 [267] and (Me3Si)2(H) CAl{[(l-O)CH2]2[AlC(H)(SiMe3)2]2}2 [268]. Aluminium-bonded alkoxyenolate and alkoxyfuran moieties support the metal centres in MeAl[(l-O)C(H) Me2C(H)¼CMe(l-O)AlMe2]2 [269] and (Al{[(l-O)CH2-C4H7O]2AlCl2}2)+ AlCl4) [253] (see above), respectively. The latter aggregate has an Al4 analogue in the mixed methoxyfuran/ethoxide system {[(l-O)CH2-C4H7O]AlCl (l-OEt)[(l-O)CH2-C4H7O]AlCl2}2 (Fig. 23) [253]. Just as salen-type complexes have been noted in bis(aluminium) organooxide systems [251, 258], similar ligands have been seen to afford spirocyclic arrays in tris(aluminium) species. Hence, N,N¢-bis(2-hydroxybenzyl)-1,2-diaminoethane, -propane and -benzene will all react with AlMe3 to yield trimetallic spirocycles which fall into two categories; either mixed N,O-bridging of the metal centres, or else Al(l-O)2Al and Al(l-N)2Al fragments are seen [270]. This last motif is noted also in the product of reaction between N,N¢-bis(2-hydroxybenzyl)-1,2-diaminoethane and Al(i-Bu)3 [271]. Finally, the presence of aluminium-bonded halogens {already shown to be efficient bridging centres in [ArO(Me)Al(l-Cl)]2

127

The Synthesis and Structural Properties of Aluminium Oxide

Me O N O

O

Al

O N Al

O

O HO H2O OH2 (a)

OH2 OH2 O N

Al

HO

Al O

O

N

O

OH OH HO

HO

Al

OH N

HO

Al O HO

OH

O

OH

Al

Al

OH2

O

N

O

H2 O (b) Fig. 22. (a) The di- and (b) hexametallic components of a co-crystalline aluminium clathrate

which incorporates a spirocyclic {[RO2Al( Æ OH2)(l-OH)2]2Al(l-OH)}2 component in the solid state [265] Cl Et

O Al Cl

O O

O

O Al Cl Cl Al O

O

O O

Et

Al

Cl O

Cl Fig. 23. {[(l-O)CH2-C4H7O]AlCl(l-OEt)[(l-O)CH2-C4H7O]AlCl2}2 [253]

128

D. J. Linton Æ A. E. H. Wheatley

Cl

Cl Al Et

O O

Et Et Cl Al

Et

Al

O O

O O

Cl Al

Et Et Cl

Cl

Fig. 24. Spirocyclic Al[(l-OEt)2AlCl2]3 has a hexa-coordinate Al-core [274]

(Ar ¼ aryl) [129, 132]} has been shown to facilitate spirocycle formation in the tetranuclear aggregate [ArO(Me)Al(l-OAr)2Al(Me)(l-F)]2 [Ar ¼ 2,6-(i-Pr)2C6H3]. This species, generated by the reaction of Me2AlF with 2,6-diisopropylphenol, is based on the association of doubly oxygen-bridged ArO (Me)Al(l-OAr)2Al(Me)F dimers via a core Al(l-F)2Al ring which reveals two ˚] distinct types of aluminium-fluorine interaction [Al-F ¼ 1.812(2), 1.879(2) A [247]. Contrary to the established view that Al-F bond dissociation is excluded in the gas phase [272], the formation of this complex occurs by a fluorinearyloxide exchange process [247]. The ability of aluminium to exhibit coordination states as high as six leads to a variant on the spirocyclic theme discussed above. Trigonal arrangements containing four metal centres where one acts as a 6-coordinate core and is connected by heteroatom bridges to the remaining metals are reported for Al[(l-OR)2AlR¢2]3 R ¼ Et, R¢ ¼ Me [273]; R ¼ R¢ ¼ Et [273]; R ¼ Et, R¢ ¼ Cl [274] (Fig. 24) and have recently been both reviewed and christened [275]. Unlike the linear spirocycles noted for RAl{[l-O(i-Pr)]2AlR2}2 in which R ¼ H [205], Cl [47], trigonal Al[(l-OR)2Al(OR)2]3 is observed for R ¼ i-Pr [276]. Lastly, the alkoxythiofuran analogue of the ligand in both {[(l-O)CH2C4H7O]AlCl(l-OEt)[(l-O)CH2-C4H7O]AlCl2}2 and (Al{[(l-O)CH2-C4H7O]2 AlCl2}2)+ AlCl4)[253] affords Al{[(l-O)CH2-C4H7S]2AlMe2}3, in which the central aluminium centre is hexa-coordinated by oxygen in preference to sulfur [277]. Organooxide aggregates containing more than four aluminium centres form complex arrays of (AlO)n (n ¼ 2, 3) metallocycles. Hence, triple deprotonation of the tripodal triol N[(CH2)2OH]3 yields a tetra(aluminium) product (Fig. 25) [278]. At its core is a bis(aluminium) fragment which combines with one oxidemoiety from each of two {[N[(CH2)2O]3}3) ligands with each metal centre also stabilised by the remaining oxide centres and an N-based lone pair. Bridging the two organic trianions are two [Al(i-Bu)2]) fragments – the result being that the core Al-centres are bridged not just by two organooxide groups but also by two [O(R)]2Al fragments which each utilise one O-centre from either trianion. The observation that polyatomic, oxygen-rich groups can bridge between two Group 13 metal centres has already been discussed in the context of aluminium carbamates and carboxylates. However, a similar coordinative

129

The Synthesis and Structural Properties of Aluminium Oxide

Cl O

i-Bu

Al

N O

Al

i-Bu O O O

i-Bu

i-Bu

O

Al N

Al

Cl

Fig. 25. The Al4-product of N[(CH2)2OH]3 reaction [278]

mode to that revealed by (AlO)2[O2CN(i-Pr)2]8 [51] and {[(Me3Si)2C(H)]2Al}2(l-O2CPh)(l-H) [206] has been reported for the O-centred group in bis (carboxylate) complexes of the type (R2Al)2(l-O2CR¢)2 [R ¼ t-Bu, R¢ ¼ t-Bu, Ph, CH2Ph, CH2O(CH2)2OMe] [279] and leads to the adoption of 8-membered (CO2Al)2 metallocyclic cores by these systems. The bridging of metal centres in a bis(aluminium) complex by each of the carboxylate groups in four 2,2,6, 6-tetramethylpiperidinylcarboxylate ligands has produced as many contiguous 8-membered (CO2Al)2 rings each of which is folded about its Al    Al axis so as to share one CO2Al2 face with the next ring in a paddlewheel arrangement (Fig. 26). This complex was made serendipitously: the reaction of ClAl(TMP)2 with LiSnMe3 yielded amorphous material which analysed as MeAl(TMP)2. However, attempts to recrystallise this species resulted only in the carboxylation of Al-N bonds [44]. Just as mono(carboxylate) bridges between Al-centres have been reported in tandem with the spanning of other (nonoxygen) heteroatoms [206], so bis(carboxylate) bridging has also been noted in the tricationic component of the ethyl ethanoate-stabilised hydroxide {[(MeCO2Et)3 Æ Al]2(l-O2CPh)2(l-OH)}3+ 3(AlCl4)) [280]. Carbamates have also been shown to participate in double-bridging of two aluminium centres: the dimeric tris(carbamato)aluminium complexes {Me2Al[l-O2CN(i-Pr)2]}2 [281] and {[(i-Pr)2NCO2]2Al[l-O2CN(i-Pr)2]}2 [282] both incorporate 8-membered metallocyclic (CO2Al)2 cores, with the latter structure also revealing terminal carbamate ligands on each Group 13 metal centre. Both bridging and terminal bidentate ligands have been reported in the context of deprotonated b-diketone complexes of aluminium. As previously discussed, a recurrent feature of such complexes is the observation of isolated, hexacoordinate metal centres which are lipophilically wrapped by three [RC(O)C(H)C(O)R]) ligands [177, 179, 180, 181]. However, the reaction of Me3N Æ AlH3 with 1,1,1,5,5,5-hexafluoropentane-2,4-dione results in metallation and reduction of the b-diketone whereas the use of pentane-2,4-dione, 1,1,1-trifluoropentane-2,4-dione or 2,2,6,6-tetramethylheptane-3,5-dione yield only the tris(b-diketonato)aluminium(III) derivatives. The implication of these results is that either reduction precedes metallation with the CF3 groups activating the ketone functions even in the presence of excess metal hydride, or else that they activate the metallated species towards nucleophilic attack (as

130

D. J. Linton Æ A. E. H. Wheatley

TMP

TMP

O Me

O

Al O

TMP

O

O Al

O O

Me O TMP

Fig. 26. Serendipitously synthesised MeAl[O2C(TMP)]2 [44]

has been reported elsewhere [283]). The bis(aluminium) compound which results from the presence of 1,1,1,5,5,5-hexafluoropentane-2,4-dione incorporates two dianions which form chelate rings but which also bridge between metals through the oxygen centre next to the chiral carbon [mean ˚ , mean Al-O ¼ 1.75 A ˚ ] while a third ligand donates one Al-(l-O) ¼ 1.91 A ˚ ). The chirality O-centre (terminally) to each metal atom (mean Al-O ¼ 1.76 A of this latter ligand is reversed relative to that noted for the first two and as a consequence an enantiomeric mixture of RRS/SSR isomers is noted [284]. The higher-order tris(aluminium) tetraanionic complex [H2O Æ Al3(l-OH) (l-OR)3]4){R ¼ C[(CH2)CO2]2CO2} (the charge on which is countered by four ammonium cations) reveals extensive coordination of the metal ions by the deprotonated carboxylate oxygen centres, with three distinct citrate ligands being observed in the solid state: one straddles all three metal ions while the remainder each bridge two such centres (one organic residue employing its central and one terminal carboxylate and the other using only terminal carboxylate groups) [101]. Akin to the arrangements of asymmetric, edge-fused arrays of aluminium oxide rings discussed in Section 2.2 [64, 65], the redetermined solid-state structure of the dimeric alkoxide {Me2AlO(CH2)2N[(CH2)2O]2AlMe}2 reveals a contiguous arrangement of one 4- and four 6-membered metallocycles (Fig. 27) [285]. This complex was prepared by the 1 : 1 treatment of AlMe3 with the triol N[(CH2)2OH]3 and, when it was originally characterised [286], represented the first hexa-coordinate alkylaluminium. The employment of polydentate Schiff bases as chelating agents for aluminium centres (as discussed above [109]) has been extended to the heptadentate product (Fig. 28) of reaction between salicylaldehyde and bis (2-aminoethyl)ethylenediamine [287]. While the reaction of this Schiff base with two equivalents of AlMe3 afforded a system which NMR spectroscopy reveals to contain only two penta-coordinate aluminium centres [38, 39], the comparable 1 : 3 reaction affords a product with three distinct aluminium environments. Whereas a single AlMe2+ moiety is bonded formally to two deprotonated O-centres (on the middle and end aryloxy groups) and thereafter stabilised by two N-centres, the remaining two AlMe2+ units show differing coordination enviroments: one is O,N,N¢-chelated whereas the other

131

The Synthesis and Structural Properties of Aluminium Oxide

Me Me Al O

O N Al Me

Me

O

Al N

O O

O Al Me

Me

Fig. 27. The dimeric alkoxide {Me2AlO(CH2)2N[(CH2)2O]2AlMe}2 [285]

is O,O-stabilised and, along with the methylaluminium unit, participates in the formation of an (AlO)2 metallocycle [287].

4 Concluding Remarks Aluminium oxides, hydroxides and organooxides (both aryloxide and alkoxide) have been extensively studied both in their own right and in order to enhance their value to the preparative chemist. In this sense, significant progress has been made towards rationalising and categorising the structural properties of such homometallic systems. For aluminium oxides, the simplest motif to be noted is that of the oxide mono-bridge (R2Al)2O. However, much more common, not only for oxides but also for hydroxides and organooxides, is the formation of (AlX)n metallocycles in which n ¼ 2, 3 in di- and trinuclear complexes (X ¼ a heteroatom which is often oxygen but which can be some other electron-rich centre in the case of non-trivial organooxide ligands). While it is extremely common for such rings as these to be edge-fused in higher-order aggregates, association via a single metal centre also occurs, to yield spirocyclic motifs of two main types. Thus, linear spirocycles have been reported for certain trimetallic aluminium species [210, 265] and, in one case, the formation of multiple Al(l-H)2Al moieties has afforded a polymer [205]. More complex are the trigonal arrangements of four metal centres, one of

N Cl

N

N

N HO

OHHO Cl

Fig. 28.

Cl

132

D. J. Linton Æ A. E. H. Wheatley

which acts as a hexavalent core and is connected by heteroatom bridges to the remaining metals, reported for Al[(l-OR)2AlR¢2]3 (R ¼ alkyl, R¢ ¼ alkyl, halogen) [275]. A futher facet of research has involved the structural characterisation of aluminium complexes which incorporate polydentate salen-type ligands. These have been noted in both neutral and monocationic (ion-separated) contexts (the latter requiring that the metal centre be stabilised by an external Lewis base) [35]. While such charged systems are invariably mononuclear the same is only usually true of their neutral analogues by virtue of the sterically demanding bis(aryloxide), chelating ligand. In the context of these latter complexes, dimerisation has been noted [251] while, more recently, the employment of flexible alkyl chains between two salen-coordinated aluminium ions has enabled the observation of dinuclear compounds [160, 161]. The synthetic utility of the complexes discussed here has been shown to be of great significance – by virtue both of their stoichiometric [75–83] and of their catalytic [25, 108–111] properties. In this context, improvements in our understanding of the steric and electronic influences on these compounds will enable the future development of new materials, preparations, and synthetic methodologies.

5 References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24.

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The Synthesis and Structural Properties of Aluminium Oxide

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Insertion and b-Hydrogen Transfer at Aluminium Peter H. M. Budzelaar1, Giovanni Talarico2 1

2

Department of Inorganic Chemistry, University of Nijmegen, Toernooiveld 1, 6525 ED Nijmegen, The Netherlands e-mail: [email protected] Dipartimento di Chimica, Universita` degli Studi di Napoli ‘‘Federico II’’, Complesso Monte S. Angelo, Via Cintia, 80126 Napoli, Italy e-mail: [email protected]

Insertion and elimination of C@X bonds are two fundamental steps of organometallic chemistry. In aluminium chemistry, these reaction steps are important in olefin oligomerization and possibly polymerization, as well as in the reaction of carbonyl compounds with aluminium alkyls. The general importance of a third fundamental step, direct b-hydrogen transfer from an aluminium-bound group to a substrate, has not been recognized as widely, despite the fact that it is the key step in the reduction of ketones by alcohols (Meerwein-Pondorf-Verley reduction) and by aluminium alkyls bearing b-hydrogens. In this review we will combine experimental and theoretical results to illustrate how the delicate balance between these three reaction types determines much of the chemistry of organoaluminium compounds. Such an understanding may create new opportunities for ligand design.

1

Introduction: Insertion, b-Elimination and b-Hydrogen Transfer . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 142

2

Aluminium Alkyls and Olefins: Oligomerization and Polymerization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 144

2.1 2.2 2.3 2.4

Oligomerization – the Aufbau Reaction . . . . . . Higher Olefins, Longer and Branched Chains . . Polymerization at Aluminium? . . . . . . . . . . . . . The Chemistry of ‘‘Modern’’ Aluminium Alkyls

3

Aluminium Alkyls and Ketones: Insertion and Reduction . . . . . 154

4

Aluminium Alkoxides and Ketones: Meerwein-Pondorf-Verley Reduction . . . . . . . . . . . . . . . . . . . . . 158

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How Special Is Aluminium? . . . . . . . . . . . . . . . . . . . . . . . . . . . 162

6

Outlook . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163

7

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163

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1 Introduction: Insertion, b-Elimination and b-Hydrogen Transfer Aluminium is used in a number of important chemical transformations. Perhaps the best-known example is the Aufbau reaction, discovered in 1950 by Ziegler [1] and still in use today:

ð1Þ

Under the reaction conditions employed, the oligomerization is living; the aluminium is removed from the chain in a separate process step [2]. Further research on the Aufbau reaction led to the discovery of transition metal catalyzed olefin polymerization (Ziegler-Natta catalysis) [3]. The Aufbau reaction itself does not lead to high-molecular-weight polymers [2]. Under process conditions, the maximum degree of polymerization attainable seems to be ca. 100. Attempts to make much longer chains lead to chain transfer, i.e., loss of the ‘‘living’’ character. If propene or higher olefins are used instead of ethene, chain transfer becomes so easy that the main reaction is dimerization. The chain transfer step of the Aufbau process is generally assumed to be b-elimination, followed by reinsertion of a new molecule of monomer. Over the last years it has become clear that in transition-metal-catalyzed polymerization the main chain transfer mechanism is frequently not b-elimination but direct chain transfer to monomer [4], and recent theoretical results indicate that – depending on the system and reaction conditions – the same may be true for aluminium [5–8].

ð2Þ

Since both insertion and b-hydrogen transfer to monomer are first order in alkyl and olefin, the competition between the two does not depend strongly on the reaction conditions and hence sets an upper limit for the attainable molecular weight.

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The oligomerization Aufbau reaction has been known for a long time. Over the past decade, several papers have appeared which claim high-molecularweight polymerization at aluminium [9–13]. These reports have led to a revival of interest in organo-aluminium chemistry, in particular in cationic and coordinatively unsaturated aluminium complexes. However, the nature of the active species in the reported polymerization reactions is not very clear [14, 15]. According to both calculation and experiment, b-hydrogen transfer is even more relevant as a chain transfer mechanism for these new systems [8, 15]. The competition between insertion and hydrogen transfer is also crucial to the selectivity of the reaction of aluminium alkyls with carbonyl compounds. Aluminium alkyls, like organolithium compounds and Grignard reagents, can add to aldehydes and ketones to form secondary or tertiary alcohols, respectively. If the aluminium alkyl has a b-hydrogen, however, reduction of the carbonyl compound is a common side reaction, and can even become the main reaction [16]. Most authors seem to accept that reduction involves direct b-hydrogen transfer to ketone.

ð3Þ

Finally, b-hydrogen transfer is the key step in the Meerwein-Pondorf-Verley (MPV) reduction of ketones by alcohols, catalyzed by aluminium alkoxides and many other catalysts. In that case, competition is not an issue, since polymerization is usually not thermodynamically favourable. The accepted mechanism for this reaction is direct transfer of the hydride from alkoxide to ketone.

ð4Þ

In the present review we will concentrate on the fundamental steps of insertion and b-hydrogen transfer for these three related reactions, with an emphasis on the area of polymerization. We will try to reconcile reported experimental data with the more microscopic picture afforded by calculations, some of which have not yet been published [17]. It will be seen that this kind of perspective allows one to draw analogies between these reactions that might not be so clear otherwise. We have made no attempt to cover all of the vast experimental literature in this area but just mention representative examples. More complete coverage can be found in several review articles [2, 18–22].

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2 Aluminium Alkyls and Olefins: Oligomerization and Polymerization 2.1 Oligomerization – the Aufbau Reaction

In the Aufbau reaction [2], trialkylaluminium reacts with ethylene to give longer-chain aluminium n-alkyls. The reaction is typically carried out at 100– 120 C and 50–100 bar of ethylene pressure. Under these circumstances, the oligomerization (taken to chain lengths of ca. 10–20 units) is living, i.e., chain transfer is slow compared to the reaction time. The aluminium alkyls can then be converted into olefins by heating to 280–320 C at low ethylene pressure (a few bar), or oxidized to alkoxides and then hydrolyzed to linear alcohols. The advantage of a ‘‘living’’ oligomerization is that the resulting Poisson distribution of chain lengths is relatively narrow. Aluminium tri-n-alkyls are dimeric in solution, although – especially for higher alkyls – the fraction of monomer can become significant at higher temperatures [19]; in the gas phase they are usually monomeric. Kinetic evidence indicates that olefin insertion involves a monomeric aluminium trialkyl; this suggests a Cossee-type insertion mechanism. Kinetic data do not indicate the presence of an intermediate olefin p-complex [23]. However, if the olefin complexation energy at the p-complex stage is low, this would be expected.

ð5Þ

The monomer-dimer equilibrium, and hence exchange of alkyl groups between aluminium centres, is much faster than insertion, so all chains grow at the same rate. Trimethylaluminium forms an exception to this rule: insertion into an AlAMe bond is significantly slower than in higher Al-n-alkyl bonds, so that reaction of Me3Al with ethene produces a mixture containing many unreacted methyl groups in addition to longer-chain alkyls. The activation energies for ethene insertion in Me3Al and Et3Al have been determined as 22.5 and 17.6 kcal/mol, respectively (relative to the R3Al monomer) [24, 25], although Egger also states that he considers the difference between these values to be too large [23]. These barriers are much larger than those typically found for transition-metal-catalyzed polymerization. This explains why traces of transition metal impurities are so problematic in this kind of study and can completely skew the results. Calculations confirm the Cossee mechanism for olefin insertion [7, 26–28]. For the simple Me2AlEt model, the ethene complexation energy is only a few kcal/mol. The activation energy calculated at the highest theoretical level [7] agrees well with the experimental estimate. An advantage of theoretical methods is that they can be used to predict the geometries of transition states, which are not directly accessible from

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Fig. 1. Calculated structures for Me2AlEt, Me2Al(Et)(CH2@CH2), the insertion transition state, and the insertion transition state for H2SiCp2Zr(Et)(CH2@CH2)+ [29]

experiment. Calculated structures for relevant species on the insertion path are shown in Fig. 1. The insertion transition state (TS) has the expected 4-membered ring structure [7, 26], with nearly equal Al-C distances for the bonds being formed and broken. Compared to a typical transition-metal insertion (H2SiCp2ZrEt+ + C2H4), the ethene is much more deformed and the transition state is much later and tighter, which is probably related to the higher barrier of the reaction. Also, there is no strong a-agostic interaction in the aluminium case. There are two reasonable paths for chain termination. Under the conditions used to liberate the olefin in the Aufbau process (high temperature, low ethene pressure), the main termination reaction is b-elimination [2].

ð6Þ

The elimination step itself is the reverse of Cossee-type insertion into an AlAH bond. This insertion is much easier than in the AlAC bond, presumably because of the lack of directionality of the hydride 1s orbital involved in the reaction (compared to the sp3 hybrid of an alkyl group). Nevertheless, b-elimination has a rather high activation energy because the initial product, a terminal aluminium hydride, is very unfavourable. This initial product can

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either form a much more stable dimer, or – as in the Aufbau process – be trapped by ethene to form a new Al-Et moiety. Elimination from aluminium n-alkyls has an activation energy of ca. 33 kcal/mol [2]. Calculations produce an activation energy of ca 35 kcal/mol for b-elimination from Me2AlEt [7]. Calculated geometries for structures on the elimination path (Fig. 2) clearly show the similarity with the insertion step described above. The transition state is somewhat closer to the p-complex, as expected from the stabilities of the species involved. At lower temperatures (or in solution) and at high monomer concentration, a second chain termination process that could occur is direct b-hydrogen transfer to a second molecule of monomer. This kind of chain transfer step is now generally accepted for many transition-metal-catalyzed polymerizations, where direct b-elimination would be too much uphill to explain the observed molecular weights. For olefin oligomerization at aluminium, a similar situation applies. Since insertion and b-hydrogen transfer have an identical concentration dependence, their ratio does not depend much on the reaction conditions (except temperature) and hence limits the molecular weight attainable in the Aufbau reaction.

ð7Þ

Unfortunately, there is no direct evidence for this kind of chain transfer step, since all kinetic studies have been carried out under conditions where b-elimination would dominate. However, the statement by Wilke [30] that the Aufbau process produces chains of at most »100 units agrees well with the predicted average degree of polymerization of »70 units [7]. The calculated transition state for b-hydrogen transfer (Fig. 3) has a nonplanar 6-membered ring structure. There is no direct interaction between the aluminium atom and the hydrogen being transferred. Thus, the situation differs sharply from that for transition metal polymerization, where the transition state has some resemblance to a hydride-bis(olefin) complex, as illustrated for a typical metallocene case in Fig. 3.

Fig. 2. Calculated b-elimination transition state for Me2AlEt, and elimination product Me2Al(H)(CH2@CH2)

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Fig. 3. Calculated b-hydrogen-transfer transition states for Me2Al(Et)(CH2@CH2) and

H2SiCp2Zr(Et)(CH2@CH2)+ [29]

In fact, the transition state for b-hydrogen transfer shows more resemblance to the typically organic [1, 5] sigmatropic hydrogen migration [31, 32]. 2.2 Higher Olefins, Longer and Branched Chains

Substitution, both at the alkyl and the olefin, has clear and easily explainable effects on both insertion and chain transfer reactions. These effects are best understood in terms of the (de)stabilization of olefin and metal alkyl by substitution. Olefins are stabilized by alkyl substituents at the double bond. Every methyl group brings an additional stabilization of »4 kcal/mol [33], as illustrated by the isodesmic equation: ð8Þ Since transition states for insertion and chain transfer are all more or less central, approximately half of the substituent effect will be felt at the transition state. Thus, insertion of propene (a more stabilized olefin) has a higher barrier (by about 2 kcal/mol) than insertion of ethene: stabilization is partly lost at the transition state (Fig. 4). Similarly, b-elimination from a propyl chain is easier than from an ethyl chain, and propyl-to-ethene hydrogen transfer is easier than ethyl-to-ethene transfer (but propyl-to-propene transfer is not). This explains why higher olefins are only dimerized at aluminium: insertion becomes more difficult, while elimination (from a b-branched alkyl, leading to a strongly stabilized olefin) becomes much easier. Aluminium alkyls, like all alkyls of electropositive elements, are destabilized by substitution at the a-carbon. A destabilization of »2.5 kcal/mol [33] can be estimated from the isodesmic equation: ð9Þ

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Fig. 4. Effect of olefin stabilization on ease of insertion (exaggerated)

Again, approximately half of this effect will already be felt at the transition states for insertion and chain transfer. Thus, internal olefins are inserted less easily, a-branched alkyls are more reactive towards insertion, and secondary alkyls tend to isomerize to primary alkyls. The abnormally low reactivity of AlAMe bonds towards olefin insertion must probably be explained in this way. The high reactivity of (t-Bu)3Al towards ethene insertion, despite its significant steric hindrance, might in part be due to the same effect. 2.3 Polymerization at Aluminium?

In transition metal chemistry, ligand variation has proven to be the key to obtaining highly active polymerization catalysts. In particular, sterically hindered monocationic alkyl complexes with an empty site seem to be well suited for polymerization. The steric bulk prevents (associative) b-hydrogen transfer, while the positive charge destabilizes the free hydride and thus opposes (dissociative) b-elimination. Aluminium is much cheaper than transition metals, and aluminium oxide is non-toxic. Aluminium residues in a polymer would probably not be harmful. Thus, a catalyst based on aluminium could be extremely attractive, even if it were significantly less active than a transition metal catalyst. This has probably contributed to the continued interest in (potential) aluminium polymerization catalysts. However, such studies are difficult, as even traces of transition metal contamination may lead to erroneous conclusions. According to calculations, insertion barriers at aluminium are typically >10 kcal/mol higher than at transition metal catalysts, corresponding to a reactivity difference of 109, so

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that ppb concentrations of transition metals may significantly affect the results. Several groups have reported olefin polymerization at aluminium, but none of the systems reported to date are completely understood. Martin and Bretinger studied the oligo/polymerization of ethene with Et3Al at room temperature [9], and found a very small amount of high-MW polymer. The ‘‘active species’’ was proposed to be a monomeric aluminium trialkyl present in low concentration in the mixture. The main problem with this interpretation is that any monomeric aluminium trialkyl is in rapid equilibrium with the dimer, which would result in alkyl group exchange and uniform growth of all chains. The polymer they observed must therefore have been formed on a species that does not rapidly exchange alkyl groups with trialkylaluminium. Whether this species is a more complicated (possibly oligomeric or heterogenized) aluminium alkyl or a transition-metal impurity remains to be seen. Gibson described the synthesis of four-coordinate cationic aluminium alkyls 1 which were reported to be well-defined aluminium polymerization catalysts [12]. However, the polydispersities of the products obtained were high (2.9–6.3), showing that there is not a single well-defined active species. The experiments were carried out in metal autoclaves, and Fe and Co complexes of pyridine-diimine ligands are extremely active in ethene polymerization [34], so a transition-metal impurity does not seem an unreasonable explanation.

Sen reported that (C6F5)2AlR (2) (generated in situ) is an ethene polymerization catalyst (precursor) [13]. Moreover, the system also catalyzes copolymerization of ethene and propene. This latter fact, in particular, is remarkable, since a b-branched alkyl (formed after propene insertion) should undergo very easy b-elimination and hydrogen transfer to ethene, as discussed above. Thus, for this system to work the intrinsic chain transfer barriers of (C6F5)2AlR should be much higher than those of trialkylaluminium.

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Jordan reported that cationic aluminium alkyl derivatives of amidinates (3) [10] and aminotroponiminates (4) [11] are active in ethene polymerization. The amidinate system was later shown to be more complicated, and the nature of the active species remains unclear [14].

For the aminotroponiminate (ATI) system, the active species has been shown not to be the monomeric cationic alkyl proposed originally. Indeed, labeling experiments indicate that (ATI)AlEt+ and (ATI)Al(i-Bu)+ primarily undergo chain transfer to monomer, whereas another, as yet unknown, species

ð10Þ

in the system is responsible for the observed polymerization activity [15]. Jordan states that he has taken very precaution to exclude transition metal impurities, but also admits that, as long as no explanation of the results is available, no possibility can be excluded. Several computational studies have been carried out in attempts to understand these systems [5, 7, 8, 28]. Calculated activation energies for insertion are comparable to, or somewhat lower than, those for aluminium trialkyls, and are compatible with the reported (low) activities of the catalysts. However, the calculated barriers for chain transfer to monomer are consistently lower than those for aluminium trialkyls, to the extent that the systems reported should dimerize or possibly oligomerize, but not polymerize. Substituent effects have been investigated in some detail. Electronegative substituents, a positive charge at the metal and small ligand bite angles all favour b-hydrogen transfer over insertion. Ligand bulk is not efficient in suppressing b-hydrogen transfer because, unlike in the transition metal series, the b-hydrogen transfer transition state does not require much more space around the metal than the insertion transition state. Thus, the abovementioned ‘‘real systems’’ are all predicted to have significantly worse insertion/chain transfer balances than trialkylaluminium [8]. The amidinate alkyl cations form dimers in solution (see next section), so the possibility of dinuclear active species has also been considered. On the basis of Car-

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Parrinello studies, Meier suggested the possibility that in dimeric complexes dynamic effects might favour insertion (at a single metal centre, e.g., in 5) over chain transfer, but no free-energy data were provided to support this hypothesis [35]. Insertion over two metal centres (as in 6), as an alternative to the standard Cossee mechanism, was shown to have a comparable barrier to the standard mechanism [36, 37]. However, this alternative mechanism does not have an improved propagation/chain transfer balance and hence does not offer a better explanation for the observed polymerization activity. It appears

to be very difficult to improve on the ancient Aufbau reaction by ligand variation. Of course, it is dangerous to exclude the possibility of aluminium polymerization on the basis of calculations. Reality is invariably more complicated than the simplified models put into computers. However, in view of the uncertainties surrounding existing systems, and the doubts thrown by calculations, any well-defined aluminium alkyl claimed to be active should at least be checked, as an isolated complex, for its propensity to olefin insertion vs. chain transfer, e.g,. using the Al-i-butyl/ethene experiment reported by Jordan [15], as explained above. 2.4 The Chemistry of ‘‘Modern’’ Aluminium Alkyls

The initial reports on Al-catalyzed olefin polymerization have led to a revival of organoaluminium chemistry. Not surprisingly, electron-deficient, lowcoordinate, cationic aluminium alkyls in particular received a lot of attention. While this has not led to efficient olefin polymerization, it did provide more insight into the structure and reactivity of such unsaturated and strongly Lewis-acidic systems. Ligands studied include amidinates [10, 14, 38–40], aminotroponiminates [11, 15, 41], b-diiminates [42, 43], guanidinates [44], thioamidinates and thioureides [38], bis(iminophosphoranyl)methanide [45], etc. The strategy followed for aluminium is similar to that used for early transition metals [46]: synthesis of a neutral metal dialkyl, followed by alkyl abstraction or protonation in the presence of a poorly coordinating anion, using reagents like B(C6F5)3, PhMe2NH+B(C6F5)4) (‘‘DANFAB’’) or Ph3C+B

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(C6F5)4) [47]. The problems encountered are also similar: the cationic and unsaturated monoalkyl is extremely reactive and tries to achieve a higher coordination number in various ways, as illustrated by the examples below. NH

N

R3 Al

-

AlR2 N

N 1) RLi 2) AlCl3

2 RLi

-R

N Al

R

N

ð11Þ

N AlCl2 N

Normally, amidinates prefer chelate structures in which both nitrogens of the ligand coordinate to the same metal centre (A). Bridging structures in which the nitrogens coordinate to two different metal atoms (B) are less common; this bridging mode might be easiest to achieve if the substituent at the amidinate carbon is small [48]. Perhaps more surprisingly, a single amidinate nitrogen can also bridge between two different metal centres; this is accompanied by localization of the amidinate C@N bond and leads to structures like C and D [14]. Such a switch from 4-e to 6-e donor ligand could also be envisaged for, e.g., aminotroponiminates and b-diiminates, and the possibility should be kept in mind when designing new ligands.

ð12Þ

Yet another type of bridge observed for such complexes is an unusual type of electron-deficient alkyl bridge. Trialkylaluminium compounds form alkylbridged dimers in which each of the two two-electron-three-centre bonds is formed from two aluminium sp3 hybrids and one alkyl sp3 hybrid (E, next page), leading to acute Al-C-Al angles (e.g., 78 in Me6Al2 [49]). This kind of bridge requires a close approach of the two metal centres and is therefore impossible in complexes bearing bulky ligands like amidinates or aminotroponiminates. For such complexes, one finds nearly linear alkyl bridges in which a planar alkyl group uses its pz orbital to form a bridge between the two metal centres (F) [10, 11, 14, 15]. Ligands are not always innocent spectator groups. Cationic b-diiminate alkyl complexes of aluminium, rather than undergoing insertion in the Al-alkyl bond, show attack of the unsaturated substrate (ethene or acetylene) at the ligand (Scheme 1). Surprisingly, the C-C coupling reaction observed with ethene is reversible [42].

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ð13Þ

Anions, too, are not always innocent. The borate anion B(C6F5)4) is very popular in early-transition-metal catalyzed polymerization, where it acts as a rather inert and non-coordinating anion. Examples of decomposition of the anion by, e.g., C6F5 transfer exist but are not very common. In aluminium chemistry, transfer of C6F5 groups from B(C6F5)3, MeB(C6F5)3) and B(C6F5)4) to the metal appears to be rather easy [14, 15], and it may be that other, even more ‘‘innocent’’ anions will be required here.

Scheme 1. Coupling of b-diiminate ligand with ethene and 2-butyne [42]

The driving force for aluminium to attain a coordination number higher than 3 appears to be remarkably high, in particular in cationic complexes. If ligand and alkyl bridging are prevented by ligand constraints, and anion decomposition can be avoided, the use of the popular DANFAB to introduce the borate anion frequently leads to strongly bound dimethylaniline complexes. If no additional donor is used [e.g., through use of the alternative Ph3C+B(C6F5)4) salt] coordination of solvent (chlorobenzene) or C6F5 groups can be observed [43]. Clearly, creating an empty site at cationic aluminium is just as difficult as at early transition metals, and as a result the chemistry of cationic alkyls can be amazingly complex. Easy and well-defined olefin insertion has not yet been observed for these ‘‘new’’ cationic alkyls. However, with terminal alkynes – which are much more

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reactive towards insertion – easy and selective dimerization has been observed, initiated by (ATI)AlEt+ (Scheme 2) [15]. Note that the initiation in this reaction involves b-hydrogen transfer (from alkyl to alkyne), but the ‘‘chain transfer step’’ does not.

Scheme 2. Dimerization of t-butylacetylene by (ATI)AlEt+ [15]

3 Aluminium Alkyls and Ketones: Insertion and Reduction Since aluminium alkyls are easily available from inexpensive starting materials, their addition to carbonyl compounds could constitute a cheap alternative to the more expensive Grignard reagents or alkyllithium compounds. Unfortunately, this addition is only clean for aluminium alkyls without b-hydrogens [16, 19], which are precisely those not accessible via the convenient

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hydralumination route. Alkyls bearing b-hydrogens invariably have reduction of the ketone as a competing reaction; other side reactions include the Tishchenko reaction, enolization and aldol condensation. Even if addition does occur, it may be followed by an undesired elimination of aluminium alkoxide to give an olefin. These complications are illustrated in Scheme 3.

Scheme 3. Reactions of ketones and aldehydes with trialkylaluminium

The mechanism of ketone addition has been studied in detail for the reaction of Me3Al with benzophenone [50]. With a 1 : 1 ratio, a coordination complex Me3Al Æ Ph2CO is formed first, and this subsequently rearranges to the adduct, presumably via a 4-centered transition state. With an excess of Me3Al, a path involving two Me3Al units seems to be preferred, and a 6-centered transition state (similar to structure 6 studied theoretically for dinuclear olefin insertion [36, 37]) has been proposed for this alternative path. Studies using more complicated systems also suggest that, if a 6-centered transition state is available, this is preferred over a 4-centered transition state [51].

ð14Þ

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This preference for a 6-centered transition state might be responsible for the significant amount of reduction found with b-hydrogen-bearing aluminium alkyls. A few authors assume that reduction involves prior b-elimination, followed by carbonyl insertion in the AlAH bond. However, reduction is frequently observed even with Et3Al, for which the rate of elimination at room temperature is negligible. Thus, direct b-hydrogen transfer from alkyl to ketone appears to be a more reasonable explanation. The large oxophilicity of aluminium, which results in strong polarization of the C@O bond upon coordination to Al, might be a second factor promoting b-hydrogen transfer. With b-branched aluminium alkyls like Al(i-Bu)3, reduction is often the main reaction [52]. Ligand modification can be used to increase the amount of reduction still further, and also to control the diastereoselectivity. In this respect, the phenoxide-modified complex 7 appears to be particularly effective [53]. A recurring problem in diastereoselective reductions is that the product can epimerize through MPV reduction (see next section) of the starting material [54]. The kinetics of this complicated system have been analyzed in terms of the iso-inversion principle [55].

Unlike olefin insertion, the reaction of aluminium alkyls with carbonyl compounds has not been studied theoretically before. The calculated barriers for addition and b-hydrogen transfer in the system Me2AlEt + CH2@O are very similar (15.4 and 14.3 kcal/mol, respectively; see Table 1), in accord with the close competition between the two reaction types observed experimentally. Calculated transition-state geometries (Fig. 5) show a clear similarity with the alkyl + olefin system described in the previous section. Transition states are somewhat earlier, since these reactions are much more exothermic: this can be seen, e.g., from the noticeably different C  H bond lengths of the hydrogen transfer transition state. Table 1. Calculated insertion and b-hydrogen-transfer barriers for simple model systemsa

Fragment

CH2CH3/CH2@CH2

CH2CH3/O@CH2

OCH3/O@CH2

Insertion b-H transfer Insertion b-H transfer Insertion b-H transfer Me2Al (Me)(NH3)Mg Me2Ga

22.7 20.1 27.8

28.3 26.0 30.2

15.4 5.8 20.9

14.3 9.5 20.3

8.7 2.7 7.7

13.1 9.4 14.4

a CCSD(T)/6-31G(d)//B3LYP/6-31G(d), basis set extrapolation to 6-311G(d,p) at MP2 level, ZPE correction at B3LYP/6-31G(d) level, no thermal corrections.

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Fig. 5. Calculated structure of Me2Al(Et)(CH2@O) and its transition states for insertion and

b-hydrogen transfer

As discussed earlier, it appears to be very difficult to get a better balance between insertion and b-hydrogen transfer than that obtained with trialkylaluminium. In view of the similarities noted here, the same might well be true for the reaction with carbonyl compounds. Thus, the ‘‘complications’’ associated with b-hydrogen-containing aluminium alkyls should probably be accepted as a fact of life, that can only be circumvented by tricks like transitionmetal catalysis. However, this is not always a disadvantage. The previous section also mentioned that electron-withdrawing groups on the ligand and a positive charge on the metal promote b-hydrogen transfer. Thus, it should be possible to shift the selectivity of carbonyl reactions completely towards reduction even with alkyl groups (like Et) for which one would normally expect a significant amount of insertion. Ligand variation could then be used to tune diastereoselectivity or even enantioselectivity. The observation by Jordan that (ATI)AlEt+ cleanly reduces acetone to (ATI)Al(O-i-Pr)+ is promising in this respect [15].

ð15Þ

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4 Aluminium Alkoxides and Ketones: Meerwein-Pondorf-Verley Reduction Meerwein-Pondorf-Verley reduction, discovered in the 1920s, is the transfer hydrogenation of carbonyl compounds by alcohols, catalyzed by basic metal compounds (e.g., alkoxides) [56–58]. The same reaction viewed as oxidation of alcohols [59] is called Oppenauer oxidation. Suitable catalysts include homogeneous as well as heterogeneous systems, containing a wide variety of metals like Li, Mg, Ca, Al, Ti, Zr and lanthanides. The subject has been reviewed recently [22]. In this review we will concentrate on homogeneous catalysis by aluminium. Most aluminium alkoxides will catalyze MPV reduction.

ð16Þ

The efficiency of the reduction is determined in part by thermodynamic considerations: secondary alcohols are stronger reductants than primary ones, and aldehydes are more easily reduced than ketones. Thus, i-propanol is the most commonly used reductant. The reaction is an equilibrium that is usually driven to the right by using a large excess of reductant. Enantioselective reduction is possible, but requires stopping the reaction before equilibrium has been established, since obviously the system is also an epimerization catalyst. Enantioselective reduction using a chiral alcohol as reductant has been reported, but the enantioselectivities obtained are generally disappointing [60, 61]. Chiral diols have been used as ancillary ligands for enantioselective aluminium-catalyzed MPV reduction using achiral alcohols; again, chiral induction observed was rather modest [62]. Much better enantioselectivities have been reported in lanthanide-catalyzed MPV using similar ancillary ligands [62, 63]. Intramolecular MPV reductions, however, can have a very high diastereoselectivity [64]. In that case, the reversibility of the reaction can actually be an advantage, since the energy difference between the diastereomers can drive the reaction to the desired isomer [65]. These observations have led to strategies involving tandem reactions, where the reductant alcohol is attached to the carbonyl compound prior to the actual reduction step [66, 67]. Electronegative substituents on aluminium like trifluoroacetate [68] or phenolate [69] increase the efficiency of the MPV reduction. Maruoka reported the use of the didentate catalyst 8 for ‘‘double electrophilic activation’’ of carbonyl compounds [70], but since no comparison with monofunctional phenolates was given it is not clear whether having two aluminium centres in the same catalyst offers any special advantages. They used this catalyst to effect transfer hydrogenation between remote aldehyde and alcohol groups in the same molecule [71], but again it is not clear whether the transfer is truly intramolecular or in any way different from that of reduction by an external alcohol using 8 or a monuclear aluminium catalyst.

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In principle, one could write two mechanisms for the MPV reduction: a stepwise mechanism involving a discrete hydride intermediate, and a direct hydride transfer from alkoxide to ketone. These alternatives are similar to the two mechanisms usually assumed for transition-metal-catalyzed transfer hydrogenation.

ð17Þ

The b-elimination mechanism might apply to certain late-transition-metal systems. In the particular case of aluminium, however, hydride addition to the ketone is so exothermic (i.e., hydride elimination is so much uphill) that formation of a discrete hydride intermediate is highly unlikely. Also, the discrete hydride mechanism would not explain chiral induction through the use of chiral alcohols. Thus, direct hydrogen transfer appears to be the only realistic path and is the generally accepted mechanism [22, 69]. Ashby argued that a single-electron-transfer mechanism cannot be excluded for aromatic ketones, but for aliphatic ketones there are no indications for the involvement of radicals [72]. Interestingly, it is only recently that the first example of an aluminium aldehyde-alkoxide complex, one of the presumed intermediates in the MPV reduction, has been isolated and characterized [69]. In the complex (9) the aluminium is five-coordinate, being surrounded by two phenolate oxygens, two bridging alkoxide oxygens and an oxygen of a coordinated aldehyde. Based on this structure, the authors proposed that MPV reduction occurs – at least in their system – at a five-coordinate aluminium centre. However, they also report examples of monomeric, four-coordinate alkoxide complexes containing a triphenylphosphine oxide or HPMA ligand (which could be viewed as a model for a ketone or aldehyde), so the matter is probably not settled. Also, it might well be that different coordination numbers are involved for different systems.

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Unlike the reactions described in the previous two sections, competition between insertion and b-hydrogen transfer is usually not an issue here. Ketone polymerization is nearly thermoneutral and disfavoured by entropy. However, aldehyde insertion is thermodynamically more favourable, and the Tishchenko reaction mentioned in the previous section can plausibly be written as a sequence of insertions and b-hydrogen transfer reactions (Scheme 4).

Scheme 4. Tishchenko reaction written in terms of insertion and b-hydrogen transfer steps

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Fig. 6. Calculated structure of Me2Al(OMe)(CH2@O) and its transition states for insertion

and b-hydrogen transfer

MPV reduction has not been studied theoretically in recent years. Our calculated barrier of 13 kcal/mol is quite similar to the one predicted for b-hydrogen transfer from alkyl to carbonyl (Table 1, Section 3). Also, the relevant geometries (Fig. 6) show a pronounced similarity, although now the 6-membered transition state is symmetric, as expected for this degenerate reaction. It should be noted that for this particular case the assumption of a 4-coordinate aluminium as the reactive centre might not be correct: as noted above, catalysis in ‘‘real’’ systems might well involve 5-coordinate structures, and for small ligands even 6-coordinate intermediates cannot be ruled out. We expect that this will not make much difference. Probably, the O-Al-O angle in the transition state would be reduced, and the 6-membered ring might become more nearly planar, in such a higher-coordinate aluminium system. The analogy between MPV b-hydrogen transfer and the transfer reactions mentioned in the previous two sections raises the question whether ligand effects cannot also be used to increase the efficiency of the aluminiumcatalyzed MPV reduction. As we have mentioned earlier, electron-withdrawing substituents or a positive charge facilitate b-hydrogen transfer in the alkyl/ olefin case. Indeed, MPV reduction is also accelerated considerably by electron-withdrawing substituents. One complication with the use of ligands is that aluminium is extremely oxophilic, so that many ligands will be displaced from the metal by the alcohol used as reductant. Nevertheless, there should be some scope here for the use of strongly chelating N2 or O2 ligands. Electronwithdrawing substituents on the ligand and small bites angles should work best.

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5 How Special Is Aluminium? To put the chemistry summarized in this review into perspective, it is worthwhile to compare aluminium with other elements. In the present context, we have chosen to compare it on the one hand with the more electropositive magnesium, and on the other hand with the heavier and more electronegative element gallium, using Me2Al, Me(NH3)Mg and Me2Ga model fragments; results are summarized in Table 1 (Section 3). Experimentally, organomagnesium compounds do not readily insert olefins [73]. Computational results, however, suggest that the reason for this does not lie in higher insertion barriers. The calculated barriers for Me(NH3)MgEt are about 2 kcal lower than those for Me2AlEt; Francl also found barriers for insertion in HMgMe that were ca 2 kcal/mol lower than those calculated for H2AlMe [27]. It seems more likely that it is very difficult to generate the coordinatively unsaturated organomagnesium species necessary for olefin insertion. Organomagnesium compounds that are not saturated by external donor ligands tend to form strong coordination polymers, which on heating often decompose to olefins and (also polymeric) magnesium hydrides. The olefin complexation energy at magnesium is even lower than at aluminium, whereas electron-deficient alkyl bridges are stronger, so that it is more difficult for the olefin to compete for an empty site. However, the theoretical results suggest that in the absence of such complications magnesium would show very similar reactivity towards olefins (i.e., oligomerization at best). Threecoordinate magnesium complex 10a has been reported not to insert ethylene (80 C, 30 bar) [74]. This was suggested to be caused by steric congestion about the metal centre, which seems a bit strange given the fact that the complex forms a stable THF adduct (10b). Possibly the lowered reactivity of the MAMe bond (relative to higher M-alkyl bonds; see Sect. 2.2) also plays a role here.

Gallium, on the other hand, shows a clearly increased barrier for olefin insertion, but smaller increases for b-elimination and b-hydrogen transfer. Thus, the Aufbau reaction at gallium should be much slower and have a much lower limit on the attainable molecular weight. Organomagnesium compounds usually insert carbonyl compounds, rather than reducing them. This is borne out by calculations, which show that barriers for both processes are lower than for aluminium, but the one for insertion is lowered most. Gallium, on the other hand, has significantly higher

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barriers for both processes, with a balance between the two that is virtually unchanged from the aluminium value. Finally, for MPV reduction we find nearly equal barriers for aluminium and gallium, and a much lower one for magnesium. Again, one should be careful not to attach too much value to these numbers, since for magnesium-catalyzed MPV, in particular, the metal will probably prefer a higher coordination number than 4. The general trend seems to be that Mg is more and Ga is less reactive than Al, as expected. However, the differences between the individual reactions are not easily rationalized and probably deserve further study.

6 Outlook The insertion of double bonds into AlAC and AlAH bonds has been known for a very long time. The importance of direct b-hydrogen transfer as a general reaction category has not been recognized as clearly; we hope that the present overview illustrates how pervasive this reaction is in aluminium chemistry. The relatively high tendency of aluminium alkyls to undergo b-hydrogen transfer, coupled with olefin insertion barriers that are much higher than those for transition metals, makes it unlikely that very efficient high-MW olefin polymerization catalysts based on aluminium will be discovered. Even for the reaction of aluminium alkyls with ketones, using ligands to tune towards insertion selectivity will be difficult. On the other hand, there are clear opportunities for using ligands to promote b-hydrogen transfer even more, and to tune its regioselectivity, diastereoselectivity and even enantioselectivity. Ligands currently used for aluminium are much less sophisticated than those that have already been used for years in the area of chiral transition metal catalysis, so there is a lot of room for ligand development. It is noteworthy that b-hydrogen transfer at aluminium, and presumably at other main group metals, proceeds via geometries that differ sharply from their transition metal counterparts. This will be an important factor in ligand design. Finally, we believe that theory can play a significant role in the design of new aluminium catalysts. Aluminium compounds are more amenable to highlevel calculations than transition metal complexes. Because of the lower coordination numbers involved the catalysts tend to be smaller, which also simplifies calculations. As a result, it is frequently possible to carry out reasonably accurate calculations on a ‘‘real catalyst’’, so that theory can at times really guide or at least assist experiment.

7 References 1. Ziegler K, Gellert H (1950) Liebigs Ann Chem 567: 195 2. Mole T, Jeffery EA (1972) Organoaluminium Compounds, Elsevier, Amsterdam

164 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34.

35. 36. 37. 38. 39. 40. 41.

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Ziegler K, Holzkamp E, Breil H, Martin H (1955) Angew Chem 57: 541 Margl P, Deng L, Ziegler T (1999) J Am Chem Soc 121: 154 Talarico G, Budzelaar PHM, Gal AW (2000) J Comp Chem 21: 398 Talarico G, Budzelaar PHM, Barone V, Adamo C (2000) Chem Phys Lett 329: 99 Talarico G, Budzelaar PHM (2000) Organometallics 19: 5691 Talarico G, Busico V, Budzelaar PHM (2001) Organometallics 20: 4721 Martin H, Bretinger H (1992) Makromol Chem 193: 1283 Coles MP, Jordan RF (1997) J Am Chem Soc 119: 8125 Ihara E, Young Jr VG, Jordan RF (1998) J Am Chem Soc 120: 8277 Bruce M, Gibson VC, Redshaw C, Solan GA, White AJP, Williams DJ (1998) Chem Commun 2523 Kim JS, Wojcinski II LM, Liu S, Sworen JC, Sen A (2000) J Am Chem Soc 122: 5668 Dagorne S, Guzei IA, Coles MP, Jordan RF (2000) J Am Chem Soc 122: 274 Korolev AV, Ihara E, Guzei IA, Young Jr VG, Jordan RF (2001) J Am Chem Soc 123: 8291 Hauske JR (1991) Organoaluminum Reagents. In: Trost BM, Fleming I (eds) Comprehensive Organic Synthesis. Pergamon, Oxford, Sect. 1.3 For convenient comparison, structures of aluminium model species shown in the text have been calculated at a uniform level of theory (B3LYP/6–31G(d)), which is not always the same as that used in individual reports. Maruoka K, Yamamoto H (1988) Tetrahedron 44: 5001 Eisch JJ (1982) Aluminum. In: Wilkinson G, Stone FGA, Abel EW (eds) Comprehensive Organometallic Chemistry, Vol 1. Pergamon, Oxford, Chap. 10 Eisch JJ (1995) Aluminum. In: Abel EW, Stone FGA, Wilkinson G (eds) Comprehensive Organometallic Chemistry II, Vol 1. Pergamon, Oxford, Chap. 10 Eisch JJ (1995) Aluminum. In: Abel EW, Stone FGA, Wilkinson G (eds) Comprehensive Organometallic Chemistry II, Vol 11. Pergamon, Oxford, Chap. 6 De Graauw CF, Peters JA, Van Bekkum H, Huskens J (1994) Synthesis 1007 Egger KW (1972) J Chem Soc Faraday 1 68: 1017 Egger KW (1969) J Am Chem Soc 91: 2867 Egger KW (1969) Int J Chem Kinetics 1: 459 Sakai S (1991) J Phys Chem 95: 175 Bundens JW, Yudenfreund J, Francl MM (1999) Organometallics 18: 3913 Reinhold M, McGrady JE, Meier R (1999) J Chem Soc Dalton Trans 484 Cavallo L, personal communication Wilke G (1995). In: Fink G, Mu¨lhaupt R, Brintzinger HH (eds) Ziegler Catalysts. Springer-Verlag, Berlin, p 3 March J (1977) Advanced Organic Chemistry, 2nd ed. McGraw-Hill, New York, pp 1037– 1041 Houk KN, Li Y, Evanseck JD (1992) Angew Chem 104: 711 B3LYP//6-31G(d) values; these are not very sensitive to the level of theory employed. (a) Small BL, Brookhart M, Bennett AMA (1998) J Am Chem Soc 120: 4049; (b) Britovsek GJP, Gibson VC, Kimberley BS, Maddox PJ, McTavish SJ, Solan GA, White AJP, Williams DJ (1998) Chem Commun 849; (c) Small BL, Brookhart M (1998) J Am Chem Soc 120: 7143; (d) Britovsek GJP, Bruce M, Gibson VC, Kimberley BS, Maddox PJ, Mastroianni S, McTavish SJ, Redshaw C, Solan GA, Stro¨mberg S, White AJP, Williams DJ (1999) J Am Chem Soc 121: 8728; (e) Britovsek GJP, Mastroianni S, Solan GA, Baugh SPD, Redshaw C, Gibson VC, White AJP, Williams DJ, Elsegood MRJ (2000) Chem Eur J 6: 2221 Meier RJ, Koglin E (2001) J Phys Chem A 105: 3867 Budzelaar PHM, Talarico G. ACS Symposium Series 822 (2002) Chap. 10, p 142 Talarico G, Budzelaar PHM (2001) Organometallics 21: 34 Coles MP, Swenson DC, Jordan RF (1998) Organometallics 17: 4042 Coles MP, Swenson DC, Jordan RF, Young Jr VG (1997) Organometallics 16: 5183 Dagorne S, Jordan RF, Young Jr VG (1999) Organometallics 18: 4619 Korolev AV, Guzei IA, Jordan RF (1999) J Am Chem Soc 121: 11605

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42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74.

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Radzewich CE, Coles MP, Jordan RF (1998) J Am Chem Soc 120: 9384 Radzewich CE, Guzei IA, Jordan RF (1999) J Am Chem Soc 121: 8673 Aeilts SL, Coles MP, Swenson DC, Jordan RF (1998) Organometallics 17: 3265 Aparna K, McDonald R, Ferguson M, Cavell RG (1999) Organometallics 18: 4241 Chen EY-X, Marks TJ (2000) Chem Rev 100: 1391 Because the easy exchange of ligands and alkyl groups between aluminiums, activation of an aluminium dichloride or dialkyl complex with MAO leads to complex mixtures and to results that are difficult to interpret. Zhou Y, Richeson DS (1996) Inorg Chem 35: 1423 McGrady GS, Turner JFC, Ibberson RM, Prager M (2000) Organometallics 19: 4398 Neumann HM, Laemmle J, Ashby EC (1973) J Am Chem Soc 95: 2597 Eisch JJ, Fichter KC (1984) J Org Chem 49: 4631 Haubenstock H, Davidson EB (1963) J Org Chem 28: 2772 Iguchi S, Nakai H, Hayashi M, Yamamoto H, Maruoka K (1981) Bull Chem Soc Jpn 54: 3033 Haubenstock H (1990) Tetrahedron 46: 6633 Meyer-Stork MA, Haag D, Scharf H-D (1997) J Chem Soc Perkin Trans 2 593 Meerwien H, Schmidt R (1925) Liebigs Ann Chem 444: 221 Verley A (1925) Bull Soc Chim Fr 37: 537 Pondorf W (1926) Angew Chem 39: 138 Oppenauer RV (1937) Recl Trav Chim Pays-Bas 56: 137 Doering WvE, Young RW (1950) J Am Chem Soc 72: 631 Menicagli R, Giacomelli GP, Lardicci L (1973) J Organomet Chem 50: C15 Xianming H, Kellogg RM (1996) Recl Trav Chim Pays-Bas 115: 410 Evans DA, Nelson SG, Gagne´ MR, Muci AR (1993) J Am Chem Soc 115: 9800 Ishihara K, Hanaki N, Yamamoto H (1993) Synlett 127 Fujita M, Takarada Y, Sugimura T, Tai A (1997) Chem Commun 1631 Node M, Nishide K, Shigeta Y, Shiraki H, Obata K (2000) J Am Chem Soc 122: 1927 Aremo N, Hase T (2001) Tetrahedron Lett 42: 3637 Akamanchi KG, Varalakshmy NR (1995) Tetrahedron Lett 36: 3571 Ko B-T, Wu C-C, Lin C-C (2000) Organometallics 19: 1864 Ooi T, Miura T, Maruoka K (1998) Angew Chem Int Ed 37: 2347 Ooi T, Itagaki Y, Miura T, Maruoka K (1999) Tetrahedron Lett 40: 2137 Ashby EC (1988) Acc Chem Res 21: 414 Lindsell WE (1982) Magnesium, Calcium, Strontium and Barium. In: Wilkinson G, Stone FGA, Abel EW (eds) Comprehensive Organometallic Chemistry, Vol 1. Pergamon, Oxford, Chap. 3 Bailey PJ, Coxall RA, Dick CM, Fabre S, Parsons S (2001) Organometallics 20: 798

Compounds Containing Five-Coordinate Group 13 Elements David A. Atwood, Aaron R. Hutchison, Yuzhong Zhang Department of Chemistry, The University of Kentucky, Lexington, KY 40506-0055, USA e-mail: [email protected]

The synthesis, characterization, and applications of five-coordinate group 13 compounds are examined and surveyed. The range of compounds covered in this study includes those with five separate ligands (called ‘‘Class A’’), a single bidentate ligand (Class B), two bidentate ligands (Class C), tridentate ligand coordination (Class D), and tetradentate ligand coordination (Class E). Each class of five-coordinate compounds is further subdivided into the possible ‘‘Types’’ of structures that may be envisioned to occur. Following an extensive (although not exhaustive) search of the literature it was found that the majority of the compounds fell into a very few Types in each Class. Thus, future reviews may focus more specifically on the details within these well-populated Types. Many of the possible Types of compounds had no members whatsoever, and thus provide clear targets for future research endeavors. As a striking example, no compound involving a single five-coordinate ligand (Class F) was found. The ultimate goal of this work is to provide a systematic understanding of the types of five-coordinate compounds that have been made and that may be sought in the future. As such it is primarily a survey of the known compounds with some details about their structures. In particular, the distinction between trigonal bipyramidal and square pyramidal geometry will be discussed in terms of a literature measure known as the ‘‘tau’’ value. This value ranges from ‘‘0’’ for square pyramidal to ‘‘1’’ for trigonal bipyramidal geometries, thus providing a quantitative indication of what geometry a structure most closely adopts. Keywords: Aluminum, Gallium, Indium, Classification, Geometry, Structure, Tau value

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3.1 3.2 3.3 3.4 3.4.1 3.4.2 3.4.3 3.4.4 3.4.5 3.4.6

Five Separate Ligands (Class A) . . . . . . . . . . . . . . One (Class B) or Two (Class C) Bidentate Ligands One Tridentate Ligand (Class D) . . . . . . . . . . . . . One Tetradentate Ligand (Class E) . . . . . . . . . . . . Salen Ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . Aminotroponimates . . . . . . . . . . . . . . . . . . . . . . . Bis(amino)thiols . . . . . . . . . . . . . . . . . . . . . . . . . Porphyrins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Phthalocyanines . . . . . . . . . . . . . . . . . . . . . . . . . Tripodal Ligands . . . . . . . . . . . . . . . . . . . . . . . . .

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Open Chain Amines . . . . . . . . Thiosemicarbazones . . . . . . . . Salan Ligands . . . . . . . . . . . . . Pentadentate Ligands (Class F)

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Abbreviations and Symbols ABZ AEL amidine APA APL BAT-TM BHT bO3N BPEDA bS3N Bu Bz DA18C6 DMAE DMAMP DMAPL DMD DMIP DPM EDBP Et H2mpdta H2msdta HN18C6 HQN i-Pr-DPS i-Pr-DPT MAMP MBZL Me MEL MIP MME MMP

[(2-Me2NCH2)C6H4] 2-aminoethanol RNCR¢N(H)R (R,R¢ = H, alkyl, aryl,) [CH2CH2CH2NMe2]) 3-aminopropanol tetramethylbis(aminoethanethiol) 2,6-di-t-butyl-4-methylphenol tris(2-hydroxybenzylamine) N,N¢-bis(3-aminopropyl)ethylenediamine tris(2-mercaptobenzylamine) butyl benzyl diaza-18-crown-6 N,N-dimethyl-2-aminoethanol {bis[(dimethylamino)methyl]pyrrole} N,N-dimethyl-3-aminopropanol 2,2-dimethylpropane-1,3-diol N,N-dimethylamino-2-propanol S-a,a-diphenyl-2-pyrrolidinyl methoxide 2,2-ethylidene-bis(4,6-di-t-butylphenol) ethyl N-methyl-N¢,N¢¢-bis(diisopropyl)diethylenetriamine N-methyl-N¢,N¢¢-bis(trimethylsilyl)diethylenetriamine aza-18-crown-6 8-hydroxyquinoline tetraisopropylimidodiphosphine selenate tetraisopropylimidodiphosphine thiolate {2-[(dimethylamino)methyl]pyrrole} 2-methoxybenzyl alcohol methyl 2-methoxyethanol 1-methoxy-2-propanol 2-nercaptomethylethanol 3-nercaptoethylpropanol

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MMPH MPH MPL NOH nor (NR)3N OEP Ph3-am Ph-DPS pmdien Pr ‘‘Salen’’ salen salophen salpen S3N sqp TAP tbmSalcen tbmSalen tms-bam TPP tbp VAPOL

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2-nercaptomethylphenol 2-methoxyphenol 3-methoxypropanol pyridine-2-carbaldehyde oxime 1-norborynl tris(3-alkylaminopropylamine) 2,3,7,8,12,13,17,18-octaethyl porphyrinate triphenylamidine tetraphenylimidodiphosphine selenate N,N,N¢,N¢¢,N¢¢-pentamethyldiethylenetriamine propyl N,N¢-bis(salicylidene)alkylene- or -arylenediimine N,N¢-bis(salicylidene)ethylenediimine N,N¢-bis(salicylidene)-1,2-phenylenediimine N,N¢-bis(salicylidene)-propylenediimine tris(3-mercaptopropylamine) square pyramidal tetra(p-anisyl) porphrinate N,N¢-trans-1,2-cyclohexanediyl-bis(3-t-butyl) -5-methylsalicylideneimine N,N¢-1,2-ethylenebis(3-t-butyl-5-methylsalicylideneimine) [N,N¢-bis(trimethylsilyl)benzamidinate] tetra(o-tolyl)porphrinate trigonal bipyramidal ‘‘vaulted’’ bisphenanthrol

1 Introduction Following the pioneering work of Ziegler in the 1950s the chemistry of threeand four-coordinate aluminum alkyl compounds has seen a steady increase in activity [1]. Despite the generally high reactivity of AlAC bonds, and their susceptibility to hydrolysis, many of these compounds have found use in organic synthesis or in catalysis. Indeed, one of Ziegler’s early discoveries was the ability of AlEt3 to polymerize olefins, and this process is still in use today for the manufacture of detergents. The synthetic and catalytic use of these and other aluminum reagents generally rely upon the fourth Lewis acidic site on the metal center. However, there is growing evidence that aluminum (and by extension gallium and indium) forms five- and six-coordinate complexes during many reactions of interest. This is undoubtedly true when bidentate substrates are involved in organic synthetic reactions. As a research target, however, highercoordinate aluminum was not actively pursued until the mid 1980s when Robinson began an aggressive campaign to explore the structures of these compounds [2]. His systematic work, as well as publications from other early pioneers in group 13 chemistry has established the means by which

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five-coordinate group 13 compounds may be formed. Now a great deal of structural information is available on this class of compounds. Building on this earlier work, researchers have, over the past decade, developed an extensive series of five-coordinate group 13 complexes based upon use of the Salen ligand system [Salen = N,N¢-(alkylene or arylene)bissalicylideneimine] [3]. The Salen-group 13 compounds are relatively stable and do not dissociate. They, thus, provide a nice soluble platform upon which fivecoordinate chemistry may be conducted. This is the most extensively developed class of five-coordinate compounds. Examples are known for a wide range of aluminum, gallium and indium, halides, alkyls, amides, alkoxides and siloxides. In some cases the compounds can be used as precursors to strongly Lewis acidic six-coordinate cations. All of this work had as its foundation the work of Goedken on five-coordinate alkyls [4], that of Storr in the preparation of bimetallics [47], and that of Barron on five-coordinate alkoxides [5]. Early-on it was discovered that these Salen compounds, and the related sixcoordinate cations [6], were useful as catalysts for the polymerization of oxiranes. These applications were anticipated in the efforts of Spassky [7] and in the substantial work of Inoue [8]. Subsequently, applications of these compounds in organic synthesis have been developed [9]. Additional applications include their use in catalytic lactide polymerization [10], lactone oligomerization [11], the phospho-aldol reaction [12], and as an initiator in methyl methacrylate polymerization [13]. Based upon their ready availability, and generally low air and moisture sensitivity, it is clear that many more applications will be found for fivecoordinate group 13 compounds incorporating Salen ligands. It is also easy to conceive that many other ligand combinations might have similarly favorable properties. It would be useful, though, to understand both the range of compounds that may be accessed as well as the nature of the bonding at the fifth coordination site (the non-Salen site for example), and not just for Salen-based derivatives. The present work will provide this information in a survey of fivecoordinate aluminum, gallium, and indium compounds with various ligand combinations. The classification scheme that is presented should provide a starting point for further systematic development of each class of compounds.

2 Geometry Five-coordinate metal complexes tend to adopt either square pyramidal or trigonal bipyramidal geometries. However, the compounds rarely conform exactly to one or the other of these geometries. Rather, they tend to form distorted structures that lie somewhere in-between the ideal geometries. In the following sections, generalities will be made regarding which of these geometries are preferred with various ligand sets. Therefore, it will be convenient to have some sort of method to organize and provide a quantitative measure of the geometry of different five-coordinate complexes. An equation producing a ‘‘tau’’ value [14] has been used to determine the distortion of

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Fig. 1. Definitions of angles and equation for calculating a ‘‘tau’’ value

these metal complexes. The tau value ranges from 0 to 1, zero being perfectly square pyramidal and one being perfectly trigonal bipyramidal. The tau equation is show in Fig. 1. The substituents around the metal are labeled a–e, with ‘‘a’’ being the substituent which does not help define either of the two largest angles around the metal. If ‘‘a’’ is defined as the z axis, then alpha (‘‘c’’ to ‘‘e’’) and beta (‘‘b’’ to ‘‘d’’) are the angles that are opposite from each other in the xy plane. A list of selected tau values for a range of different compounds can be found in Table 1. Some example calculations are presented below. For Salomphen(t-Bu)AlOSiMe3: Where ) OSiMe3 is a; b and d are N2 and O1; c and e are N1 and O2. O1-Al-N2 = 153.74 and O2-Al-N1 = 144.56. s = (153.74)144.56)/60. s = 0.153. For Salen(t-Bu)AlOSiMe3: Where ) OSiMe3 is a; b and d are N2 and O1; c and e are N1 and O2. O1-Al-N2 = 160.48 and O2-Al-N1 = 132.86. s = (160.48)132.86)/60. s = 0.46. This shows that the salomphen(t-Bu) ligand prefers a square pyramidal geometry whereas the salen(t-Bu) ligand is almost exactly in-between the two geometries. It should be noted that these numbers are meant to provide a guideline for the type of structure that results. The numbers can be somewhat variable even within a single compound. For example, in the dimeric compound incorporating dimethylamino-1-propanol (= L), [LAl(Et)(Cl)]2, the tau values for the two aluminum atoms are 0.04 and 0.17 [15]. Although they are not crystallographically equivalent the two atoms are nevertheless in almost identical bonding arrangements. Thus, both tau numbers refer to aluminum atoms in almost perfectly square pyramidal geometries.

3 Classification of Five-Coordinate Compounds The structures of five-coordinate compounds can be classified theoretically into six different structural types, progressing from those containing five

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Fig. 2A–F. The six general geometries leading to five-coordinate group 13 compounds

monodentate ligands to those bound entirely by a single hypothetical pentadentate ligand (Fig. 2). This report is not meant to present a comprehensive coverage of all the known compounds. For instance, charged compounds are not included except in selected cases. Furthermore, many five-coordinate group 13 atoms appear in complicated clusters that should be considered under a different classification system and will only be briefly mentioned herein. However, even a cursory survey of the literature reveals that the majority of the known compounds incorporate either bidentate (types B or C) or tetradentate (type E) ligands. This indicates that new contributions could be found for types A, D and F, the less developed of the structural classes. For these less represented structural types a more complete coverage was attempted to ensure identification of all of the members. 3.1 Five Separate Ligands (Class A)

Class A compounds are innately limited since they rely upon having five separate, free ligands around the central metal. This necessarily restricts the range of useful ligands to those that are small. Thus, group 13 hydrides (Fig. 3a) [16], halides, azides, and trifluoromethyl compounds (Fig. 3b) provide the majority of examples. In these compounds, the covalent groups

Fig. 3a–c. Examples of five-coordinate compounds with five separate ligands

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are found in the equatorial sites and the coordinate covalent groups (the ‘‘dative’’ bonds) in the axial sites. This is a very common arrangement for tbp compounds. In the case of the hydride and azides, the two ‘‘extra’’ groups beyond that needed to satisfy the valence of aluminum help to imbue these compounds with volatility by inhibiting aggregate formation. This is important in materials synthesis where they are used as sources of Al(0) (from the hydrides) and AlN (from the azides). The compound In(CBCCF3)3THF2 [17] is an interesting structurally characterized example of a class A compound. The ligands are in an ideal trigonal bipyramidal geometry (tau = 1) with the covalent InAC bonds equatorial and the dative InAO bonds axial. The interaction of four-coordinate compounds with dioxygen, to ultimately form alkoxides, is thought to proceed through a five-coordinate intermediate [18, 19]. In an effort to determine whether this coordination occurs, the interaction of O2 with [R2Al(pyrazolyl)]2, was explored [20]. The structures involved in this process are shown in Fig. 4. It was found that increasing the size of the R group from Me, Et, to t-Bu prevented alkoxide formation. In the mechanism the inserted oxygen molecule forms a peroxo intermediate which, in some cases, can be isolated (Type BIII) [21–23]. This intermediate would qualify as a class A compound. This also might imply that such compounds could also undergo other insertion reactions, such as carbonyl insertions that are well-known for transition metal compounds. 3.2 One (Class B) or Two (Class C) Bidentate Ligands

There are two basic structures possible for five coordinate compounds incorporating bidentate chelates: either a compound with three free ligands

Fig. 4. An intermediate five-coordinate Class A compound before oxygen insertion

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Fig. 5. Three types of Class B compound containing a single bidentate ligand

and one chelate (Type B in Fig. 2) or a compound with two chelates and one free ligand (Type C shown in Fig. 2). To simplify classification, any coordination from bridging ligands in oligomeric compounds is counted as a free ligand contribution. Furthermore, bidentate chelates can be classified as neutral, uninegative, or dinegative. The possible structures, called ‘‘types’’, for a single bidentate ligand are summarized in Fig. 5. The possible structural types for two bidentate ligands are shown in Fig. 6. It is worth noting that uninegative chelates (i.e., those that form one covalent and one dative bond) are by far the most common types seen in the literature. Since many reports of class B and C compounds appear in the same literature citations they will be discussed together in this section. Additionally, the same ligand can be used to access these two classes of compound by using either one stoichiometric equivalent (for class B) or two equivalents (for class C) of the bidentate ligand. It is possible that compounds containing one neutral chelate (Type BI) were first prepared in Lewis acid-catalyzed reactions involving aluminum reagents and various substrates capable of bidentate all dative coordination to the aluminum, although the true coordination number around the aluminum atom was not at the time understood. Based upon NMR data the formation of a five-coordinate complex, R2AlX(alkoxy ketone), was thought to be an intermediate in the reduction of alkoxy ketones to alkoxy alchohols [24]. The presence of tetracoordinate, cationic, species, however was also mentioned as a possibility, in keeping with the work of others (Fig. 7) [25]. The consensus appears to be that the addition of chelating bases causes displacement of the halide and the formation of cations as the active species (Fig. 7b) [26]. This supposition is supported, in part, by observations that group 13 halide compounds, whether chelated or not, tend to form tetracoordinate cations when combined with two or more equivalents of a Lewis base [27].

Fig. 6. Types of Class C compound containing two bidentate ligands

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Fig. 7. Two possible structural types with a single neutral bidentate chelate (Class BI)

Much more common is the combination of a group 13 compound with potentially bidentate ligands containing a covalent bonding unit and a second nucleophile with a free lone pair for dative bonding (Type BII). If the stoichiometry is 1 : 1, this can lead to dimeric compounds in which the group 13 metals are five-coordinate and usually in a trigonal bipyramidal geometry (with varying degrees of distortion). The coordination sphere is comprised of two groups from the bidentate ligand, two free ligand groups (from the original compound), and one bridging group (counted as a ‘‘free’’ ligand in this classification system). Two general examples of this type of bonding arrangement are shown in Fig. 8. Geometrically, the two free ligand groups and one bridging and covalently bonded group are equatorial, while the other bridging and covalent group and the datively bound group are axial. The bridging and covalent atoms are usually oxygen. If, on the other hand, the stoichiometry is two ligands for each metal, then five coordinate structures with two chelate molecules and one free ligand around the metal form (Type CI). There is a wide variety of chelates capable of creating structures of these two types. They are most easily classified by their bonding atoms, for examples with the atom combinations, O, O; N, O; N, N; N, C; S, O; S, S; and C, C. A systemic study of compounds containing O, O; N, O; and S, O ligand sets has been conducted in which the effect of variables such as the steric bulk of the groups on the metal and on the ligand, strength of the Lewis base donor, and size of the chelate ring formed when the ligand becomes bidentate has

Fig. 8. Steric factors (a) influence the tau value more than electronic factors (b)

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been made [28]. It was found, in general, that the compounds exist as an equilibrium between the four- and five-coordinate species. The bond dissociation energies, 5.2–13.2 kJ/mol, are in the range expected for solvation rather than donor bonds. This was attributed to interligand steric repulsion, something that was not recognized for these systems previously. Ligands capable of Type B or C bonding also have E-M-E angles of not more than 100. Those with angles potentially greater than this, acetoximes for instance [29], may support dimeric compounds, but not with chelated five-coordinate metals. Compounds incorporating O, O ligands are perhaps the most common chelates for five-coordinate compounds. It was one of these ligands (MEL, 2-methoxyethanol) that was combined with trimethylaluminum to yield the first crystallographically characterized Type BII compound [30]. A general representation of the structure is shown in Fig. 8a. This compound contains several structural features common to this class, such as bridging by the covalently bound oxygen, a distorted trigonal bipyramidal geometry, and a planar (sum of interior angles equals 360) M2O2 core. The AlAO bonds within ˚ ), presumably due to the core are nearly identical (approximately 1.8 A symmetry and resonance, while the dative methoxide bond is much longer ˚ ). The chelate bite angle is 75.9. (approximately 2.3 A This ligand was also combined with AlR3 (R = t-Bu, i-Bu, Et, Me) to determine the effect of steric bulk on the potential bonding to a fifth coordination site [28]. Although the i-butyl and ethyl derivatives are almost indistinguishable structurally from their methyl counterpart, the t-butyl compound is different in that no dative bond exists between the methoxy oxygen and the aluminum; this is presumably due to the steric bulk of the t-butyl derivative. In this pseudo-four-coordinate compound the tau value is indicative of an sqp geometry (0.2). As the steric bulk is reduced the aluminum becomes more tbp as the methoxy group is allowed to bond (for R = i-Bu tau = 0.48, for R = Me 0.58). With weaker bases, such as MME (2-mercaptomethylethanol), the structures tend towards more sqp geometry with a tau value of 0.25. Thus, with this ligand set a reduction in steric bulk, and a reduction in base strength leads to tbp geometries. The effect of changing the halide is negligible in this type of compound (Fig. 8b) [31]. In a later study [32], this ligand was also combined with (BHT)2Al(H)(Et2O). Even though a 1 : 1 stoichiometry was used, the major product was [(BHT)Al(MEL)(l-MEL)]2. Somewhat surprisingly, both one BHT and a hydrogen have left the aluminum, leaving one molecule of MEL to chelate and bridge, and another as a completely separate free ligand. Although this is a different ligand arrangement than seen before, it is still Type BII. Therefore, the same ligand is bonding three different ways to the metal. As in the compounds studied earlier, the bridging and covalent chelate bonds are very ˚ vs. 1.84 A ˚ ), while the two ligand OAAl bonds (from similar in length (1.87 A ˚ for MEL, the BHT and the non-chelating MEL) are noticeably shorter (1.70 A ˚ for BHT) and the dative bond is slightly shorter than in the 1.73 A ˚ ). MEL has also been used to form a complex alkylaluminum case (2.03 A with GaMe3 [33], which was structurally comparable to the aluminum

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˚ vs. 1.8 A ˚ for analogue, except for marginally longer bond lengths (1.9–2.0 A ˚ ˚ the covalent and bridging O to Al and 2.6 A vs. 2.3 A for the dative bond) and a slightly more acute bite angle (70.3 vs. 75.9). A related ligand, MPL (3-methoxypropanol), was combined with Al(t-Bu)3 and Al(Me)3 [28]. The resulting BII compound was similar to those formed by MEL, with a dative interaction in the methyl compound but not the t-butyl ˚ vs. 2.27 A ˚ ) for the one. However, the dative bond is much longer (2.39 A propyl compound. Another related ligand is MIP(1-methoxy-2-propanol), which was combined with Ga(Me)3 [33]. This complex is nearly identical to that formed by MEL with Ga(Me)3. The literature also contains examples of O, O ligands with aromatic backbones. Two examples are the ligands MPH (2-methoxyphenol) and MPZL (2-methoxybenzyl alcohol) [33]. When MPH was combined with AlMe3, the resulting BII compound was similar to that observed with the ligands described earlier, a Type BII structure. Once again there is a planar Al2O2 core, ˚ . The bridging bond was slightly with the covalent bond being about 1.86 A ˚ ˚ . The MPH ligand was also longer, at 1.94 A, and the dative bond was 2.2 A combined with Al(Et)3 and Al(i-Bu)3 [34]. In these compounds there was a noticeable difference between the covalent OAAl bond and the bridging bond ˚ , bridging approximately 1.95 A ˚ ), while (covalent approximately 1.86 A ˚ for the ethyl complex the dative bonds were significantly shorter (2.25 A ˚ for the i-butyl one). This chelate was also combined with GaMe3, and 2.7 A yielding a complex that was comparable to the aluminum derivatives, but with ˚ for the covalent/bridging O slightly longer bond lengths (approximately 2 A ˚ bond to aluminum and 2.5 A for the dative bond). The MPZL ligand was also combined with AlMe3 [33]. The resulting compound had covalent and ˚ , perfectly in line with the bridging OAAl bonds of approximately 1.85 A ˚ ). earlier compounds, and a long dative bond (2.57 A As another example, InMe3 combines with salicyclaldehyde (L), pyridine-2-carbaldehyde [35], 2-carboxybenzaldehyde, 2-pyrrolidinone, and N,N ¢-diphenylacetamidine to give compounds with the BII classification. In the structure of [Me2InL]2 the equatorial angles are substantially deviated from ideal. The ligand O-In-O angle is 79.9(1) while the widest O(ligand)-In-O(bridging) angle is 154.7(1). The latter angle creates space into which the C-In-C angle can bend with an angle of 141.9(2). The tau values for these compounds range from 0.19 to 0.38. It is also possible for two O, O ligands to bind the same metal, resulting in a Class CI structure. For example, when two equivalents of Al(t-Bu)3 is combined with one equivalent of ethylene glycol in hexane, the primary product is an insoluble alkone [36]. However, a minor product is a discrete, five- and four-coordinate dimer [Al2(t-Bu)3(OCH2CH2O)(OCH2CH2OH]. The five-coordinate aluminum has both ethylene glycol molecules attached, plus one t-butyl group. The geometry is sqp with a tau value of 0.16. The chelates are both equatorial and the alkyl group are axial. If the same reaction is run in diethyl ether, the result is a trimer, Al3(t-Bu)5(OCH2CH2O)2, with one fivecoordinate aluminum and two three-coordinate aluminums. This is an almost linear molecule, with a four-coordinate aluminum (two t-butyl groups and a

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bridging oxygen from each glycol) on each end and the five-coordinate aluminum (both glycols and a t-butyl) in the middle. The geometry is a more distorted square pyramid (tau = 0.43). A related ligand is DMD (2,2-dimethylpropane-1,3-diol). This was also used by Barron [37] to form a number of four- and five-coordinate trimers similar to that made with ethylene glycol. The compounds involved in this reaction are shown in Fig. 9b and, more generally, in 9c. Initially, DMD was combined stoichiometrically with Al(t-Bu)3 and Ga(t-Bu)3 to generate the four-coordinate dimers M2(t-Bu)4(DMD-H)2 (M2 = Al2, Ga2 and AlGa). When Al2(t-Bu)4 (DMD-H)2 was refluxed with AlH3(NMe3), the result was the four- and fivecoordinate Type CI trimers Al3(t-Bu)4H(DMD)2 and Al3(t-Bu)4Cl(DMD)2 (from traces of AlH2Cl(NMe3) in the starting material). Refluxing with AlMe3 also yielded Al3(t-Bu)4Me(DMD)2, although Al(t-Bu)3 could not be used to make Al3(t-Bu)5(DMD)2. However, Ga2(t-Bu)4(DMD-H)2 was more reactive, combining with Ga(t-Bu)3 to form Ga3(t-Bu)5(DMD)2, with AlH3(NMe3) and AlMe3 at room temperature to form Ga2Al(t-Bu)4H(DMD)2 and Ga2Al(t-Bu)4Me(DMD)2, respectively, and reacting with Al(t-Bu)3 at reflux to give Ga2Al(t-Bu)5(DMD)2. All of these compounds were sqp (tau
0) with the central metals somewhat above the O4 plane, with Ga being further displaced than Al. Furthermore, if Ga2(t-Bu)4(DMD-H)2 is thermally decomposed in toluene, the result is the addition of a toluene molecule to create Ga3(t-Bu)4(PhCH2)(DMD)2. The exact mechanism by which this occurred is unknown, although it is worth noting that a similar compound is not formed by the thermal decomposition of Ga3(t-Bu)5(DMD)2. Pasynkiewicz also made a series of compounds similar to Barron’s by combining trimethylaluminum with DMD, 1,4-butanediol, and 1,3-butanediol, although none of these were crystallographically characterized [38]. A general depiction of the structures

Fig. 9. The use of diols in forming Type CI structures (a) and (b), and a general depiction of the resulting trimetallic compounds (c)

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obtained with various glycols is shown in Fig. 9c. They are structurally similar to other trimetallic group 13 compounds, including those incorporating Salan ligands [39]. Amino alcohols (N, O ligands) are also well-suited to provide a fivecoordinate geometry and there are many examples of their use. Thus, the ligand DMAE (N,N-dimethyl-2-aminoethanol) combines with Al(t-Bu)3, Al(i-Bu)3, Al(Et)3, and AlH3 Æ NMe3 to form five-coordinate BII compounds in all but the t-Bu reagent [28]. The steric bulk of the t-Bu group blocks the nitrogen from forming a dative bond. For the other compounds, it was clear which O was covalently bound and which was bridging simply by comparing bond lengths; ˚ (i-butyl complex) and 1.845 A ˚ the covalent bonds were between 1.833 A ˚ (H complex), while the bridging bonds ran from 1.958 A (t-butyl complex) to ˚ (H complex). The dative bond was longest for the i-butyl complex 1.897 A ˚ ), while significantly shorter in the H derivative, (2.14 A ˚ ). Also (2.34 A referenced in this paper is a complex of this ligand with Al(Me)3, which is nearly identical to the previous compounds, with a covalent OAAl bond of ˚ , a bridging bond of 1.93 A ˚ , and a dative bond of 2.13 A ˚ . In a later study, 1.83 A the ligand was used with (BHT)2AL(H)(Et2O), generating the complex [(BHT)Al(H)(l-DMAE)]2 [32]. This compound is also quite similar to the dialkyl predecessors, with all bond lengths falling within the range described earlier. DMAE has also been used to make gallium compounds, by reacting it with GaMe3 and GaH3 Æ NMe3 [40]. As expected, the gallium compounds had very similar structures to their aluminum counterparts, except for slightly ˚ for both, longer bonds. The covalent OAGa bond was approximately 1.9 A ˚ while the bridging bond was approximately 2.0 A. The dative bond was, pre˚ ) for the methyl derivative than for the gallane dictably, slightly longer (2.47 A ˚ (2.28 A). Related to DMAE are the ligands AEL (2-aminoethanol) and DMAPL (N,N-dimethyl-3-aminopropanol). AEL is simply DMAE without the methyl groups on the amine. Therefore, there is less steric hindrance for the dative bond. It is worth noting that this ligand produced the only known fivecoordinate BII complex with Al(t-Bu)3 [28]. The structure was very similar to ˚, a the earlier structures with DMAE, with a covalent OAAl bond of 1.87 A ˚ ˚ bridging bond of 1.92 A, and a dative bond of 2.13 A, perfectly in line with the lengths found for DMAE with Al(Me)3, AlH3, and (BHT)2AL(H). In a related study the DMAPL ligand, differing from DMAE only by one methylene group in the backbone, was used with AL(Et)2Cl to form a five-coordinate BII compound [15]. Interestingly, this compound did not completely follow the trends already discussed. Unlike most of the other amino alcohol complexes, there is not a ˚ ) and the large difference in bond length between the covalent bond (1.86 A ˚ ˚ bridging one (1.89 A). Also the dative bond, at 2.12 A, is somewhat shorter than expected. As with O, O ligands, there are cases in the literature of N, O ligands with aromatic backbones. One example is with the ligand HQN (8-hydroxyquinoline), which creates a five-coordinate BII aluminum through oxygen bridging groups (Fig. 10a), exactly like its aliphatic relatives [41]. The bridging group is ˚ by comparison to the covalent distinct with an Al-O distance of 2.003 A

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Fig. 10. Example (a) of BII compounds and a CI compound (b) and the cis (c) and trans (d) conformations for Salen derivatives

˚ . The geometry is distorted square pyramidal (tau = 0.45) distance of 1.879 A with one of the alkyls occupying the apical position. The O-L-N angle is 79.3. Similar dimeric compounds result when the alkyl is Et [42]. However, when the alkyl group is more bulky, such as t-butyl, monomers result. A single Salan ligand can be viewed as two dinegative bidentate (N,O) ligands connected by an alkylene or arylene linker. These ligands form a wide range of trimetallic aluminum [43] and gallium [44, 45] derivatives containing five coordinate sites that are Type CI if the linker is ignored or Type E if the linker is taken into account. These can exist as either a cis (Fig. 10c) or a trans (Fig. 10d) isomer. The cis isomer enforces a distorted sqp geometry around the central metal. This is observed in the compounds salpanAlMe(AlMe2)2 and salophanAlMe(AlMe)2 in which, for example, the tau values are essentially zero. In saleanAlMe(AlMe2)2, which is trigonal bipyramidal with bulky groups on the aluminum, as in saleanAl-i-Bu(Al-i-Bu)2, a cis orientation, with a central square pyramidal aluminum (tau = 0) can be obtained. Interestingly, based upon the bond lengths and angles tbp is the only geometry observed for gallium as demonstrated in LGaMe(GaMe2)2 with L = salean and salpan. Thus, gallium appears to prefer a tbp geometry over a sqp geometry. Square pyramidal geometries for gallium are, nevertheless, easily obtained with other tetradentate chelates as indicated below.

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As also seen with the O, O ligands, two N, O ligands can bind a single metal to form a five-coordinate complex. One example of this comes from the combinations of the ligand DPM (S-a,a-diphenyl-2-pyrrolidinyl methoxide) with AlMe3, AlEt3, and Al(1-Nor)3 to form CI compounds of the formula AlR(DPM)2, with the 1-nor derivative being the first aluminum norborynl structure published [46]. These compounds have a distorted tbp geometry, with the nitrogens axial and the remaining three substituents equatorial. The geometry shows greater distortion from ideal tbp with increasing steric encumbrance of the ligand. For example, the axial N-Al-N angle is reduced from 171.2 to 167.9 in the series. The tau values reflect this change, dropping from 0.85 to 0.75. Salicylaldoximato (SaloxH) is another example of an aromatic N,O ligand [47]. Two of these chelates react with one equivalent of GaMe3 to create the CI compound (Salox)2GaMe (Fig. 10b). The structure is almost perfectly square pyramidal (tau = 0.08), with the two chelates forming the equatorial ring and the methyl axial. This geometry appears to be due to hydrogen bonding between the amine hydroxy hydrogen and the phenyl oxygen. Beyond that the bond lengths and angles in the structure are routine. N,N ligands can also form five-coordinate CI compounds with the group 13 elements. One such ligand is MAMP, {2-[(dimethylamino)methyl]pyrrole]}). When two equivalents of LiMAMP are combined with AlCl3, the result is the five-coordinate compound (MAMP)2AlCl [48]. The structure is interesting in that the geometry around Al is somewhat ambiguous. The tau value of 0.45 argues that this compound should be classified as very distorted square pyramidal, with the AlACl bond axial (the z axis). However, if one chooses to classify the AlACl bond as defining the y axis, then the molecule can be described with a trigonal pyramidal geometry. This would place the two covalent (formerly anionic) N atoms (short AlAN bonds) and the remaining Cl in the equatorial plane and the two dative N atoms (long AlAN bonds) axial, a ubiquitous motif in coordination chemistry. It is worth noting that this latter view is the one taken by Huang and coworkers, while the current authors prefer to simply deem this compound as possessing an intermediate geometry. Another widely used class of N,N ligands are amidines, which will also react with group 13 reagents in a 2 : 1 ratio to form five-coordinate CI structures. Three such compounds, L2AlX [with L = tms-bam; X = Cl, H (as shown in Fig. 11a), or CCPh] [49] can be prepared by reducing the chloride derivative, and then having it undergo dihydrogen elimination with phenylacetylene. Thus, the ligand environment is sufficiently stable around aluminum to allow reactions to be conducted with relative ease. The structure of (tms-bam)2AlH shows that the four ligand nitrogens are in an almost planar arrangement in forming a distorted tbp geometry around aluminum (tau = 0.76). The N-Al-N¢ angles are very acute at
69. Similar CI structures can be obtained if 2 equivalents of Li[MeC(N-i-Pr)2], Li[t-BuC(N-i-Pr)2], or Li[t-BuC(NCy)2] combine with AlCl3 [50]. The compounds that result are in a distorted tbp geometry with N-Al-N angles that are consistently
67. The tau values reflect their tendency towards this geometry with values
0.70. Many of these compounds have proven useful as olefin polymerization catalysts [51].

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Fig. 11. Examples demonstrating the great diversity of CI compounds

Two sequential substitutions of the alkyl groups on GaMe3 can be achieved with 3 moles of triphenylamidine (Ph3-am) when the reaction mixture is heated in a sealed tube to 145 C for 12 h in toluene [52]. Remarkably, and in contrast to the compounds described previously, the ligands form the basal plane in an sqp, CI arrangement around the gallium. The tau value (0.05) is effectively zero. The amidine ligands here are limited to very acute N-Ga-N angles [63.43(9) and 63.98(9)]. The trisubstituted derivative may be achieved by heating the melt to 200 C. This compound may reflect the preference of gallium for sqp geometries. However, the analogous aluminum compounds with Ph3-am have not been reported. Cationic troponimates, like the amidines described above, have activity as olefin polymerization catalysts [53]. In one compound the [H2B(C6F5)2]) counter anion coordinates the aluminum cation (i-Pr2-ATI)Al(C6F5)+ to form a five-coordinate compound (Fig. 11b). In this compound both the troponimate and borate act as bidentate ligands to aluminum, making this a CI structure. The geometry is a near perfect square pyramid, with a tau of 0.02. This compound serves as a good example of the range of compounds which may be accommodated in this classification system. Chelates utilizing nitrogen and carbon as their bonding atoms are also known. A common example of this type of ligand is APA (-CH2CH2CH2NMe2). Two equivalents of LiAPA will combine with AlCl3 to form the CI compound (APA)2AlCl [54]. The distorted tbp structure (tau = 0.67) has the nitrogens

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axial and the carbons and chlorine equatorial. This reacted with sodium azide to form (APA)2AlN3, which was structurally similar to its chloride predecessor. ˚ ) AlAN dative All of these compounds contained the slightly long (
2.2 A bonds characteristic of these ligands. Schumann also prepared five-coordinate CI indium compounds by combining two equivalents of LiAPA with InBr3 and InCl3 and (APA)2InBr was structurally characterized [55]. The geometry around the indium was intermediate between trigonal bipyramidal and square pyramidal, with a tau of 0.37. It is interesting to note that the InABr bond is ˚ ), due to a definite interaction between the bromine unusually long (2.644 A ˚ , this and the indium of an adjacent molecule. With a distance of 3.975 A interaction is longer than expected for a bridging group, but it does have an ˚ and 2.51 A ˚) effect on the bonding of the system. Also, the InAN bonds (2.53 A are among the longest InAN bond lengths known. The halides in these compounds were labile and could easily be replaced by nucleophiles such as NaO2CCF3, LiOC6F5, Li-t-Bu, ethyl Grignard, and i-propyl Grignard reagents. Further studies proved that other nucleophiles, including LiCN, LiSCN, LiCBCCF3, LiCBCMe, LiCBCsiMe3, LiOC6F5, LiOC(O)CF3 and LiI, would also readily replace the chlorine on In [17]. The iodide derivative, made via the reaction with LiI, showed the same halide quasi-bridging observed for the bromide. It also possessed a trigonal bipyramidal geometry. Efforts to place trifluoroalkyl groups on the In by direct substitution failed. However, the reaction of Cd(APA)2 with solvated RInBr2 (R = CF3, i-C3F7) yielded RIn(APA)2, although the compounds proved to be stable only at low temperatures. Two equivalents of LiAPA will also react with GaCl3 to form (APA)2GaCl, which was also structurally characterized. It is quite similar to the In counterpart, with a distorted trigonal monopyramidal geometry in which the amines are axial. With the goal of creating a six coordinate structure, three equivalents of LiAPA were also combined with GaCl3. Although the hoped for Ga(APA)3 was synthesized, it proved to be a fivecoordinate CI compound, with the three carbons equatorial and two of the N atoms axial, while the third does not coordinate. NMR studies showed that there is a rapid exchange between amine N atoms, demonstrating that the specific N atoms bonded to Ga vary. Fischer has also published work showing that (APA)2InBr could be converted to (APA)2InN3 and used as a precursor to indium nitride [56]. ABZ {)[(2-Me2NCH2)C6H4])} is an excellent and often used example of an aromatic N,C ligand. Two equivalents of this chelate will combine with AlCl2X (X = Cl or Me) to form CI compounds of the formula (ABZ)2AlX [54]. Of these, (ABZ)2AlMe was structurally characterized. The geometry around Al was distorted trigonal pyramidal (tau = 0.84), with the carbons equatorial and the nitrogens axial, which is the common arrangement for these compounds. ˚ ), although still within Interestingly, the NAAl bonds were fairly long (
2.2 A the accepted distance for a dative bond. When (ABZ)2AlCl is combined with sodium azide, a simple salt elimination occurs and the chloride is replaced by an azide [57]. The structure is comparable to the chloride derivative, with the azide replacing the chloride in an equatorial position. The azide, while itself linear, is characteristically bent in its attachment to Al (Al-N-N angle 139.5).

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ABZ has also been combined in a 2 : 1 ratio with GaCl3 and InCl3 to form fivecoordinate structures of the formula M(ABZ)2Cl [17]. Both compounds had a distorted trigonal bipyramidal geometry, with the N atoms axial and the two carbons and Cl equatorial. Three equivalents of ABZ were also combined with GaCl3 in the hope of generating a six-coordinate structure. The coordination sphere of gallium proved to be too small for this to occur, however, and the result was the compound Ga(ABZ)3. This structure is similar to the earlier one, except that the chlorine has been replaced by the carbon of a third ABZ molecule, while the amine associated with this third substituent does not bond. Once again, there was a rapid exchange among the amines. The geometry was a distorted trigonal bipyramidal. Closely related to N, C chelates would be P, C chelates, such as Li[(2-Ph2PCH2)C6H4], three equivalents of which will react with AlCl3 to form the five-coordinate CI compound Al[(2-Ph2PCH2)C6H4]3 [58] (Fig. 11c). As with APA and ABZ three molecules of chelate bind to the metal, two actually chelating and the third acting as a free ligand. The structure is distorted trigonal bipyramidal (tau = 0.66), with the three carbons equatorial and the phosphorus atoms axial, while the third chelate phosphorus does not bind the metal. A third common type of ligand, MME (2-mercaptomethylethanol), has an S,O bonding scheme [28]. When combined with Al(t-Bu)3, Al(i-Bu)3, Al(Et)3, and Al(Me)3 five-coordinate BII compounds were produced for all but the t-Bu derivative. In the methyl variant the covalent and bridging OAAl bonds were ˚ vs. 1.89 A ˚ ). The dative bond, however, was very similar in length (1.84 A ˚ noticeably longer (2.95 A) than normal. This is presumed to be partly due to the larger size of a sulfur atom compared to an oxygen and also due to the slight difference in hybridization of a datively bonding O (nearly sp2 and therefore almost planar) compared to a datively bonding S (sp3 and pyramidal), a difference that results in greater steric interactions between the alkyl on the sulfur and the groups on the aluminum. MMPH (2-mercaptomethylphenol) is an example of an aromatic S,O ligand [59]. This can be combined with Al(i-Bu)3 and Al(Me)3 to form two fivecoordinate BII compounds. The two complexes were very similar structurally and quite close to the aliphatic S,O compounds. There was, however, a noticeable difference between the bridging and covalent OAAl bonds (covalent ˚ , bridging 1.96–1.98 A ˚ ). Also, the dative bond was approximately 1.86–1.87 A ˚ ) than that seen in Barron’s compound, this is slightly shorter (2.7–2.8 A probably due to resonance from the aromatic ring giving sulfur increased sp2 character. Along with S,O ligands, the related S,S and Se,Se ligand combinations are also known. The deprotonated form of the chalcogen ligands i-Pr-DPT (tetraisopropylimidodiphosphine thiolate), i-Pr-DPS (tetraisopropylimidodiphosphine selenate) and Ph-DPS (tetraphenylimidodiphosphine selenate), when combined with InCl3 form the CI compounds InCl[i-Pr-DPT]2, InCl [i-Pr-DPS]2, and InCl[Ph-DPS]2 [60]. In all three cases, the structures were distorted trigonal bipyramidal (tau is 0.81, 0.77, and 0.69, respectively), with the chloride and one chalcogen (the covalently bonded one) from each chelate

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equatorial and one chalcogen (the dative one) from each chelate axial. Not surprisingly due to selenium’s larger size, the Se-In bonds were slightly longer than the S-In ones, but beyond that the three structures were very similar. One example is shown in Fig. 11d. There is also a unique five-coordinate CI structure involving a C, C chelate, the cyclopentadienyl ligand [61]. When AlMeCl2 is combined with mangocene in a 1 : 1 ratio, the resulting compound is AlMe(Cp)2. This is a five coordinate structure, with both Cp groups bonding through their g2 mode, a somewhat unusual occurrence in its own right. The compound has a distorted square pyramidal geometry, with a tau value of 0.18. The AlAC bonds varied ˚ ) bond and significantly in length, with each Cp having one longer (
2.18 A ˚ ˚ for the one slightly shorter, more variable length bond (2.113 A and 2.163 A two Cps, respectively). A similar compound, with an ethyl in place of the methyl on Al was also synthesized but no structure was reported. An example of a Type BIII compound is obtained by combining the EDBP (2,2-ethylidene-bis(4,6-di-t-butylphenol)) ligand with AlMe3 [62]. The fourcoordinate dimer [(l-EDBP)AlMe]2 results. This was then combined with benzyl alcohol to yield [(EDBP)Al(l-OBz)]2, which further consumed 2 equivalents of benzaldehyde to give the five-coordinate BIII species [(PhCHO)Al(EDBP)(l-OBz)]2. This compound has two covalent OAAl bonds from the chelate, two bridging bonds from the BzO groups, and a dative bond from benzaldehyde. Although this is significantly different in ligand arrangement from the compounds discussed above, the bond lengths are very similar to those found in the [(BHT)Al(MEL)(l-MEL)]2 system [63]. For the two ˚ , one bridging bond covalent chelate bonds, the distance is approximately 1.7 A ˚ ˚ , and the purely (the more covalent) is 1.8 A, the other (more dative) is 1.9 A ˚ dative benzaldehyde-Al bond is 2.0 A, at the short end of the spectrum for the dative bonds so far studied. This compound also is active as a catalyst for the ring opening polymerization of lactones. Another example of a BIII compound is offered by the ligand VAPOL (‘‘vaulted’’ bisphenanthrol), which combines with AlCl3 to form a threecoordinate complex, with no dative interactions. This compound will catalyze Diels-Alder reactions [64]. A five-coordinate compound, with two dative carbonyl bonds attached to the aluminum, has been postulated as the active intermediate in the catalytic reaction. 3.3 One Tridentate Ligand (Class D)

Although relatively more rare than the ubiquitous compounds observed between group 13 elements and bidentate ligands, there are many example of five-coordinate metal complexes formed from tridentate ligands. These ligands can largely be divided into two varieties: crown ether ligands which have highly variable coordinating ability and traditional bidentate ligands with a third bonding atom appended at the periphery of the ligand. Both types are well established and tend to follow the same general trends seen for the bidentate ligands above. That is, with low steric bulk on the chelate or the

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Fig. 12. The types of compounds which may be envisioned to incorporate a tridentate ligand

metal, the compounds are monomeric. With high steric bulk, and/or narrow chelate bite angles, dimeric compounds may result. The possible ligand combinations giving the tridentate types are shown in Fig. 12. There are a number of reports in the literature of simple crown ethers forming five-coordinate structures with group 13 metals. For example, the first neutral ligand to stabilize the fragment AlMeþ 2 was 18-crown-6 [65] (Fig. 13a). A depiction of this cationic compound is shown in Fig. 13a. The DI compound was derived from the reaction of one equivalent of Cp2TiCl2, one equivalent of 18-crown-6, and two equivalents of AlMe3 and shows the metal cation chelated by three of the crown ether oxygens. The geometry around aluminum is distorted trigonal bipyramidal (tau value could not be calculated with the data provided), with one oxygen and the two carbons equatorial and the remaining two oxygens axial. As a means of assessing the distortion in the geometry, the C-Al-C angle is
140. The equatorial OAAl bond is shorter ˚ ) than the axial ones (2.181 A ˚ and 2.435 A ˚ ), the latter being one of the (1.929 A longest OAAl bonds on record. Another derivative having the same fivecoordinate aluminum cation [AlMe2(18-C-6)(AlMe3)][AlMeCl3] was subsequently reported [66]. An important, more populous, subclass among the crown ether compounds are aza-crown ethers, which incorporate nitrogens in place of all or some of their oxygen atoms. One common diaza-crown ether is DA18C6 (diaza18-crown-6). This ligand will react with three equivalents of GaMe3 to form GaMe2DA18C6(GaMe3)2, the first structurally characterized organogalliumheterocrown ether compound [67]. This compound contains three gallium centers, two four-coordinate and one five-coordinate. The four-coordinate galliums are clearly Lewis acid/base adducts between the GaMe3 molecules and the lone pairs on the heteroatom. The five-coordinate DII gallium is structurally very similar to the crown ether structures described above. The geometry is an intermediate between trigonal bipyramidal and square pyramidal, although the tau value of 0.37 would seem to indicate the square pyramidal geometry is favored. However, a careful study of the structure could

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Fig. 13. Examples of compounds containing tridentate, Class D, ligands

also lead one to argue (as Robinson did) that the covalent bonds (two GaAMe and GaAN) form an equatorial plane and the dative (GaAO) bonds are axial, in line with a trigonal bipyramidal geometry. Two other points are worth ˚ and 2.450 A ˚) mentioning. The GaAO bonds are exceptionally long (2.278 A and the crown ether ring has been forced to assume an endodentate conformation, rather than the more common exodentate. A distinctly different compound formed, however, when this ligand was combined with AlEt2Cl in a one to six ratio [68]. The resulting compound* is a the cation/anion pair [(EtAl)2DA18C6]2+ 2[EtAlCl3]2), in which the cationic section contains two five-coordinate aluminum atoms. Both are in a nearly perfect square pyramidal geometry (tau = 0.09), with two oxygens datively bonded to each and the two nitrogens bridging between them. These four bonds form the equatorial plane, with the remaining ethyl axial. The authors suggest that this * Technically tetradentate Class E, but included here because the ligand is already under discussion.

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unusual structure resulted from disproportionation of the AlEt2Cl prior to reaction with the crown ether. Another five-coordinate structure, AlMe2N18C6AlMe3, very similar to GaMe2DA18C6(GaMe3)2, was made by combining HN18C6 (aza-18-crown-6) with 2 equivalents of AlMe3 [69]. Once again there are two distinct metal environments, a four-coordinate aluminum and a five-coordinate aluminum, with the four-coordinate being merely a Lewis acid/base adduct of the N. The five-coordinate aluminum is in an intermediate geometry, with a tau value of 0.49. In effect, this compound is the aluminum equivalent of half of the GaMe2DA18C6(GaMe3)2 compound. An early example of a bidentate-type ligand with an extra nucleophile in the backbone is pmdien (N,N,N¢,N¢¢,N¢¢-pentamethyldiethylenetriamine) [70]. This ligand combined with H3AlNMe3 to form the cation/anion pair [pmdienAlH2]+ [AlH4]). That it can be classified as a DI type of structure is apparent from examining Fig. 13b. The cationic portion is five-coordinate, with a distorted trigonal bipyramidal geometry (tau = 0.57). The hydrogens and backbone nitrogen are equatorial, while the pendant amines are axial. The H-Al-H angle of 132 reflects the modest distortion from tbp geometry. As expected, the ˚ vs. 2.01 A ˚ ). axial AlAN bonds are longer than the equatorial one (2.16 A Another tridentate ligand is H2msdta [N-methyl-N¢,N¢¢-bis(trimethylsilyl) diethylenetriamine] which normally forms four-coordinate monomers when reacted with AlCl3 or GaCl3 in the presence of strong base [71]. However, when combined with InCl3 under similar conditions, the resulting compound is a five-coordinate DIV dimer [(msdta)InCl]2. Unlike most of the compounds in this study, the geometry around the In atoms is a distorted square pyramid (tau = 0.18), with four nitrogens (two covalent, two bridging) equatorial and ˚ and the Cl axial. The two ‘‘covalent’’ NAIn bonds are slightly long (2.28 A ˚ 2.29 A), but still within reason for this system. When H2msdta is combined with AlMe3, the initial product is the five-coordinate monomer (Hmsdta) AlMe2, characterized only by NMR. Under reflux, this compound converted to the four-coordinate species (msdta)AlMe. When the related four-coordinate complex (mpdta)AlCl (mpdta is msdta with isopropyl groups in place of the trimethylsilyl groups) was combined with HCl, the result was also a fivecoordinate DIII compound (Hmpdta)AlCl2. This compound possesses the more common distorted trigonal bipyramidal geometry (tau = 0.84), with the two terminal N atoms and a Cl equatorial and the backbone N and second Cl axial. This geometry is distinct from the other structures studied in that one of the formally ‘‘dative’’ nitrogen atoms is equatorial and a formally ‘‘covalent’’ chloride is axial. This change is due to the compound’s creation from an existing complex, rather than from a metal trichloride. The original fourcoordinate compound had a trigonal pyramidal geometry, leaving open one axial site that was occupied by the chlorine. This compound was predicted to be a strong Lewis acid and to readily bind a fifth substituent [72]. Every bond around the metal is a different length, with the equatorial bonds being generally ˚ , the shorter, as expected. Thus, the equatorial ‘‘covalent’’ NAAl bond is 1.819 A ˚ , and the axial ‘‘dative’’ equatorial ‘‘dative’’ NAAl bond is longer at 2.040 A ˚ . The equatorial AlACl bond is also slightly NAAl is longer yet, at 2.130 A shorter than the axial one.

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A related ligand is DMAMP {bis[(dimethylamino)methyl]pyrrole)}, which is simply the ligand MAMP discussed earlier with a second methylamine arm attached. When combined with AlCl3, LiDMAMP will form the fivecoordinate DII structure (DMAMP)AlCl2 (Fig. 13c) [48]. As expected, the structure has an intermediate geometry comparable to its bidentate predecessor (tau = 0.46). This compound will further react with two equivalents of MeLi to form its methyl-aluminum counterpart, (DMAMP)AlMe2. Structurally, this compound was very comparable to the chloride derivative, with an ambiguous geometry (tau = 0.40) and a similar bonding arrangement. When comparing the two structures, it is interesting to note that the Cl-Al-Cl angle of 111.17 is narrower than the Me-Al-Me angle of 116.54, as predicted by Bent’s rule. The chloride derivative will also react with LiAlH4 in a 1 : 2 ratio to form a unique tetranuclear bridged hydride compound [AlH2(DMAMP)AlH3]2, the first of its type reported for aluminum [73]. The deuterium version of this compound was also synthesized and structurally characterized. The structure reveals that DMAMP is now bridging between two aluminums, with one aluminum bonding to the pyrrole (covalent) N and one methylamine N and the second methylamine arm attached to a different aluminum.* The geometry around each aluminum is trigonal bipyramidal (tau = 0.76, 0.82), with a methylamine and a bridging deuterium as the axial ligands and either three more deuterium or two deuterium and covalent N as the equatorial substituents. It is also interesting to note that the two of the bridging deuterium atoms (the ones between the aluminums also bridged by the same DMAMP) are symmetric, with the deuterium approximately halfway between the two metal centers. The other two bridging deuteriums (between aluminums attached to different DMAMP molecules) are asymmetric, with the deuterium lying closer to the aluminum with the pyrrole NAAl bond. This unique structure was quite moisture sensitive and, upon contact with water, would decompose into the five-coordinate monomer (DMAMP)AlH2, which was structurally very similar to the chloride and methyl derivatives. Another unique five-coordinate structure is formed when mercaptoethyl ether reacts with AlMe3 [74]. The dithiol undergoes an interesting rearrangement, in which an OAC bond breaks and is replaced by a SAC bond, resulting in a sulfur and the oxygen appearing to switch places in the molecule, the thiol becomes a thioether and the ether becoming a terminal oxide. This then forms a dimeric DIII type of compound (see Fig. 13d) after combining with AlMe3. The oxygens bridge between the aluminum atoms, the backbone sulfur forms a dative bond to Al, and the terminal sulfur forms a covalent bond. Unsurprisingly, the backbone S-Al distance was significantly longer than the ˚ vs.
2.6 A ˚ ), while the AlAO bond lengths terminal S-Al distance (
2.2 A were as expected for bridging alkoxides. A tau value of 0.57 (average for both aluminum atoms) indicates a close adherence to tbp geometry. * While the authors recognize that this structure is technically a mixed Type A and Type BII structure in our classification scheme, its intimate connection to two tridentate compounds led us to include it here.

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A different class of potentially tridentate ligands are thiosemicarbazones, of which gallium thiosemicarbazones are of interest as nuclear imaging or anticancer agents [75]. A series of these compounds were prepared by utilizing acetylpyridinethiosemicarbazones [RNHC(S)NHN@CMe(o-pyridine) where R = Me, i-Pr, or Ph] in conjunction with AlMe3 and GaMe3. Of the resulting six compounds, the one resulting from addition of i-Pr-thiosemicarbazone to two equivalents of AlMe3 was structurally characterized. The structure shows two different environments around the aluminum atoms: one is fivecoordinate DIII, bound by a thiosemicarbazone nitrogen and sulfur and the pyridine nitrogen, and the other is four-coordinate, bound by the other two thiosemicarbazone nitrogens. The five-coordinate aluminum possesses the expected distorted trigonal bipyramidal geometry (tau = 0.62), with the two methyls and the thiosemicazone nitrogen equatorial while the sulfur and pyridine nitrogen are axial. It is interesting to note that the sulfur, formally a covalent binder in this compound, is axial, while the covalent molecules tend to lie in the equatorial plane. This appears to be due to the geometry of the ligand, rather than any unusual electronic effects. Another five-coordinate thisosemicarbazone compound (Fig. 13e) results from the combination of PhNHC(S)NHN@CMe(o-pyridine) with three equivalents of InMe3 [76]. This structure is unusual in that it is a trimer with three five-coordinate indium centers. Each thiosemicarbazone binds one indium with three nitrogens (the pyridine nitrogen, one backbone nitrogen, and the terminal nitrogen) in a Type DIII manner, while both ligand’s sulfur and other backbone nitrogen bind a third aluminum in a CI manner. The geometry of the central (CI) indium is 0.31, while the other two (DIII) indiums are 0.03. This type of ligand has also been found to produce five-coordinate compounds in which it acts as a tetradentate ligand. 3.4 One Tetradentate Ligand (Class E)

3.4.1 Salen Ligands Five-coordinate salen group 13 alkyl complexes are conveniently prepared by combining the ligand with a trialkyl group 13 reagent in non-oxygenated solvents at ambient temperatures. Alkyl derivatives of aluminum [77–79] gallium [80], and indium [81, 82] are known. The related group 13 halide compounds, SalenMX, are prepared similarly, but with R2AlCl as the group 13 reagent [83, 84]. There are many derivatives of these two types of starting materials, including siloxide [78, 85], alkoxides, azides, amides, and bridgedoxide compounds [86]. The geometry of the compounds is dictated by the nature of the ligand backbone and occasionally by steric factors (as discussed below). Specific details of the structures of the group 13-Salen compounds were reported in a recent review and so will not be repeated here. Rather, a basic understanding of the structures that are observed, in relation to the tau values, will be addressed.

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In general the compounds contain a five-coordinate group 13 atom (Al, Ga, In) that is either tbp (tau value
0.7) or sqp (tau value
0) depending on the nature of the connection between the two nitrogens of the ligand (the ligand ‘‘backbone’’). With more flexible backbones, (CH2)n n > 2, a tbp geometry is obtained probably due to the fact that a sqp geometry would cause the methylene hydrogens to be eclipsed. In the tbp geometry they are staggered. With an ethyl, (CH2)2, or o-aryl backbone the metal adopts a sqp geometry. This ligand-geometry correlation is a fundamental property of the Salen ligands and is observed in all of the other derivatives (see below). The compounds having square pyramidal geometries are fairly consistent with tau values near zero. This is observed, for instance, in salomphen(t-Bu)AlMe3 [13]. However, there is a significant amount of deviation in the trigonal bipyramidal geometries. Such deviations, may be observed in the representative structures of salpen(t-Bu)AlCl (tbp) and salen(t-Bu)AlMe (distorted tbp). Salpen(t-Bu)AlCl contains a central Al in a distorted tbp geometry (tau = 0.77) (Fig. 14a). The axial atoms are N1 and O2 with the equatorial atoms, Cl1, N2, O1. The greatest deviations from ideal geometry occur in the N2-Al1-O1 angle, 126.3(3) and the N2-Al1-Cl1 angle, 112.3(3). The N1-Al1-O2 angle is 172.3(3). The difference in the distances to equatorial versus axial ˚ and Al1-O2ax = 1.815(6) A ˚ ] are in keeping atoms [eq. Al1-O1eq = 1.754(6) A with the trend for transition metal complexes in which the axial distances are notably longer than equatorial (to the same type of atom). An unusual distorted tbp geometry is observed in salen(t-Bu)AlMe. Although the general precedent is for the Salen ligands to enforce an sqp geometry, as observed in the salcen, salophen, and salomphen complexes [87], it may be best to consider the geometry as an approximation of tbp. This is borne out by a tau value of 0.47. There are some serious deviations from the ideal angles. The N(1)-Al(1)-O(2) angle is 158.7(5) while the N(2)-Al(1)-O(1) angle is 130.6(5). The remaining angles range from 76.4(4) for N(1)-Al(1)-N(2) to 116.9(5) for O(1)-Al(1)-C(33). The axial atoms, N(1) and O(2) form some˚ and 1.831(9) A ˚ , respectively] than the what longer bonds to Al(1) [2.069(10) A ˚ ˚ , respectively]. equatorially disposed N(2) and O(1) [2.025(11) A and 1.779(9) A This also occurs in tbp compounds. In a previous review it was predicted that ‘‘the reactivity of five-coordinate complexes would certainly be different than that of the four-coordinate derivatives [88]’’. However in the limited number of reactions of four- and fivecoordinate aluminum alkyls with silanols there appears to be little difference. 3.4.2 Aminotroponimates The alkylene-bridged aminotroponiminate ligands (TP) are very similar to the salen ligands and lead to distorted square pyramidal geometries around the group 13 metals [89]. In one example, (TP)InCl (shown in Fig. 14b), the ˚ . A signifitau value is 0.40. The In-N distances are fairly consistent at
2.2 A cant deviation in the sqp geometry occurs in the ‘‘front’’ N-In-N angle, at 112.

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Fig. 14a–g. Examples of tetradentate, Class E, ligands

This is very obtuse by comparison to the angle formed by the two nitrogens connected by the tether, 90. 3.4.3 Bis(amino)thiols Replacement of the OH groups of the salen ligands with SH produces a potentially tetradentate ligand, BAT-TM [tetramethylbis(aminoethanethiol)]

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[90]. The resulting group 13 compounds, (BAT-TM)M(X) (with M, X = Ga, Cl; In, Cl; In, SCN) adopt distorted sqp geometries analogous to the SalenMX derivatives described above. An example of this type of compound is shown in Fig. 14c. The tau values for these three compounds are all close to zero. The primary difference in these structures is a lengthened M-S distance (M = Ga, ˚ ; In,
2.4 A ˚ ) by comparison to the M-O distance in the Salen
2.3 A compounds. 3.4.4 Porphyrins Many group 13 compounds have been prepared with porphyrins. The majority of these compounds were created to serve as models for the active sites of enzymes, such as cytochrome c oxidase. Additionally, the gallium compounds are more robust than their iron counterparts. To a much lesser extent the Por-AlX compounds are used as oxirane polymerization catalysts [8]. The five-coordinate Por-MX compounds are in sqp geometries, with the porphyrin occupying the basal plane. Examples of this geometry include, OEP-AlMe (OEP = octaethyl porphyrinate) [91], and TPPGa-R (R = allyl, 2-methylallyl, TPP = tetra(o-tolyl) porphrinate) and TAPGa-vinyl [TAP = tetra (p-anisyl) porphyrinate] [92]. The latter is depicted in Fig. 14d. There is not a great deal of variability in the bonding for these compounds. Typical M-N ˚ . If the fifth site is occupied by groups capable of distances are around 2.0 A existing as separate anions, six-coordinate cationic Por-M(base)2 compounds form in the presence of coordinating solvents such as water and THF [93]. The tau values for these compounds are effectively zero. The greatest structural difference for the group 13 congeners is the displacement of the metals above ˚ , Ga
0.58 A ˚ , In
0.6 A ˚ ). the N4 plane (Al
0.5 A 3.4.5 Phthalocyanines A square pyramidal geometry around a central group 13 metal is also enforced by phthalocyanines (Pc). These are of the general formula PcMX (where M = Al, Ga, In, and X = alkyl or halide). In the structure of PcInI, the indium ˚ above the N4 plane (Fig. 14e) [94]. atom is 0.74 A 3.4.6 Tripodal Ligands In combination with the group 13 elements, tripodal (3)) tetraamine ligands (sometimes called ‘‘atranes’’) form a unique class of compounds named ‘‘azatranes’’ [95]. These ligands enforce an almost ideal tbp geometry around a central aluminum or gallium atom. The fifth, apical, coordination sites in these compounds are occupied by a bridging nitrogen from a second molecule. The axial angles for the Al and Ga derivatives are similar, 161.9 and 162.7(1),

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respectively. Interestingly, however, the non-bridging nitrogens of the Al derivative are nearly planar while for the Ga derivative they are pyramidal, as expected. The difference was attributed to some form of p bonding between the Al and N atoms. An indium derivative was prepared with a similar ligand, S3N, having the -NR groups replaced with S: [In(S3N)2] [96]. It adopts a distorted tbp geometry. Other tripodal ligands having benzyl groups connecting the central nitrogen to either oxygen (O3N-benzyl) or sulfur (S3N-benzyl) have secondary Lewis base ligands in the apical site [97]. Two examples are Ga(O3N-benzyl)(DMF) and In(S3N-benzyl)(DMF). Being monomers, as opposed to the ligand-bridged dimers described above, they feature more ideal tbp angles. For example, the axial angles are
177. This is also reflected in the tau values of 0.92 and 0.88 respectively. 3.4.7 Open Chain Amines Robinson’s systematic early work on open-chain amines set the stage for the subsequent studies on higher-coordinate group 13 compounds, which had not previously been examined to any appreciable extent [1]. As one example, the combination of BPEDA with trimethylaluminum results in a compound having a central five-coordinate -AlMe unit and a peripheral four-coordinate -AlMe2 unit [98]. The central atom is tbp and it was noted that previous examples with rigid macrocycles featured sqp aluminum [99, 100]. Thus, it appears that when the ligand is flexible enough, a tbp geometry is preferred. Similar observations were made for crown ether-aluminum compounds [101]. A monomeric ionpair, [LAlH2][AlH4] forms when H3Al-NMe3 is combined with PMDIEN [102]. The tridentate amine forms a distorted tbp around the aluminum atom. The axial N-Al-N angle is 166. The H-Al-H angle is 132. 3.4.8 Thiosemicarbazones Many examples of five-coordinate group 13 compounds have been discovered adventitiously or during studies having other targets. For example, interesting aluminum [103] and indium [76] derivatives were discovered while developing the use of bis(thiosemicarbazones) as medical imaging agents. The resulting tetradentate indium compound is shown in Fig. 14g. There is a difference in bonding for the two metals. Indium appears to be in a distorted sqp geometry [most obtuse angle = 132.4(2) tau value = 0.14] while the aluminum derivative is almost perfectly sqp [most obtuse angle = 141.1(1) tau value = 0]. Both the S-M-N angle angles are very acute, 64.9(3) for indium and 69.7(1) for aluminum. Although the difference in the two structures is almost negligible, it appears that indium has some preference for sqp geometries rather than tbp. It is interesting to note that in reactions involving gallium the possibility of a similar coordination environment leads inevitably to fourcoordinate derivatives [104].

Compounds Containing Five-Coordinate Group 13 Elements

195

3.4.9 Salan Ligands Tetrametallic dimeric Salan compounds, for example, [salopmhanAl {Al(i-Bu)}]2 are prepared by the thermolysis of the LH(AlR)(AlR2)2 starting material [105]. The compound contains two central aluminum atoms in trigonal bipyramidal geometries. The same type of structure is observed for [LAl[Li(THF)2]2 where the [AlR2]+ units are replaced with [Li(THF)2]+ (L = salpan and salomphan) [106]. These last two sets of compounds are good examples of how ligands possessing a variety of electron-rich heteroatoms can lead to five-coordinate geometries for the group 13 elements. 3.5 Pentadentate Ligands (Class F)

Two the best of our knowledge, the literature contains no recent examples of a pentadentate ligand fully coordinated to a group 13 metal.

4 Conclusions The synthesis, characterization, and applications of five-coordinate group 13 compounds have been surveyed. The compounds were found to fall into the following Classes: Class A: with five separate ligands, Class B: containing a single bidentate ligand, Class C: with two bidentate ligands, Class D: tridentate ligand coordination, and Class E: tetradentate ligand coordination. The Class F with pentadentate ligand coordination were not found to have any representation. Also, one could think of a Class G, a little understood class of five-coordinate compounds where the metals appear in clusters. Brief mention is made of these compounds as a means of making note of this broad and difficult to categorize area. Each class of five-coordinate compound was further subdivided into the possible ‘‘Types’’ of structures that may be envisioned to occur. The majority of the compounds fell into a very few Types in each Class, and these will be noted separately below. These Types should provide a starting point for future reviews that can focus on the details of these compounds. Of key importance are the Types that had no members whatsoever. These provide clear targets for future synthetic efforts. The majority of the compounds reported in the literature contained bidentate ligand coordination and formed Class B or C compounds. These were related in that the same ligands could be used to prepare both classes. The most popular ligands were uninegative with one covalent bonding group and one coordinate covalent (‘‘dative’’) group. As a result, Types BII and CI predominated. Representatives of all three B Types were discovered. Most of the BII compounds were dimeric with bridging groups. Type BI compounds appear to have the most applicability, being invoked as the primary Lewis acid-base complexes forming during organic transformations. Only Type CI compounds were found in the C class.

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Table 1.

References

Compounds

s

[45] [102] [90]

SaleanAl-i-Bu(Al-i-Bu2)2 (MeAl){CH2[C(CMe)-NNC(S)NMe]2}(AlMe2)2 (1) [GaCl(BAT-TM)] (1) [InCl(BAT-TM)] (2) [InSCN(BAT-TM)] (3) InPcI [(i-Pr2-ATI)Al(C6F5)(l-H)2B(C6F5)2] (3) [Ga3(t-Bu)5(neol)2] (7) [Ga2Al(t-Bu)5(neol)2] (10) [Et(Cl)AlOR] OR = dimethylamino-1-propanol (1)

0 0 0 0 0.2 0.02 0.02 0.03 0 Al(1) Al(2) Al(2) Al(4) 0.05 0.06 0.04 0.08 0.09 0.11 0.20 0.44 0.17 0.16 0.43 0.18 0.19 In(1) In(2) 0.20 0.48 0.36 0.25 0.26 0.43 0.22 0.25 In(1) In(2)

[94] [53] [37] [15]

{[MeAl(OR*)2AlMe2]+ [MeAlCl3]) OR* = Dpm (4) [52] [43] [47] [68] [80]

[36] [61] [87] [35] [28]

[82]

[76]

[55] [67] [48] [89] [41] [31]

[GaMe(PhNCPhNPh)2] (3) Salpan(AlMe)(AlMe2)2 (4) Salophan(AlMe)(AlMe2)2 (5) Methylbis(salicylaldoximato-O1,N)gallium(III) [(EtAl)2.diaza-18-crown-6]2+ Salophen(t-Bu)GaCl (7) Salomphen(t-Bu)GaMe (9) Salen(t-Bu)GaEt (6) Salophen(t-Bu)GaEt (10) [Al2(t-Bu)3(OCH2CH2O)(OCH2CH2OH)] (1) [Al3(t-Bu)5(OCH2CH2O)2] (2) [(C6H5)2AlCH3] (1) Salomphen(t-Bu)AlMe (5) [InMe2(ON = CHC5H4N)]2.1/2 C6H6 [(t-Bu)2Al(l-OCH2CH2OMe)]2 (1) [(i-Bu)2Al(l-OCH2CH2OMe)]2 (2) [Me2Al((l-OCH2CH2CH2OMe)]2 (7) [(t-Bu)2Al(l-OCH2CH2SMe)]2 (13) Salen(t-Bu)InCl (1) Salpen(t-Bu)InBr (6) Salophen(t-Bu)InEt (11) Salen(t-Bu)InMe (13) {(Me2In)2[NC5H4Cme-NNC(S)NC6H5]2}(InMe) Æ THF (3) {((THF)Me2In)2[CH2(Me-CNNC(S)NMe)2]} (InMe) (4) [(CH3)2N(CH2)3]2InBr (1) [Ga(CH3)2][C12H25N2O4][Ga(CH3)3]2 AlCl2{C4H2N(CH2Nme2)2-2,5} (1) AlCl{C4H3NCH2NMe2)-2}2 (2) Al(Me)2{C4H2N(CH2NMe2)2-2,5} (3) [{(i-Pr)TP}InCl] (2) [(i-Bu)2Al(l-O-8-C9H6N)]2 (5) [(l-OCH2CH2OPh)AlCl2]2 (1) [(l-OCH2CH2OPh)AlBr2]2 (2)

0.04 0.17 0.28 0.21

0.19 0.38

0.31 0.03

In(1) 0.14 0.37 0.37 0.40 0.45 0.40 0.40 0.45 0.46 0.45

197

Compounds Containing Five-Coordinate Group 13 Elements

References

Compounds

s

[78]

[70] [30] [74]

Salen(t-Bu)AlMe (1) Salomphen(t-Bu)Al-i-Bu (10) Salen(t-Bu)AlOSiPh3 (11) Salpen(t-Bu)AlOSiPh3 (12) Me2Al[N18C6]AlMe3 [(BHT)Al(H)(l-OCH2CH2NMe)]2 (2) [(BHT)Al(OCH2CH2OMe)(l-OCH2CH2OMe)]2 (6) [H2Al(pmdien)]+[AlH4]) (5) [Me2Al(l-OCH2CH2OMe)]2 [Al2{l, g3-(OCH2CH2SCH2CH2S)}2(CH3)2] (8)

[95]

[Al(MeCH2CH2)3N]2 (12)

0.47 0.14 0.40 0.74 0.49 0.50 0.36 0.57 0.58 Al(1) 0.58 Al(2) 0.56 (Al1) 0.59 (Al2) 0.58 (Ga) 0.71 (Ga¢) 0.70 0.62 0.66 0.71 0.65 0.67 0.71 0.86 0.72 0.76 Al(1) 0.76 Al(2) 0.82 0.77 0.18 0.46 0.15 0.40 0.74 0.74 0.17 0.17 0.55 0.37 0.81 0.77 0.69 0.84 0.18 0.84 0.85 0.83 0.75 0.92 0.88 1

[69] [32]

[Ga(MeCH2CH2)3N]2 (14) [75] [58] [50] [57] [62] [49] [73] [85]

[60] [71] [54] [46] [97] [17]

(Me2Al)[NC5H4(MeNNC(S)NC3H7](AlMe2) (3) Al[o-(Ph2PCH2)C6H4]3 (2) {MeC(N-i-Pr)2}2AlCl (9a) (t-BuC(N-i-Pr)2}2AlCl (10a) [Me2N(CH2)3]2AlCl (1a) [Me2N(CH2)3]2Al(N3) (1b) [2-(Me2NCH2)C6H4]2Al(N3) (2b) [(PhCHO)Al(EDBP)(l-OBz)]2 (2) [PhC(NSiMe3)2]2AlH (2) {AlD2[C4H2N(CH2NMe)2-2,5]AlD3}2 (2D) Salpen(t-Bu)AlCl (2) Salomphen(t-Bu)AlCl (3) Salen(t-Bu)AlOSiMe3 (9) Salomphen(t-Bu)AlOSiMe3 (11) Salen(t-Bu)AlOSiPh3 Salpen(t-Bu)AlOSiPh3 Salpen(t-Bu)AlOSiPh3 Salophen(t-Bu)AlOSiPh3 Salomphen(t-Bu)AlOSiPh3 (12) Salcen(t-Bu)Al{N(SiMe3)2} (13) (tbumSalcen)AlCl InCl(N{i-Pr2PS}2)2 (1) InCl(N{i-Pr2PSe}2)2 (2) InCl(N{Ph2PSe}2)2.CH2Cl2 (3) {[(i-PrNHCH2CH2)(i-PrNCH2CH2)NMe]AlCl2} (15) {[(Me3SiNCH2CH2)2NMe]InCl}2 (8) Bis{2-[(dimethylamino)methyl]phenyl}methylaluminum (5) [MeAl(Dpm)2] (1) [EtAl(Dpm)2] (2) [1-NorAl(Dpm)2] (3) [In(S3N)(DMF)] [Ga(O3N)(DMF)] [In(CBCCF3)2(THF)2] (20)

198

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Selected tau values are listed in Table 1. They range from zero for sqp compounds (Class E, for example) to one for tbp compounds (Class A, for example). There is no pattern to how these tau values vary between these two extremes. Rather, there is a continuum of values between 0 and 1. However, a value of 0.50 does not represent an ambiguous case. Even tau values of 0.40 appear more tbp than sqp. An example of this is given in Fig. 15. Also note that electronic factors (the difference in Me or Cl) have no effect, the ligand controls the geometry.

Fig. 15. Identical tau values are observed for both the Me and Cl derivatives of this DI compound

The ultimate goal of this work was to provide a systematic understanding of the types of five-coordinate compounds that have been made and that may be made in the future. As can be surmised from the foregoing, there is still a great deal of work that needs to be done both in understanding how these various compounds form, and in preparing the ‘‘missing’’ members of the various Classes. It is hoped that this review will provide a background and a stimulus for these continued efforts.

5 References 1. Wilke G (1995) In: Fink G, Mullhaupt R, Brintzinger HH (eds), Ziegler Catalysts. Springer, Berlin, Heidelberg, New York, chapter 1 2 Robinoson GH (1993) In: Robinson GH (ed), Coordination Chemistry of Aluminum. VCH, New York, chapter 2 3. Atwood DA, Harvey MJ (2001) Chem Rev 101: 37 4. Dzugan SJ, Goedken VL (1986) Inorg Chem 25: 2858 5. Gurian PL, Cheatham LK, Ziller JW, Barron AR (1991) J Chem Soc Dalton Trans 1449 6. Atwood DA, Jegier JA, Rutherford D (1995) J Am Chem Soc 117: 6779 7. Vincens V, Le Borgne A, Spassky N (1989) Makromol Chem Rapid Commun 10:623; Le Borgne A, Vincens V, Jouglard M, Spassky N (1993) Makromol Chem Macromol Symp 73: 37 8. Aida T, Inoue S (1996) Acc Chem Res 29: 39 9. Miller JA (1993) In: Downs AJ (ed), Chemistry of Aluminum, Gallium, Indium and Thallium. Chapman and Hall, London, chapter 7 10. Ovitt TM, Coates GW (1999) J Am Chem Soc 121: 4072; Radano CP, Baker GL, Smith III MR (2000) J Am Chem Soc 122: 1552 11. Taden I, Kang H-C, Massa W, Okuda J (1997) J Organomet Chem 540: 189

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57. Muller J, Fischer RA, Sussek H, Pilgram P, Wang R, Pritzkow H, Herdtweck E (1998) Organometallics 17: 161 58. Muller G, Lachmann J, Rufinska A (1992) Organometallics 11: 2970 59. Hendershot DG, Barber M, Humar R, Oliver JP( 1991) Organometallics 10: 3302 60. Darwin K, Gilby LM, Hodges PR, Piggott B (1999) Polyhedron 18: 3729 61. Fisher JD, Wie MY, Willett R, Shapiro PJ (1994) Organometallics 13: 3324 62. Ko B, Lin C (1999) Macromolecules 32: 8296 63. Francis JA, Bott SG, Barron AR (2000) J Organomet Chem 597: 29 64. Heller DP, Goldberg DR, Wulff WD (1997) J Am Chem Soc 119: 10551 65. Bott SG, Alvanipour A, Morley SD, Atwood DA Means CM, Coleman AW, Atwood JL (1987) Angew Chem Int Ed Engl 26: 485 66. Atwood JL, Bott SG, Harvey S, Junk PC (1994) Organometallics 13: 4151 67. Lee B, Pennington WT, Robinson GH (1990) Organometallics 9: 1709 68. Self MF, Pennington WT, Laske JA, Robinson GH (1991) Organometallics 10: 36 69. Pajerski AD, Cleary TP, Parvez M, Gokel GW, Richey HG Jr (1992) Organometallics 11: 1400 70. Atwood JL, Robinson KD, Jones C, Raston CL (1991) J Chem Soc Chem Commun 1697 71. Emig N, Nguyen H, Krautscheid H, Reau R, Cazaux JB, Bertrand G (1998) Organometallics 17: 3599 72. Hutchison A, Atwood DA. Inorg Chem Commun (in review) 73. Chang JC, Hung CH, Huang JH (2001) Organometallics 20: 4445 74. Janas Z, Jerzykiewicz LB, Przybulak S, Richards RL, Sobota P (2000) Organometallics 19: 4252 75. Paek C, Kang SO, Ko J, Carroll PJ (1997) Organometallics 16: 2110 76. Paek C, Kang SO, Ko J (1997) Organometallics 16: 4755 77. Dzugan SJ, Goedken VL (1986) Inorg Chem 25: 2858 78. Atwood DA, Hill MS, Jegier JA, Rutherford D (1997) Organometallics 16: 2659 79. Leung W-H, Chan EYY, Chow EKF, Williams ID, Peng S-M (1996) J Chem Soc Dalton Trans 1229 80. Hill MS, Atwood DA (1998) Eur J Inorg Chem 67 81. Atwood DA, Jegier JA, Rutherford D (1997) Bull Chem Soc Jpn 70: 2093 82. Hill MS, Atwood DA (1998) Main Group Chem 2: 191 83. Rutherford D, Atwood DA (1996) Organometallics 15: 4417 84. Davidson MG, Lambert C, Lopez-Solera I, Raithby PR, Snaith R (1995) Inorg Chem 34: 3765 85. Munoz-Hernandez M-A, Keizer TS, Wei P, Parkin S, Atwood DA (2001) Inorg Chem 40: 6782 86. Gurian PL, Cheatham LK, Ziller JW, Barron AR (1991) J Chem Soc Dalton Trans 1449 87. Wang Y, Parkin S, Atwood D (2002) Inorg Chem 41: 558 88. Oliver JP, Kumar R (1990) Polyhedron 9: 409 89. Burgstein MR, Euringer NP, Roesky PW (2000) J Chem Soc Dalton Trans 1045 90. Zheng YY, Saluja S, Yap GPA, Blumenstein M, Rheingold A, Fracesconi LC (1995) Inorg Chem 35: 6656 91. Guilard R, Zrineh A, Tabard A, Endo A, Han BC, Lecomte C, Souhassou M, Habbou A, Ferhat M, Kadish KM (1990) Inorg Chem 29: 4476 92. Arasasingham RD, Balch AL, Olmstead MM, Phillips SL (1997) Inorg Chim Acta 263: 161 93. Serr BR, Headford CEL, Anderson OP, Elliott CM, Spartalian K, Fainzilberg VE, Hatfield WE, Rohrs BR, Eaton SS, Eaton GR (1992) Inorg Chem 31: 5450 94. Janczak J, Kubiak R (1999) Inorg Chim Acta 288: 174 95. Pinkas J, Wang T, Jacobsen RA, Verkade JG (1994) Inorg Chem 33: 5244 96. Rose DJ, Zubieta J, Fischman AJ, Hillier S, Babich JW (1998) Inorg Chem Commun 1: 164 97. Motekaitis RJ, Martell AE, Koch SA, Hwang J, Quarless DA Jr, Welch MJ (1998) Inorg Chem 37: 5902

Compounds Containing Five-Coordinate Group 13 Elements

98. 99. 100. 101. 102. 103. 104. 105. 106.

201

Robinson GH, Sangokoya SA, Moise F, Pennignton WT (1988) Organometallics 7: 1887 Robinson GH, Sangokoya SA (1987) J Am Chem Soc 109: 6852 Robinson GH, Moise F, Sangakoya SA, Pennington WT (1988) J Cryst Spect Res 18: 387 Atwood JL (1993) In: Robinson GH (ed), Coordination Chemistry of Aluminum. VCH, New York, chapter 6 Atwood JL, Robinson KD, Jones C, Raston CL (1991) J Chem Soc Chem Commun 1697 Paek C. Kang SO, Ko J, Carroll PJ (1997) Organometallics 16: 1503 Paek C. Kang SO, Ko J, Carroll PJ (1997) Organometallics 16: 2110 Atwood DA, Remington MP, Rutherford D (1996) Organometallics 15: 4763 Atwood DA, Rutherford D (1995) Inorg Chem 34: 4008

Author Index Volumes 101-105

Aldinger F, see Seifert HJ (2002) 101:l-58 Anitha S, Rao KSJ (2003) The Complexity of Aluminum-DANN Interactions: Relevance to Alzheimer's and other Neurological Diseases. 10479-98 Atwood DA, see Conley B (2003) 104:181-193 Atwood DA, Hutchison AR, Zhang Y (2003) Compounds Containing Five-Coordinate Group 13 Elements. 105167-201 Berend K, van der Voet GB, de Wolff FA (2003) Acute Aluminum Intoxication. 1041-58 Bohrer D, see Schetinger MRC (2003) 10499-138 Budzelaar PHM, Talarico G (2003) Insertion and 8-Hydrogen Transfer at Aluminum. 105141-165

Conley B, Atwood DA (2003) Fluoroaluminate Chemistry. 104181-193 Fruhauf S, see Roewer G (2002) 101:59-136 Haubner R, Wilhelm M, Weissenbacher R, Lux B (2002) Boron Nitrides - Properties, Synthesis and Applications. 1021 -46 Herrmann M, see Petzow G (2002) 102:47-166 Herzog U, see Roewer G (2002) 101:59-136 Hopfl H (2002) Structure and Bonding in Boron Containing Macrocycles and Cages. 103:l-56

Hutchison AR, see Atwood DA (2003) 105:167-201 Jansen M, Jaschke B, Jaschke T (2002) Amorphous Multinary Ceramics in the Si-B-N-C System. 101:137-192 Jaschke B, see Jansen M (2002) 101:137-192 JLchke T, see Jansen M (2002) 101:137-192 Linton DJ, Wheatley AEH (2003) The Synthesis and Structural Properties of Aluminium Oxide, Hydroxide and Organooxide Compounds. 105:67-139 L u x B, see Haubner R (2002) 1021-46 Mahalakshmi L, Stalke D (2002) The group 13 Organometallic Fragment Chelated by P-centered Ligands. 103:85-116 Morsch VM, see Schetinger MRC (2003) 104:99-138 Muller E, see Roewer G (2002) 101:59-136 Oshiro S (2003) A New Effect of Aluminum on Iron Metabolism in Mammalian Cells. 10459-78

204

Author Index Volumes 101-105

PatoEka J, see Strunecka A (2003) 104:139-180 Petzow G, Hermann M (2002) Silicon Nitride Ceramics. 10.247-166 Power P (2002) Multiple Bonding Between Heavier Group 13 Elements. 103:57-84 Rao KSJ, see Anitha S (2003) 104:79-98 Roewer G, Herzog U, Trommer K, Muller E, Fruhauf S (2002) Silicon Carbide - A Survey of Synthetic Approaches, Properties and Applications. 101:59-136 Schetinger MRC, Morsch VM, Bohrer D (2003) Aluminum: Interaction with Nucleotides and Nucleotidases and Analytical Aspects of ist Determination. 104:99-138 Schubert DM (2003) Borates in Industrial Use. 1051-40 Schulz S (2002) Synthesis, Structure and Reactivity of Group 13/15 Compounds Containing the Heavier Elements of Group 15, Sb and Bi 103:117-166 Seifert HJ, Aldinger F (2002) Phase Equilibria in the Si-B-C-N System. 101:l-58 Stalke D, see Mahalakshmi L (2002) 103235-116 Strunecka A, Patocka J (2003) Aluminofluroride Complexes in the Etiology of Alzheimer's Disease. 104:139-180 Talarico G, see Budzelaar PHM (2003) 105141-165 Trommer K, see Roewer G (2002) 101:59-136 Uhl W (2003) Aluminum and Gallium Hydrazides. 10541-66 van der Voet GB, see Berend K (2003) 104:l-58 Weissenbacher R, see Haubner R (2002) 102:l-46 Wheatley AEH, see Linton DJ (2003) 10567-139 Wilhelm M, see Haubner R (2002) 10231-46 de Wolff FA, see Berend K (2003) 104:l-58 Zhang Y, see Atwood DA (2003) 105167-201

Subject Index

Adamantanoid structure 78 Adhesives 19 Agostic interactions 79,145 Agriculture 33 Aksaite 9 Alanes 69,70 Aluminum 162 -,pentacoordinate 109 Aluminum alkyls 151 Aluminum ketones, reduction 155 Aluminum nitride (AIN) 42,60 Alumoxanes 76 Amidinates 150 Amines, open chain 194 Aminotroponiminates 150 Ammonioborite 13 Anisochronicity 79 Antibiotics 20 Aplasmomycin 20 Aufiau reaction 143,144 Azaindoles 70,71,103 Azides 172,173 Azobenzene 57 Barron 170 Bayerite 104 Bidentate ligands 42,169,172-175,195 Bioessentiality 20,21 Biostat 33.34 53 Bisaminothiols 192 Boehmite 73 Bond dissociation energies 80,176 Borate esters 5,6,17-19 Borate minerals 7-10,13,14,26,29-31 Borates, metamorphic 31 Borax 25,26 Boric acid 5,8,13,26-29,34 Boric oxide 27

Boromycin 20 Boron 2-5,16,21-24 Borosilicates 3,15,17,31 Boroxol ring 15,16 tert-Butylhydrazine 43,44,56 C,F, transfer 153 Cage compounds 48,50,59 Cdixarenes 98,102,124 Carbamates 129 Carbodiimides 105 Carbonyl compounds, addition 154 Carbonyls, saturated, activation 79 Carboxylates 73,105 Cell membrane 22,23 Cell wall 2 1,22 Cellulose insulation 3,34 Ceramic glazes 3,15,33 Citrates 130 Clathrates 126 Colemanite 9,10,26,29 Cossee mechanism 144 CRAMPS 73 Crown ether 185,188 Cyclohexene/carbon dioxide copolymerization 82 Cyclohexene oxide, polymerization 82 Cyclopentadienyls 94 Deprotonation, multiple 95,100,124 Detergents 3,36,37 Di(tert-butyl)aluminum chloride 44 Di(tert-butyl)aluminum hydrazide 45 Di(tert-buty1)aluminum hydride 45,52 Diaspore 73 Diastereoselectivity 158 Diethylgallium chloride 56 P-Diketones 99 Dilithium bis(trimethylsily1)hydrazide 44 Dimerization 147

Subject Index Dimethylaluminum chloride 55 Dimethylaminoalane 51 Dimethylgallium chloride 56 Dimethylhydrazine 42 1,l-Dimethylhydrazine 47,49-51 Dioxygen 173 1,l-Diphenylhydrazine 47 Disodium octaborate 28,34 Double electrophilic activation 158 j3-Elimination 145 Enantioselectivity 158 Epimerization 158 Ethene, polymerization, with Et3AI 149 Ethylene, polymerization 81,95 EXSY 70 Fertilizer 3,33 Fiberglass 3,30-32 Fire retardants 3,29,34 Fischer 183 Fluorescence 72 Flux 15 Frit 33 Frovolite 7,8 Gallium 162 Gallium nitride (GaN) 42,48,56,60 Gel 19 Gibbsite 104 Glass 3,15-17,32,33 Goedken 170 Halides 172 Helicates 125 HETCOR 70 Hexaborate 9,10 1,5-Hexadiene, polymerization 70 Hungchaoite 9 Hydrazine 42,43,45 Hydrides 94,172,173 Hydroalumination 59,61 Hydroboracite 30 Hydrogen bonds 43,s 1 P-Hydrogen transfer 142,143,146, 163 Hydrolysis 69,70 Hyperconjugation 56 Hyperconjugative interactions 53,55

Inorganic oxides 69 Inoue 170 Kernite 9,10,26,30 Lactone oligomerization 170 Ladder-type structure 52 Lewis acids, bidentate 8 1 Ligands, salen-type 71,96,106,124,132 -,tethered 98 -,tripodal 106 Lithium alanate 53 Lithium dimethylhydrazide 56 Lithium diphenylhydrazide 56 Lithium trihydridohydrazidoaluminate 53 Lithium trimethylsilyl-tert-butylhydrazide 55 Macrocycle 49 MAD 79 Magnesium 162 Malonates 100 Mannitol 21 MAS-NMR 78,82 Meerwein-Pondorf-Verley reduction 81, 142,143,156,158 Mesocates 125 Metaboric acid 8,26 Metal disorder, solid solutions 99 Methylaluminum bis(2,6-di-tert-butyl-4methylphenoxide) 79 Methylalumoxane 69 Mitsubishi 128 Monodentate 172 Monomer-dimer equilibrium 144 MPV reduction 81,142,143,156,158 Nobleite 10 Nonaborate 13,14 Nuclear quadrupole coupling 82 Olefin B-complex 144 Olshanskyite 8 Oppenauer oxidation 158 Orthoboric acid 26 Oxide-alkoxides 73,74,77 Oxide-hydroxides 73 Oxygen, active 36

207

Subject Index

Paper pulp 37 Pentaborate 5,8,9,13 Pentahydroborite 7 Perboratelperoxoborate 35-37 Phenylhydrazine 4447 Phosphines 69 Phospho-aldol reaction 170 Phthalocyanines 193 Pinnoite 7,8 Polyborate species 6 Polymerization, with Et,Al 149 Polyvinyl alcohol 19 Porcelaine enamels 15,33 Porphyrins 193 Preobrazhensite 14 Preservatives 29,34 Propylene oxide, oligomerization 81 Pulp 37 Pyridine-diimine ligands 149 Pyridyloxides 95 Quantum-chemical calculations 56,62 Radical anions 79 Rhamnogalacturonan 21 Robinson 169,194 Rubber products 34 Salen ligand 170,190 Salicylaldehyde 130 Sborgite 8 Schiff bases 126,130 Semiconductors 42 Siloxane aluminates 104 Single-electron-transfer mechanism 159 Sodium perborate 36 Solvation 176 Sorbitol 21

Spirocycles 131 Starch adhesives 19 Tartralon B 20 Tau value 170-189,194,198 Tetraacetylethylenediarnine (TAED) 37 Tetraborate 5,9,13,14,25,26 Tetrahydroxyborate 5 1,1,4,4-Tetramethyl-2,3-diazabutadiene59 Tetramethylhydrazine 44 Thermolysis 59 Thiosemicarbazones 194 Tincal 29 Tishchenko reaction 155,160 Transformer water 12 Triazenides 94 Triborate 5,8 Trihydroxyboroxine 8 Trimethylalane 49 Trimethylaluminum 42,44,47 Trimethylgallium 42 Trimethylsilylhydrazine 43 Triphenylphosphine 69 Tripodal ligands 193,194 Tris(trimethylsi1yl)hydrazine 43

Vapor deposition, chemical 42 Vimsite 7,8 Wood composites 29,34 Ziegler 169 Ziegler-Natta catalysis 69,101 Zinc borate 14,28,29,34 Zirconocenes 70