Handbook of Lithium and Natural Calcium Chloride

Handbook of Lithium and Natural Calcium Chloride Their Deposits, Processing, Uses and Properties Donald E. Garrett Sali...

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Handbook of Lithium and Natural Calcium Chloride Their Deposits, Processing, Uses and Properties

Donald E. Garrett Saline Prucmurs, Inc. Ojai, California

2004

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Contents

Preface ...................................................................................................

vu

Acknowledgmell/s ..................................................................................

xi

Part I Lithium

Geology ..................................................................................................

1

Brine Deposits ...................................................................................... Ore Deposits ........................................................................................

5 47

Processing ............................................................................................... 98 History o f the Lithium Industry .......................................................... 98 Brine Processing; Solar Ponds ............................................................ 100 Ore Processing ............... ..... .... ... ................ ................ .... .... ....... ..... ...... 146

Toxicology .............................................................................................. 179 Uses ........................................................................................................ 180

G·lass .................................................................................................. 181 Ceramics .............................. .............................................................. Aluminum ............ ............................... ............. .................................. Balteries ............................................................................................. Grease ................................................................................................ Othe r Uses .........................................................................................

185 188 190

196 196

Industry Sta tistics ................................................................................... 200 Chemis try, Phase Data. Physical Propenies .......................................... 202

References .............................................. ................................................ 223

VI

ConteniS

Part 2 Calcium Chloride

Geology .............. .................................................................................... 237 Calcium Chloride Dolomitization Brine ............................................. Calcium Chloride in Oil and Gas Field Brines ................................... Calcium Chlo•·ide in Geothermal Brines ............................................. Calcium Chloride Lakes ...................................................................... Calcium Chloride Groundwater .......................................................... Tacbybydrite Deposits ......................................................................... Calcium Chloride Brine io Mineral Deposits ...... ...... .... .... ................. Reactions of Calcium Chloride with Minerals ...................................

238 263 266 280 311 318 332 335

Processing ............................................................... ................................ 33 7 Michigan Dolomitization Brines ......................................................... 338 Bris tol and Cadiz Lakes ...................................................................... 345 General Processing Technology .......................................................... 352 Uses of Calcium Chloride ...................................................................... 358 De1cmg .............................................................................................. Dust Control .......... .... ........ .... .... .. ...... .... .... ............ .... ..... .. ...... ...... .... . Soil Compaction or Stabilization ...................................... ................ Concrete ..................... .................... ................................................ .... Oil and Gas ........................................................................................ Ballasting ........................................................................................... Food Processing ................................................................................ Industrial ............................................................................................ Miscellaneous ....................................................................................

358 365 365 368 369 372 373 374 376

Toxicology and Safety ........................................................................... 378 Product ion Statistics .......... .. .................... .... .. .... ................ .. ....... ....... .... . 379 Phase Data and Physical Properties ....................................................... 382 Phase Data .... ..................... ............................................... ................. 382 Physical Properties ............................................................................ 406 References .......... .............. ...................................................................... 424 Index ... ................................................ ............................................... ..... 459

Preface

Lithium is one of the most interesting of the industrial minerals, occurring primarily in the unusual lithium pegmatites or in the very few high-lithium brine deposits. Many of the lithium pegmatites contain separate masses of different minerals (i.e. they are highly zoned), with the crystal size often being very large, such as up to 1 –14 m long. There may be as many as 13 distinct and separate massive zones of single predominant minerals in the pegmatite, including up to six different lithium mineral zones, and several zones with other rare elements such as tantalum, niobium, tin, tungsten, cesium and rubidium. The pegmatites appear to be the final magma that had been forced up into fractures of previously crystallized granite-type rock, considerably enlarging the fractures, and then slowly cooling and fractionally crystallizing its lithium and granitic components. Presumably the lithium pegmatites resulted from a flowing magma that had slowly cooled with some mixing, allowing the initial crystallization of the less soluble compounds such as iron and manganese silicates. This left a granitic-type melt containing the more soluble and lower-melting lithium silicates, as well as any rare metals that happened to be present in concentrations too small to be initially crystallized. Most of the minerals in the lithium pegmatites can be easily separated by selective mining and conventional mineral dressing techniques, often allowing a number of products to be recovered. For example, in one deposit separate concentrates of the lithium minerals lepidolite, petalite, spodumene, amblygonite, eucryptite and bikitaite have all been sold, along with tantalite (tantalum ore), beryl (beryllium), pollucite (cesium), cassiterite (tin) microlite, both K- and Na-feldspar, mica and quartz. With the large deposits this makes the mining and purification of lithium mineral concentrates relatively inexpensive, and they can be used directly in some lithium applications, such as in ceramics, various formulations of glass, and in producing aluminum. Most of the high-lithium brines have been found in a few playas (dry lakes) in the high plateau zones of the central Andes, or in southwestern China (Tibet). In these volcanically active areas they were formed primarily from high-lithium geothermal springs flowing into closed basins, and then over a very long period evaporating and depositing and/or concentrating the spring’s other salts. Most lithium salts are very soluble, so none of the lithium crystallized (although some of the lithium was vii

viii

Preface

adsorbed onto clays or other rocks), allowing the high-lithium brines to be formed. These brines can be processed by pumping them to solar evaporation ponds to remove additional water and progressively crystallize most of the other salts that are present. Then, because of these regions very low humidities the final solutions can almost reach the saturation point of lithium chloride, and have a low content of the other salts. This final brine can next be processed to remove its boron and small amounts of sulfate and magnesium, and then be precipitated with soda ash. The area’s very unusual concentrating ability thus allows a pure lithium carbonate to be inexpensively recovered, and if desired, several other by-products can be produced directly from the brine (such as any or all of lithium sulfate, lithium chloride, potash [KCl], potassium sulfate, boric acid, magnesium sulfate and magnesium chloride. In some brines bromine, iodine, cesium and rubidium are also potential by-products). In one lower altitude location a brine with about one-tenth the original lithium concentration as these high-altitude brines is also being solar evaporated, but because of processing limitations and the area’s higher humidity it is only evaporated to about one-tenth of lithium chloride’s saturation. Brines now supply most of the world’s lithium chemicals and metal, and a high percentage of the total lithium demand. The uses for lithium are as unusual as its deposits, primarily because of its very small molecular size. It is the third smallest element in the periodic chart of elements (hydrogen, helium and then lithium), making it the lightest metal, the smallest metallic ion, the most electropositive element, and many other distinctive features. Its small ions can often fit within the molecular structure of other compounds, thus lowering the melting point of glass and ceramics, making them cheaper to produce, giving them greater strength and a lower thermal expansion coefficient (and thus much improved temperature change stability). Its high electronegativity gives it the highest electric output per unit weight of any battery material, and thus it is preferred in both conventional and rechargeable batteries. This feature also allows it to reduce the over-voltage in aluminum cells (thus reducing costs), and makes it the most important cation in high-performance greases. Its organic compounds make unusually active catalysts, and the carbonate or acetate are almost magical psychiatric drugs. It also has a wide variety of other uses. The geology of calcium chloride is also somewhat unusual. There are only two lakes in the world with a strong, fairly pure calcium chloride brine, and only a few with a more dilute or a mixed calcium-magnesium chloride content. Small amounts of the calcium chloride hexahydrate mineral (antarctite) occur at the two lakes, and massive amounts of a hydrated calcium-magnesium salt (tachyhydrite) only occur in three unusual potash deposits. In contrast to this scarcity on the surface or in crystalline form, there are a great many deposits of calcium chloride brine under ground. It is common near potash deposits, in some oil and gas brines, in deep-sea vents, some coastal aquifers, some geothermal springs and related formations, and in the occlusions of many minerals.

Preface

ix

Calcium chloride is a very large-tonnage industrial chemical that aids in the removal of ice or snow from highways during very cold weather, or prevents it from accumulating on roads or industrial products (it lowers water’s freezing point, and has a high heat of solution). It also is extensively used to minimize dust on country roads, and for many related uses based upon its ability to absorb moisture (i.e. air conditioning, drying other products, etc.). It accelerates the curing of concrete, and has many applications in the food and beverage industries. It is used for ballasting tires, and for a wide variety of other uses. This book will review the geology, mining, processing, uses, industry statistics, phase data and physical properties of these two important industrial minerals. Lithium and calcium chloride are not related, other than having a few common brine sources and uses, and are presented together merely for convenience. Neither material has a sufficiently extensive literature base to justify being the subject of a separate book, so the two subjects have been combined in this volume as separate chapters. The manner of presentation will be the same for both minerals. This book will be the last in a sequence of books on saline minerals by the author: Natural Soda Ash, Potash, Borates, Sodium Sulfate, and now Lithium/Natural Calcium Chloride.

Acknowledgments

The author wishes to thank the many individuals and companies who have helped with the preparation of this book. The Tantalum Mining Corporation of Canada Limited, and Bill Ferguson, its Manager graciously allowed a visit to their spodumene mine at Lake Bernic, Manitoba, Canada. Peter Vanstone, Chief Geologist provided a great deal of information about the mine and lithium pegmatite geology, while Carey Galeschuk, Project Geologist conducted a most interesting mine tour. SQM S.A. and Patricio de Solmeniat, Executive Vice President also allowed a tour of their facilities at both the Salar de Atacama and Antafagasta, Chile. Sebastian Moura, Investor Relations arranged the tour and helped conduct it, while Rigoberto Arqueros P., Jefe de Operaciones conducted the lithium carbonate plant tour. Maria Virginia Ramirez A., Superintendente Procesos y Calidad provided interesting processing details, and Juan Pablo Etchart B., Ingeniero conducted the tour at the Salar de Atacama solar ponds. Carlos Nakousi S., Gerente de Operaciones made corrections and helpful additions to the SQM text, while he and Patricio Vargas M., Head of Investor Relations authorized the use of the many excellent SQM photographs. Dr. Jurgen Deberitz of Chemetall GmbH graciously authorized the use of many pictures from his 1993 book on Lithium, and provide other pictures. Bart Loundagin, Manager, Chemetall Foote Corp., Silver Peak, Nevada also provided some very helpful previously published literature on Chemetall and the Clayton Valley lithium operation. The Dow Chemical Company allowed a visit to their Ludington, Michigan calcium chloride plant, and a very interesting tour was conducted by Douglas Busch, Business Excellence Leader and Doug Dunklow, Improvement/Resource Leader. Mr. Busch also provided details on the Filer brine, and gave permission to reproduce pictures of the plant and pages from a number of Dow’s excellent brochures on calcium chloride. The National Chloride Company of America also allowed a visit to their Bristol Lake facilities. Tom Beeghly, Manager, conducted a most interesting tour of this very unusual lake, brine gathering and solar pond operation, and allowed pictures to be taken. David Morrow, Manager of Tetra Chemicals’ Bristol and Cadiz Lakes calcium chloride operations also provided some details of their production methods, while Tetra and Hill Brothers Chemical xi

xii

Acknowledgments

Co.’s sales departments furnished various informative product brochures. Many thanks are also given to the publishing companies and individuals who gave permission to reproduce figures and pictures from their previously published technical articles.

Part 1

Lithium

GEOLOGY Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. The average amount in the earth’s upper crust has been estimated to be about 20 ppm by Vine (1980), although others have quoted values as low as 7 ppm (Bach et al., 1967; igneous rocks 6 ppm, sedimentary rocks 11.5 ppm) and as high as 60 ppm (Deberitz, 1993; 27th in rank of elemental abundance). Even with these small numbers, however, there are a fairly large number of both lithium mineral and brine deposits, but only comparatively a few of them are of actual or potential commercial value. Many are very small, others are too low in grade or located in remote areas, or too expensive to recover and process. Some of the better-known deposits are roughly located in Fig. 1.1, and estimates of the lithium reserves of various deposits (or countries) are listed in Table 1.1. The deposits have been formed because of lithium’s higher solubility than most other cations, so it sometimes has concentrated in flowing and cooling magma and/or its accompanying aqueous fluids, as well as in evaporating brines. Thus, its minerals are generally found in the latter stages of alkaline magma flow, intrusion and crystallization, as occurs in pegmatite formations. There are about 145 minerals containing lithium as a major component (. 200 with . 0.002% Li2O), and about 25 contain over 2% Li2O (Deberitz, 1993). Forty-three of the better known of these minerals are listed in Table 1.2. The high-lithium brines usually have obtained most of their lithium from geothermal waters, with perhaps some of the lithium coming from surface leaching of volcanic ash, clays or other rocks. However, lithium is very difficult to leach from the lattice structure of all rocks and minerals, so little is dissolved unless the water is very hot. Experimental studies have shown that at ambient temperatures, only 55– 170 ppb dissolves from extended contact with granitic rocks, but at 275– 6008C 0.25– 2.4 ppm Li can be extracted in the same agitated, long contact-period (Dibble and Dickson, 1976). Analyses of cores into deep-ocean rift or subduction zones 1

2 Part 1 Lithium

Figure 1.1 Location and reserve estimate of some of the world’s lithium deposits (Anstett et al., 1990; reserves: R1 proven, R2 probable).

Geology

3

Table 1.1 Estimated Lithium Reserves of Various Lithium Deposits, 1000 mt Li Reserves Brines Salar de Uyuni Salar de Atacama Salar de Hombre Muerto Clayton Valley Zabuye Salt Lake, China Qinghai Lake, China Smackover oilfield brine Great Salt Lake Searles Lake Salton Sea Dead Sea Total Ore deposits Africa (other) Bikita, Zimbabwe Mali Manono-Kitotolo, Zaire Namibia Argentina Australia (Greenbushes) Austria Brazil Canada (total) Bernic Lake, Manitoba Ontario, Quebec China Portugal Russia United States (other) North Carolina

5000a 3000b(4300 –4600) f 800c 30.4a(115 –382)g,h 1000c 1000c 1000c 526c 31.6a 1000c 2000c 14,718

. 0.3a 23b 26d 309a 9.8e 0.2a 150b 10a 3.3a 240.5a 73a 139a 500d 10c 130a 44.3a 71c

a

Anstett et al. (1990). USGS (2002). c Garrett (1998). d Lloyd (1981). e Kesler (1960). f Ide and Kunasz (1989). g Kunasz (1994). h Dillard and McClean (1991). b

have shown that lithium is adsorbed on, rather than leached from near-surface rocks (up to 1.8 km in depth), and only significantly leached from deeper rocks at temperatures greater than 300 –3508C. Based upon the isotopic analyses of lithium in the upper rocks there is some exchange by adsorption and simultaneous leaching in

4

Part 1 Lithium Table 1.2 Formula and Group of Some of the Lithium Minerals (Vine, 1980)

Name and formula

Mineral group or series

Amblygonite (Li,Na)AlPO4(F,OH) Bertossaite (Li,Na)2(Ca,Fe,Mn)Al4(PO4)4(OH,F)4 Bikitaite LiAlSi2O6·H2O Bityite Ca(Al,Li)2[(Al,Be)2Si2(O,OH)10]·H2O Brannockite KLi3Sn2Si12O30 Cookeite (Li,Al4)Si3AlO10(OH)8 Cryolithionite Li3Na3Al2F12 Eckermannite Na3(Mg,Li)4(Al,Fe)Si8O22(OH,F)2 Elbaite Na(Li,Al)3Al6B3Si6O27(OH,F)4 Ephesite Na(LiAl2)(Al2Si2)O10(OH)2 Eucryptite LiAlSiO4 Ferghanite LiH(UO2/OH)4(VO4)2·2H2O Ferri-sicklerite (Li, Fe3þ,Mn2þ)PO4 Gerstleyite (Na,Li)4As2Sb8S17·6H2O Hectorite Na0.33(Mg,Li)3Si4O10(F,OH)2 Holmquistite Li2(Mg,Fe2þ)3(Al,Fe3þ)2Si8O22(OH)2 Hsianghualite Ca3Li2Be3(SiO4)3F2 Lepidolite K(Li,Al)3(Si,Al)4O10(F,OH)2 Liberite Li2BeSiO4 Lithiophilite Li(Mn2þ,Fe2þ)PO4 Lithiophorite (Al, Li)MnO2(OH)2 Lithiophosphate Li3PO4 Manandonitea LiAl4(AlBSi2O10)(OH)3 Montebrasite (Li, Na)Al(PO4)(OH,F) Nambulite LiNaMn8Si10O28(OH)2 Natromontebrasite (Na, Li)Al(PO4)(OH,F) Palermoite (Li, Na)2(Sr,Ca)Al4(PO4)4(OH)4 Petalite LiAlSi4O10 Polylithionite KLi2Al(Si4O10)(F,OH)2 Rankamite (Na,K,Pb,Li)3(Ta,Nb,Al)11(O,OH)30 Regularly interstratified montmorillonite-chlorite Sicklerite Li(Mn2þ,Fe3þ)PO4 Sogdianite (K,Na)2Li2(Li,Fe,Al,Ti)2Zr2(Si2O5)6 Spodumene LiAlSi2O6 Swinefordite Taeniolite KLiMg2Si4O10F2 Tavorite LiFe3þPO4OH Tosudite Triphylite Li(Fe2þ,Mn2þ)PO4 Virgilite Zinnwaldite K(Li,Al,Fe)3(Al,Si)4O10(OH,F)2

Amblygonite — Zeolite (?) Margarite — Chlorite — Amphibole (asbestos) Tourmaline Margarite — — Sicklerite — Smectite Amphibole — Mica — Lithiophilite–triphylite — — Chlorite Amblygonite — Amblygonite

More recently determined lithium mineralsb Diomignite Li2(B4O6)O7? Liddicoatite (end member of the group) CaLi2Al7B3Si6O27(OH)4 Chlorite, boron-bearing (end member) Li2Al5BSi2O14·4H2O a b

Manandonite is listed by some as the same formula with 2H2O. Garrett (1998).

— Mica — — — — Pyroxene Smectite Mica — — Triphylite –lithiophilite — Mica — Tourmaline Chlorite

Geology

5

the temperature range of 50– 3508C, but little net change (James et al., 2003; Chan et al., 2002b,c; You and Gieskes, 2001). Other rocks or higher temperature leaching conditions must allow a greater amount of lithium to be removed, since some geothermal springs have lithium values of 6 – 50 ppm, but in all cases the lithium concentration is still very low. When these dilute geothermal waters are concentrated by the evaporation occurring in arid climate, closed, reasonably impervious basins, comparatively strong lithium brines have resulted in a few large playa deposits. Many large medium-concentration lithium brines have also been formed in various oil or gas field waters, potash deposit end-liquors (seawater only contains about 0.17 ppm Li), and a number of miscellaneous sources such as the Salton Sea geothermal brine and end-liquors from various commercial solar pond operations. Brine Deposits As with minerals, many brines and waters contain some lithium, but as noted above it is usually found in extremely low concentrations. There are a few exceptions, but as of 2003 only three brine sources had become actual commercial operations, and each had comparatively high levels of lithium (although one only contained , 160 ppm Li), appreciable lithium reserves and good solar ponding conditions (the Salar de Atacama, Chile; Salar de Hombre Muerto, Argentina; and Clayton Valley, USA). Their brines were obtained from the porous strata under the surface of playas, and each appears to have lithium-containing hot springs as their principal source of the lithium. By-product lithium has also been recovered from Searles Lake, but its concentration in the lake brine is only 50– 80 ppm Li. Because of the very dilute lithium concentration in even the best of brine deposits, they all owe their value to the availability of solar evaporation ponds to inexpensively further concentrate the lithium. Again, as an exception to this, Searles Lake used plant evaporation, but even with multiple products it became too expensive, so their lithium recovery has ceased. At the Salar de Hombre Muerto alumina-adsorption may be used to first fairly selectively separate the lithium from the other salts in the brine, but then solar evaporation would still be needed to concentrate the eluted dilute lithium solution. The projected reserves of lithium in the world’s few potential or actual commercial brine deposits has been roughly estimated as about 15 million metric tons of Li (Table 1.1), but the practical recovery of the lithium from many of these deposits would be very difficult. The more important lithium-brine deposits are separately discussed in the following sections. Clayton Valley (Silver Peak), Nevada This relatively small 83 km2 (O’Neill et al., 1969; 100 km2, Davis et al., 1986) dry lake (playa) is about 16 km long and 6.4 km wide, and has a drainage basin area of about 1300 km2 (Fig. 1.2). It is located in central Nevada about 87 km southwest of Tonopah, 274– 282 km from Reno and Las Vegas, and 40 km east of the Nevada – California border. Its elevation is 1300 m, and in the porous strata under its

6

Part 1 Lithium

Figure 1.2 Location map of Clayton Valley and its surrounding mountains (Davis et al., 1986).

surface there is a fairly concentrated sodium chloride brine with comparatively high amounts of potassium and sulfate, but very little magnesium and other ions (Table 1.3). It also has a fairly high content of lithium in a brine pool with about a 54 km2 area and an average deposit depth of about 460 m. Originally the central area contained 100 –800 ppm Li, and the discovery well at 229 m depth contained

Table 1.3 Various Analyses of the Clayton Valley Brine, wt.% (or pppm) Barret and O’Neill (1970) First wella

Anon. (1966)

Feed

Product

Kunasz (1974)

Wellb

Springsp

Brown and Beckerman (1990)

Garrett (1996)

8.16 1.17 678 533 407 — — 50 — — — — 0.7 13.11 1.12 233 70 — 0.7 28 — — — 1.180 — — 229

7.50 1.00 400 600 500 — — — — — — — — 11.70 0.75 — — — — — — — — — — — —

6.20 0.80 400 400 500 — — — — — — — — 10.10 0.71 — — — — — — — — — — — —

7.80 4.80 5000 70 40 — — — — — — — — 16.10 2.90 — 2000d — — — — — — 1.25 — — —

6.62 0.77 380 600 560 — — — — — — — — 9.50 1.11 — — — — — — — — — — — —

6.37 0.80 230 360 450 179 211 20 42 18 7 5 — 10.00 0.66 650c 90 90 57 31 29 11 — 1.079 19.5 7.1 158, 218

9010 892 36 51 372 — — — — — — — — 13,850 545 609c — 4.2 — — — 46 2105 — 36.5 7.3 —

6.20 0.80 200 200 200 — — — — — — — — 10.06 0.71 — 50 — — 20 — — — — — — 213

4.69 0.40 163 190 450 — — 21 — — — — — 7.26 0.34 74 67 — — 23 — — — 1.058 — — —

Geology

Na K Lip Mgp Cap Srp Tip Rbp Mnp Znp Nip Vp Asp Cl SO4 COp3 Bp Fp Pp Brp NOp3 SiOp2 d2H Density Temperature pH Depth (m)

Davis (1986)

a Garrett b

(1960). And others of about the same year, including Davis and Vine (1979). Brown and Beckerman (1990).

7

c HCO3. d

8

Part 1 Lithium

678 ppm when pumped at 450 gpm (Garrett, 1960). The average brine analysis when commercial production of lithium carbonate began in 1966 was about 400 ppm. Since that time it has been slowly declining, and in 1998 the concentration was about 100 –300 ppm Li (averaging 160 ppm, Harben and Edwards, 1998). In 1990 the average depth of the brine production wells was 213 m. The basin has a fairly limited rainfall (89 – 127 mm/yr) and a very high (water) evaporation rate of 760 –1370 mm/yr (Dillard and McClean, 1991). The sediments in the basin are primarily Quaternary alluvial gravel, sand, silt and clay (partly derived from the alteration of volcanic glass or pumice; Vine, 1980), with some gypsum and calcite and several halite layers. At a 35 m depth the sediments’ age is 26,900 years (Fig. 1.3), and there is one zone of 6.9 million year old volcanic ash. It is one of the major brine aquifers (Davis et al., 1986). The sediments tilt, and have several fault lines, which apparently trap the more concentrated lithium brine (Fig. 1.4). The general structure of the sediments in Clayton Valley compared to two other lithium brine deposits is shown in Fig. 1.5 (Vine et al., 1979; Gadsby, 1967). The lithium (and potassium) in the deposit apparently originated from currently flowing hot springs along the Silver Peak Fault, with the current brine composition being a blend of evaporated water from these springs and surface and ground water that drains into the basin. During several periods of evaporation halite was crystallized, forming occasional beds in the alluvial, ash-fall or stream-carried insoluble sediments. The springs flow at 10 – 20 gpm and contain 9280– 10,000 ppm Na, 786 – 826 ppm K and 24 – 43 ppm Li (Table 1.3). The unusually high temperature of the brine in some areas of the deposit (up to 448C at fairly shallow depth [25.5 m]) would also tend to support the theory of a geothermal origin for the lithium. In the playa sediments there are limited beds of the lithium clay hectorite

Figure 1.3 Several drill core sections from Clayton Valley, and their possible age correlation with Searles Lake (Davis et al., 1986) (see Fig. 1.47 for the core locations).

Geology

9

Figure 1.4 General cross section of Clayton Valley with its fault planes and possible groundwater flow (Davis et al., 1986).

Figure 1.5 The salt structure in three basins containing high-lithium brine (Kunasz, 1980; reprinted with permission of the Northern Ohio Geological Society).

10

Part 1 Lithium

(analyzing up to 1700 ppm Li) that appear to have been formed by a reaction of the brine’s lithium with volcanic ash. The brine is highly supersaturated with hectorite (assumed from its solubility product), indicating that the lithium was deposited from the brine, and not leached from the mineral (Anon, 1981, 1979; Barrett and O’Neill, 1970). There is one very small Tertiary high-lithium pegmatite dike (with lepidolite; Kunasz, 1974), a zone of moderately high-lithium exposed ancient lake-bed sediments, and some Tertiary extrusive volcanic rocks in the playa’s drainage basin (Davis and Vine, 1979; O’Neill et al., 1969). However, such rocks have not created lithium playas elsewhere in the world from normal run-off water leaching, and in the relatively brief lifetime of the Clayton Valley brine it is unlikely that either hectorite, volcanic ash, the old lake-bed sediments or the pegmatite contributed much lithium to this basin by normal-temperature water leaching (see the temperature-solubility data noted above). Waters accompanying the pegmatite, or very hot rock-leaching during vulcanism could have left a high-lithium brine, but it is doubtful that it could have survived the 6– 50 million year period from their occurrence to the present. The Silver Peak Fault geothermal waters, however, do appear to have leached deep, hot, high-lithium containing rocks. The total lithium reserves in Clayton Valley were originally estimated to be about 115,000 mt Li, along with 30 million tons of potash (Kunasz, 1994), but have also been estimated as 382,000 mt of Li (Dillard and McClean, 1991). Salar de Atacama, Chile This Salar is stated to be the third largest playa in the world (Jordan et al., 2002), and is located in Northern Chile about 200 km east of Antofagasta (Figs. 1.6 and 1.9; 280 km by road and rail) in an upper plateau on the western slope of the Andes mountains (Fig. 1.7). The playa has an area of 3000 km2, it is 85 km long with a maximum width of 50 km, and in its southern central section is a 1700 km2 “halite nucleus” of massive salt (Fig. 1.8; Jordan et al., 2002). Its drainage basin area is about 11,800 km2, and its elevation 2300 m. The Cordillera de Domeyko with its extensive Tertiary halite and gypsum outcrops (the Cordillera de la Sal) borders the Salar on the west, the high Andes mountains closely border its east side and more distantly the north. A tall (5,200 m), active volcano (Volcan Lascar; also called Cerro Miniques, with smoke frequently emitting from its cone top) and lower hills border the south. Mud flats dominate the northern part of the salar (Fig. 1.9), and there is a small perennial lake where the Rio San Pedro enters from the north. Several other small lagoons or wet areas are usually present in the mud flats further south and along the eastern side of the salar. This northern area acts as a settling, precipitating (calcite, gypsum, borates, etc.) and brine concentrating area for the Salar, allowing fairly pure halite to crystallize in the south. The large, highly mineralized El Tatio geyser field lies further to the north, and discharges brine containing up to 47 ppm Li (Table 1.4) into the Rio Salado, which joins the Rio San Pedro (Fig. 1.9) and then flows into the Salar. Some small

Geology

11

Figure 1.6 Map showing some of the Andean high-lithium salars and their connecting roads.

intermittent streams enter the Salar from various canyons on its eastern side, but there are essentially none from the western mountains with their salt and gypsum outcrops. The average flow of the San Pedro River is 1000 liter/sec, and the intermittent eastern rivers which enter the Salar primarily through alluvial fans (and not on the surface) have maximum spring flow rates (as liter/sec) of: Vilama 218, Aquas Blancas 134, Tulan 59, Honar 50, Peine 17, and Camar 3 (Ide and Kunasz, 1989). The southern “nucleus” of the playa consists of massive salt, and in areas that are frequently flooded (the eastern side) the surface salt has a beautiful clear white-topink color, and is uniformly fractured into polygonal patterns (usually pentagonal or

12

Part 1 Lithium

Figure 1.7 The sequence of elevations of the Andean salars, and their highly faulted plateaus (Vila, 1990; reprinted from Stratabound Ore Deposits in the Andes by permission from Springer-Verlag GmbH & Co. KG).

hexagonal) , 0.6 m on a side and 1.5 –2 m across. All of these salt polygons have developed expansion cracks at their edges, and then capillary evaporative crystallization of additional salt in the cracks has caused an uneven heaving of the polygon edges up to 0.3 m in height. The salt in most of the central area, however, appears to only have been flooded in rare events (such as 50– 100 year

Figure 1.8 Typical salt surface of the central section of the Salar de Atacama (SQM, 2001, courtesy of SQM S.A).

Geology

13

Figure 1.9 Surface structure of the Salar de Atacama, and its adjacent rivers and streams (Ide et al., 1983; reproduced with permission of the Salt Institute).

storms). It has a silty brown color and its surface is fairly flat in profile (Fig. 1.8), but composed of very rough, jagged and sharp salt. There are no flat surfaces, making walking on the salt very difficult, and driving impossible, although the salt can be bulldozed to a smooth surface. The salt is very porous (30, 20, 15 and 5% at 0– 0.5,

14

Table 1.4 Typical Analyses of Several Lithium-Containing Geothermal Brinesa (ppm)

Na K Mg Ca Fe Mn Sr Zn HN4 Te As Li Ba Pb Rb Cs Cu Ag Sb U Cl CO3

Salton Seaa,b

From power plant

From salt pond

El Tatio Springs, Chilec

Paradox Basin, CO

CO

NV

Yellowstone, Norris, WY

50,000 –70,000 13,000 –34,200 700– 5700 22,600 –39,000 1200–3700 1000–2000 540 –2000 500 –700 504 –650 520 312 100 –400 200 90 –210 25 –100 24 0.5– 20 0.5– 2 — — 142,000 –209,000 —

8700 1700 — 400 0.94 0.64 15.7 0.2 — — 1.20 16 9.73 — 9.4 3.5 0.12 — — — 15,610 —

70,000 36,000 — 9400 0.4 1.8 — 1.7 — 4.0 6.5 393 17.0 6.7 — — 0.6 0.8 3.3 31.0 159.000 —

4460 523 — 15.4 — — — — — —

25,200 26,700 30,900 43,500 1380 (260) 1300 50 (1090) — (20) 110d — 6 95 16 8 A1 (66) — — 201.100 (800)

682 103 0.3 6.8 0.3 — — — — — 2.7 7.1 — — — — — — 0.5 — 952 —

653 71 0.8 5.0 — — 1.0 — ,1 — 2.7 7.6 — — — — — — 0.4 — 865 0

439 74 0.2 5.8 — — — — 0.1 — 3.1 8.4 — — — — — — 0.1 — 744 0

46 — — 6.6 15.5 — — — — 8050 —

Typical volcanic springs 815 101 — — — — — — — — — 9.4 — — — — — — — — 1255 17 (continues)

Part 1 Lithium

Steam Boat Springs

Cerro Prieto

Table 1.4 (continued) Steam Boat Springs

Cerro Prieto

HCO3 SO4 B Br Si S F I PO4 TDS Density pH Temperature (8C)

Salton Seaa,b

From power plant

From salt pond

— 42–50 400–500 109–200 40 15–30 4.6 –10 0.5 1.5 — 1.18–1.26 4.6–5.5 100–400

— — 12 20.5 — — — — — — — — —

— — — — 2.1 — — — — — 1.250 — —

El Tatio Springs, Chilec 45 32 179 — 102 — — — — — 7.4 85

Paradox Basin, CO

CO

(882) (227) 1690 1960 (10) — (25) (264) (1000) 359,000 — (6.2) —

246 125 67.5 — 145 6.9 — — 0.8 2,500 — — —

NV

Yellowstone, Norris, WY

Typical volcanic springs

305 100 49 0.2 137 4.7 1.8 0.1 — 2360 — 7.9 89.2

27 28 11.5 0.1 247 0 4.9 ,0.1 — 1890 — 7.45 84

177 53 36 — — — 7.2 — — 2850 — 8.1 93

Geology

( ) Limited number of analyses. Reprinted from Borates: Handbook of Deposits, Processing, Properties and Use, Table 5.6, pages 248–249 q 1998, with permission from Elsevier. a Garrett (1998). b Christopher et al. (1975). Also, as ppm: Ce 10, Mo 10, Zr 8, Ta 6, As 3, Se 2.5, Ti 2.5, Cr 2, Ge 1, Cd 0.9, Al 0.6, Ga 0.5, Ni 0.5, V 0.3; 500 –3000 m depth. c Cusicanqui et al. (1975). Located 80 km east of Calama; elevation 4250 m; erupting brine near the boiling point. Power production wells at 550– 1800 m depth; maximum temperature (2638C) at 800–1000 m. Surface source water dD 274 to 278; d18O 210.5 to 211.0. d Li range 66 –173 ppm.

15

16

Part 1 Lithium

0.5 –2, 2 –25 and . 25 m depth, respectively; average , 18% for the upper 25 m), and brine-filled from about 0.6 – 35 m. The porosity then decreases rapidly to nearly zero after about 35– 40 m. The salt’s average depth is about 800 m (with one area up to 1400 m), but the depth varies greatly. Based upon a few drill holes and fairly complete seismic data it appears that the Salar has experienced major faulting (Fig. 1.10) during its basinfilling period, forming large escarpments that have later been filled with additional salt. The major fault (SFS) starts at the lower end of the Salar in about its center and extends in a NNW direction through much of the halite nucleus. The escarpment that it formed resulted in a deep zone to the east of the fault, and made the halite in the east be on average about 240 m thicker than in the west, or about 640 –960 m thick (average , 720 m) west of the SFS Fault, and 1400– 620 m thick (average , 960 m) to the east (Fig. 1.10). The age of the basement rock in the basin appears to be about 5.1 ma (million years old), and the lower , 500 m of sediments do not contain halite. They are about 2 5.1 ma old based upon uranium – thorium age dating of ignimbrites (consolidated ash flows, welded tuff or recrystallized ash) within or at the edges of the deposit. A second major fault (the Peine Fault) in the Salar occurs near its eastern edge, running in an NNE direction. Detailed stratigraphic sections of three halite cores, age dating and the estimated environment of the Salar during each period are shown in Fig. 1.11 (Jordan et al., 2002; Bobst et al., 2001). Brine can be pumped from the Salar’s near-surface salt mass at relative high rates, such as . 31.5 liter/sec (500 –1000 gpm) without appreciable draw-down, although such high pumping rates would hasten the short-circuiting of brine from nearer the surface and from other areas of the Salar. The brine is saturated with salt, and contains variable concentrations of lithium, potassium, magnesium, sulfate and borate in different locations in the Salar (Tables 1.5 and 1.6; Fig. 1.12). The lithium concentration varies from about 1000 –4000 ppm, and averages over 1500 ppm for the two commercial operations on the Salar. The total lithium

Figure 1.10 Thickness of the halite on both sides of the Salar de Atacama’s Salar Fault (Jordan et al., 2002, courtesy of the Geological Society of America).

Geology

17

Figure 1.11 Stratigraphy and age dating of three halite cores in the Salar de Atacama (Jordan et al., 2002, courtesy of the Geological Society of America).

reserves have been estimated at 4.3 –4.6 million mt of Li (Anon., 1998, 1995, 1981; Kunasz, 1994; Coad, 1984). Brine from the El Tatio geyser field contains 28 –47 ppm Li (Cusicanqui et al., 1975), and is probably the major source of the Salar’s lithium, boron and potassium (and perhaps the magnesium; Tables 1.4 and 1.13), while the majority of the salt and sulfate must have come from the Cordillera salt –gypsum mass. The ratio of lithium and potassium in the Salar are roughly the same as in the El Tatio run-off waters, and the mineralization from this geyser field alone, based upon current surface flow rates, could have supplied the Salar’s lithium and potassium in 250,000– 360,000 years. Since the Salar’s salt mass and arid evaporating climate has existed for much longer than that period, it would appear likely that the geyser’s flow and mineral

18

Part 1 Lithium Table 1.5 Various Salar de Atacama Brine Analyses, wt.% or pppm

Minsala

Garrett (1998)

Vergara-Edwards CORFO and Parada-Frederick (1981) (1983)b

Na 6.50 9.10 8.00 K 3.13 2.36 1.84 Mg 1.30 0.965 0.93 Lip 2420 1570 1500 Cap 530 450 300 Cl 17.30 18.95 15.90 SO4 0.80 1.59 1.70 Bp 556 440 600 Brp — — — HCOp3 600 230 — Density 1.227 — — pH — — —

7.60 1.79 1.00 1600 245 15.66 1.90 685 — — 1.226 —

Brown and Beckerman (1990)c Brine

Product

7.17 770p 1.85 190p 0.96 1.29 1500 63,000 310 530 16.04 34.46 1.46 166p 400 7300 50 — — — — 1.250 — 6.50

Orrego et al. (1994)d product 570p 160p 1.92 60,000 — 35.10 220p 6270 — — 1.252 —

a

Estimated. Km-20 brine. c Patent assigned to Foote. d Final solar pond brine. b

input might also have existed for that long. Also, subsurface flow from El Tatio through the region’s many aquifers and faults must have been appreciable, and would have greatly reduced this formation time. Alonso and Risacher (1996) have estimated that 63.4% of the Salar’s annual water input at present enters through subterranean flow, but that 85.2% of the yearly lithium addition enters in this underground flow. If these values are correct it would reduce the estimated Salar’s lithium input time to as little as 44,000 –68,000 years. Since a considerable amount of lithium must have been adsorbed on clay particles while traveling in the river or underground (as occurred in the Owens River feeding Searles Lake), the true age of lithium accumulation is probably somewhere between these two estimated ranges. These authors also estimated that a considerable lithium input came from the small streams flowing from the eastern mountains. However, considering the major uncertainties in their annual flow rate, the amount flowing underground, and the loss by adsorption (which has been very high in studies on deep sea vents, the Owens River and Mono Lake), their lithium contribution was proabably fairly small. One of the mysteries with this formation theory, however, is what has happened to the borate content of El Tatio’s brine? Based on the Salar’s lithium and potassium analyses only about one-tenth of the El Tatio boron is now in the Salar’s brine. As a related question, why are there not beds of ulexite, colemanite or other calcium borates in the Salar as there are in all of the Puna regions’ other lithium-containing,

Table 1.6 Typical Analyses of Various Lithium-Rich Brines in Northern Chile, ppm (Garret, 1998) pH

Dissolved solids

Na

Aguas Calientes Ascoton

7.7 7.8

Atacama

6.6

10.4

81,436 153,600 47,022 370,000 310,000 190,000 73,000 62,000 40,100 170,300

25,460 45,000 13,870 91,000 85,800 45,100 18,220 14,840 10,280 50,000

6.0 6.8 7.1 8.6 — 7.5 7.8 7.7 —

150,100 390,000 271,900 89,298 102,138 167,200 4357 152 —

38,000 126,800 86,000 28,500 28,160 54,000 1210 23 2.2

Bellavista, Pintados Huasco Lugunas Punta Negra Pujsa San Martin Surire Hot Springsa Riversb Soilc

K

Mg

Ca

Li

1183 3500 1670 23,600 13,000 9000 4220 2900 1690 5403

1361 5125 827 9650 6350 5330 1810 1930 750 3665

2538 920 1195 450 1100 900 360 1080 1160 5935

152 186 82 1570 940 520 290 190 130 85

10,000 14,280 10,000 1295 2614 8700 200 2.8 0.3

1750 3630 2620 653 6252 1250 28 4.3 1.8

840 110 2080 375 1566 750 135 1.0 13.5

130 412 320 137 187 340 8.3 0.1 65

SO4

HCO3

B

46,690 70,000 24,000 189,500 163,900 83,780 36,750 27,500 20,300 100,600

3154 25,000 4693 15,900 8540 18,170 3430 7900 2160 2720

0 2900 0 230 280 240 320 100 92 178

474 783 595 440 360 360 100 88 61 225

83,600 176,600 164,500 27,660 60,050 79,800 1905 22 1.2

13,600 47,770 4480 28,110 2490 20,300 534 20 22.9

Cl

— 406 — 0 625 90 150 33 3.1 CO3

Reprinted from Borates: Handbook of Deposits, Processing, Properties and Use, Table 4.3, pp. 201, q1998, by permission of Elsevier. a Also SiO2 129, NO3 7.4. b Also SiO2 45, c South-center of playa, average 6.1 m depth. Also As 158 ppm.

2200 979 2230 675 426 1820 47 0.7 3.1

Geology

Salar

19

20

Part 1 Lithium

Figure 1.12 Isopach map showing the lithium, potassium and sulfate concentrations in the Salar de Atacama (After Ide and Kunasz, 1989; CORFO, 1985).

geothermally fed salars? The borates are usually found as nodules (like potatoes) buried in the playas’ near-surface muds to a depth of 0.2 –1.3 m, although some nodules can be much deeper, and some playas have formed layered deposits. Most of these playas have been commercially mined for their borates in the turn of the century, and quite a few still are (Garrett, 1998). However, no large-scale borate deposits have ever been noted in the near-surface of the Salar de Atacama. This probably implies, since the El Tatio geyser field almost certainly supplied most of the Salar’s lithium content, and the Salar’s sediments are so old that there probably are deeply buried borate deposits in the northern non-halite zone of the Salar. Most of them would have deposited prior to the last ice ages 10,000 –21,000 years ago, and then been buried by the sediments carried into the Salar with the massive ice age water flows. This would make their depth below 18 – 20 m (Fig. 1.11), and indicate that there should be over 114 million tons of ulexite (NaCaB5O9·8H2O; 15.2 MMmt B) buried in the Salar’s San Pedro de Atacama mudflats. The ulexite would have slowly formed as the Rio San Pedro water met the high-calcium sulfate run-off water from the Cordillera de la Sal, since ulexite is much less soluble than gypsum (but also

Geology

21

much slower to crystallize). Based upon the hydrodynamics of the other salars most of it would be in the mud flats close to the Cordillera. This reaction, in turn would have liberated the sulfate from the Cordillera’s dissolved gypsum, and would account for the high sulfate content in the salar’s brine. Ide and Kunasz (1989) have reported that there are zones of high-lithium sedimentary rocks (up to 470 ppm Li) in the basin’s run-off area, and considerable masses of low-lithium volcanic ash and rocks. They made some tests indicating that the rocks could be rapidly leached by water at ambient temperatures to yield 1– 15 ppm Li solutions. However, no one else has ever reported such simple lowtemperature lithium leaching, as the leach solutions at ambient temperatures from other researchers contained less than 0.01 –0.04 ppm Li, and temperatures . 3008C were required to leach that much lithium from the rocks tested. Further, if lowtemperature leaching could generate such lithium solutions from volcanic and sedimentary rocks, there should be far more high-lithium brine deposits of this type, since there are many playas in the Andes and elsewhere in the world with similar rocks. However, the only similar (but smaller and/or less concentrated) lithium deposits occur where there are known high-lithium hot springs feeding into closed basins to form playas (Garrett, 1998). The evaporation rate of water at the Salar is relatively high, even with its high altitude and cold winters and evenings, since there is usually some-to-moderate wind, and the humidity is very low (usually only 5– 10%). These conditions allow even the Salar’s very hygroscopic MgCl2 or LiCl solutions to evaporate and crystallize salts. The evaporation rate for water is in the range of 3200 mm/yr, compared with 2300 at Hombre Muerto and 1800 at the Great Salt Lake and Clayton Valley. The area’s average rainfall is 10– 50 mm/yr (average 10 – 15; Hombre Muerto 55 – 70, Clayton Valley 230 and Great Salt Lake 330 mm/yr; Harben and Edwards, 1998). Ide and Kunasz (1989) list the Salar de Atacama’s rain at 10– 30 mm/yr (average near 10 mm/yr), the solar radiation 630 langleys/day (6.3 £ 106 cal/m2/day), the temperature range from 2 58C (winter) to 358C (summer), and the brine level generally 50 –70 cm from the surface. They also state that there are two confined aquifers in the mud flat area of the Salar that cause some springs to form in both the mud and the edges of the halite zones. There are also a number of smaller, but similar fairly high-lithium salars throughout the Puna region of Chile, Argentina, Bolivia and Peru as seen in Tables 1.6 –1.8, and Fig. 1.6. For example, the Salar de Surrie in Chile contains a brine with 389 ppm Li, 1334 ppm B and 1120 ppm K. It has an area of 150 km2, an altitude of 4480 m, 5 and 188C average winter and summer temperatures, and has an 808C hot spring in its southeast corner (Garces, 2000). In this high Andean plateau area with its many volcanoes and geothermal springs each of the salars appeared to have originated from the region’s hot springs. As the groundwater from rain and snow percolates through the region’s many faults it would be heated by contact with rock still hot from the recent or active vulcanism, and in some cases leach both lithium and boron (always an accompanying mineral in these salars). This is the

22

Brine Analyses at the Salar de Hombre Muerto (ppm or wt.%) (Garrett, 1998)

Na

K

Ca

Mg

Li (ppm)

Cl

SO4

B2O3(B) (ppm)

Total solids

Initial brine in the 15 76 210 72 2.1 900 1100 25 (7.77) 3500 Catal Lagoon (ppm) Catal Lagoon 9.45 0.55 0.02 0.16 930 15.8 1.06 1400 (435) 28 brine at NaCl saturation (wt.%) Average brine 10.1 0.519 0.088 0.054 521a 16.0 0.846 750 (233) 27.8 in the top 1 m of sediments (wt.%) Range (wt.%) 9.9– 0.24– 0.068– 0.018– 190– 15.8– 0.53–1.14 260–1590 27.2–29.4 10.3 0.97 0.121 0.141 900 16.8 (87–535)

Density (g/cc)

pH

1.001

7.5

—

1.22

7.2

—

1.204

6.9

1.199–1.212 6.5–7.2

Reprinted from Borates: Handbook of Deposits, Processing, Properties and Use, Table 5.1, p. 232 q1998, by permission of Elsevier. As well as 29 ppm Rb and 33 ppm Cs.

a

Conductivity (mmho/cm)

1.74

1.68–1.80

Part 1 Lithium

Table 1.7

Geology

23

theory for the much-studied El Tatio geyser field, and it probably holds true for all of the region’s high-lithium salars. Salar de Hombre Muerto, Argentina This medium sized (565 km2) playa (Fig. 1.13) lies about 240 km SE of Antofagasta, Chile (as the crow flies; 565 km by road and rail), 395 km from Salta, Argentina and 1300 km NW of Buenos Aires, in the remote altiplano area of the Andes Mountains. The nearest large town is Salta in the Andes foothills to the northeast (Fig. 1.6). The Salar is at an altitude of 4300 m (4100 and 3964 m have also been reported) with a small mountain-peak island in its western center, and its surface is always partially flooded (from the Catal Lagoon in the southeast corner; Fig. 1.14) and partially dry. The Salar’s surface near the Lagoon often floods in the winter from the perennial Los Patos river, and thus after the water evaporates is usually covered with smooth white salt, while most of the other areas have a dirty,

Figure 1.13 Surface structure of the Salar de Hombre Muerto (Garrett, 1998; reprinted from Borates; Handbook of Deposits, Processing, Properties and Use, Fig. 5.1, p. 228, q1998 by permission of Elsevier).

24

Part 1 Lithium

Figure 1.14 Areal view of the Salar de Hombre Muerto’s Catal Lagoon, with the northwestern section of the Salar in the background.

very uneven salt –clay surface as in most of the Salar de Atacama (Fig. 1.8). A . 50 m thick massive salt body underlies much of the Salar, and contains a highlithium brine in its porous upper section. Its lithium content averages 521 ppm, but varies from 190 –900 ppm (Fig. 1.15 and Table 1.7), and the brine is estimated to contain 800,000 mt of lithium, 1.1 million mt of B2O3 and 80 million tons of potassium. The Salar’s average air temperature has a high of 78C and a low of 2 68C, with a maximum and minimum of 13 (also noted at 288C) and 2 328C, respectively. The average rainfall is 60 – 80 mm/yr, and the evaporation rate about 1500 mm/yr. The algae in the Catal Lagoon support a large colony of pink flamingoes, and wild burros and domesticated llamas graze on the bunch grass near the Salar. The brine in the Salar is usually within 20 cm of the surface, and it appears to be relatively constant in composition with depth to at least 15 m. The analyses listed above were taken from samples mostly at a depth of 0.7 – 0.9 m, and in these holes the average amount of insolubles in the salt was 1– 11%. The average porosity of the salt to 15 m was about 15%, there appeared to be some circulation of the brine in the salt mass, and there was capillary evaporation from the surface. Most of the calcium, and some of the magnesium and sulfate in the Los Patos river water precipitated as it advanced to the edges of the Catal Lagoon and into the main brine body. About 32% of the Salar’s near-surface crust contained on average 12.7 cm of cotton ball ulexite

Geology

25

Figure 1.15 Brine analyses across the Salar de Hombre Muerto, g/liter (Garrett, 1998; reprinted from Borates; Handbook of Deposits, Processing, Properties and Use, Fig. 5.4, p. 231, q 1998 by permission of Elsevier).

26

Part 1 Lithium

(much of it as large nodules like potatoes), starting at an average depth of 80 cm and extending to a depth of several meters. The total ulexite would appear to be about 7 million tons of B2O3, and in the early 1900s some of it had been commercially harvested (Garrett, 1998). In a 40 m core sample uranium – thorium age dating on layers of ash indicated that the bottom was about 82,000 years (82 kyr) old. At a 6.4 m depth the age was about 8000 years. The amount of mud in the halite from 8– 26 kyr indicated that the weather had been comparatively dry, as at present, followed by a brief period of wetter weather. During the period from 64 –82 kyr the Salar appeared to have been a saline lake, and the weather was much wetter (Lowenstein et al., 1998). Salar de Uyuni, Bolivia The Salar de Uyuni is located in central Bolivia fairly near its western border with Chile, about 190 km from Iquique, Chile and the Pacific Ocean (660 km by rail; Fig. 1.6). It lies within the Puna region of the Andes Mountains, a very high, large and arid valley region within the Andes that extends from central Argentina to Peru. The Salar is the world’s largest saline playa, with its surface area having been estimated at 9000– 10,500 km2 (its longest dimension is 120 km; Fig. 1.16), and the

Figure 1.16 Map of the Salars de Uyuni, Coipasa and Empexa, and the ancient Lago Minchin (Ericksen et al., 1978; reprinted from Energy, Vol. 3, No. 3, q 1978 with permission of Elsevier).

Geology

27

Figure 1.17 Brine concentration map of Salars de Uyuni, Empexa and Coipasa (Ericksen et al., 1978; reprinted from Energy, Vol. 3, No. 3, q 1978 with permission of Elsevier).

smaller but quite similar Salars de Coipasa and Empexa are adjacent to it (Fig. 1.17). Its altitude is 3653 m, and much of its northern surface is hard, smooth and flat, similar to the Bonneville salt flats in Utah (used for very high speed racing). There are occasional algal reefs up to 75 m high, as well as much lower algal terraces, indicating that the playa was once a much larger and deeper lake (called Lago Minchin). It apparently started to reduce to its present size about 10,000 years ago, and the evaporation was completed about 3520 years ago based upon the age of near-surface organic matter. The Salar’s average depth is 121 m, and it has a 0.1– 20 m thick salt mass (average 3– 6 m) in its central area in the form of 11 porous (20 –30% void space) halite beds separated by layers of mud and sand. In the southern section there is a small perennial lake, the surface crust is much more irregular and the near-surface sediments contains significant amounts of ulexite (NaCaB5O9·8H2O; Garrett, 1998). The Rio Grande de Lipez river flows into the Salar from the south, forming an extensive delta area and the perennial lake. The Salar floods from 0 – 75 cm deep during the rainy season, but most of the playa usually dries completely in the summer. However, beneath the surface the Salar is filled with brine, and it is always within 5 – 20 cm of the surface. The brine contains 80 to 1150 ppm of

28

Part 1 Lithium

lithium, and averages 321 ppm Li, but one limited area in the southeast corner averages 625 ppm (Table 1.8 and Fig. 1.17). There are many thermal springs feeding into the Rio Grande river (and thus the Salar) that have a high lithium (4 – 30 ppm) and boron content, and they are probably the major source of the Salar’s minerals. As an example of the Salar’s concentrating effect on the springs’ water, the Rio Grande river enters the Salar with 11 ppm boron (B is about 2/3 the Li value), but by the time it has passed through the Salar’s delta system its concentration is 520 ppm B. The surface crust in this southern area is being mined for its ulexite content (in 1996 at a rate of 5000 mt/mo), and salt has been mined from the central area since the 1500’s. It has been estimated that the Salar contains 13 km3 of brine with 5.5 million mt of lithium, 110 million mt of potassium and 3.2 million mt of boron. The average rainfall in the area varies from 20– 50 cm/yr, while the evaporation rate is about 150 cm/yr. There are a number of smaller highlithium salars at a higher elevation (4000 – 5000 m) and to the south of Uyuni (Table 1.8; Garrett, 1998). Searles Lake, California Searles Lake is a medium sized playa located about 200 km north of Los Angeles in the Mojave Desert. Its surface area is about 100 km2, the center of which consists of massive halite that is about 8 m thick. The halite is 3500 years old at the surface and 6000 years old at its base (Fig. 1.3). Beneath the halite are two zones consisting of many saline minerals (halite, trona, hanksite, borax, thenardite, etc.), with a 5 m thick clay layer in between (Fig. 1.18). Both salt masses are very porous (, 35%) and filled with a high-density (, 1.30 g/cc) brine (Table 1.9). In the central section of the Lake the brine contains an average of 50 – 80 ppm Li, which grades to about 10 –70 ppm Li near the edges of the deposit. The massive, porous halite in the center grades to an overburden consisting of clay at the edges. Under the halite is about 13 m of the massive Upper Salt (1.05 km3 of salts), the clay layer, and finally 14 m of the Lower Salt (with alternate layers of salts and mud; , 0.5 km3 of salts) that extends to a depth of about 40 m. Beneath the Lower Salt is about 30 m of clay, and then a mixed zone which has layers of clay alternating with layers of varying purity sodium carbonate and sodium sulfate minerals with halite. The source of the various salts in Searles Lake, including the lithium has been fairly positively determined. First, as seen in Fig. 1.19 Searles Lake received most of its water and minerals from the Owens River. The Lake was usually the lowest and the final basin to receive Owens River water after it had accumulated the runoff and spring water from about 420 km along the east side of the Sierra Nevada Mountains. Water could overflow from Searles Lake into the Panamint Valley, and from there into Death Valley, but apparently when this occurred during the peak of the glacial periods the lake had stratified, and only fresh water overflowed. During the interglacial periods the water could evaporate, and for two periods it crystallized the upper salts. This succession of dry and then pluvial periods lasted

Table 1.8 Average Analyses of Various Lithium-Rich Bolivian Brines wt.% or ppm (Garrett, 1998) Pastos Grandes Salara Salar de Empexab,c

Salar de Uyuni

Na K Mg Ca (ppm) Li (ppm) Sr (ppm) Cl SO4 HCO3 (ppm) B (ppm) Br (ppm) F (ppm) SiO2 (ppm) pH Density

Salar de Coipasac

Hot Springs d

1

2

3

2

4

2

4

Edge

Playa

Subsurface brine

8.2 0.66 0.64 456 321 — 14.8 1.08 — 187 — — — — —

8.72 0.72 0.65 463 349 14 15.71 0.85 333 204 49 10 7 7.25 1.21

7.06 1.17 1.25 306 625 — 5.0 — — 525 — — — 7.3 1.19

5.4 0.27 0.68 209 172 — 9.7 2.8 347 176 — — — — —

6.7 0.52 0.32 410 253 — 14.8 — — — — — — — —

7.51 1.10 1.36 156 350 17 15.10 2.46 747 786 142 33 10 7.23 1.231

71 1.21 1.36 227 243 — 16.5 — — —

317 43 4.5 10.2 5.2 — 560 36.5 106 2.0 — — 39 6.55 1.001

0.448 0.05143 67 ppm 212 69 — 0.876 0.014 539 26 — — 37 6.30 1.013

7.72 0.891 0.174 1440 1800 — 15.67 0.932 608 376 — — 7.1 7.14 1.194

— — — —

Geology

1; an average of 40 Salar de Uyuni samples, 2; a more detailed individual sample analysis, 3; an average of the eight highest-lithium samples, found in a narrow band in the southern half of the eastern lobe of the salar, 4 is an average of four Salar de Empexa or Coipasa samples. Reprinted from Borates: Handbook of Deposits, Processing, Properties and Use, Table 4.2, p. 198, q 1998 by permission of Elsevier. a This salar is slightly southwest of Uyuni, in Bolivia. There is also a Pastos Grandes in Argentina. b Total dissolved solids (TDS), 19.3. c Adjacent to Uyuni. d All analyses as ppm.

29

30

Part 1 Lithium

Figure 1.18 Summary of stratigraphic units in Searles Lake evaporite sequence (Smith, 1979).

Table 1.9 The Analyses of Several Lithium-Containing Lakes or Brines, wt.% or pppm Searles Lakea,b Upper

Lower

Evaporated

11.08 2.53 — 16p 54g 144 22 14 5 — — — 12.30 4.61 2.72 — 2990 846 330 300 54 12

11.84 1.57 — — 60 — 32 — — — — — 10.81 4.44 3.84 — 4120 537 1560 190 20 20

8.32 10.62 — — 139 3480 — — — — — — 13.55 1.06 3.56 — 14810 7390 2840 2400 — 360

Southc 3.7–8.7 0.26– 0.72 0.50– 0.97 0.026– 0.036 18h 138 — 184 — 156 46 23 7.0–15.6 0.94– 2.00 5p 600 18h 55 — — 1.7 —

Bonnevilleb

Dead Sea

Mg Plantd

GSLe

Sea f

Conc.f

Sua Panb

Brine

Conc.

0.5 0.8 7.5 — 600 — — — — — — — 20.3 4.4 — — 540 — — — — —

0.118 0.058 8.55 50p 1160 — — 10 — — — — — 2.46 — — 700 2120 — — 50 —

3.01 0.56 3.09 1.29 12 — — — — — — — 16.10 0.061 — 190 30 3760 — — — —

0.38 0.22 7.01 2.65 23 — — — — — — — 25.60 — — — — — — — — —

6.00 0.20 — — 20 — — — — — — — 7.09 0.83 1.17 0.62i — 200 — — — —

8.3 0.5 0.4 290p 57 — — — — — — — 14.0 — — 60 — — — — — —

0.2 0.3 8.3 88p 980 — — — — — — — 23.0 — — — — — — — — —

Geology

Na K Mg Ca Lip Asp Wp Fep Sbp Cup Znp Crp Cl SO4 CO3 HCOp3 Bp Brp S¼ p Pp Fp Ip

Great Salt Lake

(continues)

31

32

Table 1.9 (continued)

Upper SiOp2 Np Density a

— — 1.29

Lower — — 1.30

Great Salt Lake

Evaporated — — 1.34

Southc

Mg Plantd

— 1.8 ,1.1

— — —

Bonnevilleb

Dead Sea GSLe — — 1.344

Sea f 11 — 1.198

Conc.f — — —

Sua Panb

Brine

Conc.

— — 1.124

— — ,1.2

— — ,1.3

Gale (1945). Garrett (1996) (Searles Lake, Upper and Lower Structure brine; Sau Pan, Botswana; Bonneville Salt Flats, Utah, brine and solar pond end liquor). c Strum (1980) (South arm of lake; dilute values unless a range is given). d Toomey (1980) (product from the magnesium plant’s solar ponds). e Nelli and Arthur (1970) (maximum evaporated brine from the GSL solar ponds). f Tandy and Canfy (1993) (Sea brine or end liquor from potash plant’s solar ponds). g Other authors list 70–80 or 150 ppm Li. h North arm 42 ppm Li and B. i Wt.%; brine to the soda ash plant’s solar ponds. b

Part 1 Lithium

Searles Lakea,b

Geology

33

Figure 1.19 Location of Searles, Bristol and Cadiz Lakes, showing the flow paths to and from the lakes (Ver Planck, 1957, courtesy of the California Division of Mines and Geology).

for about 50,000 years, flushing the concentrated brine and deposited salts from the intermediate basins (Owens Lake, China Lake, etc.) during each wet cycle into Searles Lake, with the last flushing occurring 3500 years ago. This is the age of the surface salts in Searles Lake, and the age of all of the salts in Owens Lake.

34

Part 1 Lithium

It appears that most of the minerals in the lake, other than much of the halite, came from a cluster of geothermal springs in the Owens Valley area (primarily in the Long Valley). After the Los Angeles Metropolitan Water District (MWD) purchased the water rights to the Owens River in the early 1920s its water initially killed several orange groves in the Los Angeles area because of its high boron content. This caused the MWD to analyze the various water sources into the river, and they found that the problem was the hot springs near the Long Valley Caldera, and some hot springs in the White Mountains on the eastern side of the Owens Valley. Using the MWD’s data on flow rate and analysis for a 10 year period from these springs, Garrett and Carpenter (1959) made a material balance for the annual tonnage of minerals from the springs, the tonnage of salts in Searles Lake, and the age of the deposit. There was a very close balance for most of the major ions (carbonate, borate and sulfate), as well as the minor ions (tungsten, iodine and fluorine), but the springs were low in Na, Cl and lithium. Later Smith (1976) made a much more detailed material balance using Owens River water for a later period, and each of the two salt beds separately. His results were similar, but generally indicated somewhat less contribution from the River. However, both calculations would appear to definitely indicate that these springs did supply most of Searles Lake’s salts except halite, and that there had been considerable loss of lithium from the springs, presumably because of it reacting with, or being adsorbed onto clays and volcanic rocks that the water contacted as it flowed to Searles Lake. This loss is similar to the apparent reaction of volcanic ash or smectite clay to form hectorite (a lithium clay) in the Clayton Valley, and the near-surface adsorption of lithium from the brine in deepocean rift or subduction zones. The large scale recovery of products from Searles Lake began in 1916, with brine being pumped to large plant evaporators, followed by a succession of processing facilities to recover most of the brine’s individual components. In 1936 facilities were installed to remove the very fine crystals of dilithium phosphate that crystallized with the burkeite in the evaporators, and became an impurity in the soda ash and salt cake products. These “licons,” containing about 20% Li2O were at first sold to Foote Minerals, but in 1951 a plant was built to produce lithium carbonate and phosphoric acid from them. When the plant to process the burkeite was closed in 1978, the lithium operation was also discontinued (Garrett, 1998; Vine, 1976). Mono Lake is one of the lakes that drained into Searles Lake during the ice age periods, and it has been studied for its lithium isotopic composition in a series of papers (Tomascak et al., 2003, 2001, 2000). The current flow of water into the lake comes from streams (75%) and springs, although almost all of the lithium enters from a few thermal springs. The lake water contains 10 ppm Li with a d7Li of 19.5, while groundwater and stream water d7Li is variable but up to 31 and 29, respectively, and the thermal spring water d7Li is 8.4. The volcanic rocks in the area have a d7Li of 3.8, indicating that their high-temperature leaching is the source of the thermal spring’s lithium.

Geology

35

Potential Deposits Great Salt Lake, Utah The Great Salt Lake is one of the larger inland lakes in the world (Fig. 1.20), and contains a seawater-type brine. Far more than 5.7 km2 of solar ponds have been constructed in the north end of the lake in the broad Bear River mud flat estuary, and

Figure 1.20 Map of the Great Salt Lake showing distribution of most common types of sediments (Gwynn and Murphy, 1980; reprinted from The Great Salt Lake by permission of the Utah Department of Natural Resources).

36

Part 1 Lithium

Figure 1.21 The limited-permeability railroad causeway across the Great Salt Lake, Utah (Gwynn, 1980; reprinted from The Great Salt Lake by permission of the Utah Department of Natural Resources).

on the west side of the lake to produce potassium sulfate, salt (NaCl), magnesium chloride brine and previously sodium sulfate. Here the brine can be drawn from the lake’s more concentrated zone north of the semi-permeable railroad causeway (Fig. 1.21). Solar ponds have also been constructed in the southern part of the lake for salt and magnesium metal production. The solar ponds produce an end-liquor which is much more concentrated in lithium than the original lake brine (Table 1.9), and thus has been considered for commercial lithium recovery. Although the lake’s brine concentration varies widely with the climate cycle, the southern brine has contained from 18 to 43 ppm Li, and the northern brine from 40 to 64 ppm Li (Whelan, 1976). The magnesium plant’s solar pond end-liquor often contains about 600 ppm Li, while the potassium sulfate plant has achieved values from 700– 1600 ppm Li (Toomey, 1980; Nelli and Arthur, 1970). If the latter end-liquor’s lithium were recovered it could amount to over 41 mt/yr of Li. Extensive tests have been conducted on both solvent extraction (with a ferric chloride – organic solvent mixture), and a plant evaporation-selective crystallization process to recover this lithium. However, neither appeared to be sufficiently economical to be competitive with other deposits. The Great Salt Lake is estimated to contain 526,000 mt of lithium. The subsurface brines of the nearby Bonneville Salt Flats are of a simpler type with very little sulfate and bromine, and contain 20– 60 ppm Li (Table 1.9). The Bonneville brine is also commercially evaporated in solar ponds to produce potassium chloride (potash) and a concentrated magnesium chloride brine that is sold for road de-icing and other uses. This end-liquor has been further concentrated

Geology

37

than that produced at the Great Salt Lake, so it has a higher lithium content of 2000– 4000 ppm, and averages 3000 ppm Li. It might be capable of yielding up to 20 mt/yr of lithium (Whelan, 1976). Salton Sea Geothermal Brine, California In the most southern part of California, in about its central section is the large Salton Sea. South of it and extending into Baja California, Mexico is a very large geothermal brine field in the porous sediments at a depth of about 500– 3000 m. The brine is a hot (100 –4008C) concentrated solution of predominately sodium and calcium chloride, with a very large array of metals and other uncommon ions, including from 100– 400 ppm Li (200 ppm average, Vine, 1980; Table 1.4). Geothermal power is obtained from these brines in both the USA and Mexico, and many studies have been made on the possible recovery of lithium from the plants’ effluent brine. At the southern end of the Salton Sea district, Cerro Prieto, Mexico contains equally hot, but much more dilute brine with a lithium content of 5– 100 ppm. Pilot plant solar ponds have been operated that concentrated the brine to about 400 ppm Li after a potash recovery process, but no attempt was made to recover the lithium (Garrett, 1996; Vine, 1980; Berthold and Baker, 1976). The origin of the Salton Sea geothermal brine is believed to be meteoric water (based upon dD and d18O values) from the nearby Chocolate Mountains flowing through fault lines deep into the earth, where it is heated by hot rocks or magma. The brine composition might indicate that the descending water dissolved highmagnesium potash salts that had been formed in this former seawater estuary area (from the present Sea of Cortez [Gulf of Baja California]). Then this brine underwent the very common dolomitization reaction (see Chapter 2) by reacting with calcite and converting most of its calcium content to magnesium. Later, when the brine was heated it became highly corrosive to other rocks, and dissolved the wide array of metal ions that it now contains. There is no present indication of a buried potash deposit or of a high temperature rock source, but it is on the very active San Andres fault line and the plunging Pacific Plate, so the heat source might be at considerable depth. The brine could be a concentrated rift vent brine, but its low bromine content, the terrestrial dD and d18O values, and the lack of dolomite and gypsum beds in the formation or nearby makes this very unlikely (Garrett, 1996). Dead Sea, Israel and Jordan The Dead Sea is one of the world’s largest and lowest inland lakes, containing a concentrated calcium – magnesium – sodium – potassium chloride brine, with about 10 ppm Li (Table 1.9) and reserves of about 2 million tons of Li. The brine is commercially evaporated in large solar ponds to produce potash in both Israel and Jordan, and their pond end-liquors often contain about 30 ppm Li. Some of this brine is processed for bromine and magnesia recovery, but most of it is merely returned to the sea. Because of its ready availability and potential value several laboratory studies have been made on lithium recovery from it, but without economic success. The source of the Dead Sea appears to be a blend of hot end-liquor dolomitization brine (such as is found in many springs near the Dead Sea) that has traveled along

38

Part 1 Lithium

fault lines from potash deposits under the Red Sea (Folle and Beutel, 2000), and surface water entering from the Jordan River. Neither seawater nor the river source has a high lithium content, making the lithium concentration unusually low for such a strong brine. However, the Dead Sea is very large, as are the two solar pond potash operations, making the total reserves very large. Even though the Dead Sea is located in one of the lowest valleys in the world, and in a very hot desert, the wind and humidity conditions are not sufficiently favorable to evaporate the brine to crystallize calcium and magnesium chloride salts as at the Salar de Atacama. If they were, then further solar evaporation could appreciably concentrate the lithium, and make its recovery more practical (Epstein et al., 1981). Chinese Dry Lakes In the high mountainous region of Tibet (Xizang) there are more than 57 highlithium playa lakes similar to those in the Puna or altiplano region of Argentina, Bolivia, Chile and Peru (Fig. 1.22). Geothermal springs with a high lithium content flow into most of the playa lakes, and it has been stated that 37 different ionic species are found in them, along with 27 different minerals in the playas (including many forms of borates). Several of the individual lake analyses are listed in Tables 1.10 and 1.11, and the shape and surface structure of others is shown in Fig. 1.23. Another group of these lakes is in the Qaidam Basin, containing brine with an average of

Figure 1.22 Map showing the location of some of the high-lithium lakes in the Qinghai-Xizang (Tibet) Plateau (Dapeng and Bingxiao, 1993; reprinted from the Seventh Symposium on Salt [ISBN 0444891439], Vol. 1, p. 178, Fig. 1, q 1993, with permission from Elsevier).

Table 1.10 Chemical Composition of the Brines in Various Alkaline Lakes in the Xizang (Tibet) Plateau, China (Garrett, 1998) Concentration (mmol/l unless noted) Lakes

Salinity (g/l)

pH

Na

K

Mg

Ca

Li

SO4

CO4

HCO3

B2O3

Cl

282.63

8.3

4065.42 7.29 10.66

543.18 1.66 3.83

1.37 26 ppm —

3.39 106 ppm —

90.34 489 ppm 660 ppm

361.98 2.71 2.19

698.92 3.27 3.75

— — 0

182.94 0.99 0.45

3449.21 9.53 12.30

Zabuye Caka Surface brinea wt.% Interstitial brine

434.47

Bangkog Cuo Lake I surface brine Lake II interstitial brine Lake III surface brine

68.52 119.20 221.88

8.7 8.6 8.7

1130.74 1562.62 2696.38

109.08 99.88 200.57

5.45 3.07 4.53

— 0.01 —

14.98 35.30 18.30

183.90 994.16 18.87

76.90 124.49 221.64

26.55 69.90 201.51

38.28 45.31 78.66

675.39 561.25 1118.63

Guogaling Cuo Surface brine Interstitial brine

114.17 125.97

— 8.8

1761.35 2217.99

87.32 200.85

0.41 2.64

— —

16.14 21.61

892.14 769.51

207.42 396.29

53.50 3.07

42.70 106.54

459.94 1165.32 Geology

Reprinted from Borates: Handbook of Deposits, Processing, Properties and Use, Table 5.3, pages 238–239, q 1998, with permission from Elsevier. Na2CO3 (After subtracting the equivalent Ca, Mg, and Li) 5.36 wt.%; Na2SO4, 4.01 wt.%; Na2B4O7, 1.11 wt.%; KCl, 3.15 wt.%; NaCl, 13.51 wt.%.

a

39

40 Part 1 Lithium

Table 1.11 Analyses of Several Lithium-Rich Lakes in the Qaidam Basin, China, wt.% or ppm (Garrett, 1992) Cl

SO4

HCO3

CO3

B2O3

310 182

13.42 14.16

3.41 2.04

0.06 0.21

0.02 —

0.20 0.26

0.08

38

12.14

3.57

0.004

—

0.96

0.07

51

10.84

2.33

—

0.79

0.01

5.79

2.35

0.68

Na

K

Mg

Ca

Da Qaidam Intercryst. Surface

5.63 7.77

0.44 0.36

2.02 1.17

0.02 0.03

Kiao Qaidam Surface

5.43

0.13

0.39

8.08

0.16

3.93

0.16

Mahai Intercryst. a

Qinghai Lake a

Li (ppm)

0.84

Br (ppm)

Total salts

pH

Density (est.)

58 80

25.05 25.68

7.3 7.4

1.234 1.240

0.19

16.6

21.76

—

1.203

—

434 ppm

—

22.38

—

1.208

0.52

15 ppm

1.5

14.23

—

1.133

Also (ppm): Si, 0.93; P, 0.50; Al, 0.26; Cr, 0.12; Ni, 0.092; Fe, 0.067; U, 0.042; Sr, 0.04; Ba, 0.02; Cu, 0.016; Mn, 0.016; Ti, 0.01; I, 0.004; Zn, 0.0021.

Geology

41

Figure 1.23 Several of the high-lithium Chinese Lakes, and their surface composition (Dapeng and Bingxiao, 1993; reprinted from the Seventh Symposium on Salt [ISBN 0444891439], Vol. 2, p. 179, Fig. 1, q 1993, with permission from Elsevier).

42

Part 1 Lithium

(as ppm if not noted): 320 Li, 542 B, 6.2% Na, 0.66% K, 0.47% Mg, 159 Ca, 6.7 Sr, 5.4 Rb, 1.9 Cs, 1.7 As, 9.2% Cl, 2.8% SO4, 0.12% CO3, 948 HCO3 and 113 F. The age of the lakes is estimated to be 5600– 20,000 years. Several of these lakes are quite large, and the Qinghai playa was being prepared for commercial potash production in 2001, with its 8– 13 ppm Li brine to be sent to solar evaporation ponds. The final potash end-liquors with about 120 ppm Li were being considered for the production of lithium (Qian and Xuan, 1983). Lake Zabuye Caka has an especially complex brine, somewhat resembling Searles Lake, but with higher concentrations of all of the alkali metals (K, Li, Rb and Cs; Tables 1.10 and 1.12). The lithium content in its brine varies from 500 to 1000 ppm, and the brine has been extensively studied for potential multiple mineral production. The brine is saturated with both salt and potassium sulfate, and during solar evaporation the lithium starts to crystallize at about a two-fold concentration (Table 1.12; Garrett, 1998, 1992). This lake has also been reported as being developed for lithium production (USGS, 2001). Other Geothermal Brines Brines in the Reykjanes, Iceland geothermal field contain modest concentrations of lithium (7.4 ppm), and several studies have been made on its potential recovery. At current flow rates the production of perhaps as much as 500 mt/yr of lithium (as lithium compounds) might be possible. Similar brine and studies have been conducted at the Wairakei, New Zealand geothermal area. Their brine contain 12 ppm Li, and their flow rate of 3.785 million l/hr might allow the production of 2400 mt/yr of lithium carbonate. Studies have also been made on the potential recovery of lithium from the geothermal brines of Cesano, Italy; Cronembourg, Alsace, France; and the Hatchobaru and Othake areas of Kyushu, Japan (Pauwels et al., 1990). Many other geothermal waters have a lithium content in the range of 1 –10 ppm (White et al., 1976), and a few have slightly higher values as indicated in Tables 1.4, 1.13 and 2.6. The East Pacific Rise deep-ocean thermal vents (Fig. 1.11) contain 7.2 ppm Li, have a pH of 3.5, and most of the other ocean thermal vents have a similar relatively high-lithium content. These vents have been extensively studied to determine their lithium source, including the extent of leaching from the rift zone rocks, or the adsorption of lithium onto the clay and other sediments on the ocean floor. For instance, the East Pacific Rise had a volcanic eruption at 9 –108C N latitude in 1991. It was noted that the vents’ brine greatly decreased in their lithium content near the eruption for several weeks, presumably because of the formation of new and shallower (and thus cooler and briefer) travel paths for the seawater to circulate in the Rise’s fracture patterns. Then the lithium slowly increased over the next 6 years (Bray, 1998). The North Fiji Basin vents also have a relatively high lithium content, along with appreciable methane and hydrogen sulfide. These vents are cold, and appear to be seawater-diluted hot vents that have had most of their heavy metal content precipitated by the hydrogen sulfide (Koschinsky et al., 2002). Laboratory studies have been made to simulate the rock-leaching by the Juan de Fuca Ridge

Table 1.12 Solar Evaporation of Zabuye Caka Lake Brine (g/l) (Garrett, 1992) Sample no.

Density

pH

Total salts

K

Na

Li

Rb

Cs

CO3

SO4

Cl

Br

l

B

Conc. ratio

1 Wt.% 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16

1.3047

10.86 — 10.83 10.38 10.67 10.75 — 10.92 10.98 11.11 11.36 11.22 11.12 11.22 11.08 11.13

50.00 3.83 55.15 59.00 65.10 62.50 62.50 62.50 59.00 59.00 55.15 55.15 56.00 63.00 56.00 56.00 93.35

139.11 10.66 142.44 142.44 146.15 146.15 155.06 158.77 162.25 156.91 160.18 173.60 173.61 173.61 181.02 181.02 178.13

0.87 667p 1.08 1.26 1.44 1.48 1.50 1.51 1.63 1.65 1.50 1.82 1.62 1.56 1.64 1.46 1.26

0.000 69p 0.000 0.090 0.140 0.140 0.180 0.180 0.240 0.240 0.240 0.360 0.480 0.720 0.880 1.240 1.800

0.034 26p 0.043 0.054 0.066 0.069 0.081 0.085 0.100 0.100 0.104 0.113 0.180 0.225 0.290 0.450 0.633

42.13 3.25 40.03 52.33 70.33 80.20 82.33 88.78 76.93 105.02 97.34 116.35 111.44 116.30 106.82 113.66 107.02

28.52 2.19 29.65 25.56 23.09 22.19 20.99 20.21 19.76 18.54 19.55 19.86 18.11 19.92 20.02 19.76 19.76

160.54 12.31 156.12 151.67 144.55 139.81 132.09 126.16 121.41 121.69 118.15 107.17 106.93 108.47 105.39 103.73 99.09

0.965 790p 1.04 1.29 1.49 1.78 2.07 2.07 2.34 2.41 2.56 3.33 4.92 5.84 6.78 9.35 13.37

0.0013 1.0p 0.0024 0.0030 0.0021 0.0021 0.0028 0.0011 0.0024 0.0025 0.0021 0.0026 0.0038 0.0064 0.0071 0.0074 0.0144

5.90 1.46 6.81 8.40 9.87 10.81 13.01 13.25 15.22 15.89 16.66 15.97 15.34 16.27 18.59 22.08 22.08

1.0

1.3226 1.3333 1.3573 1.3630 1.3703 1.3762 1.3880 1.3851 1.3987 1.4168 1.4120 1.4155 1.4214 1.4367 1.4409

428.56 32.85 433.01 442.58 462.82 465.13 469.82 473.54 456.88 481.45 473.74 491.73 488.63 505.92 497.44 507.76 524.24

1.25 1.52 1.67 1.76 2.18 2.23 2.73 2.77 2.85 3.66 5.31 7.31 9.16 13.51 19.31 Geology

Concentration of Mg, Ca, Sr, HCO3 all , 0. pppm. Reproduced from Natural Sada Ash: Occurrences, Processing, and Uses, Table 5.12, pp. 184–185, q1992 with kind permission of Kluwer Academic Publishers.

43

44

Part 1 Lithium Table 1.13 Analyses of Various Geothermal Brines Cesanoa (mg/l) Cronembourgb (mg/l) Wairakeic (ppm) El Tatiod (ppm) Puga Valleye (ppm)

Li Na K Mg Ca As Fe Cs Rb NH4 Cl SO4 HCO3 B F SiO2

350 63,570 21,370 12 43 — 0.7 — — 12 37,010 91,010 1900 13,800 100 130

220 32,200 3978 145 4600 — 5.2 — — — 61,415 508 305 — 4.6 235

12 1200 185 — 18 4.5 — 2.5 2.5 — 2100 32 18 28 — 560

38 3620 357 2.2 252 45 0.13 11.3 4.2 — 6470 36 46 145 2.9 184

5.9 588 57 2.1 7.3 — — 10.5 0.9 — 375 128 816 135 12 —

a

Italy (Pauwels et al., 1990). Alsace, France (Pauwels et al., 1990). c New Zealand (Rothbaum and Middendorf, 1986). d Chile, average of 12 samples. Also: Sb 40, Sr 1.13, Mn 0.38, Pb , 0.06, Ti , 0.06, Ni , 0.03, Cu , 0.02, Ag , 0.005, Zn , 0.004, S22 18, CO3 5.8, pH 7.2, temperature 84.68C (Cusicanqui et al., 1975). e India (Garrett (1998)). Also: Ba 20, Cu 2.0, Sr 0.22, TDS 2202, pH 7.9. b

thermal vents, indicating that lithium is not removed from basaltic rocks below about 1508C, and based upon d7Li data, at that temperature some lithium is adsorbed by the basalt, and some is leached. Terrigenous sediments can have some of their lithium leached and vent brine lithium re-adsorbed at 508C, but most of their lithium is not leached until the temperature is over 3508C (James et al., 2003). An ocean drilling program at the Costa Rica Rift showed that the upper , 1.8 km of sediments were enriched in lithium (5.6 – 20.8 ppm Li) due to the adsorption of lithium onto the clays that were present, and that this lithium had a heavier d7Li value (6.6 – 20.8) because of that zone’s relatively higher seawater circulation. At greater depth the rock’s lithium content was depleted to 0.6 ppm, and its d7Li values were 2 0.8 to 2.1, typical of basaltic or magmatic rocks (Chan et al., 2002b,c). Other laboratory studies have been conducted on the leaching or adsorption of lithium from rocks obtained by the Nankai Trough drilling program. Some of the rock’s lithium began to be leached at a temperature of 1508C with considerable isotopic fractionation, and the leaching was nearly completely at temperatures over 3008C. At higher temperatures there was considerable albitization of calcic plagioclase (You and Gieskes, 2001). Various studies have also been made on the lithium

Geology

45

isotopic distribution of unaltered island arc lavas and Mid Ocean Ridge basalt. The Central American arc lava has a d7Li value of 2 4.5, and the Kurile arc (eastern Russia), Sunda arc (Indonesia) and the Aleutian arcs averaged from þ 2.1 to þ 5.1, thus indicating no partitioning of lithium in these lava flows (Tomascak et al., 2002). Table 1.14 lists the lithium isotopic data for various other lithium rocks or waters. Oil Field Waters A few of the world’s oil field waters have a medium – high lithium content, with limited areas of the extensive Smackover brines in the US Gulf Coast states perhaps being the highest. One zone in both Arkansas and Texas has high-sodium and calcium chloride brines with lithium contents of 50 –572 ppm. The Texas brine has an average of 386 ppm Li, and the Arkansas brine averages 365 ppm Li (Table 1.15). All of the Smackover oil field brines appear to be concentrated seawater dolomitization brines (because of the high Ca and Br; see Chapter 2), and the high-lithium areas must have had additional geothermal input to supply the Li, B and many of the other trace ions. The brine is found at depths of 1800 – 4800 m, and its formation thickness is up to 213 m of oolitic limestone with an average porosity of about 5%. Smackover brines are commercially processed to recover bromine in Arkansas, and several studies have been made on the potential recovery of lithium from them. Various other oil field brines have medium– high lithium values such as is indicated in Table 1.15 (e.g., some Texas Cretaceous reservoirs have 132– 333 ppm Li (Table 2.5), North Dakota Devonian formations 100 –288 ppm Li and the German Altmark gasfield 263 ppm; Holdorf et al., 1993; Chan et al., 2002a; Burkowsky et al., 1991; Collins, 1976). Other Lithium Brines There are many other medium – high lithium brines in the world, with one type being most of the high-calcium chloride brines, which includes those in the Salton Sea geothermal field, the Dead Sea, some groundwaters (i.e., in Germany with 290 ppm Li) and some of the oil field and geothermal brines noted above. Most of the marine potash deposit end-liquor brines have a high to medium –high lithium content, such as the Angara-Lena basin, Russia’s 1600 – 1900 ppm, the Paradox Basin’s 66– 173 ppm Li (Tables 1.4 and 2.2), the Michigan Basin’s Sylvania Formation’s 36– 72 ppm (Table 2.4), and the English Zeichstein Formation’s 7– 65 ppm, etc. (Table 2.12). However, some end-liquors have only a nominal lithium contents, such as from the Saskatchewan, Canada potash deposits (Bottomley et al., 1999). A few calcium chloride lakes also have medium – high values, such as the Don Juan Pond’s 235 ppm, Bristol Lake’s 30 –108 ppm, Cadiz Lake’s 20 –67 ppm, and Lake Vanda’s 27 ppm (Tables 2.9– 2.11). There are other potential high-lithium brine sources that were initially mediumlithium brines extensively evaporated to recover other minerals (such as at the Great Salt Lake, Bonneville Salt Flats, the Dead Sea and the Qinghai playa noted above). The Sua Pan in Botswana (Fig. 1.24), for example contains brine with about 20 ppm Li (Table 1.9), and it is evaporated in solar ponds to produce soda ash. The endliquors should contain from 200 to 400 ppm Li, and could be further concentrated as

46

Part 1 Lithium Table 1.14 The d7Li and d6Li Values of Various Lithium-Containing Materialsa

Sea water Potash end liquor Oilfield brine, Israel River water Lakes, fresh water Great Basin closed lakes Walker Lake, Nevada Great Salt Lake, Utah Lakes, brine (Mono Lake) Springs, fresh water, Mono Lake basin Salt crusts, Mono Lake basin Springs, thermal, Mono Lake Basin Deep ocean vent brine Costa Rica rift Marine clastic sediments Near sea floor sediments Marine carbonates Orinoco River sediments Volcanic Sedimentary Volcanic rocks, Mono Lake watershed Mid Ocean Ridge basalt OIB basalt BABB basalt Costa Rica Rift dyke complex Aleutian Arc lava Central American Arc Arcs, Costa Rica to Guatemala Metabasalt, Canadian Shield Glass inclusions in olivine, Sicily Meteorites a 7

d6Li values

d7Li values

231,b 232,c 232.3d (232.1 to 236.3)d (226.3 to 217.9)c (232.2 to 26.0)b — — — — —

þ 32e — — (þ11 to þ29) f þ 32e (þ16.7 to þ 23.7)g þ 24e þ 16e þ 19.5g, þ 20 f þ31e, (þ7.4 to 31) f þ 32 f þ8.4 f — (20.8 to þ 2.1)h — (þ6.6 to þ 20.8)h —

— — (211.5 to 22.6)b — (214.7 to 20.9)b — (240 to þ12)b (230 to 222)i (222 to 27)i (24.7 to 23.4)b (26.6 to 24.7)b (22.9 to 20.7)b — — 24.5h (24.5 to 26.4)h (214.7 to 215.6)d 210.0b

— — þ3.8 f — — — (21.7 to þ7.9) j (þ 2.1 to þ 5.1) j — — — (23.4 to þ1.2)k þ 16l

d Li ¼ 1000 ([7Li/6Li]sample/[7Li/6Li]standard 2 1), while d6Li has the 6Li and 7Li reversed. The standard is 92.48% 7Li and 7.52% 6Li, so [7Li/6Li]standard is 12.29787, and the inverse is 0.081315. Thus if a sample had 6Li ¼ 7.3038%, the d7Li would be þ 32.0, while d6Li would be 231.0, making the two values of opposite sign and somewhat different. Since most other isotopic data is based upon the higher molecular weight atom, it would appear that that d7Li shoud become the satandard notation for lithium. b Kuidong and Shaoyong (2001). c Chan et al. (2002a). d Bottomley et al. (1999). e Tomascak et al. (2001). f Tomascak et al. (2000). g Tomascak et al. (2003). h Chan et al. (2002b). i Huh et al. (2001). j Chan et al. (2002c). k Guerenko and Schmincke (2002). l Ustinova (1998).

47

Geology Table 1.15 Various Analyses of Smackover Oilfield Brines, wt.% or ppmp High lithiuma

Lip Na Ca Mg K Sr Fep Mnp Bap Rbp Csp Cup Cl Br SOp4 SiOp2 Bp Ip Fp Density No. of Samples

Collins (1976)

Dow (1984) b

Texas

Arkansas

146 5.69 2.91 0.29 0.24 0.16 35 25 19

170 6.70 3.45 0.35 0.28 0.19 41 30 20 — — 1.1 17.17 0.313 450 200 134 25 — — —

386 5.49 — — 0.59 — — — — 21c 21c — — — — — — — — 1.171 6

365 5.98 — — 0.51 — — — — 11.2 6.1 — — — — — 366 — 4.6 1.229 11

— — 0.9 14.45 0.263 375 — 123 21 — 1.180 71 Li; 64 –284 others

a

Collins (1976). Maximum Li 572 ppm. Several patents on lithium recovery ion exchange resins. c Only one sample. b

other salts crystallize. Lake Abijdata in Ethiopia has similar brines and solar ponds, but is a much smaller soda ash operation. The Sebka El Adhibate, Tunisia has about a 16 ppm Li concentration in a seawater-type brine, and after solar evaporation for potential potash production the end-liquor would contain 250 – 340 ppm Li (Hamzaoui et al., 2000). There are several other solar evaporation or mineral recovery projects throughout the world with end-liquors that might also be considered for potential lithium recovery. Ore Deposits Theory of Origin The predominant type of lithium mineral formation is that of high-lithium pegmatites, although a few micas, clays and other minerals have been reported with a comparatively high lithium content. “Pegmatites are an exceptionally coarse-grained

48

Part 1 Lithium

Figure 1.24 Location of the Sua Pan, Botswana Soda Ash Plant, and a sketch of its solar ponds (Low et al., 2000; reprinted from the Eighth Symposium on Salt [ISBN 0444500650], Vol. 1, p. 523, Fig. 1, q2000, with permission from Elsevier).

igneous rock with interlocking crystals, normally found as irregular dikes, lenses or veins, especially at the margins of tremendously large bodies of (solidified) magma (flow). Most grains are 1 cm or more in diameter, and the pegmatite’s composition is usually that of granite” (-type rocks; Bates and Jackson, 1976). If a large mass of

Geology

49

magma were induced to flow, such as by the extensive shifting of crustal plates or to balance large lava eruptions, and it traveled into a zone where it could slowly cool and crystallize as it flowed, some fractionation of its components would usually occur. During any crystallization of multi-component mixtures where there is at least some circulation (such as by mixing, thermal gradient currents or diffusion) the least soluble compounds will crystallize first, followed by a sequence of minerals crystallizing based upon their solubility. For example, in the evaporation of seawater calcite, gypsum, halite, epsomite, kainite, etc., crystallize somewhat separately and in a sequence as listed. Upon cooling seawater there is a crystallization sequence of ice, hydrohalite, mirabilite, schoenite, etc. When magma was slowly cooled with some mixing the high manganese, iron and magnesium silicates would usually crystallize first and deposit predominately at the upper, cooler surface, or partially segregate as a more viscous fluid that was more available for lava flow. This would tend to leave a granitic-type composition in the remaining magma. The salts dissolved in any accompanying fluids such as super critical water, carbon dioxide or methane would also crystallize in a sequence, but the solubility of salts above the fluids’ critical point does not change much during cooling, and the critical point of water (and most gases) is much lower in pressure and temperature than the melting point of the magma constituents. Thus the accompanying aqueous fluid would usually not crystallize most of its salts until toward the end of the magma’s solidification. Further slow cooling of the low-iron, manganese and magnesium magma would form the large crystals characteristic of granite-type rocks, and if there was continued flow and/or circulation the lithium and other “rare elements” that might be present would concentrate in the final (liquid) magma and aqueous fluid (if any were present). The need for these favorable circumstances and the scarcity of lithium has resulted in only a small fraction of the world’s pegmatites having a high lithium content, but there are so many pegmatites throughout the world that there are still a large number of lithium pegmatites. Lithium is one of the more soluble cations in any magma, and its silicate compounds also have comparatively low-melting temperatures and greater fluidity (i.e., they have a lower viscosity) than the Na, K, Al, etc., silicates that would make up the bulk of the granitic magma. Also, most of the other rare metals (if present) would not be sufficiently concentrated to crystallize early in the cooling process. Thus, as the granitic melt continued to cool and crystallize with some internal mixing, the near-final melt (and the aqueous fluid, if any) would become enriched with lithium and many of the other rare elements that may have been present in that particular magma. This remaining fluid would tend to be in the warmest lower or central section of the flowing magma. Then, as the previously solidified granitictype rock suffered contraction or seismic fractures, or as the flowing magma fractured other overlying rock, because of the greater pressure with depth much of the remaining fluid would be forced into these fractures to form pegmatites. For that portion of the final magma that had experienced sufficient fractionation to concentrate the lithium and other rare metals (if present) the pegmatites would be of

50

Part 1 Lithium

the lithium type. The various components of the magma forced into the pegmatite would then crystallize, and it would usually be in a sequence or zoned pattern depending upon their concentration, solubility and the other components in the magma (and aqueous fluid). If there had been less lithium fractionation, and/or the magma cooled more rapidly in the pegmatite (as at a much shallower depth) the number of zones would be fewer, or it could even be non-zoned. Many of the lithium pegmatites occur in a scattered cluster of barren pegmatites with only a small fraction of them containing lithium, and occasionally some of the other pegmatites would contain predominantly other rare metals. The depth (pressure), temperature and amount of water present in the lithium pegmatites as they intruded into the overlying fractures and cooled is not known, although considerable laboratory experimental work has been done to help determine these factors. Since spodumene has the highest melting point (14238C; eucryptite 13978C and petalite 13568C; quartz is , 14258C) of the commercial lithium compounds, it might be expected to be the first lithium mineral to have crystallized and be the most common, as it is. This appears to have been the case with the “non-zoned” deposits and those containing massive primary spodumene. However, for the secondary spodumene deposits London (1984) has established phase data with quartz, slightly impure minerals, very high pressures and considerable water (9%) that indicates that spodumene does crystallize first above about 4 kbar pressure and 7008C, but that b-spodumene crystallizes below that pressure (and at 700– 9008C). The data further indicates that the b-spodumene would then undergo a “solid-state” conversion to petalite as it cooled below 6808C, and that petalite, in turn, could decompose to a-spodumene and quartz as the temperature dropped to 3208C and the pressure was reduced from 4 to 1.6 kbar. Below that temperature and burial depth the data showed that petalite could transform into eucryptite and quartz (Fig. 1.25). These pressures correspond to depths of 25,500 – 10,200 m, respectively, assuming an overburden density of 2.5 g/cc. Normal (low) thermal gradients would reach these temperatures. At least the petalite conversion aspect of this data appears to be confirmed in some deposits, since after the entire pegmatite was solidified much of the petalite appeared to have been transformed into spodumene and quartz, and occasionally some eucryptite and quartz. The spodumene was in the form of psuedomorphs after petalite, and its very low iron content was similar to that of petalite, and much lower than that of most primary spodumene. However, this phase data and these reactions do not appear to have general applicability, since: (1) solid phase reactions do not occur without the solid being remelted or dissolved, at least on a molecular scale, and then recrystallized after the reaction has taken place. (2) The data shows bspodumene being stable at 6808C, whereas at atmospheric pressure the change in form from a to b-spodumene only occurs in ore concentrates above about 10008C. Also, b-spodumene has never been found in a lithium deposit, and a-spodumene’s melting point is shown as 7008C compared to its atmospheric value of 14238C. (3) Changes in pressure normally have only a relatively minor effect upon a solid’s

Geology

51

Figure 1.25 Experimental phase diagram for the petalite–spodumene–quartz fields at high pressure and 9% water (with 89% petalite; London, 1984; reprinted from the American mineralogist, courtesy of the Mineralogical Society of America).

melting point, and can not cause chemical reactions, so pressure alone is unlikely to be the reason for these “solid-state” changes, or for the much lowered temperatures. Furthermore, it is hard to visualize how the deposit would be remelted in the long intervening period (presumably with the water having escaped) as the pressure was reduced by the deposit being raised or the surface eroded. (4) The pressures indicated by these tests would require an extreme burial depth of 10,200 –31,900 m (33,500 – 83,700 ft), which is very hard to reconcile with so many deposits now near or outcroping on the surface. (5) The laboratory tests were made with 9% supercritical water (water becomes critical at pressures greater than 0.22 kbar and temperatures above 3748C), but no physical evidence of water’s presence has been reported in any of the spodumene-to-petalite deposits. The materials used in the tests beside the water were 45.5% petalite (89% purity) and 45.5% quartz of unknown purity. The impurities should have only slightly lowered the melting points, implying that the super critical water reduced the melting or transformation temperatures dissolved enough of the surface of the minerals to let the solid phase reactions take place. It then slowly worked its way into all of the crystals and back out and if this is what happened, perhaps the presence or absence of water explains why only a few petalite deposits converted to spodumene, and others did not. (6) The temperatures are unbelievably low for any flowing magma.

52

Part 1 Lithium

Most of the pegmatite deposits, including the much-studied Tanco deposit in Canada have a very low water content, and yet the 9% water in these tests is a substantial amount. There are no highly crystallized adjacent cracks or passage ways, even though the escaping magma water would contain many salts, and should have deposited at least some of them. When presumably similar magma water cools or evaporates it deposits large amounts of solids, such as the many veined mineral deposits, and the skarns that formed large massive ore bodies. Even the Salton Sea geothermal brine deposits silica, silver, barium, boron and many other elements as its pressure is released. It has been speculated that aqueous fluids reacted with the wall rock and outer pegmatite zones to form the modest amounts of tourmaline at Tanco, but the tourmaline more likely crystallized from the melt, since almost all lithium sources contain some-to-considerable boron, and very little has been reported elsewhere in the deposit. Also, the very small amounts of water found in lithium mineral inclusions do not predict that there had been a major aqueous phase present during crystallization, since similar inclusions are rather common in many other mineral deposits (see the Calcium Chloride chapter). Finally, other experimental data has shown that with water and lithium minerals heated to 275– 6008C only 0.25 – 2.4 ppm solutions result, making the presence of water at reasonable pressures (and burial depths) very unlikely to have caused the petalite-to-spodumene reaction. As an alternate theory to the super-pressure, low temperature, high water content hypothesis, perhaps a re-heating of the deposit (caused by a nearby magma or pegmatite flow, intruding dikes, or even the heat of crystallization from an adjacent mineral phase) allowed sufficient softening and localized melting of the few transformed deposits for the solid phase reactions to occur (without the need for much water to have accompanied the pegmatite). Also, with this theory the magma would have been at a temperature somewhat above its atmospheric melting point, such as are observed in modern volcanic lava flows (often , 10008C or higher). Without assuming the need for water and high pressures, the high-lithium magma flow more logically would have been at a much shallower depth. This would much better explain why so many thousands of lithium pegmatites are now at, or very near the surface. Structure In most lithium pegmatites there are several zones rich in different lithium minerals, and zones with varying amounts of quartz, feldspar, mica, albite, apatite and other granitic minerals, as well as zones with other rare minerals such as those containing tantalum, niobium, tin, tungsten, cesium, rubidium, boron, fluorine and other elements. As many as 18 zones have been reported in the Bikita lithium pegmatite, with each rich in specific minerals. The lithium minerals are usually found in the intermediate-to-late (core) zones of the pegmatite, and pegmatites vary widely in the number of zones and different mineralization. As an example of the various minerals that may occur in separate zones of a pegmatite, in the early days of its production the Bikita deposit produced the wide array of minerals

Geology

53

Table 1.16 Mineral Production from the Bikita lithium deposit from 1952–1960 (Symons, 1961) Mineral products

S. tons

Approximate range of contents

Lepidolite Petalite Spodumene Amblygonite Eucryptite Beryl Tantalite Microlite Cassiterite

393,000 177,000 24,000 5000 2000 565 29 9 5

3.5–4.3% Li2O 3.6–4.7% Li2O 4.0–4.7% Li2O 7.5–9.5% Li2O 5.5–6.5% Li2O 10–12% BeO

70% Sn

listed in Table 1.16. The Big Whopper mine in Kenora, Ontario, Canada in 2002 anticipated the production of (as mt/yr): 21,200 high grade petalite concentrates (the ore contains about 1.3% Li2O), 25,400 petalite – feldspar mix, 46,200 sodium feldspar, 9400 potassium feldspar, 8400 spodumene, 22,000 mica and 12.6 tantalum minerals (Saller and O’Driscoll, 2000). By contrast, an example of a sparsely-zoned (even called non-zoned) lithium pegmatite is the North Carolina Tin Belt pegmatites which contain predominantly a fairly uniformly mixture of , 20% spodumene, 41% feldspar, 32% quartz, 6% muscovite and 1% various rare minerals, with no significant amount of secondary minerals or phase transformations (Anstett et al., 1990). There is a zonation of grain sizes in these pegmatites (Table 1.25), but very little of separate minerals. All of the commercial pegmatite deposits initially owed their value to the minerals being susceptible to hand-sorting because of their large crystal size and/or distinctive color, and now they are dependent upon the selective separation ability of flotation, heavy media and magnetic force to concentrate and purify the lithium minerals. The first commercial lithium flotation process (on spodumene) began in the late 1930s (Singleton, 1979; Manser, 1975). There are literally thousands of lithium pegmatites throughout the world, but most of them are small and not of potential commercial quality, making the actual number of large, developed deposits relatively small. There are also many lithiumcontaining minerals of potential commercial interest as seen in Table 1.2, but again, only a few have ever been mined on a large scale. Most of the lithium pegmatites are quite old, such as in Precambrian shield areas that are from 600– 3000 million years of age (Vine, 1980). Several of the largest lithium deposits are: the very large and complex lepidolite, petalite and other lithium minerals deposit at Bikita, Zimbabwe; the Gwalia Greenbushes primary spodumene deposit near Bunbury, Western Australia; the pegmatite zone that includes Tanco’s low-iron spodumene deposit at Bernic Lake, Manitoba, Avalon Ventures Ltd.’s Big Whopper petalite deposit at

54

Part 1 Lithium

Kenora, Ontario, and Emerald Field Resources Big Mack petalite deposit near Kenora, Canada; and the now-closed Tin-Spodumene Belt of Foote and FMC in North Carolina, USA. There are large commercial deposits in China, Russia and Zaire (the latter with limited lithium production), and medium-sized ones in Brazil, Namibia, Portugal, Finland and Afghanistan (the latter two not yet mined in 2002; Saller and O’Driscoll, 2000). In the past, and a few at present, of the smaller deposits throughout the world have had limited mining, such as in Rwanda, South Africa and Europe. The major commercial lithium minerals in these deposits are described in the following section. Commercial Lithium Minerals Spodumene The most abundant of the lithium minerals is spodumene (LiAlSi2O6), a lithium pyroxene containing up to 3.73% Li (8.03% Li2O), with high-grade deposits usually ranging from 1.35 to 3.56% Li (2.9 – 7.7% Li2O) and 0.007 – 0.03% Fe2O3, and the lower-grade deposits 0.5 – 1.0% Li (1.0 – 2.2% Li2O) and 0.6 – 1.5% Fe2O3. Spodumene has been classified into three types. (1) Secondary that has been formed by the conversion of petalite to spodumene ðLiAlSi4 O10 ! LiAlSi2 O6 þ 2SiO2 Þ: It is comparatively fine-grained, usually high grade and white, and has a very low iron content (0.01 – 0.04% Fe2O3). (2) Primary and zonal, with variablesized laths of large crystals in well-zoned pegmatites. It is very pure, comparatively low in iron and usually white. (3) Primary and unzoned or phenocrystic, where the crystals are usually relatively small and the spodumene low in grade, high in iron, and fairly uniformly mixed with other minerals. Its color is commonly greenish, and the deposits may be quite large (Table 1.17). Spodumene is often intermixed (or intergrown) with quartz (Fig. 1.26) and sodium or potassium feldspar, most of which can be removed by flotation and/or gravity separation to produce higher grade lithium concentrates with a lower iron content (in the US 2.3 –2.8% Li (5.0 –6.0% Li2O) and about 0.1% Fe2O3; Bach et al., 1967). The commercial product is sold at grades ranging from 2.2– 3.5% Li (4.5 – 7.5% Li2O), and with 0.01 –0.1% Fe2O3. Spodumene’s hardness is 6.5 – 7 on the Mohs scale, and its density 3.13 – 3.20 g/cc. Its crystals are monoclinic prisms with a pronounced longitudinal (110) cleavage, causing it to form lath-shaped particles upon being broken. Giant crystals have been found up to 14 m in length, but the dominant size range is 1.3 –30 cm long and 0.3 –5 cm wide. Its color is usually pale green to white, but it varies from nearly clear white to dark green when it has a high iron content. Some clear spodumene crystals are considered to be gem stones in colors of deep green (hiddenite) or yellow and lilac (kunzite). It decomposes by surface weathering to kaolinite and/or hydrous mica. Several ore analyses are listed in Table 1.18, and some of the larger spodumene deposits are in: Greenbushes, Australia; Ontario and Manitoba, Canada; North Carolina, USA; Bikita, Zimbabwe; Minas Gerais, Brazil; the Chita Region, Russia; and the Altai Mountains, China (Heinrich et al., 1977; Kesler, 1960).

Geology

55

Table 1.17 Examples of Various Types of Spodumene Deposits (Heinrich et al., 1977) Iron content Location

Color

FeO (%)

Fe2O3(%)

Spodumene’s description Pegmatites are essentially unzoned. Disseminated uniformly in fine-grained quartz–microcline– muscovitea Well-zoned pegmatite. Giant crystals in quartz-rich intermediate zonea Complex horizontally zoned pegmatite. Laths with quartz directly below quartz corea Complexly zoned pegmatite. In lepidolite–cleavelandite– quartz–tourmaline rockb Well-zoned pegmatites. Very coarse-grained, with quartza Complex horizontally zoned pegmatite. Spodumene–quartz with zones of lepidolite and some amblygoniteb

Kings Mountain, NC

Greenish, semi-vitreous

0.08–0.10

0.70–0.80

Etta mine, Black Hills, SD

Buff, dull

0.02

0.03

Harding mine, Dixon, NM

White-pale buff, dull-vitreous

0.01

0.01–0.03

Strickland Quarry, Portland, CT Pala, CA

Pale pink, vitreous

0.02

0.03

Pink, gem quality

0.10

0.02

0.01–0.02

0.02–0.04

Bernic Lake, White to pinkish, Manitoba (Tanco) vitreous

a b

Primary spodumene. Secondary spodumene.

Figure 1.26 Typical crystals of lithium ore: left, zinnwaldite; right, spodumene; both with quartz (Deberitz, 1993, courtesy of Chemetall GmbH).

56

Table 1.18 Analyses of Various Lithium Ores, wt.% Spodumene

b

China Indiaa Li2O K2O Na2O Al2O3 Fe2O3 SiO2 MgO CaO MnO Rb2O Cs2O F P2O5 As2O3 CuO H2O Acid Insol. Volatile, 4008C

3.70 10.69 1.12 33.17 2.90 47.57 — — — — — — — — — — — —

Ore

Conc.

South Dakotac

Tancod

1.63 2.98 5.12 18.91 0.17 68.84 0.01 0.06 0.12 0.33 — 1.86 0.49 — — — — —

4.65 8.35 1.13 23.64 1.29 55.33 — — — 1.18 0.20 1.11 — — — — — —

3.76 7.00 1.24 29.14 0.37 49.18 — — — — — 2.30 — — — 1.35 — —

7.28 0.15 0.15 26.00 0.045 — — — 0.03 — — 0.015 0.20 — — — — —

Petalite

Li2O K2O Na2O Al2O3

Green.e 2.53 1.50 2.57 13.94 0.94 73.74 0.16 0.45 — 0.32 0.04 — 0.20 79 ppm — — — —

Wek.f

Afg.g

Port.h

6.60 1.50 0.27 26.70 1.29 64.60 0.07 0.02 0.14 0.24 0.04 0.23 — — — — — —

7.56 0.05 0.16 27.42 0.24 64.39 0.06 0.06 0.05 — — — 0.01 — — 0.06 — —

7.65 — 0.01 27.20 0.23 65.09 — — 0.01 — — — 0.06 — — — — —

Eucryptite

Brazil i

BigWhopp.j

Port.h

Bikita.k

Montebrasiteh

Liconsl

4.35 ,0.01 ,0.01 17.06

4.67 0.01 0.01 16.24

4.49 — 0.01 16.75

11.49 , 0.05 0.05 40.40

8.87 — 0.01 35.07

20.30 0.59 22.13 —

(continues)

Part 1 Lithium

Lepidolite

Table 1.18 (continued) Petalite

Fe2O3 SiO2 MgO CaO MnO Rb2O Cs2O F P2O5 As2O3 CuO H2O Acid Insol. Volatile, 4008C

Eucryptite

Brazil i

BigWhopp.j

Port.h

, 0.01 78.37 , 0.01 , 0.02 , 0.01 — — — , 0.01 — — 0.13 — —

0.01 77.93 0.00 0.02 0.01 — — — — — — — — —

0.05 78.91 — — 0.03 — — — 0.05 — — — — —

Bikita.k 0.08 47.92 — 0.09 , 0.01 — — — , 0.01 — — 0.23 — —

Montebrasiteh

Liconsl

— — — — — — — 0.04 49.75 — — 0.55 — —

0.14 — — 0.54 — — — — 46.40 0.41 0.01 3.39 0.08 2.25

a

Lepidolite from Bihar, Rajasthan, India (Vyas et al., 1975). Lepidolite from Yichun, Hunan, China. Conc. (Xu et al., 1998); Also in the ore: FeO 0.15%, TiO2 0.06%; as ppm: Ta 137, Nb 68.9, Zr 18.7, Sr 6.9, Ce 2.52, Nd 1.63, Sm 0.95, La 0.94, Gd 0.84, Eu 0.11 (Yin et al., 1995). The concentrate has had one stage of flotation and roasting (reducing its fluorine). c Lepidolite from the Black Hills, South Dakota (Page, 1953). d Spodumene concentrates, Tanco, Canada (Burt et al., 1988). e Spodumene ore from Greenbushes, Australia; Also: TiO2 0.19%, as ppm: Sn 179, Ta2O5 86, Be 74, Nb2O5 61, Sr 34, Ni 12, Zr 11.5, U 6.2, Th 3.1 (Partington et al., 1995). f Spodumene from Wekusko Lake, Manitoba (Dresler et al., 1998). g Spodumene from Nuristan, Afghanistan (London, 1984). h Petalite from Minas Gerais, Brazil; also 0.02% TiO2 (London, 1984). i Spodumene, petalite and montebrasite from the Covas de Barroso district, Portugal (Charoy et al., 2001). j Petalite from the Big Whopper deposit, Separation Rapids, Ontario, Canada (Pearse and Taylor, 2001). k Eucyrptite from Bikita, Zimbabwe; also ,0.01% TiO2 (London, 1984). l Licons from Searles Lake, California (Stenger, 1950). b

Geology

57

58

Part 1 Lithium

Petalite (LiAlSi4O10) has a monoclinic crystal habit with a framework silicate structure, and its cleavage is in two planes 1148C apart. It has a density of 2.4 g/cc, and a hardness of 6. Its color is white, grayish white and more rarely pinkish, with a theoretical lithium content of 2.27%(4.88% Li2O), while the commercial deposits vary from 1.4– 2.2% Li (3.0 –4.7% Li2O; Kesler, 1960) and the standard sales grade is 2.0% Li (4.3% Li2O). The petalite crystal does not accommodate very much iron, so its deposits have a very low iron content. In some deposits it has been transformed to quartz and spodumene, as noted above. Various larger deposits of petalite occur in: Bikita, Zimbabwe; Kenora, Ontario, Canada; Karibib, Namibia; Aracuai, Brazil; Londonerry, Australia; the Transbalkin area of Russia; and at Uto, Sweden. Lepidolite [K2(Li,Al)5-6(Si6 – 7Al1 – 2O20)(OH,F)4] or [K2Li2Al4Si7O21(OH,F)3] or [KLiAl2Si3O10(OH,F)3] is a mica with a complex and variable formula. Its lithium concentration ranges from 1.39% (3.0% Li2O) to a theoretical maximum of 3.58% Li (7.7% Li2O). Its density is 2.8 – 3.0 g/cc, it has lamellar cleavage, and its crystals have a book-type structure. The books range in size from microscopic to about 5 cm in thickness. Normal lithium concentrations in commercial deposits range from 1.4 –1.9% Li (3.0 – 4.1% Li2O), although in the early days of shipments from Rhodesia (now called Zimbabwe) the ore was hand-sorted (based upon its beautiful violet-colored crystalline clusters), and the lithium percentage was at or higher than theoretical. This was probably because of the presence of some of the higher grade range of the series such as zinnwaldite (Fig. 1.26), polylithionite or protolithionite (Kesler, 1960). Lepidolite also contains a high, but variable concentration of potassium, rubidium and cesium, which made a valuable by-product for the former American Potash & Chemical Co. plant at San Antonio. Several analyses of lepidolite samples are shown in Table 1.18. The major commercial deposits of lepidolite are in: Bikita, Zimbabwe; Bernic Lake, Manitoba, Canada; Karibib, Namibia; Mina Gerais, Brazil; and Sociedad Mineria de Pegmatites, Portugal. Amblygonite (LiAl[PO4][F,OH]) is the fluorine-rich end member of a lithium aluminum phosphate group, and montebrasite is the hydroxyl-rich end member. Amblygonite’s color is generally white or creamy, although it can vary from colorless to many other pale colors. It resembles potassium feldspar, but with a bluish or grayish tint instead of feldspar’s cream or salmon tint. It is usually found as fine anhedral and compact crystals, but it is sometimes found as short prismatic, tabular or equant (lath-shaped) crystals. It can also occur with polysynthetic twinning in two directions at 908C. It has a vitreous, greasy or pearly luster on its cleavage planes, and cleaves in all four directions with non-right angles. The crystals have a density of 2.98– 3.11 g/cc and a hardness of 5.5 –6. Its theoretical lithium content is 4.76% (10.2% Li2O), but most commercial ores contain 3.5– 4.4% Li (7.5 – 9.5% Li2O). It has been mined in Canada, Brazil, Surinam, Zimbabwe, Rwanda, Mozambique, Namibia, South Africa and the Black Hills and Pala Districts in the United States (Kesler, 1960).

Geology

59

Eucryptite (LiAlSiO4) has a theoretical Li content of 5.53%, (11.84% Li2O) and its ores average 2.1 – 3.0% Li (4.5 – 6.5% Li2O). The only large deposit is at Bikita, Zimbabwe with an average grade of 2.34% Li (5.0% Li2O), and much of the impurity is quartz (Kunasz, 1994). In the early days of the lithium industry eucryptite and amblygonite were the favored minerals, since the lithium could be leached directly (without roasting) by strong acids. However, their deposits are fairly uncommon, and those that were initially worked were quite small. Several of the more well-known present or formerly operated lithium deposits are listed in the following sections. Gwalia Consolidated Ltd. (Greenbushes), Western Australia The Greenbushes lithium pegmatite deposit is located about 300 km south of Perth (220 km, Tambourakis et al., 1990) and 80 km southeast of the port of Bunbury in the center of a 100 km2 pegmatite field. The main lithium zone of the pegmatite is 3.3 km long and up to 230 m wide in a 7 km long north– south band of pegmatites that are up to 1 km long (Figs. 1.27 and 1.28). The only lithium mineral is a low-iron primary spodumene (Table 1.18), but there are also major amounts of tin, tantalum (stated to contain half of the world’s tantalum reserves; Partington et al., 1995) and niobium, as well as kaolin in the deposit’s weathered overburden. The pegmatite has intruded into a gigantic granitoid mass, with the contact minerals

Figure 1.27 Top and side views of the Greenbushes pegmatite (Partington et al., 1995; reprinted with permission from Economic geology, Vol. 90:3, p. 620, Fig. 3, Partington, G. A., McNaughton, N. J. and Williams, I. S., 1995).

60

Part 1 Lithium

Figure 1.28 Schematic plan of the zoning in the Greenbushes pegmatite (Partington et al., 1995; reprinted with permission from Economic Geology, Vol. 90:3, p. 621, Fig. 4, Partington, G.A., McNaughton, N. J. and Williams, I.S. 1995).

being foliated greenstone and dolerite. The hanging wall dips to the west 40 –508C, and the footwall 55– 608C west. The spodumene is relatively pure (about 50%, or 4.01% Li2O), and has an unusually high rubidium content. The main accessory mineral is quartz (about 49%), together with minor amounts (, 1%) of feldspar, mica, tourmaline, apatite and beryl. The pegmatites were formed as a series of linear dikes, varying greatly in their width and length (Fig. 1.27). The main lithium zones are in the hanging and footwalls (Fig. 1.28), with the hanging wall lithium decreasing in thickness to the south and eventually disappearing. The hanging wall lithium is generally richer with up to 5% Li2O(2.32% Li; 60 –80% spodumene), but the footwall is laterally more continuous. At the top of each zone the spodumene consists of coarse-grained euhedral crystals (many over 5 cm in width) intergrown with quartz, and with a lustrous white or pink color. The crystals are finer-grained in the center of each zone, and are intergrown with quartz and potassium feldspar. The central area also has some lenses of quartz-albite or microcline, and all of the spodumene contains various impurities of apatite, tourmaline, muscovite, beryl and tantalite. There are also some zones of spodumene in the centers of albitic pegmatite dikes and pods to the north of the main lithium pegmatite. It is speculated that the footwall spodumene crystallized soon after the initial pegmatite intrusion 2527 MyrBP (million years before the present; Archean age), while the hanging wall crystallized

Geology

61

slightly later (perhaps as much as 7 Myr). The granitoid host rock may have formed about 90 Myr before the main pegmatite intruded into it. London’s (1984) phase data (Fig. 1.25) would indicate that foot and hanging wall temperatures would be 770 and 6908C, and the pressure 5.5 and 5 kbar, respectively. However, this would mean burial depths of over 32,000 m and a thermal gradients of only 2.148C/m (1.188F/ft) despite the magma flow being a high temperature event. More logically the temperatures were much higher than this, and the pressures (and depth) much lower. The pegmatite was later cut by a swarm of near-vertical, generally east – west trending dolerite dykes (Partington et al., 1995; Tambourakis et al., 1990). The Gwalia mine on this deposit is the largest producer of lithium mineral concentrates in the world, with ore reserves to a 220 m depth estimated to be 42 million tons averaging 1.36% Li. The mining zone in 1993 had reserves of at least 7.9 million mt of ore containing 2.02% Li (4.35% Li2O) and 0.12% Fe2O3, as well as 3.63 mm mt of 1.58% Li (3.40% Li2O), 0.27% Fe2O3 ore. The cut-off point for high grade ore is above 4.0% Li2O (i.e., about 50% spodumene) and 0.1% Fe2O3, and the low grade ore averages about 3.0 – 4.0% Li2O (37.5 – 50% spodumene; Flemming, 1993). Other much smaller occurrences of lithium ores in Australia have also been noted, including the Coolgardie District with petalite, spodumene and amblygonite; Ravenstrope’s spodumene; Wodgina’s lepidolite; and Euriowie, New South Wales’ amblygonite (Harben and Edwards, 1998; Kunasz, 1994). Bikita Minerals, Zimbabwe A large pegmatite area occurs about 75 km east (69 km ENE; Cooper, 1964) of Fort Victoria and 64 km NE of Masvingo, Zimbabwe, extending for 1700 m and varying in width from 30 –70 m (average 64 m; Figs. 1.29– 1.31). It generally dips from 14 –458C east, but at the 1800 £ 60 m outcrop it strikes north, and dips 358C. Because of the dip the true deposit thickness is about 23 m, and in 1964 it had been worked to a depth of 60 m and drilled to 152 m. The deposit’s age is about 2650 million years, and thus from the Archean period. The pegmatite area is divided into the Al Hayat, Bikita, Southern and Nigel sectors, and it is distinctly zoned. Table 1.19 lists the general order of zoning in the two sectors, and compares them with a general zoning model for many of the North American pegmatites that had been explored during the same period. The deposit has an unusual variety and tonnage of commercial lithium minerals, as well as tantalum, tin, beryl and pollucite (a cesium mineral). Its proven reserves were 6 million ton at 1.35% Li in 1961 (Cooper, 1964), 12 million mt with an average grade of 1.4% Li in 1979, and 23 million mt in 2002. The pegmatites outcrop into the area’s granite and ironstone (laterite) rocks, and the first commercial operations were for tin, tantalum and microlite from their weathered surfaces, but these alluvial deposits were soon depleted. The wall

62

Part 1 Lithium

Figure 1.29 Map of the Bikita pegmatite and the surrounding area (Symons, 1961).

rocks of the lithium pegmatites are massive Precambrian greenstone (finely banded quartz –amphibolite schist) containing occasional zones of black tourmaline (some with up to 7.5 cm crystals) and biotite. The Li Cs minerals in the pegmatites are distinctly zoned and originally contained in order of abundance: petalite, lepidolite, spodumene, pollucite (H2O·2Cs2O·2Al2O3·9SiO2), beryl ðBe3 Al2 ðSiO3 Þ8 Þ; eucryptite (LiAlSiO4), amblygonite and bikitaite (H2LiAlSi2O7).

Geology

63

Figure 1.30 Geology on 900 ft level horizon, Al Hayat and Bikita sectors (Symons, 1961).

The accompanying minerals include quartz, muscovite, microcline and albite, with very little tourmaline being present. As an example of this highly zoned (and thus easily recoverable) mineralization, Table 1.16 lists the products shipped from the deposit between the years 1952 –1960. The largest deposits of petalite occur in the Al Hayat sector (Fig. 1.30) where very large crystals of low-iron (0.03% Fe2O3) petalite occur with massive microcline (potassium feldspar) in a matrix of fine-grained albite, muscovite and lesser amounts of quartz. Often there are roughly equal amounts of petalite, microcline and the matrix. The petalite occurs as laths up to 1.8 m long and 46 cm wide with a pronounced platy cleavage, and the laths are often oriented at right angles to the walls of the pegmatite. The laths size and quantity decrease toward the footwall of the deposit, but all of the petalite crystals are quite pure, with essentially no feldspar or other intergrowths. There are also zones of petalite– feldspar in which both minerals occur in giant crystals up to 2.4 m in length in a matrix of albite, quartz and muscovite. The microcline could be white, cream colored or grey, with occasional pale lilac flakes of lepidolite. Typical analyses of the petalite and other lithium minerals found in the deposit are shown in Table 1.20. The Bikita sector has dimensions of about 427 £ 29 and 64 m deep, and lepidolite was originally its dominant lithium mineral (most of it has now been mined), but there were also major amounts of spodumene, petalite and amblygonite. The lepidolite was associated with quartz (but no feldspar) in three distinct zones: (1) A lepidolite – quartz shell on the upper side and the ends of the deposit with about

64 Part 1 Lithium Figure 1.31 Mineral zonation in the 900 and 1000 ft levels in the Bikita pegmatite (Cooper, 1964). Reproduced with permission of the Geological Society of South Africa.

65

Geology Table 1.19 Typical Lithium Pegmatite Zones at Bikita and in North America (Cooper, 1964)

A. Zones of the Al Hayat sector Hangingwall greenstone H.W. contact Border zone

Selvedge of plagioclase, quartz, muscovite Mica band Hangingwall felspar zone Petalite-felspar zone Spodumene zone Felspathic lepidolite zone Pollucite zone All-mix zone (felspar, quartz, muscovite, lepidolite, etc.) Footwall felspar zone

Wall zones Intermediate zones

Wall zone

F.W. contact Footwall greenstone B. Zones of the Bikita sector Hangingwall greenstone H.W. contact Border zone Wall zones Intermediate zones (upper)

Core zones

Intermediate zones (lower) Wall zone

Selvedge (plagioclase, quartz) Muscovite band Hangingwall felspar zone Petalite and felspar Spodumene (i) massive (ii) mixed Pollucite Felspar-quartz “All-mix” zone Massive lepidolite (i) high grade core (ii) near solid Lepidolite-quartz shell (i) lepidolite (ii) amblygonite “Cobble” zone Felspathic lepidolite Footwall felspar zone (i) rhythmically banded beryl zone (ii) muscovite band (iii) “spotted dog” F.W. contact Footwall greenstone (continues)

66

Part 1 Lithium Table 1.19 (continued)

C. The American classification of zones Border zones Relatively fine-grained selvedge, generally only a few inches thick. Wall zones Coarser and much thicker. Both wall and border zones are more continuous and more constant in thickness than those which follow. Intermediate zone Any zone between the wall zones and the core is termed intermediate. Such zones are very variable in shape, size, and continuity. Core zone The innermost or central zone, commonly elongate or as a series of disconnected segments (or lenses). Reproduced with permission of the Geological Society of South Africa.

40% lepidolite, which was in 0.3– 3.1 m diameter masses separated by barren quartz layers. (2) A zone with 60– 70% coarse grained lepidolite (with quartz) that was easily hand-sorted. And (3) a central core of almost pure (. 90%) fine-grained lepidolite that was a beautiful mauve color, and very dense and tough. It often contained small amounts (less than 1%) of microscopic-sized topaz. Near the ends of the lepidolite deposits there were giant (0.3 – 6.2 m across) masses of white, irregular crystalline amblygonite in a grey quartz. It had sharp contacts to the lepidolite, and contained no intergrowths. In addition to the lepidolite core there were other lepidolite-containing zones such as the “all mix zones” in which microcline, coarse lepidolite and quartz occurred in spectacular arrangements, sometimes surrounding laths of microcline, petalite or spodumene. Lepidolite could also occur in a “cobble zone” where fairly pure boulders or lenses of lepidolite were found in a matrix of fine albite. In a “felspathic lepidolite” zone the mixture consisted of fine-grained albite, quartz and variable amounts of disseminated lepidolite. In the “wall zone”, and particularly the footwall there were rhythmically banded layers of fine grained lepidolite and albite, with irregular and wedge-shaped beryl crystals lying across and through the banded layers. The layers varied from 2.5 mm – 10 cm in thickness, and often the lower edge of the lepidolite had a sharp boundary with the albite, while its upper edge graded into the albite. The layers undulated and were occasionally contorted. In the Bikita sector spodumene occurred in both massive and mixed zones as acicular intergrowths with quartz, and in the form of blocks or laths from 5 cm to 5.5 m in length. In the more massive zones the blocks were separated by 2 – 5 cm of

Table 1.20 Typical Analyses of the Lithium Minerals in the Bikita Pegmatite, wt.% (Cooper, 1964) Petalite

Eucryptite

Spodumene–quartz intergrowth

Lepidolite

Amblygonite

Pollucite

SiO2 Al2O3

76.79 16.85

73.98 18.15

76.50 17.09

56.24 24.65

1.62 33.36

47.09 17.41

Li2O Rb2O Cs2O

4.36 0.00 0.00

4.98 0.00 0.00

4.12 0.00 Trace

3.64 2.71 0.31

8.60 ND 0.00

K2O Na2O CaO MgO

0.00 0.46 0.31 0.21

0.23 0.51 0.23 0.07

0.71 0.94 0.36 Trace

7.20 0.26 ND ND

0.20 1.00 0.00 0.72

P2O5 F

0.00 0.02

0.00 0.02

0.08 ND

ND 5.10

43.97 3.48

Fe2O3 TiO2 MnO

0.05 0.00 0.04

0.06 0.00 0.12

0.017 0.00 ND

a

a

a

ND ND

ND 0.03

0.00 ND

BeO Loss on ignition

0.006 0.65

0.01 1.56

ND 0.64

ND 1.12

ND 6.24

ND 2.20

99.75

99.92

100.46

ND, not determined. Reproduced with permission of the Geological Society of South Africa. a These samples contained metallic iron introduced during sample preparation.

101.23

99.22

2.00 3.02 0.00 0.00 0.30 ND

99.94

Geology

Total

0.41 0.91 26.60

67

68

Part 1 Lithium

a lepidolite – quartz mixtures. The lepidolite appeared as radiating aggregates with their centers near the walls of the spodumene blocks. In the mixed spodumene zone the percentage of spodumene – quartz laths was lower, and they were usually accompanied by darker lepidolite, along with albite and cleavendite. The pollucite zone of the deposit was massive and had little apparent crystal faces or other indication of grain size, although some cleavage planes indicated crystals 15 cm or so in width. The pollucite was randomly honeycombed by veins of lepidolite 6 mm or so in width that divided the mass into a mosaic pattern with each section 7.5 –30.5 cm in size. The lepidolite veins stopped abruptly at the boundaries with other minerals. The cesium content of the pollucite was about 25% Cs2O, compared to pure pollucite with about 30% Cs2O. The impurity, other than lepidolite appeared to be very fine particles of quartz and some other minerals. The pollucite ore body enclosed a mass of petalite 1.5 m wide, a similar mass of microcline, and a 6.2 m wide lens of lepidolite. In each case the boundaries were sharp between the different minerals. Other minor lithium mineralization in Zimbabwe occurs in the Insizia, Matobo, Mazoe, Mtoko, Salisbury, Umtali and Wankie districts (Kunasz, 1994; Cooper, 1964; Symons, 1961; Kesler, 1960). Tantalum Mining Corp. (Tanco); Bernic Lake, Manitoba, Canada This high-lithium pegmatite deposit also contains tantalum, cesium, rubidium and beryllium. It occurs about 180 km NE of Winnipeg (Fig. 1.32) in the Cat LakeWinnipeg River district, with the plant on the northeastern shore of Bernic Lake, and most of the deposit under the Lake (Fig. 1.33). The 2.55 –2.65 billion year old pegmatite has intruded into the Archean age Bird River greenstone belt in the Canadian Shield of southeast Manitoba, and is zoned in a complex manner with several lithium ores, tantalum, cesium and other industrial minerals (Tables 1.21 and 1.22). It is lens-shaped with a maximum length of 1990 m, a width of 1060 m and it is up to 100 m thick (Fig. 1.34; 1440 £ up to 820 m by . 100 m depth, Crouse et al., 1984; 2400 –3,000 long £ 820 m wide, dipping west 7– 128C, Cerny and Lenton, 1995). In 1984 it was probably the second largest known complex-zoned lithium deposits after the Bikita pegmatite, with an unusually large number of minerals (about 100) having been identified in the deposit (Table 1.23). The primary lithium mineral is spodumene, occurring in two separate zones, but there is also considerable petalite, lepidolite and amblygonite (typical analyses are given in Table 1.24). The proven spodumene reserves were estimated to be 7.4 million mt of ore in 1984 with an average of 1.34% Li (2.88% Li2O), and the three lepidolite areas had lithium concentrations ranging from 0.87 to 1.31% Li (1.87 – 2.82% Li2O; Vanstone et al., 2002; Harben and Edwards, 1998; Kunasz, 1994; Burt et al., 1988; Crouse et al., 1984). The Tanco pegmatite appears to have been formed in a series of faults in the overlying granite or basalt that allowed the intruding pegmatite magma to spread and lift the overlying rock, thus forming its somewhat horizontal, lens-like shape

Geology

69

Figure 1.32 Location of the Tanco and Separation Rapids pegmatites (Pearse and Taylor, 2001; figure first published in the CIM Bulletin, Vol. 94, No. 1049. Reprinted with permission of the Canadian Institute of Mining, Metallurgy and Petroleum).

(Vanstone et al., 2002). It is composed of nine different mineral zones (Table 1.23), with the different ores of commercial interest, tantalum, spodumene, cesium and rubidium, each occurring primarily in separate zones. Most of the spodumene is in the Upper Intermediate Zone and to a lesser extent the Lower Intermediate Zone (Fig. 1.34). The Upper Intermediate Zone is a lens up to 24 m thick, and overlying the central portion of the pegmatite. It appears to have initially crystallized as coarse-grained petalite and potassium feldspar, with crystals of each up to 13 m long. There were also some coarse-grained primary spodumene blades in quartz, albite and other minerals. After the original cooling process most of the petalite and potassium feldspar was transformed into intergrowths of fine-grained spodumene and quartz, in the form of psuedomorphs after the original minerals. Very little of the petalite or potassium feldspar remained unchanged. Since petalite crystals appear to not be able to incorporate much iron, the resulting spodumene has an unusually low iron content (less than 0.05% Fe2O3; Table 1.24). Small amounts of eucryptite and several other minerals appear to have also formed from a more extensive transformation of the petalite –potassium feldspar mixture or from the secondary spodumene. Only the Border and Wall Zones of the Tanco deposit occur as concentric shells around the entire pegmatite, although when combined the Lower and Upper

70

Part 1 Lithium

Figure 1.33 Areal view of the Tanco pegmatite showing the extent of the Upper Intermediate Zone (Burt et al., 1988) (this figure appeared in Industrial Minerals No. 244, January 1988, p. 54. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. q Metal Bulletin plc 2003).

Intermediate spodumene zones also form a fairly uniform shell. The normally central Quartz Zone is only central in the western portion of the pegmatite, while the Border Zone is relatively thin (2 – 30 cm) and composed mainly of fine-grained albite and quartz. The hanging Wall Zone contact is relatively sharp and characterized by the presence of brick-red perthite and schrol, with considerable tourmaline and other minerals, as well as some unaltered petalite –potassium feldspar. The Upper Intermediate Zone is in contact with the hanging Wall Zone, and usually the Lower Intermediate Zone is in contact with the up to 35 m thick lower Wall Zone. In the upper 2 m of the Upper Intermediate Zone the potassium feldspar is a pinkish color, decreasing in intensity from the contact to about 2 m depth. As it mixes with spodumene the spodumene attains a greenish color, which decreases in intensity with depth. Most of the spodumene is in the form of “SQUI” (or squi; spodumene –quartz intergrowths that are psuedomorphs after primary petalite with a very large crystal size). The zone also contains some large blades of primary spodumene (up to 7 –13 m long), as well as areas of quartz and amblygonite. The up to 25 m thick Lower Intermediate Zone has a footwall contact that is more gradual, with a decrease in grain size and the amount of SQUI, although some crystals are up to 2 m long. There is an increasing content of other minerals as the contact is approached, including some 0.5– 2 m quartz pods with spodumene and

Table 1.21 Zonation and Location of the Economic Minerals in the TANCO Pegmatite (Crouse et al., 1984)

Zone Exomorphic unit (1) Border zone

Main constituents Biotite, tourmaline, holmquistite albite, quartz

(3) Aplitic albite zone

Albite; quartz (muscovite)

(4) Lower intermediate zone

Microcline–perthite, albite, quartz, spodumene, amblygonite Spodumene; quartz;

(5) Upper intermediate zone

amblygonite

(6) Central intermediate zone

Microcline – perthite

(7) Quartz zone

Quartz

quartz, albite, muscovite

Tourmaline, apatite, (biotite) (beryl, triphylite) Beryl (tourmaline)

Muscovite, Ta; oxide minerals; beryl, (apatite, tourmaline, cassiterite) (ilmentie, zircon, sulfides) Li-muscovite, lithiophilite (lepidolite, petalite, Ta-oxide minerals) Microcline–perthite, pollucite, lithiophilite (albite, Li-muscovite), (petalite, eucryptite, Ta-oxide minerals) Beryl, ðTa-oxide mineralsÞ; (zircon, ilmenite, spodumene, sulfides, lithiophilite, apatite, cassiterite) (Spodumene, amblygonite)

Fine-grained reaction rims and diffuse veins Fine-grained layers

Geochemically important major and (minor) elements K, Li, B (P, F) Na (B, P,Be,Li)

Medium-grained, with some giant K-feldspar crystals Fine-grained undulating layers, fracture fillings, rounded blebs, diffuse veins Medium- to coarse-grained, non-homogeneous

K, Na (Li, Be, F)

Giant crystal size of major and most of the subordinate minerals

Li, P, F (K, Na, Cs, Ta)

Medium-to coarse-grained

K (Na, Be, Ta, Sn, Zr, Hf, Ti)

Monomineralic

Si (Li)

71

Albite, quartz, muscovite, Li-muscovite, microcline–perthite

(Arsenopyrite)

Textural and structural characteristics

Geology

(2) Wall zone

Characteristic subordinate (accessory), and (rare) minerals

Na (Be, Ta, Sn, Zr, Hf, Ti)

K, Na, Li, P, F (Ta)

(continues)

72 Part 1 Lithium

Table 1.21 (continued)

Zone

Main constituents

(8) Pollucite zone

Pollucite

(9) Lepidolite zone

Li-muscovite; lepidolite; microcline–perthite

Characteristic subordinate (accessory), and (rare) minerals Quartz, spodumene (petalite, muscovite, lepidolite, albite, microcline, apatite) Albite, quartz, beryl, (Ta-oxide minerals; cassiterite), (zircon)

Textural and structural characteristics

Geochemically important major and (minor) elements

Almost monomineralic

Cs (Li)

Fine-grained

Li, K, Rb, F (Na, Be, Ta, Sn, Zr, Hf, Ga)

Underlined minerals occur in economic quantities in the zones indicated. Table published in the Geology of Industrial Minerals in Canada. Reprinted with permission of the Canadian Institute of Mining, Metallurgy and Petroleum.

Table 1.22 Estimated Distribution of the Major Minerals in the Tanco Pegmatite, wt.% (after Cerny et al., 1998) Zonesa

(1)

Biotite

0.1

Quartz Albite K-feldspar Muscovite Li-Muscovite Lepidolite

(2)

(3)

(5)

(6)

28.7 66.0

36.0 40.7 15.0

27.0 67.0

38.0 25.0 24.0

29.0 7.0 25.0

15.0 20.0 50.0

1.0

3.0 3.0

3.0

1.0 2.0 0.5

0.1 0.1 0.1

12.0

6.4

33.0 2.0

0.1

0.1

Apatite Lithiophilite Amblygonite

2.0 0.1

Beryl Pollucite Tourmaline

0.2 2.0

(7)

(8)

(9)

Bulkb ,0.01

Spodumene Eucryptite 0.4 0.2 0.1

0.2 0.5 1.5

0.1 1.0 1.0

0.2 0.5

0.5

1.0

0.1 0.5 0.1

0.1 1.0 0.1

1.0

0.02

0.01 0.02

0.5

0.01

0.02 0.02 0.02

5.5 5.0 2.5

10.0 8.0 10.0

36.07 25.56 22.05

70.0

2.97 3.00 0.27

2.0

0.1

0.1 0.1 0.5

1.0

94.9 0.1 2.0

7.0 2.6

1.7 0.5

0.4

0.5

6.41 0.26 0.1 0.2 0.5

0.15 0.38 0.77

0.5

0.37 1.28 0.38

75.0 0.1 0.02 0.02 0.01

0.02

0.01 0.01 ,0.01 (continues)

Geology

Cassiterite Rutile Ferrotapiolite

(4)

73

74

(continued) Zonesa

(1)

Columbite group Wodginite group Microlite group Simpsonite

(2)

(3)

(4)

(5)

(6)

0.02

0.03 0.05 0.02

0.02

0.02 0.02 0.02

0.01 0.05 0.03 0.01

0.005 0.01

0.005 0.02

Uraninite Zircon Totals

0.02 100.00

99.93

99.48

Courtesy of the International Mineralogical Association. See Table 1.21 for the zone designations. b Percent of total pegmatite. a

0.02

99.76

99.71

99.18

(7)

100.00

(8)

(9)

Bulkb

0.03 0.01

0.01 0.03 0.02

0.02 0.01 0.01 ,0.01

0.005 0.02

,0.01 ,0.01

99.74

99.41

100.00

Part 1 Lithium

Table 1.22

Geology

75

Figure 1.34 Three representative north–south cross sections, and a longitudinal east– west section of the Tanco pegmatite (Crouse et al., 1984). Figure published in the Geology of Industrial Minerals in Canada. Reprinted with permission of the Canadian Institute of Mining, Metallurgy and Petroleum.

amblygonite, and some cloudy pink primary spodumene columns up to 18 cm long. In both spodumene zones there is very little disseminated tourmaline. The Lepidolite Zone is up to 18 m thick in the form of two elongated sheets within part of the Central Intermediate Zone. It is in contact with the Upper Intermediate Zone, and consists of fine-grained lithian muscovite and lepidolite with an unusually high rubidium content. It also has a moderate amount of tantalum minerals, for which it is mined (Table 1.22; Vanstone et al., 2002; Cerny et al., 1998; Burt et al., 1988; Crouse et al., 1984). North Carolina Tin-Spodumene Belt; Chemetall GmbH (Foote Minerals originally, and Cyprus Foote until 1998) The combined spodumene reserves in different sections of this 60 km long, . 1.6 km wide belt (Fig. 1.35) make it one of the world’s large lithium deposits, containing 185,000 mt of Li that was proven, and perhaps twice that in total ore averaging 0.7% Li (1.51% Li2O). All of the ore had a fairly high iron content, such as often about 0.6– 0.9% Fe2O3 (Henderson, 1976). Numerous pegmatites had intruded to the surface in this area, with the lithium pegmatites up to 1000 m long, 90 m wide, and more than 200 m deep. They usually did not have the typical zoned structure, and spodumene was essentially the only lithium mineral except for a little

76

Table 1.23 Various Minerals Found in the Tanco Pegmatite (Cerny et al., 1998)

Lead Bismuth Arsenic Copper (?) Antimony Stibarsen

PbS Bi As Cu (Sb»Bi) SbAs

Sulfides and sulfosalts Galena Sphalerite Hawleyite Pyrrhotite Pyrite Marcasite Arsenopyrite Molybdenite Cosalite Gladiate Pekoite Gustavite Tetrahedrite Freibergite Bournonite Dyscrasite Pyrargyrite Miargyrite

PbS (Zn,Cd)S (Cd,Zn)S Fe12xS FeS2 FeS2 FeAsS MoS2 PbBiS2 CuPbBi5S9 CuPbBi11S16 Pb3Ag3Bi11S24 (Cu,Fe,Ag)12Sb3S13 (Ag,Cu,Fe)12Sb3S13 PbCuSbS3 Ag3Sb Ag3SbS3 AgSbS2

Oxides Cassiterite Rutile Ferrotapiolite Ferrocolumbite Manganocolumbite Manganotantalite Wodginite Ferrowodginite Titanowodginite Ferrotitanowodginite Lithiowodginite Simpsonite Stibiotantalite(?) Microlite Uranmicrolite Cesstibtantite Rankamaite – Sosedkoite Ilmenite Uraninite Manganite

SnO2 (Ti,Fe,Ta,Nb)O2 FeTa2O6 (Fe . Mn)(Nb . Ta)2O6 (Mn . Fe)(Nb . Ta)2O6 (Mn . Fe)(Ta . Nb)2O6 Mn(Sn . Ta,Ti,Fe)(Ta . Nb)2O8 (Fe . Mn)(Sn . Ta,Ti,Fe)(Ta . Nb)2 (Mn . Fe)(Ti . Sn,Ta,Fe)(Ta . Nb)2O8 (Fe . Mn)(Ti . Sn,Ta,Fe)(Ta . Nb)2O8 LiTaTa2O8 Al4Ta3O13(OH) SbTaO4 (Na,Ca)2Ta2O6(O,OH,F) (Na,Ca,U)2Ta2O6(O,OH,F) (Sb,Na)2Ta2(O,OH)6(OH,Cs) (Na,K)32xAl(Ta,Nb)10(O7OH)30 (Fe,Mn)TiO3 UO2 MnO(OH) Phosphates

Fluorapatite Carbonate Hydroxyapatite

(Ca,Mn)5(PO4)3(F) Ca3(PO4,CO3)5(OH)

Carbonates Calcite Rhodochrosite3 Dolomite Zabuyelite

CaCO3 MnCO CaMg(CO3)2 Li2CO3 Borates

Diomignite

Li2B4O7 Silicates

Quartz Albite Microcline Sanidine (adularia) Rb-feldspar Biotite Muscovite Lithian muscovite Lepidolite Illite Montmorillonite Cookeite Eucryptite Spodumene Petalite Foitite Schorl

SiO2 Na(AlSi3O8) K(AlSi3O8) K(AlSi3O8) (Rb,K)(AlSi3O8) K(Mg,Fe)3(AlSi3O10)(OH)2 KAl2(AlSi3O10)(OH)2 K(Al,Li)2(Al,Si)4O10(OH,F)2 (K,Rb)(Li,Al)2(Al,Si)4O10(OH,F)2 (K,H2O)Al2(AlSi3O10)(OH, H2O)2 (Na,Ca)(Mg,Al)2(Si4O10)(OH)2·n H2O LiAl4(AlSi3O10)(OH)8 LiAl(SiO4) LiAl(Si2O6) Li(AlSi4O10) Fe2þ 2 AlAl6(Si6O18)(BO3)5(OH)4 NaFe2þ 3 Al6(Si4O18)(BO3)5(OH)4 (continues)

Part 1 Lithium

Native elements

Table 1.23 (continued) Sulfides Cubanite Chalcopyrite Stannite Kesterite Cerny´ite

CuFe2S3 CuFeS2 Cu2FeSnS4 Cu2ZnSnS4 Cu2CdSnS4 Halides

Fluorite

CaF2

Phosphates Lithiophosphate Lithiophilite Amblygonite Montebrasite Tancoite Whitlockite Fairfieldite Crandallite Overite Dorfmanite Switzerite

Li3PO4 Li(Mn . Fe)PO4 LiAlPO4(F,OH) LiAlPO4(OH,F) LiNa2HAl(PO4)2(OH) Ca3(PO4)2 Ca2(Mn,Fe)(PO4)2·2H2 CaAl3H(PO4)2(OH)6 Ca3Al3(PO4)5(OH)6·15H2O Na2HPO4·2H2O (Mn,Fe)3(PO4)2·7H2O Sulfates

Barite

Silicates Elbaite Rossmanite Feruvite Dravite Beryl Topaz Pollucite Cesian analcime Holmquistite Zircon Thorite Coffinite (?) Garnet (?)

NaLi1.5Al1.5Al6(Si6O10)(BO3)5(OH)4 LiAl2Al6(Si6O10)(BO3)5(OH)4 CaFe2þ 3 Al5Mg(Si6O18)(BO3)5(OH)4 NaMg3Al6(Si6O10)(BO3)5(OH)4 Be3Al2(Si6O10) Al2SiO4(F,OH) (Cs, Na)(AlSi2O6).nH2O (Na,Cs)(AlSi2O6).nH2O Li2Mg3Al2(Si6O22)(OH)2 (Zr,Hf)(SiO4) ThSiO4 USiO4·(OH)4 (Mn,Fe)3Al2Si3O12

BaSO4

Courtesy of the International Mineralogical Association.

Geology

77

78

Table 1.24 Typical Analyses of Lithium Minerals from the TANCO Mine, wt.% (All from the Upper Intermediate Zone, Except as Noted) (Cerny et al., 1998)

a

SQI

Petaliteb

Amblygonitec

Montebrasited

Lithiophilitee

Eucryptitef

Lepidoliteg

Lithian muscoviteh

7.87 63.45 27.40 0.114 0.038 0.16 0.012 0.053 — 0.002 0.001 — 0.02 0.41 —

4.48 77.08 16.57 0.253 0.048 0.09 0.015 0.13 — 0.012 0.008 — 0.05 0.37 —

4.55 77.83 16.58 0.054 0.050 0.008 0.028 0.007 — — — — — 0.41 —

9.90 — 32.86 0.047 0.004 0.132 0.002 — — — — 6.30 49.26 3.27 —

9.52 — 35.10 0.139 0.005 0.085 0.003 — — — — 1.40 49.11 5.32 —

9.13 — — 0.05 0.01 1.00 0.51 8.83i 34.39 — — — 44.95 — 0.07

11.03 45.18 43.79 — — — — — — — — — — — —

4.70 47.85 26.02 0.16 8.52 — — 0.11 0.55 4.29 0.93 4.54 — 1.76 —

4.57 43.89 36.01 0.23 9.85 — — 0.40 0.15 1.92 0.25 0.38 — — —

Log-shaped white crystals. SQI is a spodumene, quartz and feldspar mixture resulting from the decomposition of petalite and feldspar; white crystals. Grey or white crystals. c White crystals. d Lower Intermediate Zone; secondary; brownish crystals. e Grey at the wall contact; pale brown or orange-pink below the wall. f Pale pink crystals. g Lepidolite zone; plus 0.04% TiO2. h Lepidolite zone; plus 0.06% TiO2. i Analyzed as FeO, not Fe2O3. j The OH is included in the H2O analysis. b

Part 1 Lithium

Li2O SiO2 Al2O3 Na2O K2O CaO MgO Fe2O3 MnO Rb2O Cs2O F P2O5 H2O j Insol.

Spodumenea

Geology

79

Figure 1.35 Map of the Kings Mountain, North Carolina pegmatite area (Kesler, 1976).

amblygonite as microscopic crystals. The deposit’s average composition was about 15– 20% spodumene as fine-grained crystals, 41% feldspar (27% albite and 14% microcline), 32% quartz, 6% muscovite (as small flakes) and 1% rare minerals such as beryl (finely disseminated crystals; often , 0.5% of the pegmatite), cassiterite, columbite, lithiophilite and tantalite. There was also minor amounts of apatite, pyrite, rhodochrosite and sphalerite. The spodumene crystals reached a maximum length of 0.91 m (3 ft), but averaged less than 0.15 m (6 in.; Table 1.25). The spodumene was fairly evenly distributed in the ore, making it comparatively easy to mine and process (Kesler, 1976, 1960). The Foote Mineral deposit is located about 1.6 km SW of the town of Kings Mountain and had been mined periodically since the early 1900s, but only on a large scale since Foote purchased the deposit and completed exploratory drilling in 1956. Their mining was from a cluster of eight nearly vertical pegmatite dikes that could be recovered from a single open pit. More extensive drilling later outlined ore reserves in the northern part of this area of 29 million mt with an average grade of 0.7% Li. Inferred reserves in the remainder of the area might be about 14 million mt (Kunasz, 1994). The deposit is comparatively young for a lithium pegmatite, with its age initially estimated to be about 260– 375 million years

80 Part 1 Lithium

Table 1.25 Typical Grain Size of the Minerals in the North Carolina Lithium Pegmatite (Spanjers, 1990) Rock type

Spodumene, % (grain size)

Beryl, % (grain size)

Mediumgrained pegmatite

20 (0.2–25 cm)

0–2 (0.05–2 cm)

Coarsegrained pegmatite Aplite (spod.-rich)

30 (1–40 cm)

Aplite (spod.-poor)

Microcline, % (grain size)

Albite, % (grain size)

Muscovite, % (grain size)

Quartz, % (grain size)

15 (0.2–25 cm)

35 (0.01 –1.0 cm)

5 –10 (0.01 –1.0 cm)

25 (0.01–2 cm)

0

40 (1–50 cm)

5 (0.01 –1.0 cm)

5 (0.01 –1 cm)

20 (0.01–2 cm)

20–30 (0.01–0.5 cm)

0–2 (0.01–0.2 cm)

0–10 (0.02–0.2 cm)

20–50 (0.01 –0.2 cm)

5 –10 (0.01 –1 cm)

20– 40 (0.01–0.1 cm)

0–2 (0.01–0.5 cm)

0–2 (0.01–0.2 cm)

0–5 (0.02–0.2 cm)

50–80 (0.01 –0.2 cm)

5 –10 (0.01 –0.1 cm)

15– 35 (0.01–0.1 cm)

Comments Most prevalent rock type; dominates major ore bodies and in peripheral dikes Between aplite (at wall rock contacts) and medium-grained pegmatite One end member of layered aplite; occurs locally within medium-grained pegmatite Other end member of layered aplite; often occurs adjacent to wall rock contact

Geology

81

(Vine, 1980), and later to be middle or late Paleozoic, or only about 250 MYBP (Spanjers, 1990). North Carolina Tin-Spodumene Belt; FMC In 1969 Lithium Corporation of America (Lithco or LCA; later purchased by FMC) began to mine their Cherryville, Hallman-Beam spodumene pegmatite deposit in the Long Creek area of North Carolina (Fig. 1.36). They had begun operation in Minnesota processing ore from South Dakota, but moved to Bessemer City, North Carolina in 1954 to utilize the ore purchased from other local operators, and later Foote Minerals’ nearby Kings Mountain mine. The LCA deposit had proven and probable reserves of about 27.5 million mt of ore recoverable by open pit mining, containing an average of 0.70% Li (Kunasz, 1994). The ore on average contained 20% spodumene, 33% albite, 25% quartz, 15% potassium feldspar, 6% muscovite and 1% other minerals (apatite and beryl, with rare cassiterite, columbite and tantalite). It was rhythmically zoned with four components: (1) low-spodumene apalite, which usually occurred adjacent to the pegmatites’ wall rock; (2) coarse-grained high-spodumene ore next to the low-grade apalite; (3) medium grained high-spodumene rock, the dominant component of the ore and (4) high-spodumene apalite that occurred locally within the medium-grained spodumene (Table 1.25). The deposit contained about 80% of the large and mediumgrained spodumene, and the structure also had several large (30 –90 m thick) adjacent spodumene pegmatite dikes, with a swarm of small spodumene dikes in various orientations filling the joints of the central section. Numerous barren dikes and a few ore dikes from 0.2 –10 m thick occurred along the periphery of the deposit, and appeared to be of the same age. The spodumene and microcline crystals were usually perpendicular to the apalite borders in the coarse-grained ore, and more randomly orientated in the medium-grained ore.

Figure 1.36 Top view of the Lithium Corporation’s North Carolina Pegmatite (Spanjers, 1990).

82

Part 1 Lithium

It has been speculated that flowing magma that was being cooled and mixed by density and viscosity gradients was slowly being fractionated by solubility differences, and reached this area with a relatively low lithium content. Some of this magma was forced into existing fracture planes in the overlying brittle amphibolite and rapidly cooled, forming thin, fairly contorted low-lithium pegmatite dikes. The remaining magma, which by then had fractionated to a higher lithium content, was finally forced in a series of pulses with a little of the remaining apalite into the now weakened and fractured host rock. This resulted in much more extensive displacement, and made much larger dikes. There was a longer period between some of the pulses, allowing slower cooling and larger crystals in the larger dikes. The variation in pulses helped to form the various lithium zones, but in all cases the cooling of the pegmatites was comparatively rapid to prevent further fractionation and more complex zoning. The temperature and pressure of the pegmatites may have also been influenced by saturated, super critical water that possibly accompanied the molten magma (Spanjers, 1990). Smaller United States Deposits One of the first lithium pegmatites operated in the United States was from 1900 to 1927 at the Stewart mine in the Pala district, San Diego county, Southern California. This very small, but high-grade lepidolite deposit (Fig. 1.37) also contained some amblygonite and small amounts of the unusual kunzite variety of spodumene (Kesler, 1960). Slightly earlier (1898), and extending for much longer the Keystone and other districts of zoned pegmatites in the Black Hills of South Dakota became the country’s major lithium producers. The limited early US demand for lithium increased considerably during World War I when the use of lithium in alkaline storage batteries required larger supplies. The demand for lithium products greatly increased again during World War II for rescue devices (inflated by

Figure 1.37 Side view of the Stewart Mine’s lepidolite pegmatite, San Diego County, California (O’Neill et al., 1969).

Geology

83

hydrogen from lithium hydride), lithium hydroxide for the absorption of carbon dioxide in submarines and for multi-purpose greases. The Black Hills deposits filled most of these needs, with the largest operation being the Etta mine (after 1908 owned by the Maywood Chemical Co.). It supplied spodumene somewhat intermittently for almost 50 years, and there was additional production from many smaller operators. The Black Hills lithium district consists of a very large number of small pegmatites that intruded into Precambrian rocks in a 710 km2 area about 1715 million years ago. There are over 200 pegmatites in the main 16 km2 lithium area (Fig. 1.38), and over 1500 in an adjacent 34 km2 area with less lithium. However, of these pegmatites only 43 were known to contain from traces (R) to abundant ðA ¼. 15%Þ lithium minerals, with two containing abundant spodumene (33 Rare[R] –A), two abundant lepidolite (5 R – A) and five with fairly common (C) amblygonite (24 R – C). Many other pegmatite minerals such as beryl, cassiterite, columbite– tantalite, feldspar, mica, microlite, pollucite and quartz were also mined in the district, but not in the lithium mines. The lithium production to 1944 totaled: 7127 mt of lepidolite, 7678 mt of amblygonite, 22,619 mt of spodumene, and 100 mt of triphylite (LiFePO4). The spodumene in the Etta deposit had unusually large crystals, with some up to 14 m long, and the average length was about 3 m. It, like most of the Black Hills lithium pegmatites was distinctly zoned, with as many as 12 zones, but among these there were only three lithium mineral zones. Most of the lithium occurred in the intermediate or core zones, accompanied or surrounded by perthite, quartz and albite. The ore grade of spodumene varied widely from 0.5 – 3% Li, with the amblygonite containing the most, 1.18 – 2.15% Li (after the initial period it was usually mined as a by-product). In several mines lepidolite was the dominant mineral, with the Bob Ingersoll Mine being by far the largest producer. Their ore averaged 5 – 20% lepidolite, 40 – 50% quartz and 40 – 45% cleavlandite, but the center of the core contained massive, fairly pure lepidolite with high concentrations of the other alkali metals (Table 1.18; Tuzinski, 1983). Its crystals were usually small lilac or purple flakes, about 3 mm wide, and arranged in felted aggregates. Amblygonite was found in an intermediate zone, often in light grey spherical masses about 1.8 m in diameter, but one mass measured 9 £ 1.2 £ 100 m. For each of the minerals hand sorting had originally been employed to obtain a fairly high-grade product, but later flotation was used to concentrate each of the ores. Figure 1.39 shows the mineralization in one of the larger South Dakota lithium pegmatites (Page, 1953). Sirbescu and Nabelek (2001) studied quartz inclusions from the Tin Mountain, South Dakota lithium pegmatite, and found indications of both solid and aqueous (with CO2 and NaCl) phases. Based upon the average boiling point of the aqueous phase, they came to the conclusion that the deposit formed at 3408C and 2.7 kbar pressure. This estimated depth of . 11,000 m and that low temperature would

84

Part 1 Lithium

Figure 1.38 Outline of several of the South Dakota lithium pegmatites (Page, 1953).

appear to be unlikely, and implies that the inclusions formed from supercritical water and carbon dioxide. Other areas in the US with small lithium pegmatite deposits include the western part of Arizona from Phoenix to Lake Mead. In the White Picacho district distinctly zoned pegmatites range up to 610 m in length by 60 m in width, with zones

Geology

85

Figure 1.39 Geologic maps and sections, northwest part of the Beecher Lode pegmatite, Custer County, South Dakota (Page, 1953).

of 8 –23% spodumene, amblygonite or lepidolite. The Harding and Pidlite mines operated for a brief period in New Mexico, and there are several smaller lithium pegmatite areas in Colorado, Connecticut, Maine, Massachusetts, New Hampshire, Utah and Wyoming (Kunasz, 1994; Vine, 1980; O’Neill et al., 1969; Kesler, 1960). The Big Mack and Big Whopper Deposits, Ontario, Canada The Separation Lake metavolcanic belt of northwestern Ontario (Fig. 1.32) contains the Separation Rapids pegmatite group, with its large number of lithium and rare metal pegmatites. In this group there are more than 37 major rare metal pegmatites and many smaller ones, including two of the world’s very large lithium deposits, the Big Mack and Big Whopper. The Separation Rapids formation (Figs. 1.32 and 1.40) is to the east and adjacent to the Winnipeg River –Cat Lake formation in Manitoba that contains the Tanco deposit (60 km to the west of the Big Whopper), and has a very similar age (2646 compared to 2740 –2844 Ma), geological setting and mineralization. This makes it appear that they are closely related to each other, although the Separation Lake lithium pegmatites are primarily of the petalite-type while spodumene predominates in Manitoba. There is some evidence that the Separation Rapids pegmatite group originated from the Treelined Lake granitic complex in the adjacent English River Subprovince to the north. The Separation Lake area also contains pegmatites of the beryl-type, along with a few containing mostly lepidolite. Various tantalum and other rare metals are common in most of these pegmatites (Tindle and Breaks, 2000).

86 Part 1 Lithium Figure 1.40 Some of the lithium pegmatites in the Separation Lake area, Ontario, Canada (Tindle and Breaks, 2000; Zoning: H ¼ homogenous, O ¼ oscillatory, P ¼ patchy; courtesy of the Ontario Geological Survey).

Geology

87

The Big Whopper deposit consists of five pegmatite lenses that are 6 £ 56– 12 £ 122 m in size, and swarms of 1 – 10 m thick pegmatites adjacent to the lenses (Fig. 1.41). The area is 1350 m long and up to 160 m wide, with the largest lens, the Big Whopper, 350 m long and up to 60 m thick. The pegmatites are nearly vertical, with an unknown depth, and composed of two sections: (1) the majority of the deposit contains alternating zones of potassium feldspar with on average 37% petalite, and lenses of fine-grained muscovite aplite. (2) contains 2 –4 cm thick pink petalite layers in orange garnet –quartz –muscovite – albite aplite. The majority of the deposit appears to have been modified by ductile deformation and recrystallization, forming a polygonal net-like mosaic of medium to coarse-grained petalite crystals. They are a translucent light pink color, and very pure, containing about 4.74% Li2O with very little iron (Table 1.18; Tindle and Breaks, 2000). The 842 ha (2080 acre) Big Whopper deposit is owned by Avalon Ventures, and located 60 km north of Kenora on the English River. It has estimated reserves of 11.6 million mt of ore averaging 0.62% Li (1.33% Li2O, or 25 ^ 5% petalite), and one 5.6 million mt area contains 0.46 –0.66% Li ore (0.99 – 1.42% Li2O; Saller and O’Driscoll, 2000). Pearse and Taylor (2001) describe the deposit as being 1.5 km long and usually 10 –80 m wide, with 3.2 mt of ore in a 55 £ 400 m zone that is amenable to open pit mining. The rock in which the pegmatite intruded is dark green to black amphibolites, and the entire formation appears to have undergone north– south compression. This has produced foliation, folding and ductile shear zones in the deposit, and made it more homogenous. Some zoning remains, however, with usually a thin albitite or lepidolite – albite –petalite wall zone, and the later minerals also occur in separate thin dykes parallel to the main pegmatite. There is one Rb – Kfeldspar, low-petalite zone, and the other zones are less distinct, being based upon petalite’s color (white, blue, grey or pink), as well as variations in the content of Rb– K-feldspar, albite, mica and quartz. Pearse and Taylor (2001) characterized the deposit as containing about 25% petalite (LiAlSi4O10), 11% albite (NaAlSi3O8), 10– 15% Rb – K-feldspar ([K,Rb]AlSi3O8), up to 15% muscovitic mica, 3 –5% spodumene, and lesser amounts of lepidolite (with up to 4% Rb2O), tantalite minerals (predominately columbite –tantalite), cassiterite (0.04% Sn) and many other minerals. The Big Mack pegmatite system lies to the west of the Big Whopper, and is owned by the Emerald Fields Resource Corp. It consists of a swarm of beryl and petalite-type pegmatites within an area 3 km long and an average width of 200 m. The swarm contains the Big Mack pegmatite which is about 30 £ 200 m in size, nearly vertical and of an unknown depth. Its mineralogical features are similar to the Big Whopper, consisting primarily of K-feldspar and petalite, with minor amounts of quartz and muscovite. It contains a thin border zone of fine-grained garnet – quartz – albite aplite with some blue cordierite, and a small 2 –3 cm thick core zone of mafic metavolcanic rock. In addition to the petalite the deposit contains small amounts of spodumene, eucryptite, bikitaite and many other minerals (Tindle and Breaks, 2000). The initially estimated ore reserves were greater than 300,000 mt of

88

Part 1 Lithium

Figure 1.41 The Big Whopper lithium pegmatite (Pearse and Taylor, 2001; figure first published in the CIM Bulletin, Vol. 94, No. 1049. Reprinted with permission of the Canadian Institute of Mining, Metallurgy and Petroleum).

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petalite, and consideration was being given to the production of 15,000 mt/yr of this product (USGS, 2000). Other Canadian Lithium Deposits In Ontario the area from Beardmore to Lake Nipigon, including an area 140 km south of Georgia Lake also contains many zoned and unzoned pegmatites, with at least 6 million mt of predominantly spodumene ore. Other areas in Ontario with lithium pegmatites are Dryden, Falcon Lake, Fort Hope, Gull Lake, Lac La Croix, O’Sullivan, Pakeagama Lake, Quetico Park, Root Lake, Superb Lake, Wekusko Lake and Tashota Lake (Kunasz, 1994; Kesler, 1960). The mineral and chemical analysis of an outcrop sample of the unzoned, very small-crystal Wekusko Lake spodumene deposit, averaging 0.79% Li (1.70% Li2O) is shown in Table 1.18. A similar outcrop sample from the Pakeagama Lake pegmatite analyzed 0.62% Li (1.34% Li2O; in its spodumene), 0.90% Rb2O, 0.20% Cs2O, 0.038% Ta2O3, 0.014% Nb2O3, 0.0076% Sn, 0.0047% Ga and 0.0029% Tl (www.hustonlakemining.com 1999). In 1960 there were considered to be nine lithium pegmatite areas in Ontario, extending from Georgia Lake (137 km NE of Port Arthur) to SE of Lake Nipigon. Both zoned and unzoned spodumene deposits had been located (Kesler, 1960). Quebec; Quebec Lithium Corp In the Preissac-Lacorne district of western Quebec, half way between the towns of Val d’Or and Amos there is a lithium-rich pegmatite district with dikes containing primarily spodumene and small amounts of lepidolite and other lithium micas. There are also other non-lithium ores such as beryl, pollucite, molybdenite and minor amounts of columbite and tantalite. The district is in a “T” shape extending about 33 km east –west and 20 km north –south. The lithium pegmatites appear to have intruded into this Precambrian Shield area perpendicular to a greenstone – granodiorite contact at shallow depth, and are richest near the contact zone (Fig. 1.42). The largest deposit was owned by Quebec Lithium Corp. with reserves of 15 – 20 million mt of ore averaging 0.56% Li (0.51 – 0.93%; 1.21% Li2O), and they produced spodumene from the deposit from 1955 – 1965. Many smaller deposits have also been found in this area, including in Montainier and Delbreuil Townships, and within the Abitibi greenstone belt in Steele and Lowther Townships (Flanagan, 1978; Kunasz, 1976; Kesler, 1960). In the Abitibi region Raymor Resources was considering the development of a spodumene operation on their La Motte property in 1997. The deposit was stated to have reserves of 4.55 million mt of ore containing 0.50% Li (1.07% Li2O) to a 100 m depth, and ore with . 1.16% Li (. 2.5% Li2O) below that. The pegmatite was amenable to surface mining (USGS, 2000, 1997). In Manitoba the East Braintree-West Hawk, Gods Lake, Herb Lake and Cat Lake – Winnipeg River (see Tanco) areas all contain lithium pegmatites. Spodumene is the principal lithium mineral, with some petalite, lepidolite and amblygonite. The small Buck and Pegli lithium deposits were also found in the latter area (Cerny and Lenton, 1995). In 1958 the lithium ore reserves were about 8 million tons in the

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Figure 1.42 Side view of the Quebec Lithium Corporation’s spodumene pegmatites (Flanagan, 1978; reprinted from Energy, Vol. 3, No. 3, q1978 with permission of Elsevier).

Winnipeg River area, and 2 million tons in the Herb Lake area, 137 km east of the town of Flin Flon. In the Northwest Territories there are also many pegmatites in this Precambrian Shield area, such as an extensive area (13,000 km2) north of the Great Slave Lake and the Yellowknife-Beaulieu area. However, of the 500 pegmatites examined by 1960, only 30 contained lithium. Spodumene was the principal lithium mineral, although amblygonite was a prominent constituent of some ore zones. The spodumene was often high grade with a large crystal size, and accompanied by tantalite –columbite, beryl and cassiterite. Both the Moose No. 2 and Best Bet openpit mines have produced lithium ore in this area (Kesler, 1960). The age of these deposits is about 2.2 billion years (Lasmanis, 1978). Other Countries In Brazil commercial lithium-bearing pegmatites have been operated for many years in the Jequitinhonha River basin of the Aracuai –Itinga area in northern Minas Gerais province. As an example of this, in 1978 1600 mt of petalite, 1200 mt of lepidolite, 1000 mt of amblygonite (for conversion to lithium chemicals) and 800 mt of spodumene were mined for domestic consumption, and 2500 mt of petalite were exported. Some cassiterite, tantalite and beryl that accompanied the lithium in the pegmatites was also recovered, with the reserves in the Precambrian Brazilian Shield region of that state at that time being about 25,000 mt of Li (Afgouni and Silva Sa, 1978). In 2001 Brazil was the sixth largest lithium producing country.

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Small amounts of spodumene and amblygonite have also been mined in the states of Paraiba, Rio Grande do Norte and Ceara (Solonopole and Quixeramobim). The pegmatites dip steeply, are up to 300 m in length and 30 m wide, and distinctly zoned. At least eight contained appreciable spodumene, and 20 some amblygonite (Kesler, 1960). In the province of Goias a number of lepidolite pegmatites have been prospected, with one area about 1 km long and 100 – 150 m wide. The ore varied from 3 – 6% Li2O, and was colorless to several shades of pink (Petruk and Sikka, 1987). As the demand for lithium minerals increased (particularly petalite) further exploration has resulted in the discovery of other important petalite pegmatites, such as one that contained 100,000 mt of petalite with a 2% Li grade. Spodumene reserves in 1990 were estimated at 300,000 mt, but the richer lepidolite deposits in the country appeared to be nearly exhausted. The most important mining company was Arquena de Minerios e Metais Ltd., mining all of the above-listed minerals. They also supplied spodumene to Companhia Brasileria do Litio, who in 1991 produced about 1500 mt/yr of lithium carbonate with some government assistance in a nearby town. In Argentina there are lithium-bearing pegmatites in the western part of Sierras Pampaneas province in the San Luis, Cordoba and Catamarca areas. The pegmatites are zoned and contain about 18,000 mt of spodumene reserves (Kunasz, 1994). In Bolivia there are also lithium pegmatites in the Bolivian tin belt. Zaire One of the largest lithium pegmatite deposits in the world occurs in the Manono and adjacent (2.4 km away) Kitotolo deposits in Zaire. Each deposit is 5 km long, 120– 425 m wide, proven to a 125 m depth (up to 50 m from the surface is kaolinized), and appear to be zoned. From 1929 to at least 1991 only columbite and cassiterite were being mined, and that from the weathered surface rock of the Manono pegmatite. The near-surface ore of the zoned pegmatites contain 10 –25% spodumene in microcline, albite and mica. The deposits face the almost insurmountable problem that the nearest shipping port is at Lobito, Angola, about 2200 km away. The construction of a 10 million lb/yr lithium carbonate plant had been planned for 1980, but was not built (USGS, 1991; Kunasz, 1976; Kesler, 1960). Namibia, Other African Deposits Lithium minerals have been periodically produced by a number of different companies from the Karibib-Omaruru district, Namibia, approximately 190 km from Walvis Bay. The area has several strongly zoned pegmatites containing lepidolite, petalite and smaller amounts of amblygonite. A flotation plant had been installed to process the 20– 50% lepidolite ore, based upon reserves of about 1 million mt of ore. There are also an estimated 200,000 mt of petalite ore in the deposit (Kunasz, 1994; Kesler, 1960). In the Alto Ligonha area of Mozambique lepidolite, amblygonite and spodumene occur in small zoned pegmatite bodies. The lepidolite has been commercially mined, but the amblygonite is rare, and the spodumene has been altered to kaolinite. In Rwanda large amblygonite masses have been found in the pegmatite district west of the capital city Kial. This district is 5.6 km long by 46 –760 m wide, and though mined primarily for tin, it is said to

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contain 10 –25% spodumene and other lithium minerals. The Bougouni area of Mali had limited production of spodumene and amblygonite concentrates from 1956 – 1970, with reserves of about 266,000 mt of 3.02% Li2O ore. In the Noumas and Norrabees areas of the Cape Province in South Africa spodumene pegmatites occur that contain about 30,000 mt of lithium. Other reports of lithium minerals have been noted in the Ivory Coast, Malagasy Republic, Sudan and north of Kampala, Uganda (Kunasz, 1994; Anstett et al., 1990; Kesler, 1960). In China lithium minerals (such as petalite) and lithium chemicals are produced from the Altai Mountains in Mongolia, 600 km north of Urmchi, and several other pegmatite deposits in the Hunan, Sichuan and Xinjiang-Uygur provinces. In the latter province the large zoned Koktokay pegmatite has produced spodumene and lepidolite (as well as tantalum, niobium and cesium) since 1946. Its reserves are stated to be over 5 million mt of 0.7% Li ore, and from 2110– 7332 mt of Li2CO3 were produced from the deposit each year from 1984– 1988. A major low-grade deposit is also operated in southwestern Jiangxi Provence, South China, where the Yichun open pit mine produces lithium, tantalum and niobium. The ore is characterized as a lepidolite – topaz granite, consisting of lepidolite (Table 1.16), albite and quartz, with lesser amounts of topaz, potassium feldspar and amblygonite. The accessory minerals are zircon, monazite, pollucite, columbite – tantalite, microlite and tantalum-rich cassiterite (Fig. 1.43). The pegmatite has intruded through the Yashan granite batholith in the form of a small sheet-like body, which now outcrops and has been weathered to an unknown extent. The lepidolite –topaz granite contains from 15 – 20% lepidolite (0.74 – 1.63% Li2O) as up to 3 mm crystals in a 40– 60% albite lath structure, with Ta– Nb mineralization disseminated throughout the mass. The lepidolite contains relatively high amounts of fluorine, phosphorous and rubidium, and the deposit’s age is estimated as 178 MyrBP (Yin et al., 1995). There are many small lithium pegmatites in Europe, and in the early days of the industry several of them were commercially operated. Significant spodumene pegmatites have been found in southern Austria by Minerex, and exploratory underground mining and separation procedures have been conducted. The large Koralpe spodumene deposit is estimated to contain 10 million mt of 0.77% Li ore. A large lithium pegmatite has also recently been discovered in Finland, and in 2001 petalite production from it was being considered (USGS, 2002). Previously only the small Viitaniemi, Eraejaervi, central Finland lithium pegmatite containing mostly lithium phosphate minerals had been discovered (Volborth, 1954). Lithium pegmatites also occur in France, but none appear to be of sufficient size or grade to be commercial deposits. At Beauvoir in the Allier province of the French Massif Central a lithium phosphate pegmatite containing lepidolite and amblygonite, with tin and tantalum mineralization has been studied (Raimbault et al., 1995). One of the minor lithium pegmatites that has also been extensively studied is the Varutrask pegmatite in Sweden. It is in the form of a “C” with a shallow dip, and is zoned somewhat as the Tanco pegmatite in Canada. The Inner Intermediate Zone contains amblygonite, spodumene – quartz pseudomorphs after petalite, and petalite

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Figure 1.43 Simplified geologic map and cross section of the Yashan Batholith (Yin et al., 1995). Reprinted with permission from Economic Geology, Vol. 90:3, p. 578, Fig. 1 Yin, L., Pollard, P. J., Shoux, H. And Taylor, R. G. 1995.

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remnants. Some of this zone is replaced by lepidolite, and in the eastern limb of the “C” it also extends into the Outer Intermediate and Wall zones. The deposit was mined periodically to 1950, and very little of the reserves remain (Heinrich, 1976). In Portugal there are many lithium pegmatites, with the Barroso-Alvao area of northern Portugal being perhaps the most studied (Fig. 1.44). This area has at least 10 good-sized lithium pegmatites containing spodumene with some petalite, amblygonite and montebrasite. Most of the spodumene consists of laths up to 30 cm long, but there are zones where it is in the form of poeciltic aggregates. The major component of the deposits is aplite, and the spodumene is primary and comparatively uniform (similar to the North Carolina, USA deposit). When petalite is present it is usually in the form of a thin continuous coating on spodumene crystals, or as blades with quartz in pseudo-vugs. There is some alteration of the spodumene to albite (^ muscovite), and of the petalite to K-feldspar and/or eucryptite (but not spodumene). The bulk samples of the ore vary from 0.80 – 1.65% Li2O, and trace amounts to 0.81% Fe2O3. Crystals of spodumene contain about 0.23% Fe2O3 (Table 1.16; Lima et al., 1999). Charoy et al., (2001) have further studied three of the pegmatites in the eastern part of this district, the Covas de Barroso area, and found primary spodumene and small amounts of the secondary minerals noted above. However, they also noted that in the western area the deposits contained only primary spodumene. Their lithium pegmatites outcropped in a swarm of dykes over a 0.1– 1 km distance, and were 2 –10 m wide with a gentle to steep dip. Most of the spodumene occurred as long blades of euhedral to subhedral crystals, pearly in color, up to 15 cm long and averaging about 5 cm. It contained almost no inclusions, and usually petalite could only be detected by microscope in thin sections of the spodumene’s edges, except in one deposit where there were a few angular voids in the spodumene containing needles of petalite with quartz. Some of the petalite had been altered to brownish microcrystals of eucryptite. Aplite was the dominate matrix material, with some Narich plagioclase and quartz, and very little K-feldspar and muscovite. The Li2O content of the pegmatites averaged 1.55%, and the composition of three of the lithium minerals in the deposits is given in Table 1.18 (Charoy et al., 2001). In the Guarda area of Portugal a lepidolite and amblygonite deposit has been mined on a small scale for many years (Anstett et al., 1990). At Goncalo, Beira Alta province there is also a lepidolite deposit. India has a number of small lithium pegmatites (the lepidolite in one is listed in Table 1.18), including the lithium phosphate deposit in the Sewariya batholith of the South Delhi Fold. This leucogranite structure may have been intruded by a volatilerich magma at about 4848C and 500 – 700 bars pressure, based upon the phosphate minerals that are present (Pandit and Sharma, 1999). In the Govindpal area of Madhya Pradesh there is a zoned pegmatite containing lepidolite, zindwaldite, amblygonite, spodumene and tantalite of unknown size and grade. Sizable lithium deposits exist in Russia, with the Zavitaya, Chita Oblast spodumene district in the Ural Mountains having received the most attention, and in 1979 was Russia’s largest producer. In the

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Figure 1.44 Map of the Covas de Barroso lithium pegmatites (with names), and other pegmatites in Northern Portugal (Charoy et al., 2001, courtesy of the Canadian Mineralogist). (1) Syntectonic two-mica granites; (2) post-tectonic (post-D3) biotite granites. Insert: regional geology for location of the Re´gua-Verin fault system.

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Ukrainian Shield’s Archean Kirovograd Block there are also a number of lithium pegmatites, including fine-grained petalite in the Polohov deposit, and spodumene and petalite in the Lipnyazhka pegmatite field. The ores are accompanied by microcline, albite, apatite, muscovite, etc. (Eremenko et al., 1996). In the Kola Peninsula’s Baltic Shield an unusual lithium – cesium deposit occurs in an andalusite –biotite –quartz schist that is separated from, and may not be part of nearby pegmatites (Nagaytsev and Belyaev, 1995). A number of other Russian lithium pegmatites have also been mentioned, as well as high-lithium granitoids in the Altai (Fig. 1.45) and Transbaikal regions. These Li – F granites also contain appreciable amounts of various rare metals. In the western Kirovogard Block of the Ukrainian Shield there are lithium pegmatites high in petalite, with no minerals containing volatile components, little quartz, and highly disordered K-feldspar. Inclusion evidence indicates that the pegmatites formed at greater than 11008C and 87 MPa pressure, in the presence of CO2-rich fluids. Some petalite later recrystallized in the presence of CO2-rich fluids at 6808C and , 46 MPa (Voznyak et al., 2000).

Figure 1.45 Location of rare-metal Li- and F-rich granitoids in the Altai. (l) Li-F-type granitoids: (1) vicinity of Ust’-Kamenogorsk, (2) Kalguta pluton, (3) Dzhulalyu pluton, (4) Alakha pluton, (5) Belokurikha pluton, (6) Ust’-Tulatin pluton, (2) Late Paleozoic and Early Mesozoic Li-rich granitoids: (A) Kalba, (B) Kalguta, (C) Chindagatui, (D) Belokurikha, (E) Karakol, (F) Shchebetin, (G) Sinyushin and Kolyvan, (H) Tigirek; (3) granitoids of the Kalba and Monastyr complexes with unspecified Li and F contents; (4) other granitoids; (5) major faults (Dovgal et al., 1998).

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Other lithium production has been obtained from the Tro San Zhen district of the former Russian province of Kirgiztan, which is stated to contain significant spodumene and petalite reserves (USGS, 1979). In Uzbekistan the Chatkal-Kurama region of central Tien-Shan contains a cluster of pegmatites that includes the Shavaz lithium deposit, the Sargardon tungsten deposit and the Shabrez fluorite deposit (Akhundzhanov, 1997). In Afganistan’s Hindukush Mountain Range in the NilauKulam and Parun areas there is a potentially large spodumene deposit where much of the ore mined to 1990 was of gem quality (Anstett et al., 1990). Lithium pegmatite occurrences have also been noted in many other countries, including Bolivia, Chile, Czechoslovakia (lepidolite in the Rozna area in Moravia province), Ireland, Japan (Nagatare, Fukuoka Prefecture), Korea, Sicily (Guerenko and Schmincke, 2002) and Spain (Kunasz, 1994; Deberitz, 1993; Anstett et al., 1990). High-Lithium Clays (Hectorite); Other Rocks The range of the lithium content in igneous rocks is often about 6 –28 ppm Li, but it can vary widely from zero to much higher values, such as the exceptionally Li-rich igneous rock found in the McDermitt Caldera in Nevada and Oregon containing up to 0.35% Li. Sedimentary rocks often contain high lithium values, such as from 10– 53 ppm Li, but again the range is considerable, from zero to the medium-rich shales (often 20 –100 ppm) to the lithium clays (Table 1.26). The latter includes the highmagnesium-lithium end member of the smectite group hectorite [Na0.33(Mg,Li)3 Si4O10(F,OH)2] with a lithium concentration range of from 0.24 –0.53%. It is found in a large deposit at Hector, California, 120 km east of the large geothermal

Table 1.26 Physical and Chemical Properties of Some Lithium Clays (Vine, 1980) Li

Mineral name Hectoritea Li-Stevensiteb Li-smectitec Li-bearing highalumina clayd a

Al

Mg

F

(percent or relative amount) 0.5 0.4 0.7 0.5

Trace Low Medium High

Major Major Low None

5.0 5.0 Low None

Relative dispersion of clay in water

Relative Li solubility in water

Very high Variable Poor Very poor

Low High Low Very low

Type material from Hector clay pit, San Bernardino County, California. Representative material from Kirkland area, Yavapai County, Arizona. c Representative material from Montana Mountains, Humboldt County, Nevada. d Representative material from Missouri fire-clay district, Mint Hill area, Osage County, Missouri. Probably a mixture of kaolinite, cookeite, and possibly diaspore. b

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spring-formed borax deposit at Boron (that also contains high-lithium shales; 0.2 – 0.4% Li). At Hector it occurs as a massive deposit with some travertine-like calcite, and associated rocks of sandy to bentonitic clay and volcanic ash. It is partly covered by a younger basalt flow, leading to the suggestion that it may have been formed by a reaction with the existing volcanic ash and high-lithium thermal spring waters that accompanied the basalt. A large deposit of bentonitic clay 90 km further east containing 0.2% Li perhaps indicates some relationship between these three high-tomedium lithium content clay deposits (Kesler, 1960). Smaller amounts of hectorite have also been found in Clayton Valley, McDermitt, and the Lake Mead area, Nevada; near Wickenburg, Arizona; Socorro, New Mexico; Lincoln, Montana; and a number of other locations with lithium contents of up to 0.11 –0.19 (Anon., 1979). Hectorite is usually white, with swelling characteristics that at one time gave it some use in cosmetics as an adsorbent for facial oil, as a paint thickener and a beer clarifier. Extensive tests on the commercial processing on hectorite have indicated that the recovery of lithium is both expensive and difficult. The US hectorite reserves have ben estimated as 15.1 million tons of Li (Kunasz, 1994). In a limited number of locations several other types of high-lithium clays have also been found, such as the flint clays in Missouri, Kentucky and Pennsylvania with up to 0.5% Li (Vine et al., 1979; Anon., 1979). An estimate of the average abundance of lithium in some of the other comparatively high-lithium members of common rocks and waters has been made by Vine (1980) and White et al. (1976), as ppm: shales 60 – 66, pelagic clay 57, granite 30 – 40, basalt 10 –17, sandstone 15, carbonates 5 and geothermal water 1 –10. Seawater is 0.17, and Vine (1980) estimates the earth’s average crust at 20 ppm Li.

Lithium Isotopes Since the late 1990s there has been a growing literature on lithium isotopes, as indicated in Table 1.14. Both d6Li and d7Li have been reported, since a standard nomenclature has not yet been established. This is quite unusual in isotopic work, and makes the values of different investigators difficult to compare. Nevertheless, the studies have been quite useful in determining the origin of some brines, the amount of lithium adsorbed on clays or other rocks, and on various other physical or chemical changes. PROCESSING History of the Lithium Industry Lithium was discovered by the Swedish geologist, Arfvedson in 1817, who separated it from petalite found in Sweden’s Uto pegmatite, and named it after the Greek word lithos, meaning stone. It was first isolated as a metal in trace amounts by Sir Humphrey Davy and Brande in 1818, in larger quantities by Robert Bunsen and

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Mathiessen in 1854, and the metal was first prepared as a commercial product in 1925 by Metallgesellschaft in Langelsheim, Germany. They utilized zinndwaldite recovered from the Zinndwald, Germany base metal tailing dumps, and this company has remained in continuous production of lithium compounds from purchased ore concentrates or lithium carbonate since 1923. Lithium products were first commercially produced in the US by Maywood Chemical Co. in New Jersey in 1927, and then by the Foote Mineral Co. in the late 1930s, both using their own South Dakota or purchased ore concentrates. LCA later started operation near Minneapolis, Minnesota, and the US Government sponsored the Solvey Process Co. in the late 1940s to initiate mining at the large Kings Mountain, North Carolina spodumene deposit. Foote bought the mine in 1951, and in 1953 started lithium hydroxide production from this ore at Sunbright, Virginia. Also in 1953 the US Government let three contracts for the purchase of lithium hydroxide to Foote, LCA and American Potash, with the contracts expiring in 1960. The latter two companies built new plants in Bessemer City (1955 using their own North Carolina ore) and San Antonio (1956 using purchased Southern Rhodesia [Zimbabwe] lepidolite), respectively. In 1960 the industry’s over-capacity only allowed operation at about 20% of capacity, and Maywood ceased the production of lithium compounds. LCA closed their Minneapolis plant in 1959 (and canceled their long-term ore contract with Quebec Lithium), while American Potash & Chemical Co. closed their Texas plant in 1963. American Potash’s Searles Lake lithium operation had started in 1951 and closed in 1978. Quebec Lithium in turn started producing lithium chemicals, but closed their plant in 1965. The production of ore from South Dakota stopped in 1969, and sanctions against Southern Rhodesia (still one of the world’s major suppliers) curtailed their ore imports from 1965 –1980. Foote’s Clayton Valley brine operation commenced in 1966, and LCA started mining spodumene in North Carolina in 1968. Foote’s Salar de Atacama operation started in 1984, while SQM’s started at the Salar de Atacama, and FMC’s (formerly LCA) at the Salar de Hombre Muerto in 1997. Both of the North Carolina mines closed after their brine operations had been well established, and FMC essentially closed their Hombre Muerto plant in 1998 because of SQM’s greatly reduced lithium carbonate pricing. Some of the first large-scale uses of lithium were in lithium batteries, lithium hydroxide in lithium greases, the absorption of carbon dioxide in submarines in World War II, and the filling of balloons with hydrogen made from lithium hydride. During 1955 – 1960 lithium began to be used much more extensively in ceramics and glass, and 6Li was employed to produce tritium for hydrogen bombs. In the late 1950s lithium bromide began to be used in air conditioning on a large scale, and in 1961 n-butyl lithium began to be used as a catalyst for synthetic rubber. Lithium carbonate began to be employed in aluminum reduction cells in the second half of the 1960s, and from 1972 onwards many new uses for lithium were developed.

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Brine Processing; Solar Ponds Because of the very dilute concentration of lithium in even the most favorable brines, and the expense of directly recovering the lithium in a processing plant, solar evaporation of the brine to further concentrate it has been a necessary first step in all of the world’s current lithium brine operations. Intrinsically solar evaporation is a very simple and inexpensive operation, but it does have a number of demanding requirements, and can have many problems. Since it relies upon solar radiation for the energy to evaporate water from the brine, the pond areas must be very large and the land comparatively flat and inexpensive. Most of the highest-lithium brines are found in dry lakes (playas), so this requirement is easily met, but it would not be with most of the oilfield, geothermal, potash deposit end-liquor or similar brines. The pond area must also have good evaporating conditions, which again is usually the case with most dry lakes but not with most other brines. The rate of evaporation depends upon the amount of solar radiation (sunlight), the humidity, wind and temperature, and these conditions vary widely. This effects the pond size, the final brine concentration, the cost of the ponds and their operation, and the final brine treatment methods and cost. The ponds require careful construction, operation and control, with adequate provisions for product brine storage during the winter and periods of unusual weather. To be most cost-effective the ponds must be divided into many segments to maximize the overall evaporation rate (the rate decreases with the brine concentration), so that each of the salts in the brine may crystallize in separate ponds, and that ponds may be taken out of service to periodically remove (harvest) these salts (after the entrained brine has been thoroughly drained) without undue disruption to the entire system. There must be as much gravity flow between the ponds as possible, and the banks must be protected from wave erosion. Then, the most important of the pond design factors is for them to be reasonably free from leakage. If the ponds are constructed of soil a careful soil survey must first be made of the entire pond area to be sure that there is a continuous layer of adequately impermeable clay under the pond, at least at a reasonable depth. If there are occasional zones of permeable soil it must be removed and back-filled with clay. If the clay layer occurs at depth, then the outer pond walls must be cut with a trench and back-filled with clay to prevent the upper porous zones from leaking laterally. Meandering former sandy stream beds must also be sealed with these clay “cut-off walls.” In cases where the soil permeability is of border-line value canals may be built adjacent to the outer walls and filled with a more dilute brine at a higher level than in the ponds to form a “hydraulic seal.” In this case the ponds are usually operated with the feed brine at the outer edges of the pond system, and the product brine in the center. With most of the world’s large playas sufficient areas of impermeable soil can be found to form a solar pond system with a reasonably low amount of leakage. However, without the soil testing and the cut-off walls the ponds may leak excessively, as did at least the initial ponds at Clayton Valley.

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As an alternative to clay-sealed ponds they can be lined with an impermeable plastic membrane. For the most demanding ponds (or pond areas) one of the most popular membranes is 25 mil reenforced Hypalon. It is highly puncture-resistant, has a very long service life (.25 years) with low-to-moderate concentrations of magnesium chloride in the brine, and it is quite resistant to ultra-violet oxidation. For the bulk of the ponds a much less expensive, but less rugged membrane might be used such as 20–40 mil PVC (polyvinyl chloride) or polypropylene. Although theoretically the membranes can make a leak-free pond, in practice this is often far from the case. The problem is that the plastic sheets can only be produced with a limited width in the factories, and they may have some-to-many small holes in them as they are formed. Also, the membranes are very heavy, further limiting the size of each piece. This means that strips of membrane must be joined together (sealed) to from the desired pond size when actually laid out on the floor of the pond. The Hypalon sheets are joined with a glue, while the polypropylene, PVC and most other membranes can be heat-sealed. Both jobs are difficult, and the chance for imperfect seals is always present, as well the possibility of punctures from underlying rocks, or tears as the membrane is stressed. Careful visual inspection, as well as some electric testing can find some leaks at this time, and after the membrane has been formed the ponds may initially be filled with water, with moisture or conductivity sensors in the soil, and other leaks found and repaired before the ponds are placed in service. There are a number of articles on solar pond design and operation, such as by Garrett (1966). Because of the low concentration of the lithium in the original brine, leakage prevention, just as the recovery of as much as possible of the entrained brine from harvested salts, is very important. Many methods have been suggested to determine leakage in both clay and membrane-lined ponds, but they have had very limited success. The simplest is that shown in Fig. 1.46 utilizing the difference in pressure between the brine in the pond and the moist soil underneath (Lee and Cherry, 1978), while various electrical measurements can also be used to detect moisture in the soil. Piezometers (small open-ended tubes to measure the hydrostatic pressure) are sometimes employed, but the most positive type of detector is the use of small porous tubes placed under the surface of the ponds, where actual samples of the leaked brine can be withdrawn by vacuum. In all cases, once a leak is detected it is still very difficult to find the exact hole, and then repair that area of the membrane. However, despite these problems, if very carefully constructed and maintained, membrane linings can provide a most satisfactory brine retention barrier, and even though relatively expensive, in general be superior in performance to clay-lined solar ponds. The various commercial operations, and other suggested processing methods for lithium recovery will be reviewed in the following sections. Clayton Valley (Silver Peak), Nevada; Chemetall Production of lithium from this deposit was initiated in 1966 by the Foote Mineral Co. in a $2 million plant with a capacity of 14 million lb/yr of lithium

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Figure 1.46 An example of a solar pond leakage detection meter (Low et al., 2000; reprinted from the Eighth Symposium on Salt (ISBN 0444500650), Vol. 1, p. 524, Fig. 1.4, q2000, with permission from Elsevier).

carbonate (that could be extended to 18 million lb/yr for a 2-year period). Foote was acquired by Cyprus Minerals Co. in 1988 (and then called Cyprus Foote Mineral), and it in turn by Chemetall in 1998. The Clayton Valley operation, as are all of the lithium brine deposits, is dependent upon solar evaporation to concentrate the brine from the playa (as discussed above) to a value where lithium can be precipitated with sodium carbonate. In the ponds (Figs. 1.47 and 1.48) as the brine (Tables 1.3 and 1.27) evaporates it first crystallizes small amounts of calcite and gypsum (Garrett, 1960), and then salt at a rate to deposit a layer about 0.3 m thick per year. The evaporation rate of water in the area usually varies from 760 –1200 mm (30 – 40 in.)/yr, and the rainfall is often less than 130 mm (5 in.)/yr. In 1969 the operation pumped 100– 300 gpm of brine from each of 30 0.3 m (12 in.) diameter gravel-packed wells, perforated for their entire depth, and surrounded by 15 cm (6 in.) of gravel. They were 90– 240 m (300 –800 ft) deep with multi-stage centrifugal pumps at their base, powered by 50 HP engines on the surface. The brine level was sometimes as low as 15 –76 m (50 –250 ft) beneath the playa floor, and it contained an average of 400 ppm Li. The brine was sent through transite pipes to the initial 308 ha (760 acre) pond. The total area of the nine solar evaporation ponds in use at that time was 642 ha (1587 acre), and brine was advanced from pond to pond as it concentrated (Fig. 1.49) to minimize the entrained brine lost in the crystallized salts. Salt did not crystallize in the first pond, but it did in ponds 2, 3 and 4 (520, 90 and 86 acre, respectively). Slaked lime was added to the brine leaving the fourth pond after it had been evaporated for about 10 months, and it

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Figure 1.47 Map of Clayton Valley and its early solar ponds and wells (Davis and Vine, 1979; reprinted courtesy of the Rocky Mountain Association of Geologists).

Figure 1.48 Photograph of the more recent Clayton Valley solar ponds in 1991 (courtesy of Rocky Mountain PAY DIRT).

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Part 1 Lithium Table 1.27 USGS Logs for the Clayton Valley Drill Holes Shown in Fig. 1.47 (Vine, 1980) Maximum lithium content (ppm)

Well number

Elevation at surface (m)

Total depth penetrated (m)

Brine

Sediment

CV-1 CV-2 CV-2A

1302.4 1303.0 1304.5

120.4 120.4 83.8

60 55 100

310 930 390

CV-3 CV-4 CV-5 CV-5A

1304.2 1301.5 1301.5 1301.5

187.5 242.3 146.3 224.0

160 190 110 640

640 1700 770 960

Comments Gravel below 300 ft (91 m) Gravel below 235 ft (72 m); bottom hole temperature 448C Maximum 338C at 415 ft (126 m) Mostly in gravels Penetrated thick sequence of salt beds

precipitated gypsum and hydrated magnesia (reducing the Mg content of the brine to 2 –3 ppm) in the fifth 19 ha (46 acre) pond (also listed as two 12 ha [30 acre] ponds). The resultant magnesium hydroxide and calcium sulfate were periodically dredged (Fig. 1.50) from the ponds and sent to a sludge-containment reservoir. Since the brine was then basic, much of the remaining calcium precipitated with absorbed carbon dioxide in pond 6 (17 ha; 41 acre) with a mixture of sodium and potassium chloride (sylvinite). The sylvinite was harvested and stockpiled separately for possible future potash recovery. Salt and glaserite (K3Na[SO4]2) crystallized in pond 7 (17 acre), while in ponds 8 (5 ha; 13 acre) and 9 (6 ha; 14 acre) salt, potassium chloride and glaserite all precipitated. The final brine to be sent to the plant contained 5000 ppm Li, and it was stated that if the final concentration were over 6500 ppm Li, lithium-potassium sulfate would crystallize and be lost (Gadsby, 1967). Because of lithium carbonate’s appreciable residual solubility when precipitated in the plant, this rather low lithium concentration required that a high percentage of the end-liquor brine from the plant be recycled back to the ponds. From 75 –90% of the pond evaporation occurred during the months of April through October, so the final pond held enough brine to service the plant throughout the year and to help smooth-out the yearly weather fluctuations. Brine was initially pumped primarily from an unconsolidated volcanic ash aquifer, and later some was also pumped from porous halite (although it tended to dissolve, causing the upper sediments to cave-in) and sand or gravel beds (Gadsby, 1967). In 1970 it was noted that the average brine concentration had dropped to 300 ppm Li, the well field covered an area of 5.2 km2 (2 mi2), and that 10 of the 30 wells, and

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Figure 1.49 (a) Solar pond and pumping station at Clayton Valley (Deberitz, 1993, courtesy of Chemetall GmbH). (b) Pumping brine between ponds at Clayton Valley (Dillard and McClean, 1991, courtesy of Rocky Mountain PAY DIRT).

the No. 9 pond had been added during the first expansion in 1967. Twenty new wells were to be drilled from 1970– 1971 in a second expansion. The new wells were spaced on 610 m (2000 ft) centers, not perforated in their upper 12.2 m (40 ft) to minimize the entry of more dilute near-surface water, and the wells pumped at

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Figure 1.50 Dredge removing loose salts from the Clayton Valley solar ponds (Dillard and McClean, 1991, courtesy of Rocky Mountain PAY DIRT).

50 –500 gpm. The brine in the ponds varied in depth from 0.9 –1.2 m for the larger ponds to about 0.3 m for the smaller ponds. The 23 km (14 mi) of pond dikes were constructed of granular dolomitic limestone or a gravel – silt – clay mixture, and had an inner clay core to minimize leakage. There were also 27 km of access roads and 16 km of power lines. Pumping was required between ponds 1 and 2, 5 and 6 (Fig. 1.49), and from 9 to the plant, but all of the other interpond flow was by gravity, since the land sloped at about 0.19– 0.38 m/km(1 – 2 ft/mi). By 1970 pond 5 was full of gypsum and magnesium hydroxide, so it was abandoned and the lime slurry added to pond 6 (Anon., 1970; Barrett and O’Neill, 1970). In 1991 there were 40– 60 wells that were 150– 300 m deep (average 213 m), and they pumped several million gallons per day of brine from the playa. There were 22 ponds covering 1620 ha (4000 acre), subdivided into about 30 sections. The well construction was a variation of typical water wells, and both submersible and turbine pumps were used. The salt was only periodically removed (Fig. 1.51) from 61 ha (150 acre) of the smaller ponds as they became too full, and the dikes were raised on the larger ponds instead of removing the salt. The final ponds were plastic-lined (Fig. 1.48) to improve the pond efficiency, since the initial earthen ponds suffered considerable leakage. The entire evaporation process took 16 –24 months time, compared to 12 – 18 months in 1966– 1970 (Dillard and McClean, 1991). By 1993 the piping had been changed to PVC, and in 2001 the initial brine averaged 160 ppm Li, the final brine 6000 ppm Li, and there were 16 km2 of solar ponds (Kunasz, 1994).

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Figure 1.51 Hauling salt from Clayton Valley solar ponds with a scraper-carrier (Dillard and McClean, 1991, courtesy of Rocky Mountain PAY DIRT).

From the final pond the concentrated brine (Table 1.3) with a density of about 1.25 g/cc was pumped nearly 4.8 km (3 mi; 1.5 mi in 1967, Gadsby, 1967) to the processing plant in the town of Silver Peak. The plant had been converted from a silver ore cyanide-leach plant that had operated there from 1864 –1961. In the conversion all of the tanks and settlers were rubber lined to reduce iron contamination in the product, and considerable new equipment was added. The solar pond brine was first reacted with lime to remove most of the residual magnesium and some of the sulfate and borate ions, and then a small amount of soda ash was added to precipitate most of the calcium from the lime reactions. The slurry from these operations was settled and filtered, and the overflow solution sent to storage tanks. From there the brine was pumped through filter presses to be totally clarified, and then heated to 938C (2008F; lithium carbonate has an inverse solubility) and reacted with dry soda ash and hot wash and make-up waters to precipitate the lithium carbonate product. Extra water was added to prevent salt from crystallizing, since the pond brine was saturated with salt. The lithium carbonate slurry was thickened in a bank of cyclones, and the underflow fed to a vacuum belt filter where it was washed and dewatered. The cyclone overflow and filtrate were

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returned to the solar ponds, since they still contained at least half of the feed brine’s lithium. The filter cake was sent to a stainless steel rotary steam-tube drier, and the , 99% pure product was then air-conveyed to a storage bin. The product was considered a “commercial” grade since it still contained about 400 ppm B, and its major impurities were sulfate, sodium and potassium, with lesser amounts of calcium, magnesium and moisture. These impurities made it unsuitable for metal production or other demanding uses (Brown and Beckerman, 1990). From Silver Peak it was shipped by truck 89 km to the rail station at Mina, Nevada in either bulk or bags as a white, granular, free-flowing product. In 1970 plans were being considered to ship some of the product in a pellet form for the aluminum industry. Part of the lithium carbonate was also converted into lithium hydroxide at the plant (Anon., 1970; Gadsby, 1967). Foote hired contractors (Target Construction Co. in 1991) for the maintenance of the solar pond’s dikes and the 320 km (200 mi) road system. For this work Target used four 12 yard dump trucks, five 30 yard bottom dumps, a 6 yard loader, and employed nine people. In 1991 Foote employed 62 workers with a payroll of about $2 million and combined taxes of about $1 million (Kunasz, 1994; Dillard and McClean, 1991; O’Neill et al., 1969). In 1981 their capacity was 8000 mt/yr, and in 1997, 5700 mt of lithium carbonate were produced from the deposit. The output from the plant was shipped to Germany or their conversion plants in Pennsylvania, Tennessee and Virginia (USGS, 1997; Lloyd, 1981). Salar de Atacama, Chile Two of the world’s four commercial lithium brine recovery operations are located on the Salar de Atacama (Fig. 1.52). The mineral rights to the Salar are owned by the Chilean government, and in the late 1970s to mid-1980s their development agency, Corporacion de Fomento de la Produccion (CORFO) and their contractor, Saline Processors of the USA conducted exhaustive tests on the Salar to explore its mineral reserves and to develop economic methods of recovering the minerals. A brine sampling and drilling program initially established the halite’s porosity and permeability, and then the area, depth and composition of the Salar’s brine. This allowed isopach maps to be made of each of the important ions in the brine, as shown in Fig. 1.10 for lithium, potassium and sulfate. There is a considerable shifting in the ratios and concentrations of these ions, as well as boron and magnesium, within different areas of the Salar. Detailed meteorological data was gathered, including solar evaporation rates for different concentrations of brines to establish solar pond sizing. Typical evaporation rates for the brine as it concentrates are listed in Table 1.28. The average evaporation rate of water was about 3000 – 3300 mm/yr, the rainfall about 10– 25 mm/yr, the average relative humidity about 10%, and there was frequently a moderate wind. Temperatures ranged from a minimum of 2 208C in June to 9– 288C in January, but because of the low humidity and the wind there was excellent evaporation capability even in the

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Figure 1.52 Map of Northern Chile showing the location of the Salars de Atacama and Carmen (Harben and Edwards, 1997; this figure appeared in Industrial Minerals No. 353, February 1997, p. 29. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc 2003).

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Part 1 Lithium Table 1.28

Typical Brine Analysis in the Salar de Atacama Experimental Solar Ponds (g/liter) (Garrett, 1998)

Cl SO4 H3BO3 (B) Na Mg K Li H2O Density Evaporation Rate (mm/day) Salts cryst.

a

Original brinea

To sylvinite pond

To sulfate pond

192.0 23.3 4.4 (0.77) 93.2 12.3 22.0 1.96 873 1.227

205 45 9.2 (1.61) 72.0 23.7 46.8 3.66 856 1.258

195 88 18 (3.15) 40 46 37 7.07 860 1.284

Summer Winter 8.0 4.0 Halite

Summer Winter 7.3 3.7 Sylvinite

From carnallite pond 292 23 50 (8.74) 4.0 92 4.0 8.9 867 1.323 Summer Winter 5.7 3.0 Sulfates, Carnallite

Also containing, as wt.%: Ca 0.03, NO3 0.012, CO3 0.003, I trace; KM-20 brine.

winter with such normally hygroscopic salts as bischoffite. The periphery of the Salar was also examined for impervious clay that could be used to locate and construct inexpensive solar ponds. Only two areas were found that were adequate in size and impermeability, but the southwestern area had the unusual feature of frequent tunnels as might be formed by small rodents or moderately sized roots that had totally decayed. However, the mid-salar area appeared to have reasonably uniform and impermeable clay (Fig. 1.53). Finally, laboratory and solar pond studies were made to determine the phase chemistry of the different brine types upon being evaporated (Crozier, 1986; CORFO, 1985; Saline, 1985). These studies allowed CORFO to establish the most economical processes for the production of each of the potential products from the Salar’s brine: lithium carbonate, lithium chloride, lithium sulfate, potassium chloride, potassium sulfate, boric acid, magnesium chloride and magnesium sulfate. In most of the experimental studies brine was taken from the location Km-20 (Fig. 1.53), which contained somewhat of an average of the Salar’s brine composition. When solar evaporated in the summer a sequence of salts crystallized as the brine concentrated (Table 1.28), initially being salt (halite, NaCl); then halite and sylvite (KCl that forms a mixture with the NaCl called sylvinite); then halite, sylvite and potassium lithium sulfate (KLiSO4); then halite, kainite (KCl·MgSO4·2.75H2O) and lithium sulfate (Li2SO4·H2O); then halite, carnallite (KCl·MgCl2·6H2O) and lithium sulfate; then primarily bischoffite (MgCl2·6H2O), and finally primarily bischoffite and lithium carnallite (LiCl·MgCl2·7H2O; Vergara-Edwards and Parada-Frederick, 1983).

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Figure 1.53 Location of the Chemetall Salar de Atacama Lithium and Potash Leases and Buffer Zone (Coad, 1984). (This figure appeared in Industrial Minerals No. 205, October 1984, p. 28. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc 2003.)

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Because of the unusually low humidity and the prevailing winds at the Salar the brine can be evaporated to precipitate all of these compounds, in contrast to most places in the world where carnallite would be the final salt crystallized in the solar ponds. With the above-noted salts deposited in separate solar ponds the harvests could be halite, then sylvinite to produce potash (KCl, using a flotation separation from halite), then carnallite to produce coarse potash (Garrett, 1975), and then the potassium, lithium and sulfate salts could be harvested together. They could be converted to lithium and potassium schoenite (K2SO4·MgSO4·6H2O), these salts removed from the halite by flotation, and then the schoenites converted into potassium sulfate. The lithium salts would dissolve in the schoenite conversion liquor and be sent back to the ponds to join the end-liquor from the potash and lithium salts pond. This process is in commercial operation on the Great Salt Lake (but without the lithium salts; Garrett, 1970, 1967), and its equal viability with mixed potassium – lithium salts was demonstrated by CORFO. Later the same separation and conversion with salts quite similar to those formed at the Salar de Atacama has been conducted in China and at the Salar de Uyuni (Ramos and Kirigin, 2000). All of the brine’s lithium and boron could then be simply recovered from this end-liquor (Vergara-Edwards et al., 1985, 1983; Garrett, 1985; Pavlovic-Zuvic et al., 1983; Garrett and Laborde, 1983). Salar de Atacama, Chile; Chemetall The company Sociedad Chilena del Litio (SCL) was formed in 1982 as a 55% Foote Mineral (now Chemetall) joint venture with the Chilean government agency CORFO to produce lithium and potash from the Salar de Atacama. They received a 30 year concession from CORFO that would be renewable for 5 year periods until 200,000 mt of lithium equivalent had been produced, which might take 40 – 45 years. The concession covered 16,720 ha (41,315 acre), with a 6850 ha strip of land adjacent to the concession guaranteed not to be leased to any one else (Fig. 1.53). They would also have an exclusive right to recover lithium from the Salar for 10 years. SCL then constructed a solar pond system and lithium carbonate plant costing $56 million, and with 14 million lb/yr (6350 mt/yr) of lithium carbonate equivalent (LCE) capacity. The pond location at the Salar de Atacama was called Chepica del Salar, and the concentrated brine from the ponds was shipped to a lithium carbonate (Li2CO3) plant on the Salar de Carmen near Antofagasta and the small town of La Negra. Production started in 1984, and in 1986 Foote purchased CORFO’s share of the operation. The original plant capacity was soon raised to 24 million lb/yr of LCE, and in the early 1990s potash also began to be recovered as a by-product from the sylvinite harvested from their solar ponds. By 1991 they shipped 11,800 mt of lithium carbonate, and in 1995 they were exporting 28 million pounds of LCE, and selling their byproduct potash to SQM’s large Chilean potassium nitrate operation. The process that Sociedad Chilena de Litio Ltd. uses is presumably similar to that used by Foote at Clayton Valley, except that the brine initially had about

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a 1500 –1900 ppm Li content and a much higher magnesium to lithium ratio (Tables 1.5, 1.6 and 1.28). Foote located their brine wells in a position on the lake where the sulfate content was comparatively low (see Figs. 1.10 and 1.53), and yet there was still a high lithium concentration. Based upon their initial announcements they then planned to precipitate most of the sulfate and part of the borate from the brine leaving the halite ponds with calcium chloride (Anon., 1984a,b). The resulting low-sulfate brine would produce a larger crop of sylvinite, followed by halite– carnallite and then halite– bischoffite to remove most of the magnesium. As the evaporation proceeded to about 4% Li, there would finally be a period when bischoffite and lithium carnallite crystallized together. To obtain the highest yields this mixture could be leached to dissolve the lithium and leave most of the bischoffite, with the leach-brine recycled to the bischoffite ponds. The mixture could also be harvested and sold, or discarded as desired. In any case the final brine from the ponds would contain up to 6% Li and be a nearly saturated lithium chloride solution. It would have a comparatively low magnesium, sodium, potassium and sulfate content. As far as is known, the Salar de Atacama, a few other Andean salars and the Tibetan region of China are the only places in the world where the humidity is low enough to allow bischoffite and lithium salts to be crystallized in solar ponds on a commercial scale. At the Foote operation the initial solar pond system had an area of 89 ha (220 acre; 100 ha, Coad, 1984), but was soon expanded to 130 ha (320 acre). In 1993 there were 1.5 km2 (150 ha; 371 acre) of solar ponds (Fig. 1.54; Deberitz, 1993). Initially there were nine ponds varying in size from 2.2– 14 ha, with three of the ponds divided into two parts to make a total of 12 sections. The ponds were constructed on a flattened and smoothed area of the Salar’s salt surface, and lined with 0.5 mm (20 mil) PVC plastic sheet. The lining was made from 1.5 m wide strips that weighed about 1 t, and they were sealed together in the ponds. This required 61 km (38 mil) of carefully-made seams that then had to be both manually, and later when the ponds contained , 20 cm of brine, electrically examined for leaks. Initially the brine in the ponds was maintained at a 25– 40 cm depth, and the brine flowed by gravity or was pumped from pond to pond. After the final pond the brine was pumped to a 0.7 ha deep storage reservoir with a floating cover to prevent further evaporation (Anon., 1984a). Their area’s evaporation rate was 1270 – 1780 mm (50 – 70 in.)/yr, and the rainfall very little most years (10 – 30 mm), but on rare occasions there were heavy storms. The solar radiation in the area was 6.3 £ 106cal/m2/day, the relative humidity as low as 5%, and moderately intense winds arose in the afternoons. Brine was initially pumped at 1000 gpm from three wells that were 30 m (100 ft) deep to fill the ponds to an average 38 cm depth (Anon., 1984a). After the halite ponds the brine was mixed with calcium chloride and end-liquor from the processing plant to precipitate gypsum and some of the boron, with the precipitate being washed to recover some of the entrained liquor’s lithium content. The salt was harvested from the halite ponds once per year and placed in stockpiles, while

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Part 1 Lithium

Figure 1.54 Chemetall solar ponds at the Salar de Atacama (Deberitz, 1993, courtesy of Chemetall GmbH).

the salts from the sylvinite ponds were harvested more frequently and processed in a plant on the Salar to produce potassium chloride. The plant at Chepica (at the southeastern end of the Salar de Atacama) employed 51 people, 32 of whom came from the nearby village of Peine, where all of the employees lived. They worked on an 11-day schedule, with 4 days off, allowing them time to travel by road or light plane to the larger towns of Antofagasta or Calama (450 km away), if they desired. The plant maintained a modern quality control laboratory (Coad, 1984; Anon., 1984a). The pond’s final brine (Table 1.5) was removed from the holding pond at a concentration of 4.3 – 6% Li, trucked about 80 km (90 km, Coad, 1984) south to a new railroad station at Pan de Azucar (Fig. 1.6; initially at a rate of 100 m3/day), and then shipped by rail cars about 170 km further to La Negra (a small town south of Antofagasta on the Salar de Carmen and the Pan American highway) for final processing (Fig. 1.55; Kunasz, 1994; Crozier, 1986; Anon., 1984a). The process employed in the plant at La Negra has not been described, but is thought to follow

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Figure 1.55 Aerial view of the Chemetall Lithium Carbonate Plant at La Negra (near Antafagasta Chile, courtesy of Chemetall GmbH).

the steps of the Clayton Valley operation and one of their patents. The large amount of boron (, 8000 ppm) in the brine must be removed to avoid serious contamination of the product, and this can be easily done by solvent extraction. In an acidic brine all long-chain alcohols in an insoluble solvent can extract boron fairly selectively with a moderate extraction coefficient, and with all brines (including very basic ones) multi-carbon diols have a more selective and much higher extraction coefficient (Garrett, 1961, 1963). Foote’s boron-removal patent follows the patent of Folkestad et al. (1974) (removing boron from strong MgCl2 solutions) and suggests using a simple 7– 12 carbon alcohol such as iso-octanol or 2-ethyl hexanol in about a 20% mixture with kerosene. The brine is first brought to a pH of 2 with hydrochloric acid, and then contacted with a ratio of about four volumes of the solvent to one part of brine. Under these conditions the solvent has an extraction coefficient ranging from 6– 14, so when mixed with the brine and settled in four stages of counter current contact the residual brine should contain less than 5 ppm of boron (Table 1.29; but also suffer a 5 –10% lithium loss in the solvent). The loaded solvent can then be stripped of its boron and lithium content with water or dilute caustic in several other mixer – settler stages, and be ready for reuse (Brown and Beckerman, 1990). The amount of solvent loss was not mentioned in the patent, but with strong magnesium chloride solutions Folkestad et al., (1974) estimated that the stripped

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Part 1 Lithium Table 1.29

Brine and Product Analyses of a Boron Extraction Process with Lithium Chloride Crystallization (Brown and Beckerman, 1990)

Chemical species

Brine

Brine after B removal

Li Na K Mg Ca B SO4 Cl

6.30 0.077 0.019 1.29 0.053 0.73 0.016 5.86

6.03 0.073 0.018 1.29 0.051 0.0001 0.019 34.46

Chemical species LiCl (dry basis) Na K Mg Ca B SO4 Cl H2 O

With B removal evap. crystallized, 90% recoverya 99.2% 0.17 0.0015 0.075 0.004 ,0.0001 0.004 82.9 0.4

Boron extraction with four parts of 20% isooctanol in kerosene to one part of brine, and four mixer– settler stages. a 90% of the lithium chloride input is recovered.

brine would contain at least 200 ppm of solvent, and the boron eluting solution 2000 ppm of solvent. The purified brine might then be treated as at Clayton Valley, with lime to precipitate most of the magnesium and sulfate, followed by a reaction with a small amount of soda ash to precipitate the remaining magnesium and calcium. After the precipitates were removed by settling and filtration the brine could be heated and lithium carbonate formed by the reaction with soda ash. When thoroughly washed and dried this would form an excellent product for most uses. Alternately, however, to form lithium chloride directly, or for higher purity lithium carbonate, the brine following the lime, minor soda ash additions and filtration steps could be evaporated at about 1108C to crystallize lithium chloride. Some of the slurry would be continuously removed, thickened, filtered, the solids washed at 1308C, and the lithium chloride dried at 1708C. The high-lithium filtrate and wash water could be returned to the solar ponds. This should produce a 99.2% LiCl product (Table 1.29). To produce a 99.9% lithium chloride product the previous crystals could be dissolved and re-crystallized, or further washed with, or dissolved in isopropanol and then recrystallized. Also, the lithium chloride could be dissolved and reacted with soda ash to precipitate a high-purity lithium carbonate. The plant at La Negra initially employed 63 people, and Foote’s lithium carbonate was either sold directly, or some of it compacted into granules (Anon., 1984a). In 1998 lithium chloride production was also initiated using lithium carbonate as the raw material (Crozier, 1986; Coad, 1984). If it is assumed that sulfate is still precipitated from the brine, the general flowsheet shown in Fig. 1.56 should roughly illustrate the Chemetall process.

Processing

Figure 1.56 A general flowsheet for obtaining lithium carbonate from Salar de Atacama brine (Wilkomirsky, 1998).

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Salar de Atacama, Chile; SQM S.A (Formerly SQM Chemicals or Minsal) In 1986 CORFO formed a consortium called Minsal where they had a 25% ownership, Amax Exploration had a 63.75% participation, and Molymet 11.25%. The objective was to establish a multi-product operation on the Salar, and they were granted mineral rights on 1586 km2 of the Salar, with water rights to 240 liter/sec. The central 820 km2 of the concession to a depth of 40 m was estimated to contain 26 million mt of K, 1.7– 1.8 million mt of Li (or 20 billion pounds of LCE), 22 million mt of sulfate, and 0.7 million mt of boron. Amax initiated additional drilling and a new feasibility study, but in 1992 sold their interest in the venture to SQM, as did Molymet in 1993. SQM had been a purchaser of considerable potassium chloride to be converted into potassium nitrate at their large nitrate operations fairly close to the Salar (at Maria Elena, Coya Sur and Pedro de Valdivia; Fig. 1.6), making it the product of their greatest initial interest. They planned to add lithium, potassium sulfate and boric acid later in a sequence of expansions using the end-liquor from the potash ponds as the feed material for the lithium, and new brine for the other products (Table 1.30). In 1995 SQM purchased CORFO’s then 18% interest in the venture for $7 million, becoming the sole owner of the company, and their 300,000 mt/yr, $55 million potash plant started production. In December, 1996 lithium carbonate also began to be produced as a by-product in a 40 million lb/yr LCE, $51 million pond system and plant. SQM, as Foote, initially selected a brine extraction location for its well field where the brine had the maximum potassium and the least sulfate for potash and lithium production, and later a location with the maximum sulfate content for potassium sulfate production (Fig. 1.57). Because of this the plants could initially use the simplest processes and have the lowest capital and operating costs. In the initial operation brine with up to 3400 ppm Li was pumped from the Salar in 40 wells, 28 m deep on a 200– 500 m grid, which delivered up to 5280 m3/hr of brine to the solar ponds. There were also 13 monitoring wells to follow any changes in the brine concentration and its depth from the surface. The ponds were lined with flexible PVC or reinforced hypalon membranes, and the brine flowed through the sections of the pond system in series. The initial salt ponds had an area of 1.16 million m2, followed by 3.36 million m2 for the sylvinite ponds, and later 1 million m2 of ponds were installed for lithium production. The plant employed 184 people, of which 120 were hired from the sparsely populated local area. Contractors were used to drill and maintain the wells, harvest the salts, transport them to their respective stockpiles, and reclaim the sylvinite to feed the potash plant’s conveyor belt. They also provided all of the miscellaneous trucking needed at the Salar, and transported the potash to Coya Sur or Maria Elena and the concentrated lithium chloride brine to the Salar de Carmen. SQM unloaded the brine and potash, and stacked the later material at its nitrate plants (Harben and Edwards, 1997). In the solar ponds (Fig. 1.58) salt (halite) crystallized immediately as the brine evaporated, and with the low-sulfate brine utilized by SQM, much of the potassium

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Table 1.30 Initial Schedule for SQM’s Salar de Atacama Plant Construction (Harben and Edwards, 1997) Capacity (mt of product) Stage 1 Potash

Capital expenditure ($MM, US)

Start-up date

300,000

55

October 1995

Stage 2 Lithium carbonate

18,000

51

December 1997

Stage 3 Potassium sulphate Boric acid

250,000 20,000

150

Total investment

January 1998

$256

This table appeared in Industrial Minerals No. 353, February 1997, p. 35. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc 2003.

Figure 1.57 Location of the various pond and plant operations on the Salar de Atacama (SQM, 2001, courtesy of SQM SA).

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Figure 1.58 Aerial view of SQM’s potash–lithium solar ponds at the Salar de Atacama (SQM, 2001, courtesy of SQM SA).

then crystallized as sylvinite. In the initial operations when the ponds were adequately full of salt or sylvinite crystals they were drained for about 1 week, the crystal mass broken and windrowed to drain for 4 weeks, the crystals then loaded and carried to the stockpiles for 4 weeks, and 1 week was spent smoothing the pond floor and re-flooding it to commence another cycle. A permanent 30 cm thick floor of crystals was maintained in the ponds to protect the membrane during harvesting, and the thickness of the deposited crystals was built-up to a minimum of 35 cm before being removed. During harvesting (Fig. 1.59) the crystal bed was first lifted from the floor and broken by Caterpillar or Rahco modified pavement-breakers

Figure 1.59 Harvesting salts at SQM’s solar ponds on the Salar de Atacama (SQM, 2001, courtesy of SQM SA).

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controlled by lasers to cut the harvested crystals exactly to the permanent crystal floor. An average pond produced 35,000 m3 (, 35,000 mt) with each harvest, using ten 20 mt trucks, two bulldozers and seven 3 m3 front-end loaders (Harben and Edwards, 1997). The harvested salt was sent to large storage piles (Fig. 1.60), or some could also be used to reinforce the solar ponds’ walls or make internal baffles in the ponds to better control the brine flow. The sylvinite was taken to the potash plant (Fig. 1.61) and crushed and ground to its sylvite liberation size (about 6 mm), and then the potassium chloride was separated from the mixture in froth flotation cells. The potash was next thickened, centrifuged, washed and trucked about 250 km to their potassium nitrate plants as a , 95% KCl (on a dry basis) slightly moist product. The salt from the flotation cells was also centrifuged, washed and sent to disposal stockpiles, while a bleed stream from the flotation brine was returned to the sylvinite solar ponds. The brine leaving the sylvinite ponds contained about 1% Li, and was sent to the lithium ponds to be concentrated to about 6% Li (38% LiCl, or essentially LiCl’s saturation point), 1.8% Mg and 0.8% B. A portion of this brine to produce the desired amount of lithium carbonate was trucked 250 km to near Antofagasta at the Salar de Carmen, and the remainder allowed to seep into the Salar for potential future use (Harben and Edwards, 1997). A very brief outline of the Salar operations is given in Fig. 1.62.

Figure 1.60 One of SQM’s salt (halite) solar ponds, with a salt disposal pile in the background (courtesy of SQM SA).

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Figure 1.61 SQM’s potassium chloride plant at the Salar de Atacama, with a solar pond in the foreground (SQM, 2002, courtesy of SQM SA).

Figure 1.62 General flow sheet of SQM’s Salar de Atacama process (Harben and Edwards, 1998). Figure published in the Forum on the Geology of Industrial Minerals. Reprinted with permission of the Canadian Institute of Mining, Metallurgy and Petroleum.

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In 2002 the brine was delivered from 92 wells that were equipped with submersible pumps powered by diesel or electric motors. All of the wells were 40 m deep, with perforated casings and no gravel-pack because of the structural strength of the halite. Their pumping rate varied considerably because of the differences in salt porosity, and the wells’ useful life depended upon how rapidly the brine composition changed to an undesirable composition. There were 17 pond strings on the Salar, with 3 for lithium production, and 14 for the potash and potassium sulfate – boric acid plants. The total pond area was 15 million m2 (1500 ha or 3710 acre), including 150 ha for the lithium ponds, and all of the ponds were lined with plastic membranes. Brine averaging 1500 – 2000 ppm Li and 1.85% K was gathered from wells along an 8 km canal (Fig. 1.63) for the potash – lithium ponds, and was then pumped from the canal to the halite ponds (Fig. 1.58). These ponds had an average size of 120,000 m2, and crystallized about 2 million tons of salt per year. The salt was periodically harvested and disposed of in nearby piles that were limited to a 10 m height (Fig. 1.60). When the brine in the halite ponds became saturated with potassium chloride it was pumped to , 100,000 m2 sylvinite ponds where it was joined by brine from a few wells that were already saturated with potash. The sylvinite ponds were also periodically taken out of service, drained and harvested, and their salts sent to the potash plant storage – drainage piles (Fig. 1.61). In 2002 the total of the potash plant’s capacity was 650,000 mt/yr, and the KCl was hauled to Coya Sur in covered

Figure 1.63 The brine canal for SQM’s potash–lithium solar ponds (courtesy of SQM SA).

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dump trucks to be converted to potassium nitrate (also with a plant capacity of 650,000 mt/yr). Brine from the sylvinite ponds next went to the carnallite ponds (Fig. 1.64), and from there to the 500,000 m2 bischoffite ponds (Fig. 1.65). These two series of lithium ponds were also periodically taken out of service to harvest predominantly carnallite from the first ponds, and bischoffite from the later ponds. These minerals were stockpiled separately, with some of the bischoffite sold as magnesium chloride (with a capacity of 450,000 mt/yr), and the carnallite saved for later conversion to potash. The six carnallite ponds were divided into two groups of three, with the higher sulfate brine directed to one group, and then its end-liquor was returned to the Salar by being flooded onto its porous surface. The final brine from the bischoffite ponds contained 6.0– 6.1% Li, and was sent to 40,000 m2, about 3 m deep holding ponds to await truck shipment to the lithium carbonate plant. The plant had a capacity of 22,000 mt/yr of Li2CO3 in 2002, to be raised to 28,000 mt/yr in 2003 (Moura, 2002; Etchart, 2002; Nakousi, 2003). The potassium sulfate and boric acid plants (Fig. 1.66; with capacities of 250,000 and 16,000 mt/yr, respectively) started production in 1998 using a separate brine supply and solar evaporation system. After the brine had left the initial halite ponds all of the potassium and sulfate salts were allowed to crystallize and be harvested together from one set of ponds. In 2002 the first processing step was to leach its halite content, and then convert the residue to schoenite with return liquor from

Figure 1.64 One of SQM’s carnallite solar ponds with a typical mild wind-rippled surface (courtesy of SQM SA).

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Figure 1.65 Bischoffite draining prior to harvesting from one of the final lithium solar ponds (courtesy of SQM SA).

the potassium sulfate crystallizers. The schoenite was next reacted with potash and converted to potassium sulfate (Fig. 1.67; Ramirez, 2002). In other potassium sulfate plants the harvested salts are first converted to fairly large crystals of schoenite, and the halite then removed by flotation. The harvest salts could also be initially floated

Figure 1.66 SQM’s potassium sulfate plant at the Salar de Atacama (SQM, 2002, courtesy of SQM SA).

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Figure 1.67 A general flowsheet for the production of potassium sulfate and boric acid at the Salar de Atacama (Harben and Edwards, 1997; this figure appeared in Industrial Minerals No. 353, February 1997, p. 31; published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc 2003).

to remove the halite, but the liberation size of most of the potassium sulfate compounds from solar ponds is very small, and the yields are poor. The final brine from the sulfate ponds is acidified with sulfuric acid to crystallize (“salt-out”) boric acid. Lithium sulfate and magnesium sulfate could also be readily recovered from the boric acid end-liquor if desired (CORFO, 1985). A flowsheet for the Salar de Carmen (Antofagasta) lithium carbonate plant in 2002 is shown in Fig. 1.68. Brine is first unloaded from tank trucks bringing it from the Salar de Atacama, and sent to storage tanks. It is a sparkling clear, bright yellow color with a boron content of about 8000 ppm B. The boron was first removed from the brine to a 2 ppm B level by contacting it in four stages with an alcohol –kerosene mixture in a liquid – liquid extraction plant. After extraction the brine was colorless, indicating that the color may have been from a metal or organic borate, since alkali borates are colorless. The boron was removed from the solvent in stripping cells with a dilute sodium hydroxide solution, and this sodium borate –lithium chloride solution was returned to the Salar. Pictures of the plant are shown in Figs. 1.69 and 1.70. A Chilean laboratory study by Orrego et al. (1994) (based upon the Folkestad et al., 1974 patent), had suggested using iso-octanol as the boron solvent in a 50 vol% mixture with kerosene. In their study the solvent contacted the acidified brine (, 0.1 N Hþ) in a one-to-one (by volume) ratio in four countercurrent liquid – liquid extraction stages to reduce the boron in the brine to less than 5 ppm B (Table 1.31). The solvent could then be regenerated by three stages of water wash or with a 0.02 N NaOH solution. The aqueous wash would contain a significant amount of boron (the equivalent amount at SQM in 2002 would be 3500 mt/yr H3BO3), and

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Figure 1.68 Flowsheet for SQM’S lithium carbonate plant at the Salar de Carmen (SQM, 2002, courtesy of SQM SA).

Figure 1.69 SQM’s lithium carbonate plant at the Salar de Carmen. View from Office to the east, Laboratory on left; Warehouse in middle right; Plant in middle left (courtesy of SQM SA).

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Figure 1.70 Second view of SQM’s lithium carbonate plant. Processing building to the right; boron extraction to the left (courtesy of SQM SA).

Table 1.31 Isotherms at 258C for Boron Extraction from Salar de Atacama Brine, and Re-extraction in Water (Orrego et al., 1994) Isoterma de Extraccio´n

Isoterma de Reextraccio´n

[Boro]ac (g/liter)

[Boro]org (g/liter)

[Boro]ac (g/liter)

[Boro]org (g/liter)

6.70 6.21 5.56 4.07 2.50 0.58 0.063 0.040 0.031 0.026 0.003

11.46 11.46 11.46 11.35 10.70 7.27 3.89 2.60 1.56 1.12 0.78

64.40 49.05 29.49 19.84 9.82 5.24 3.58 2.19 1.57

2.04 1.43 1.41 1.32 1.20 0.77 0.51 0.31 0.23

Boron extraction with one part of 50% isooctanol in kerosene to one part of brine, and re-extraction with six parts of solvent to one part of water. Reprinted courtesy of Nucleotecnica.

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could be returned to the Salar’s boric acid plant. Figure 1.71 shows the rapidity of the extraction, and indicates the high lithium loss (in these tests , 10%) with the boron based upon the brine’s greatly reduced density and viscosity. After boron extraction the brine was next sent to the “Chemical Area” building where it was first reacted with a fraction of the stoichiometric amount of soda ash needed to precipitate all of the magnesium. The resultant magnesium carbonate was then filtered on two large rotary drum vacuum filters with a traveling cloth filter-anddischarge membrane. The brine was next reacted with lime to remove the rest of the magnesium and much of its sulfate content. This magnesium hydroxide – calcium sulfate precipitate was then filtered in a bank of plate and frame filter presses. The filter cake from both precipitations was next repulped and sent to two solid bowl centrifuges. Their discharge cake contained from 30 to 40% moisture, and was hauled by truck to a local dump. The filtrates from each of the dewatering devices were returned to the brine stream. The purified lithium chloride brine was next heated and precipitated with soda ash to produce the lithium carbonate product. It was filtered on a belt filter, and washed first with wash water and then with fresh water. The filtrate still contained about 1% Li, so it was recycled to the feed brine to the extent that water and salts had been removed from the system. The lithium carbonate filter cake was next dried in a direct fired rotary dryer with a multiple cyclone dust collector. The 99.3– 99.4% Li2CO3 product then went to a three-tray screen, the oversize was ground and returned to the screen, the middle size became a product grade, and the fines were sent to a compactor to form a , 20 mm thick briquette. These granules were ground and fed to a screen to form the desired product sizes, and some of the product was further ground to form a very fine powder. In 2002 they had 18 product grades of different particle size and sulfate content. The products were shipped in 0.5 and 1 t bulk bags; 25 kg, 25 and 50 lb bags; or in plastic lined 100 kg fiber drums. The plant operated on a three shift, 7 days/week basis, had a semi-automatic control system and an analytical laboratory to insure the product quality and assist in the plant operation. There were automatic sprinklers over the liquid extraction mixer – settlers in case of fire, and any off-spec or spilled product was re-dissolved and added to the incoming brine. Only a nominal amount of pilot plant testing had been needed for the plant design and construction, and the start-up operation went very smoothly (Nakousi, 2003; Arqueros, 2002; SQM, 2002; Harben and Edwards, 1998). The initial offering price that SQM posted for its lithium carbonate was $0.90/lb, or about half of the then-existing market price. This reduced price considerably stimulated the market for lithium carbonate (sometimes at the expense of lithium ore concentrates), and caused high-cost producers to close their plants. This, in turn, has allowed the price to again slowly rise, but stay below that of higher-cost producers. The SQM plant’s nominal capacity was 40 million lb/yr of lithium carbonate in 1998, and they sold 15.4 million pounds. They expected to sell 28.6 million in 1999 and then run at near-capacity of 22,000 mt/yr of lithium carbonate (Schmitt, 1998). In 1999 SQM began selling lithium hydroxide, and later

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Figure 1.71 The rate of extraction and solvent washing, and the density and viscosity of strong lithium chloride brine before and after boron extraction (Orrego et al., 1994; reprinted courtesy of Nucleotecnica).

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lithium chloride. In 2002 they were considering installing a butyl lithium plant in Texas (SQM, 2001; USGS, 2000). Salar de Hombre Muerto, Argentina; FMC Production of lithium carbonate and chloride started at the Salar de Hombre Muerto in 1997 utilizing a new process developed by FMC. Although details of the process have not been disclosed, several of their recent patents are based upon selective lithium adsorption onto alumina. In the early stage of their development work 17 wells were drilled in the Salar’s central salt mass, from which 678 m of core were obtained to determine the salt’s porosity. Brine samples were taken at various depths, and the reserves of the deposit were estimated (Anon., 1984b). Considerable experimental work was then conducted in the laboratory, pilot plants and on the Salar. An extensive infrastructure was next constructed for this remote, high altitude location, and eventually a multiple-well brine-gathering system was installed on the Salar. In their process patents and announcements the brine to be utilized was assumed to be saturated with NaCl and contain about 600 ppm lithium. In the alumina patents brine would be sent in counter-flow through a series of columns packed with polycrystalline hydrated alumina. The flow rate and number of columns would be adjusted so that the lithium would be almost completely (and fairly selectively) adsorbed from the brine leaving the last (freshest) column. It would then be discharged from the plant and returned to an area of the Salar far from the inlet wells. After the alumina in the first (or oldest) column was nearly saturated with lithium it would be removed from brine flow circuit, and the lithium mostly removed from the alumina (eluted) by a water wash. The resultant solution would contain up to 1% lithium, and could then be concentrated in solar ponds (Fig. 1.72) to the desired concentration for further purification and/or processing. The eluted column would next be given a saturated sodium chloride wash (perhaps containing some lithium) to recover the entrained lithium and to raise the ionic concentration in the alumina.

Figure 1.72 Solar ponds at the Salar de Hombre Muerto (Chem. Week, 1995 and 1998; reprinted with permission from Chemical Week, November 22, 1995, Chemical Week Associates; reprinted with permission from Chemical Week, December 2, 1998, Chemical Week Associates).

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It would then be returned to the brine adsorption system as the last column in the lithium recovery process. In preparing alumina to have a high lithium capacity for the process it would be necessary to initially react the alumina with lithium chloride in a saturated sodium chloride solution. This would form LiCl/Al(OH)3 crystals in which the lithium chloride was present in amounts up to a 0.2 mol fraction. The alumina has a strong affinity for the lithium in high-ionic solutions, but the adduct is not stable (i.e., the lithium can be eluted) in dilute solutions. The aluminum hydroxide reacts somewhat as if the lithium chloride were a hydrate [Al(OH)3·n H2O·LiCl] molecule that was stable (could be attached or removed) depending upon the solution’s total ionic strength. If the alumina were first treated with lithium hydroxide followed by hydrochloric acid the LiCl in the LiCl/Al(OH)3 crystal could be increased up to a 0.33 mol fraction. The lithium adsorption – desorption cycle was stated to be repeatable many times before the lithium-treated alumina had to be regenerated or discarded. Eluate concentrations up to 1.1% Li with comparatively small amounts of impurities, and loadings of 3.6 g Li/liter of alumina were claimed in the process patents (Bauman and Burba, 1997, 1995). It was announced that the commercial process worked very well, reportedly at a 20% cost saving over the conventional solar evaporation-magnesium and sulfate precipitation process. The initial brine strength varied from 0.22 to 1.08 g/liter Li, and averaged 650 ppm Li. The plant had a capacity of 45 million lb/yr of LCE (although it produced fairly pure LiCl directly from the brine), and cost $68 million. The product distribution as the plant started in late 1997 was 12,000 mt/yr lithium carbonate and 9650 mt/yr lithium chloride. However, in 1999 after SQM had greatly lowered the price of lithium carbonate they partly closed the plant and contracted to purchase lithium carbonate from SQM. They announced that some lithium chloride would still be produced from the Salar, and then purified and perhaps converted to other products at Guemes, Salta province, Argentina (Fig. 1.6; Harben and Edwards, 1998). In 2002, it was reported that this production was 4729 mt LiCl and 906 mt Li2CO3 (944 mt of Li; USGS, 2002). Some of the favorable conditions for the operation were that the Argentine government had granted complete ownership of the Salar to FMC, the brine had a relatively high lithium content, there was an adequate area for solar ponds, the evaporation rate was quite high, and there was an excellent fresh water supply. However, if they only produced a crude product at the Salar (the lithium chloride brine from the solar ponds) and shipped it to another plant for final processing, the transportation costs would be very high. The trip to the nearest harbor would require first trucking the product a considerable distance (, 145 km) over a tortuous unpaved mountain road, loading it into rail cars, and then shipping by rail down the steep west side of the Andes, across the Atacama desert (420 additional km) to Antofagasta. If the product were to go to a plant near Salta, the same mountain road would have to be traveled, and the rail haul would go down the east side of the Andes

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to the plant (, 395 km). In either case supplies would have to follow the same difficult routes back to the Salar. Searles Lake, California; American Potash & Chemical Co As Searles Lake brine was evaporated in triple-effect evaporators the lithium concentration reached about 140 ppm Li (Table 1.9). During this evaporation process salt (NaCl) was crystallized predominantly as large crystals, burkeite (3Na2SO4·Na2CO3) as medium-sized crystals, and dilithium sodium phosphate (called licons; Li2NaPO4) as very fine crystals. This allowed most of the salt to be removed from the other crystals by hydraulic classification, and the licons (representing about 40% of the feed brine’s lithium) to later be removed by froth flotation. The untreated licons contained about 9.3 –10.7% Li, and typical yearly production rates were 200 mt in 1938, 522 mt in 1943 and 765 mt in 1976. Originally the licons were sold to the Foote Mineral Co., but starting in 1951 they were converted to lithium carbonate and phosphoric acid at Searles Lake. Searles Lake brine had a comparatively high organics content (called “humates”), and coconut oil was added as a froth-control agent in the brine evaporators. Much of these organics were adsorbed onto the burkeite and licons as they crystallized, and these fine salts were separated together from most of the salt crystallized in the evaporator, and then filtered and washed. Next this burkeite mixture was given a hot, partial leach, and in a second step the burkeite was totally dissolved. Most of the licons did not dissolve, leaving the remaining licons with a high proportion of the organics. In the early days of the soda products operation these licon solids contaminated the products, caused slurries to foam, they inhibited filtration, and formed a sticky scum on the top of liquid-filled tanks and on metal surfaces. To reduce this problem in 1936 they began to scrape the scum from the surface of various processing vessels, put it in a small tank and on a campaign basis add water and steam, and then filter and dry it to form their first , 20% Li product (Table 1.18) in the amount of about 200 mt/yr. In 1942 the Government asked companies to increase their lithium production, so at Searles Lake they installed a plant to recover the licons much more effectively, with the first production being in mid-1943. They took advantage of the licons being self-floating (because of their adsorbed organics) by designing a vessel to be aerated, and then collecting the foam. Initially they added some kerosene (to better control the foam) to the burkeite leach tank, and then cooled the , 0.2% licons slurry of burkeite leach liquor to about 278C in a multiple-spray cooling tower. The slurry next went to a “conditioning” tank, and then to four parallel 10,000 gal flotation vessels where air was forced at 4 – 8 psi through porous carbon plates in their base. The foam that was formed was scraped from the top surface, and the clear brine passed on to the soda products plant. Later they found that the kerosene and conditioning could be eliminated, and that instead of porous plates the air could be introduced into the suction of the pumps bringing the licons slurry to the flotation

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tanks. They also installed a second stage of flotation with a commercial foam-release separator, and a small amount of additional licons were scraped from the surfaces of several of the soda products plant tanks. The foam was then sent to agitated, heated tanks to dissolve any remaining burkeite (it would also float), filtered, washed and prior to 1951 dried in a steam-jacketed, mixing-type dryer. This product was shipped in 100 lb bags, and in 1945 sold for $256/mt, f.o.b. Trona (Rykken, 1976; Gale, 1945). In 1945 they began work on a process to produce their own lithium carbonate, following the patent of May (1952; Fig. 1.73) who noted that lithium sulfate and sodium sulfate had a quite low (, 1.4% Li) solubility in . 30 – 40% phosphoric acid. In the commercial process the licons were first roasted to burn off their organics content, and then mixed with 93% sulfuric acid at 1158C to form 45 – 50% phosphoric acid and a mixture of lithium and sodium sulfate crystals. The phosphoric acid was then evaporated to 78%, which crystallized additional salts, and reduced the lithium content to less than 0.4% Li. The mixed sulfate crystals were centrifuged, washed and re-dissolved, and then soda ash was added to the solution at

Figure 1.73 Flowsheet for the conversion of licons into lithium carbonate and super-phosphoric acid (May, 1952).

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938C to precipitate lithium carbonate. The end-liquor was next treated with a small amount of phosphoric acid and evaporated to nearly sodium sulfate’s crystallization point, precipitating trisodium phosphate that was recycled to the licons leach step. The final solution then only contained , 0.07% Li instead of its original 0.28% Li, and it was sent to the soda products plant. The operation produced about 900 mt/yr of lithium carbonate, with an overall recovery from the lithium in the brine entering the evaporators of about 30% (Rykken, 1976; Williams, 1976). The operation was terminated in 1978 after 40 years of production when the soda products plant was closed. Chinese Lakes Plans were announced in 2000 that lithium and various other products would be produced from Zabuye Salt Lake by the Tibet Lithium New Technology Development Co. The lake has a complex mineral content, including over 1 million mt of lithium, along with recoverable amounts of boron, bromine, cesium and potassium (Garrett, 1992). Experimental work had produced a lithium concentrate containing 78% Li2CO3, and a $170 –240 million, multiple-product plant was proposed to be started in 2003 (Saller and O’Driscoll, 2000). Discussions have also been made on the possible by-product recovery of lithium from the end-liquors (with about 120 ppm Li) of Qinghai Basin’s large projected potash plant at Qarhan Lake (Fig. 1.74). A joint venture called Qinghai Lithium Ltd. was formed in 2000 by Pacific Lithium Ltd. of New Zealand and the Chinese government. It was stated that the lake contains 1 million tons of lithium, 1 million tons of boron and greater than 17 million tons of potassium (USGS, 2000; Garrett, 1996).

Figure 1.74 Structure of Qaidam Basin, China (Sun and Lock, 1990). Legend: (1) mountains; (2) playa surface; (3) major potash deposits; (4) smaller potash deposits; – – –, outline of Qarhan Playa; Q, Qarhan Salt lake; D, Dabuxun Saline Lake. Reprinted by permission of Science Press (China).

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Various Proposed Processing Methods Liquid Extraction The selective extraction of lithium from brines has been a much studied but elusive target. Many organic compounds can dissolve some-to-considerable lithium chloride with some selectivity for it compared to sodium and potassium chloride (and perhaps calcium chloride), but usually not magnesium chloride. An example of this is organic alcohols, where often the lithium chloride solubility is greater the lower the molecular weight, and the selectivity improves with greater molecular weight. Based upon these properties Hermann (1966) has suggested that a dry mixture of lithium, sodium and potassium chlorides or sulfates can be dissolved in anhydrous alcohols containing 3 – 8 carbon atoms (preferably butanol; solubility 10.57 g LiCl/100 g solvent). Almost no sodium (82 ppm) and sulfate dissolves, and very little potassium, so upon partial evaporation of the saturated butanol most of the impurities will crystallize. After they are filtered the lithium chloride can be recovered by further distillation or by re-extraction with water. Brown and Beckerman (1990) more recently suggested doing the same treatment with isopropyl alcohol (solubility 12.2% LiCl, 67 ppm Na), and others have discussed a similar lithium extraction with propanol (16.22 g LiCl/100 g solvent, 152 ppm Na), isobutanol (7.3% LiCl, 113 ppm Na), pentanol (8.1% LiCl, 34 ppm Na), 2-ethylhexanol, isopentanol, amyl alcohol (9.02 g LiCl/100 g solvent), isoamyl alcohol, allyl alcohol (4.36 g LiCl/100 g solvent), tetrahydrofuran (4.6% LiCl, 42 ppm Na) and other alcohols. Morris and Short (1963) noted that 0.001 –1.84 M of lithium chloride could dissolve in tri-n-butyl phosphate, and that pure solutions had distribution coefficients (D, the concentration of lithium in the solvent/concentration of lithium in the aqueous phase after vigorous contact and settling) of 0.002 –0.16, depending upon the initial LiCl concentration. Despite the appreciable solubility of low-molecular weight alcohols in water or brine, Gabra and Torma (1978) also suggested using butanol to extract lithium from aqueous solutions of sodium, potassium and calcium. The distribution coefficients were very low, but still could allow an extraction with some purification of the lithium (Table 1.32). The use of insoluble carriers (such as kerosene) for the solvent (to lower the solvent loss in the aqueous phase) reduced the extraction to impractical levels. Holdorf et al. (1993) suggested such an extraction with amyl alcohol or fermentation alcohol (e.g., 52.1% 2-methylbutanol-1 and 47.9% 3-methylbutanol1). They preferred 2 –2.5 volumes of solvent per volume of brine, six extraction stages and two re-extraction (stripping) stages. They claimed a 95% lithium recovery from a 260 ppm Li gas field brine, with 52% of the magnesium also being extracted, and only a 0.1% solvent loss in the stripped brine. This low solvent loss is difficult to understand with such a low molecular weight alcohol. Many other organic compounds have been noted that might be able to somewhat selectively extract lithium, but essentially all of them also extract magnesium, and require a drastic pH or composition modification of the brine. For instance some

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The Extraction of Lithium, Sodium, Potassium and Calcium from a Dilute Brine Containing All Four Chlorides (Gabra and Torma, 1978) Distribution coefficient

n-Butanol sec-Butanol Isobutanol Pentanol Isopentanol 2-Ethylhexanol 2-Ethylisohexanol Octanol 2,4-Dimethyl-3-pentanol o-Chlorophenol p-Chlorophenol o-Cresol m-Cresol p-Cresol Cresol Molar sol. of phenol/benzene p-sec-Butylphenol/benzene p-1,1,3,3-Tetramethyl-(butyl)phenol/benzene p-tert-Butylphenol/benzene

Separation factor

DLi

DNa

DK

DCa

SLi Na

SLi K

SLi Ca

0.058 0.044 0.018 0.011 0.01 0.004 0.005 0.005 0.006 0.018 0.0119 0.045 0.007 0.011 0.008 0.011 0.005 0.002 0.012

0.023 0.022 0.009 0.008 0.006 0.007 0.006 0.007 0.007 0.0056 0.0049 0.019 0.014 0.01 0.019 0.003 0.017 0.003 0.0089

0.020 0.021 0.009 0.005 0.007 0.004 0.005 0.006 0.008 0.020 0.01 0.03 0.013 0.006 0.0049 0.005 0.0013 0.007 0.008

0.017 0.020 0.005 0.006 0.005 0.006 0.006 0.006 0.014 0.013 0.012 0.006 0.0008 0.0012 0.009 0.011 0.009 0.003 0.012

2.5 2.0 2.0 1.38 1.66 0.57 0.83 0.71 0.86 3.21 2.43 2.36 0.50 1.0 0.42 3.67 0.29 0.67 1.35

2.9 2.09 2.00 2.20 1.43 1.00 1.00 0.83 0.75 0.90 1.19 1.50 0.53 1.83 1.63 2.20 3.85 0.29 1.50

3.4 2.1 3.6 1.8 2.0 0.7 0.8 0.8 0.4 1.4 1.0 7.5 8.8 9.2 0.9 1.0 0.6 0.7 1.0

With n-butanol the lithium extraction was the same at pH values from 1 to 11, although pH . 8 reduced the calcium extraction. The optimum ratio of solvent to brine was 1/1, among the ratios of 1/5 to 3.2/1 that were tested. As high as a 90% lithium recovery was obtained in four mixer–settler stages from solutions containing from 30–300 g/l lithium. Reprinted from Hydrometallurgy, Vol. 3, p. 26, Table 1, q1978, with permission from Elsevier.

diketones can extract lithium and a limited amount of magnesium (to improve the Li/Mg ratio 10 to several 100-fold), but only when the solution is strongly basic. Dipivaloylmethane ([CH3]3C–CyO)2 –CH2 as a 0.7 M solution in ether (as a carrier), with equal volumes of Great Salt Lake (GSL) potash plant end-liquor (Tables 1.9 and 1.33) and the solvent, and with the brine adjusted to 3% NaOH, can extract 90% of the lithium in one stage of mixing and settling. With only 0.3% NaOH the lithium extraction was 20%, and with no caustic, or with the carrier solvent being kerosene, benzene, petroleum ether, acetyl ether, chloroform, carbon tetrachloride, and many other solvents there was no extraction. Dang and Steinberg (1978) have also hypothecated a lithium recovery process for Smackover oilfield brine based upon this solvent. Other diketone extractants such as pivaloyltrifluoroacetone have an extraction coefficient for lithium of 0.1 with and without pH adjustment, 4-methylbenzoyl trifluoroacetone’s coefficient is 0.06, and the chelating

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agent diacetylmethane 0.043, but in all cases magnesium was also extracted. Other expected lithium solvents such as tri-n-butylphosphate, bis(2-ethylhexyl)phosphoric acid and mono(octylphenyl)phosphoric acid had no extraction effect on lithium in untreated Great Salt Lake end-liquor. Others have suggested the b-diketones thenoyltrifluoroacetone and dibenzoylmethane, and the acidic organophosphorus compounds di-(2-ethylhexyl)phosphoric acid and 2-ethylhexylphosphonic acid mono 2-ethylhexyl ester as lithium extractants, but their effectiveness is limited. Many authors have studied lithium extraction processes in which they only attempted to separate lithium from other alkali metals, and sometimes calcium. For the separation from calcium chloride Goodenough and Gaska (1967) used alcohols or ketones and massive amounts of ammonia or urea. Ma and Chen (2000) studied the addition of two organic compounds that acted synergistically together (only LIX54 was a slightly effective solvent by itself) in batch experiments and on a supported liquid membrane. They used the commercial ion exchange chelating agent LIX54 (a-acetyl-m-dodecylacetophenone) and the neutral complexing agent TOPO (tri-octyl-phosphine oxide) in a kerosene solution for the batch experiments or to soak a thin membrane film of Celgard 2500 (37 –48% porosity; 0.05 – 0.19 mm pore size). With the membrane the film was dried under vacuum for 30 min, placed in the extraction cell, and a Li, Na, K solution passed over it. The extraction coefficients were only appreciable at a pH higher than 12, and in the batch experiments LIX54 had a moderate coefficient for lithium, and a lesser one for potassium and sodium. The mixed solvents had a very high coefficient when lithium was alone, but it fell to a moderate value in mixed-salt solutions, and some sodium and potassium were also extracted. The stripping (elution) solution for the solvent or membrane was water at a pH of 0.05– 2. The initial solution concentrations were 10 –100 ppm Li, 1000 – 7000 ppm Na and 30– 500 ppm K, and with the membrane tests a 90% lithium removal was obtained after 2 hr of recirculated flow, and a 70 – 75% recovery was obtained from the strip solution in a similar period. The strip solution had about the same lithium concentration as the feed solution, but the reduction in its sodium and potassium content was not noted. Kinugasa et al. (1994) studied the kinetics of this solvent pair, and noted that in their work sodium was not extracted with the lithium. Somewhat similar results have been obtained when crown ethers (such as dibenzo-14-crown-4) and lipophilic anions were incorporated into membranes, with the eluate having a considerable reduction in the divalent cations (Olsher, 1982). Lee et al. (1968) studied the synergistic use of trioctylphosphine oxide (TOPO) again as the adduct former, but with the chelating agent dibenzoylmethane. Their individual extraction coefficients from 0.02 M Li or 0.88 M KCl when in a 0.1 M KOH solution were 0.025 and 0.010, respectively, and 131 for a mixture of LiCl and KCl. When in p-xylene as a solvent carrier, D was 82.4, in carbon tetrachloride 68.5 and dodecane 49.1, while other solvents greatly reduced the lithium extraction. The selectivity factors (DLi/DNa) for Li/Na were 570 and Li/Cs 12,400 in 3 M NH4OH solutions, and the lithium extraction was only effective at pH values above about 10.

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Table 1.33 The Extraction of Lithium by Ferric Chloride in an Insoluble Solvent, wt.% (Nelli and Arthur, 1970) Feed

Raffinate

Li 0.116 0.0113 Na 0.118 0.098 K 0.058 0.042 Mg 8.55 6.19 Ca 0.005 — Fe 0.001 0.010 SO4 2.46 — Br 0.212 — B 0.070 — P 0.005 — HCl 0 0.15 Density 1.344 1.228 Pounds 1000 1385 Add to the brine or solvent, as pounds Fe 1.05

Wash liquor

Strip liquor

Strip recycle

0.305 0.120 0.024 1.72 — 6.40 — — — — 0.20 1.240 212

0.420 0.009 0.001 0.02 — 5.60 — — — — 0.30 1.180 294

0.087 0.079 0.001 0.005 — 6.20 — — — — 0.03 1.170 266

H2O 164

H2 O 241

H2O 220

HCl 6.82

Na 17.5

Solvent 626

Solvent 1540

Li 0.071

HCl 0.63

Product 0.360 5.86 0.001 0.02 0.0002 0.002 0.016 0.006 0.0027 0.001 0.03 1.130 296

The only fairly selective liquid extraction process that has been suggested for lithium solutions with a high magnesium content is that by Nelli and Arthur (1970). They employed end-liquor (Table 1.33) from the massive Great Salt Lake (Strum, 1980) potassium sulfate plant’s solar ponds. The process involved a quite elaborate series of liquid – liquid extraction steps centered around the substitution of lithium chloride into a ferric chloride complex. In the presence of strong chloride solutions and some hydrochloric acid lithium is nearly all converted into the relatively stable lithium tetrachlorferrate, which can be easily extracted by a number of solvents. In their process (Fig. 1.75 and Table 1.33) about an equal stoichiometric amount of ferric chloride is added to the lithium in the brine, along with enough hydrochloric acid to make the solution 0.04– 0.1 N in HCl. This forms the iron complex which is then extracted by 1 – 2 parts of solvent per part of brine, with the solvent being a 20% mixture of tributyl phosphate and 80% diisobutylketone. This lithium extraction is made in seven countercurrent mixer – settler stages. The exiting brine (raffinate) is depleted of about 90% of its lithium content, but unfortunately, some magnesium is also extracted. This magnesium is recycled back to the initial extraction step (unavoidably along with about 34% of the lithium) by a “washing” step of the solvent with four countercurrent mixer – settlers and just enough water (1 part/10– 11

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Figure 1.75 Flow sheet for the ferric chloride extraction process for lithium from high-magnesium brines (Nelli and Arthur, 1970).

parts solvent) to remove all of the magnesium and the minimum of lithium. Then the remaining lithium tetrachlorferrate is “stripped” with more water (1 part/3 – 7 parts of solvent) in five countercurrent stages and the solvent is ready to be recycled. The strip water is next made about 2 N with sodium chloride, and contacted with a different solvent in six countercurrent stages to remove its ferric chloride content. This solvent is an equal molar mixture of di(2-ethylhexyl)phosphoric acid (20 vol%) and tri-n-butyl phosphate (30 vol%) with benzene or a similar non-polar diluent (50 vol%), and its total volume is about twice that of the salt-adjusted strip solution. The raffinate from this step is the product, containing about 0.36% Li, 200 ppm Mg and 20 ppm Fe. The second solvent is then stripped of its ferric chloride with about 0.3 parts of water in six countercurrent stages. The solvent is recycled, and its raffinate joins the feed brine in the first stage as the ferric chloride source. Alumina Adsorption A very large number of articles and patents have been issued on methods to precipitate or adsorb lithium from brines, but by far the most common is the suggested adsorption or co-precipitation of lithium on aluminum hydroxide or alumina. When aluminum chloride is added to a neutral or basic solution containing lithium most of the lithium joins the voluminous aluminum hydroxide gel-like precipitate. In a similar manner, hydrated aluminum hydroxide can adsorb lithium, and a wide range of mixtures with aluminum hydroxide (either in a solid phase or as a co-precipitate) can act in a similar manner. This method was first proposed by

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Goodenough (1960), with further details provided by Neipert and Bon (1967), and then by numerous investigators in many combinations (i.e., such as with ion exchange resins or other co-precipitants). Many different lithium-containing brines were also tested, such as seawater, geothermal brines from Italy, France and Japan, as well as brines from the Dead Sea, Salton Sea, the Great Salt Lake and various oilfield waters. However, no commercial processes were ever developed utilizing aluminum hydroxide precipitates. One of the investigated brines was the Wairakei Geothermal power plant discharge. In 1986 this highly mineralized (As, B, Li, Si, etc.; Table 1.13) brine was being dumped into the Waikato River, causing considerable pollution. Consequently, studies were made on the recovery of many of the brine’s constituents, including an aluminum hydroxide co-precipitation process for lithium, a patented process for the recovery of high-grade colloidal silica, and the recovery of calcium silicate and arsenic. Rothbaum and Middendorf (1986) found that after silica removal a 95% yield of lithium could be obtained by adding sodium aluminate to the brine at a pH of 10, and at 308C. The lithium could be recovered from the alumina gel (which was claimed to filter well) by washing at 608C, and the residual alumina could be recycled by dissolving the gel in sodium hydroxide. However, since the gel only contained 3% solids the eluted brine was only four times stronger than the original solution, and did contain some other salts. With the fairly concentrated Italian and French geothermal brines (Table 1.13) at a nearly neutral pH, a modest excess of aluminum chloride, and temperatures from ambient to 808C there was an almost complete recovery of lithium in the aluminum hydroxide gel. However, there was no practical means to make a product from the lithium chloride – aluminum hydroxide mixture (Pauwels et al., 1990). The Goodenough (1960) and Neipert and Bon (1967) patents were based upon removing lithium from the high-calcium, medium– low magnesium dolomitization brines of the Michigan Basin’s potash end-liquors (56 ppm Li; Table 2.5), both as-is and concentrated. Their optimum temperature of precipitation was 818C, the pH 6.8 and using about 0.007 parts of AlCl3·6H2O/g of brine (with the as-is brine). However, both the optimum temperature and pH varied with the brine’s total concentration and magnesium content, so the conditions were different with concentrated or magnesium-precipitated brine. Their recovery was from 80 to 90%, and they suggested removing the lithium from the aluminum hydroxide precipitate with hot water. With Salton Sea brine laboratory tests had apparent technical success (up to 99% recovery), but were not considered to be economically practical. Under optimum conditions aluminum chloride in the amount of 3.0 times the stoichiometric amount of lithium was added to the brine at a pH of 7.5 and a temperature of 758C. It was noted that if most of the magnesium had been removed lower temperatures (25 – 508C) and different pH’s would be optimum, but that with magnesium higher temperatures (50 – 1008C) and lower pH’s were required (Berthold and Baker, 1976). Seawater, even though it only contains 0.17 ppm Li, has been studied for lithium

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recovery in many reports, such as by Kaneko and Takahashi (1990). They found that the best adsorbent was a 50 mol% mixture of aluminum hydroxide and magnesium ˚ . A 40% lithium hydroxide, precipitated together and dried to a pore size of 20 – 30 A recovery was obtained, and when eluted with 0.05 N HCl there was a sevenfold increase in the lithium concentration. With Dead Sea potash pond end-liquor Pelly (1978) also used aluminum chloride to co-precipitate the brine’s 30 ppm of lithium, adding a threefold excess of AlCl3, adjusting the pH to 6.8 – 7.0, heating the brine to about 558C, and allowing a 4 hr residence time. After the aluminum hydroxide (with about an 80% Li yield) was filtered and washed, the precipitate was dissolved with hydrochloric acid to form a 680 ppm Li solution with the composition shown in Table 1.34. For each ton of lithium recovered it required 73 t of lime for the pH adjustment, 19,000 m3 of water to wash the precipitate, and 92 t of HCl for the product leaching. Epstein et al. (1981) continued this work, but dissolved the lithium – alumina precipitate in isoamyl alcohol, and then stripped the lithium from the alcohol with water. A 1.7% Li solution was obtained, but it also contained 1.2% Ca and 0.4% Mg, with considerable loss of alcohol and lithium. Bukowsky et al. (1991) proposed concentrating Smackover oilfield brine, precipitating the lithium with aluminum hydroxide, leaching the alumina with iso-amyl alcohol, stripping the alcohol with water, precipitating the solution with soda ash, and carbonating the slurry to form a fairly pure lithium bicarbonate solution after filtration. Tests have also been made on adsorbing lithium onto alumina from strong potash end-liquor dolomitization brines, such as found in the Angara-Lena basin’s Znamenskoe deposit in the Irkutsk oblast. This brine contained, as wt.%: 25.44 CaCl2, 9.40 MgCl2, 1.49 NaCl, 0.52 KCl, 1640 –1870 ppm LiCl, and 37.31% total salts, with a density of 1.34 g/cc. Ryabtsev et al. (2002) followed the work of Bauman and Burba (1997, 1995) and described their aluminum oxide adsorbent as having the formula LiCl·2Al(OH)3·m H2O. They noted that it was formed in a layered structure that resulted in a molecular-sieve effect that could only be penetrated by lithium. They further stated that up to 40% of the lithium in this compound could be released by a water wash, and then be replaced (the adsorption reaction) by contact with a lithium-containing strong brine. They prepared their adsorbent by mixing equimolar amounts of crystalline LiCl and Al(OH)3 with 0.5 m of H2O, forming granules, and then drying them. It was stated that their capacity for lithium was 7 mg/g of adsorbent. When the granules were made by using 6 – 8% PVC as a binder, they had a capacity of 5 –6 mg lithium/g of granules. The PVC – alumina pellets were ground to various size fractions, placed in columns, and then contacted with brine at various flow rates. It was found that 1 –2 mm particles were optimum, and that over a 90% lithium recovery could be obtained from the brine. However, to have a fairly pure lithium product the column needed to be drained of brine, or displaced before the water wash started. This required as much as 2.2 bed volumes of water, and resulted in about a 20% lithium loss, but this wash water could be re-treated in the column. To obtain a higher lithium concentration during

Table 1.34 Laboratory Tests on Dead Sea Brine, g/kg (Tandy and Canfy, 1993) Phosphate process Dead Sea Brine Li Na K Mg Ca Al Cl PO4 H2O a

11.6 (ppm) 31.5 6.4 35.7 14.4 — 184.1 — —

Precipitation

End-liquor 30 (ppm) 2.5 2.6 63.5 26.9 — 244.8 — 658.6

Alumina Leach liquor 680 (ppm) 0.16 — 1.25 0.29 16.0 70.8 — —

End-liquor 7 (ppm) 3.2 2.5 63.2 25.9 — 237.7 0.02 667.5

Leaching precipitate (wt.%)

Precipitate (wt.%) 0.40 10.4 1.4 2.3 3.3 — 6.5 40.0 35.7

Originala

Final

2600 (ppm) 11.9 — 2.4 2.0 — 8.1 40.8 —

170 (ppm) 0.5 — 6.9 11.4 — — 50.8 —

1440 (ppm) 4.1 — 0.05 — — 3.4 13.7 — Processing

Same as the precipitate in the previous column, except only one run, while the previous column was an average of seven runs.

Brine

143

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the elution (water wash) stage, the more dilute portions of the eluate could be used as the elution water, allowing a lithium concentration in the product of as much as 12 – 15 g/liter. However, during their runs the average product concentration was about 6 –9 g/liter, and the most effective adsorbent capacity 3 mg/g. The decrease in efficiency with repeated adsorption – elution cycles was not discussed, nor was the amount of other ions in the product, although in one run the calcium content was about equal to that of the lithium. A possible commercial use of similar alumina technology is with FMC’s hydrated alumina –lithium chloride granules suggested to be used in countercurrent adsorbent beds for Salar de Hombre Muerto brine, as discussed above. Here the lithium brine to be processed would be maintained saturated with salt as it contacts the alumina –LiCl granules, and the lithium fairly selectively adsorbed. Then the adsorbed lithium would be removed (eluted) from the granules in a second step with low-lithium water in a similar countercurrent manner. The dilute, fairly pure lithium eluate could finally be concentrated in solar ponds, and the resulting strong lithium chloride solution purified and made into the desired products (Bauman and Burba, 1997, 1995). Other Processing Methods Pan et al. (2002) have presented a general review of various methods to recover lithium from brines, and Sprinskiy (2000) made a similar review of methods to recover lithium from Carpathian groundwater. Many other adsorbents for lithium have also been suggested, such as spinel or cryptomelane-type MnO2, or antimonates of Snþ4 or Tiþ4. Abe et al. (1993) recovered lithium from seawater (at 0.17 ppm Li) with a number of metal oxide adsorbents, and found that granules of l or (l þ g)MnO2·0.18H2O, 3.1TiO2 – Sb2O5·4.9H2O and 1.1SnO2 – Sb2O5·4.9H2O could all recover up to 99% of the lithium when seawater was slowly passed through packed beds of the oxides. The adsorption preference for the manganese dioxide and tin antimonate was in the sequence of Li . Cs . Rb . K . Na, while with the titanium Cs was preferred over Li. This allowed lithium separations from sodium of 104 –105 fold, Li from K of about 1/10th that amount, and separations from Mg and Ca only about 10-fold or less. The maximum amount of lithium adsorbed was about 0.003 g Li/g of oxides, and when eluted from the column with 1 – 5 M HNO3 the best separation was with lMnO2 and a 63% recovery. The peak strength of this eluate (as ppm) was about 6 Li, 4 K and Ca, and 2.4 Na and Mg, with the average eluate being about half that value. No testing was done on the re-use or re-generation of the adsorbents, or of re-treating the eluate. A subsequent series of reports were made on similar studies with different adsorbents, perhaps culminating with the selection of H1.6Mn1.6O4 as the preferred adsorbent. It was prepared by heating LiMnO2 to 4008C to form Li1.6Mn1.6O4, and then reacting it with 0.5 M HCl. In column tests this material was capable of loading from 34 to 40 mg of Li/g of adsorbent from seawater, along with 4.1 –6.6 Na, 0.5 – 1.4 K, 2.3 – 2.5 Mg and 2.9 –4.0 Ca mg/g. The cations could be almost totally removed (eluted) by 0.5 M HCl (along with 2.5 –3.5% of the Mn), and in a second

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adsorption cycle the recovery and loading were almost the same. The recovery efficiency from the seawater was about 60% (Chitrakar et al., 2001). Umeno et al. (2002) later added the same manganese oxide adsorbent to a polyvinyl chloride polymer to prepare an adsorbent film. Using a specially designed membrane – seawater contact box the loading was 10.6 mg/g of membrane for lithium, along with 4.3 Na, 0.4 K, 10.8 Mg, 5.3 Ca and 0.5 Sr as mg/g. It was speculated that the manganese oxide was in the form of an ion sieve with a predominant pore size small enough for lithium, but not sodium, potassium or calcium. The magnesium, with about the same ionic radius has a much higher energy of hydration, and thus needs more energy to become dehydrated and enter the pore space. The larger particle size of the manganese oxide granules in the packed bed accentuated this effect, and thus rejected more magnesium. Other adsorbents that have been suggested include Li2Cr(PO4)1.67, which was claimed to react similarly to lithium – alumina, have a capacity of 9.3 mg/g in seawater, and have a concentration factor of 3.3 £ 104. It was most effective above a pH of 6.2, but could be used down to a pH of 3 (Miyai et al., 2001). Activated carbon impregnated with sodium oleate has also been suggested for seawater, along with many types of equipment to facilitate the lithium adsorption. Precipitating lithium from low-lithium brines with sodium phosphate has also been tested, after the model of licons being precipitated from Searles Lake brine. Tandy and Canfy (1993) studied the precipitation of lithium phosphate from Dead Sea potash pond end-liquor, and found that perhaps a 70% Li recovery could be obtained. By adding over a 30-fold molar excess of disodium phosphate to the lithium in the brine, adjusting the pH to 6– 7, heating to 808C, and with a 20– 30 min residence time about 76% of the lithium would be precipitated along with dicalcium phosphate and the excess disodium phosphate. The precipitate contained about 0.3% Li, and could be leached with water to recover over 90% of the Li, with the remainder being in the residual phosphate precipitate. The filtrate contained about 1440 ppm Li in a sodium phosphate –chloride solution (Table 1.34). For complex brines where solar ponding is possible, and potash recovery is also desired, a salting-out process has been suggested by Garrett and Laborde (1983). Using the brine from the Salar de Atacama as an example, the brine could be evaporated to first crystallize salt and then the bulk of the potash salts as either sylvinite or sylvinite and potassium double salts. Then the brine could be cooled in a plant to about 2 108C to crystallize about 50% of the remaining sulfate as magnesium sulfate heptahydrate (epsomite). The residual brine would be further solar evaporated and then again cooled to crystallize additional epsomite. After its removal the brine would again be solar evaporated, and then heated to about 308C (the normal summer pond temperature) and epsomite added to salt-out much of the lithium as lithium sulfate. Boric acid could be precipitated from the residual brine by adding sulfuric acid. To recover the remaining lithium lime could be added to remove the small amount of sulfate that was still present, and the brine evaporated to form a concentrated lithium chloride solution.

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Brine from the Great Salt Lake is similar to seawater, and it has also been studied for lithium recovery. In the commercial production of potassium sulfate from the Great Salt Lake the end solar pond liquors can contain 1000 –1900 ppm Li (average as ppm: Li 1500, B 360, Br 300; as %: Mg 6.9, SO4 3.6; Toomey, 1980; Tables 1.9 and 1.33). Upon further evaporation in a processing plant at about 708C to a Li value of 8000 – 9000 ppm essentially only magnesium sulfate and sodium and potassium chloride crystallize, and upon continued evaporation bischoffite and the lithium – magnesium double salts would form. Although expensive, presumably the plant evaporation of the Great Salt Lake solar pond end-liquor could form the basis of a process similar to that accomplished in the much lower humidity Salar de Atacama solar ponds. Ore Processing Greenbushes, Australia; Sons of Gwalia The Greenbushes operation of Australia (previously called Gwalia Consolidated Ltd. and then Lithium Australia Ltd.) is the largest producer of lithium mineral concentrates in the world. Production started in 1983, and from 1998 to 2002 they had 150,000 mt/yr of capacity. In 1993 they were mining . 4.0% Li2O (1.86% Li), low iron ore in an open pit with an average 1.8/1 overburden ratio (Fig. 1.76). They excavated 7000 mt/month of ore by conventional drill-and-blast techniques, employing a contractor’s personnel and mining equipment as directed by the Greenbushes’ staff. The ore was overlain by an average of 20 m (maximum 60 m) of

Figure 1.76 Side view of the Greenbushes spodumene mine in Western Australia (Flemming, 1993a; reprinted courtesy of the Australasian Institute of Mining and Metallurgy).

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weathered clay-bearing material that generally could be removed by a front-end loader or excavator. The massive deposit was also sufficiently homogenous to be easily monitored for grade-control by blasthole sampling. The benches were 5 m high, and the blasted and transported ore was first stockpiled into 8000– 10,000 mt piles (equivalent to a single blast) to allow blending to the crushers for a fairly uniform plant feed. The most harmful impurity in the ore was iron, which was partly in the spodumene molecular structure, and partly from contamination by the greenstone and dolerite in the hanging wall, and lateritic mud from the overburden and material spilled on the roads. Great care was taken to prevent iron contamination, as well as selective mining to not mix high and low grade ores in the pit (Flemming, 1993a). In the nearby plant in 1993 the 2 800 mm (31.5 in.) ore was reduced in size at a rate of 230 t/hr by being sent through a series of grizzles, screens and two stages of jaw crushers, followed by a cone crusher with a 2 12 mm (2 1/2 in.) discharge (Fig. 1.77). This equipment was used on a campaign basis with either their spodumene or tantalum (which was also mined) ore in batches of 7000– 8000 mt.

Figure 1.77 Crushing section flowsheet at the spodumene ore treatment plant of Gwalia Consolidated Limited, Greenbushes, WA (Flemming, 1993a; reprinted courtesy of The Australasian Institute of Mining and Metallurgy).

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Figure 1.78 Spodumene concentrator flowsheet at the Greenbushes mine of Gwalia Consolidated Limited, Greenbushes, WA (Flemming, 1993a; reprinted courtesy of The Australasian Institute of Mining and Metallurgy).

The 2 12 mm spodumene ore was sent to a storage bin, and then as needed to a primary screen. From there the þ 3 mm (þ 6 mesh) ore was delivered to a wet ball mill in series with a screen to be reduced to a 2 3 mm size (Fig. 1.78). The ball mill had 75 mm balls, and ran with a slurry density of 75 –80%. The 2 3 mm slurry was pumped to a hydraulic classifier where the , 250 mm (60 mesh) particles were removed in the overflow stream. The underflow was sent to a vibrating 1700 mm by 800 mm slotted screen and its , þ 820 mm (, 20 mesh) oversize returned to the ball mill. The screen underflow then went to a second vibrating screen of the same size, and its underflow was sent to rougher spirals to remove the heavy minerals (mainly cassiterite and tantalite). The cleaned classifier overflow solids were filtered and washed on a flat bed filter, the cake sent to a fluid bed dryer, and the dry product was re-screened and passed through a low-intensity magnetic separator to become “Glass Grade Spodumene”. When the raw ore contained about 4.35% Li2O (2.02% Li), the underflow from the hydraulic classifier (and the final product) contained 4.8 –5.0% Li2O (2.23 – 2.32% Li) and , 0.13% Fe2O3.

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The classifier overflow was next sent through two stages of cyclones to remove the 2 20 mm (, 400 mesh) slimes, and the coarser particles sent to flotation cells. The reagent was 700 g of fatty acids per ton of ore, the feed slurry density 35 –40% solids, the caustic conditioning time 10– 15 min, and the pH 7.1. The rougher concentrate was reprocessed (cleaned) in two stages of additional flotation, with a 45% slurry density and a pH of 6.7. An 88% lithium recovery was obtained in the flotation cells from a feed material with 3.8 Li2O (1.77% Li; this is a fairly typical spodumene flotation circuit, Manser, 1975). The flotation tailings consisted primarily of quartz and feldspar, with 0.8% Li2O (0.37% Li), and was stored for possible future recovery. The flotation overflow was vigorously washed with sulfuric acid at a pH of 1.5 –2.0 to remove the fatty acid (which discolored the spodumene and reduced the subsequent magnetic separation efficiency) and some of the iron and apatite, and then sent to low intensity magnetic separators to remove any tramp iron (from the grinding balls). Next the ore was sent to spirals (gravity separators) to remove more of the heavy minerals (which were further processed), and finally to high-intensity magnetic separators to remove the para-magnetic minerals, which were primarily tourmaline. The remaining material was passed through a 560 mm (, 28 mesh) screen, and then filtered and washed on a belt filter. The cake was dried in a fluidized bed dryer to form their “Spodumene Concentrate” product. If desired, the dust from the dryer’s flue gas cyclone collector could be sold separately as a “Fine Grade” product. The plant could treat 22 mt/hr of ore in 1993 with an overall lithium recovery of 82% when the ratio of glass grade to concentrate was 3.5/1. The plant operation was on a 24 hr, 7 days/week basis, with two operators and a supervisor (shared with the tantalum plant) on each shift. The products were shipped in bulk from the port of Bunbury, or in bags from Fremantle, and careful quality control was maintained (Harben and Edwards, 1998; Flemming, 1993b). The product could contain up to 1– 2% tourmaline (small black specks) with a 14.4% Fe2O3 content and a melting point of 11008C, but other than the iron, being a borosilicate it did not harm the glass. Quartz was the major impurity in the spodumene products, with 0.5 –1% other minerals such as Na- or K-feldspar and tourmaline (Kingsnorth, 1988). Four grades were being produced in 1993, with the main ones being concentrate containing 3.49% Li (7.5% Li2O; about 95% spodumene, 5% quartz), 0.10% Fe2O3, a maximum of 5% þ 212 mm (65 mesh), and a minimum of 60% þ 75 mm (200 mesh) particles. The glass grade contained 2.23% Li (4.8% Li2O; about 60% spodumene, 40% quartz with 0.5 – 1% other minerals), 0.13% Fe2O3, zero þ 820 mm (, 20 mesh), and a minimum of 95% þ 105 mm (150 mesh) particles. Much of their ore concentrates went into the production of container glass, television tubes and pyroceramics (Flemming, 1993a,b). A “Chemical Grade” product was initiated in 1992 that could be made with either low, or preferably high grade ore. In 1996 they commenced production of lithium carbonate from these concentrates in a small plant, but because of SQM’s drastic

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reduction in the price of lithium carbonate the next year, and technical problems with the plant it was closed in 1997 (USGS, 1997). Bernic Lake, Canada; Tanco The Tanco operation at Bernic Lake, SE Manitoba was originally (in 1929) a tin prospect because of cassiterite (up to 6.9% Sn) in the pegmatite’s outcrop, and subsequent drilling found a large rare element deposit. Over the years the deposit was more extensively drilled, and a limited amount of amblygonite was mined in 1960. A three compartment, 93 m deep shaft was sunk into the pegmatite in 1956, it was deepened to 103 m in 1959, and 1800 m of exploratory mining was done. In 1968 the shaft was again deepened to 169 m, and in 1969 a 400 m long, 20% inclined entry was constructed. From 1969 to 1982 tantalum mining was conducted, and in 1984 a 5 t/hr spodumene pilot plant program began utilizing the idle tantalum mill. Its successful conclusion led to a $6 million, 15,000 mt/yr spodumene plant being built to begin operation in 1986. Tantalum mining re-commenced in 1988, amblygonite concentrates began to be separated in 1989, and a cesium formate plant opened in 1997 (Vanstone, 2002). The mining operation in 1988 was on the Upper Intermediate Zone, approaching within about 60 m of the floor of Bernic Lake, with the three-compartment shaft and the 3.1 £ 4.3 £ 400 m incline being used for ore hauling and entry, respectively. Room and pillar open stope mining was practiced, and since the mining followed the thickness of the ore the stopes were of variable height. The first entry was made in the upper zone of the ore, and lower entries were made as needed. The principal lithium ore was “squi” (spodumene – quartz intergrowth with considerable potassium and sodium feldspar, and albite), along with occasional laths of primary spodumene and feldspar. Masses of petalite, amblygonite and lepidolite (except when mined for its tantalum content) were avoided unless it was necessary to mine through them to obtain additional squi. The initial pillars were 16 m square, and the rooms 16 m wide, but as additional rock mechanic details were established, in 1988 the standard rooms were 15 m wide with 7.7 m pillars, and on average about 20 m (10 – 30 m) high (requiring a tall roof-scaling “giraffe” (Fig. 1.79). This allowed an 89% ore recovery, with very little dilution and the minimum of waste rock development. The roofs were arched for greater strength, and seldom needed to be bolted except for some long-term entries. When bolts were used they were 3 m (10 ft) long with a pressure-expandable seal to the rock. By 2002 the room size was increased to 22 m. In 2002 two-boom hydraulic jumbos performed all of the drilling for drifts, benches and entries, while a single-boom long-hole drill was used for pillar sizereduction. The drill holes were often up to 10 m (30 ft) deep, and it was found that with this very hard rock ordinary steel drills with an abrasive fluid provided the most economical drilling. After blasting the ore was picked up and transported to various ore passes that were located throughout the mine using 3.82, 4.59 or 5.35 m3 (5, 6 or 7 yd3) load-haul-dumps (LHDs), front end loaders or 20 t trucks. The larger ore

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Figure 1.79 Roof scaler “Giraffe” in the TANCO mine (Taylor, 2002, courtesy of the Tantalum Mining Corporation of Canada Limited).

particles were broken to a 2 300 mm size on grizzlies over the chutes by mobile or stationary pneumatic or electric rock-breakers. The ore then dropped to the 130 m level where it was loaded into 4 t side dump rail cars, and hauled to the shaft loading pockets. From there it was hoisted to the surface in 4 t skips, and stored in 450 mt ore bins before being delivered to the processing plant. The mining rate was about 300 mt/day in 1988 (Vanstone et al., 2002; Burt et al., 1988). The tantalum and spodumene ores in 2002 each had their own skip pocket and surface storage bins, and their ore was hoisted daily. However, to accommodate the

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pollucite ore, one of the pockets and bins was emptied each week, and the pollucite was hoisted and processed on a campaign basis as needed. The extensive diamond drilling of the deposit allowed the mining plan to be made well in advance, and if needed all three ores (spodumene, tantalite and pollucite) could be mined at the same time. Each ore was handled in the same primary crushing equipment (reduced to a 2 12 mm size), but each type of ore had its own fines storage bin. For the mine ventilation the air entered through the shaft and a vent rise, and exited through the inclined entry at a rate of 5300 m3 (190,000 ft3) per minute. All of the mine maintenance was conducted on-site (Vanstone et al., 2002). In the plant in 1988 (Figs. 1.80 and 1.81) the spodumene ore was initially crushed in a three-stage operation, first in a jaw crusher, and then by standard and short head cone crushers, both in closed circuits with screens. The ore was crushed to a 2 12 mm size at a rate of 90 mt/hr in a single shift operation, and then stored in a 450 mt bin. From there the crushed ore was fed to a wet screen at a rate of 14 mt/hr to remove the 2 0.4 mm particles (0.5 mm in 2002). The 2 0.4 mm particles were sent to a 150 mm cyclone, with its overflow pumped to the slimes pond, and the underflow joined the heavy media underflow ore. The coarse ore from the screens next went to a two-stage heavy media separator with an effective separation density of 2.65 g/cc. The heavy media was a 70:30 mixture of ferrosilicon and magnetite with a feed density of 2.74 kg/liter. Both of the separator’s sink and float discharge streams were drained and thoroughly washed to recover the heavy media, which was thickened in a low-intensity magnetic separator for recycling (the loss was about 0.15 kg/t of ore).

Figure 1.80 Aerial view of the TANCO mine’s surface facilities (Taylor, 2002, courtesy of the Tantalum Mining Corporation of Canada Limited).

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Figure 1.81 Flow sheet for the Tanco spodumene concentration process, Ontario, Canada (after Burt et al., 1988; this figure appeared in Industrial Minerals No. 244, January 1988, p. 58. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc 2003).

The light fraction was feldspar which was stored for potential later processing, while the underflow stream and the original 2 0.4 mm fraction was filtered and sent to a second 500 mt storage bin. From there it was fed at 8.5 mt/hr to a 2.4 £ 3.7 m ball mill with 75 mm steel balls, operating in closed circuit in 2002 with two sets of 2 mm primary screens followed by a hydroseparator that only allowed about

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2 150 mm sized (100 mesh) ore to leave the ball mill circuit. Rougher and cleaner spirals also recovered tantalum from the grinding circuit streams, and a 1.5 m diameter low intensity magnet removed the iron produced during the grinding process. The 2 150 mm ore was next deslimed in two stages with 150 and 100 mm cyclones, the final overflow sent to the tailings pond, and both cyclone underflows advanced to froth flotation cells. The heavy media overflow, the cleaner spirals’ underflow and the slime streams were processed through several additional stages of density separation by spirals, tables and Falcon separators to recover more of the tantalum as a valuable by-product from the spodumene ore. The first set of rougher and cleaner flotation cells was used to remove amblygonite, since it would contaminate the spodumene with both phosphate and fluorine, and an amblygonite – spodumene mixture called Montebrasite could be sold as an extra product. The separation was made at a pH of 9.2 with limited amounts of tall oil fatty acid and petroleum sulfonate as the collector, and starch as the spodumene depressant. The floated amblygonite stream was sent through high intensity magnetic separators to remove any weakly magnetic iron minerals, and then the Montebrasite was filtered, dried and packaged for sale. The amblygonite flotation underflow solids were next re-floated in a single stage to remove mica, and thus reduce the K2O level in the final product (Vanstone et al., 2002). The underflow stream from the mica flotation was next de-starched in two stages of cyclones, and the final underflow sent to the main flotation cells to remove the feldspar and quartz. These cells operated at the same pH as the amblygonite cells, with more of the same reagents, and at a pulp density of 35 –40%. The underflow went to scavenger cells, and the spodumene overflow to cleaner cells, and then the scavenger overflow and cleaner tails were reground in a 1.8 £ 1.4 m ball mill to be returned to the rougher cells. The flotation concentrate was next passed through a low-intensity magnetic separator, and then acid washed at a pH of 1.5 to remove any iron or it’s stains, and to dissolve some of the minor minerals in the ore such as lithiophilite (a lithium phosphate compound). Finally the ore passed through a high-intensity magnetic separator to remove most of the remaining iron and any magnetic minerals such as tourmaline. The spodumene was then washed, thickened, filtered on a belt filter to an 8% moisture content, and dried in a propane-fired rotary drier to less than 0.1% moisture. The spodumene was sent to shift bins for analysis, and when approved to three 180 mt storage bins. From there the product was trucked in bulk or bags 70 km to a rail siding at Moslon, and then by rail to the main line of the Canadian Pacific railroad. Overseas shipments were made from Thunder Bay, one day to the east (Harben and Edwards, 1998; Burt et al., 1988; Crouse et al., 1984). In 1991 Tanco produced 12,000 mt/yr of low-iron lithium ore, and 160 mt of tantalum ore, and in 2001 Tanco was the largest supplier of lithium, tantalum and cesium (pollucite) minerals to the United States. In addition it sold amblygonite (containing . 7% Li2O and 20% P2O5), and potentially could sell feldspar (some containing 0.4% Li2O, 9.5% K2O and 1.5% Rb2O), lepidolite (with an even higher

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rubidium content), quartz, mica, beryl, rubidium, gallium, and more than 80 minerals ranging from bismuth to zircon – hafnium (Harben and Edwards, 1998; Kunasz, 1994; Burt et al., 1988). Zimbabwe; Bikita Minerals The famous Bikita mine has been the largest producer of lepidolite and petalite in the world, and is also one of the oldest lithium mineral producers. The mine is located 64 km NE of Masvingo, Zimbabwe (69 km east of Fort Victoria in the former Southern Rhodesia), and is in the largest (1550 m length; 29– 64 m width) of a series of pegmatites in the Bikita tin field. The mine was first opened in 1911 to recover tin, and later mined for tantalum, beryllium and lithium. Bikita Minerals first acquired some of the deposit in 1953, and added more in 1959 to form a total area of 32.5 km2. The land varies in altitude from 1097 to 1372 m, and consists of rolling hills and flat, swampy valleys, with the climate being subtropical. The average maximum and minimum temperatures are 34 and 6.78C, respectively, and the annual rainfall 678 mm. In the early lithium operations prior to 1960 a 610 m adit was used to enter underground workings in the lepidolite deposit at the 61 m level. Cars were loaded through overhead chutes, and then hauled by diesel locomotives to the beneficiation plant 450 m from the adit. The ore was crushed to about a 2 76 mm size and screened, the waste removed by hand sorting from a conveyor belt, and the waste and fines stockpiled for potential future use. The concentrates were often of almost mineral collector’s appearance, and they were trucked to the port at Beira for overseas shipment, and later delivered by the Rhodesian Railways. Petalite was mined in an open pit, and also hand sorted, but amblygonite was only obtained by selective hand mining in outcrop areas (Kesler, 1960). Lepidolite was the first lithium mineral mined in large quantities (from 1954 to 1959; primarily for the American Potash plant), followed by petalite from a 460 by 46 m zone east of the lepidolite mine, and amblygonite in smaller quantities. Mining was later conducted by open pit operations in both of the two sectors, as shown in the 1960 cross-section drawings of Fig. 1.82. An extensive drilling and underground development program had first been conducted to better define the ore body, and mining plans were developed for both predominantly lepidolite (the Bikita pit) and petalite (the Al Hayat pit). The benches in both pits were 15.2 m high for each development level, and then subdivided into working faces 7.6 m high. Slots 7.6 m wide were initially cut at right angles to the strike, and into the ore bodies to initiate the pits. The average overburden ratio was 1/1, and the overburden consisted of weathered greenstone country rock and low-grade lithium ore, feldspar, quartz and mica, all of which were stockpiled for potential future recovery. If lenses of spodumene, amblygonite, petalite, beryl or cassiterite were encountered in the overburden they were sent to the hand-picking plant for possible recovery. Once the overburden was removed selective mining was employed to recover each type of ore. Also, the lepidolite and petalite ores were mined at different periods on

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Figure 1.82 Side view of the Bikita lithium mines in Zimbabwe (Symons, 1961).

a campaign basis so that the same mining, crushing and hand sorting equipment could be used for each ore. During this period the mined ore was separated by a complex crushing-hand sorting operation (Fig. 1.83). It was hauled to the surface of the pit in trucks, and

Figure 1.83 Flow sheet for the Bikita lepidolite hand-sorting process, Zimbabwe (Symons, 1961).

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then transported to the plant (which was about 2100 m away) in narrow gauge rail cars. The first stage of size reduction was a 91 £ 64 cm jaw crusher set with a 12.7 cm ore discharge. The crushed ore was screened at a 7.6 cm size, and the oversize was washed and sent to a hand-picking belt. Quartz and feldspar were first removed, then spodumene, and finally mixed ore which was returned to the screen undersize stream. The large ore particles that remained on the belt were rescreened with the oversize sold as a glass maker’s product, and the undersize sold for chemical processing. The 2 7.6 cm ore from the screen went to a 76 £ 45.7 cm jaw crusher set to discharge at 4.45 cm, and the crushed ore was sent to a screen with 3.18 cm openings. The underflow went to a storage pile, while the overflow was washed and rescreened at 4.45 cm and 9.53 mm in a two-deck screen. The underflow went to the reject storage pile, while both oversize fractions were separately hand sorted, and ˚ ultraviolet light which after screening the product stream was passed under a 2537 A caused the eucryptite to glow salmon-pink, allowing it to be seen and removed. Both lepidolite and petalite were handled in this manner, but often the former did not require the amblygonite screening step (Symons, 1961). Production by the hand-sorting method for the period 1952 –1960, when Bikita was the world’s dominant lepidolite producer, is shown in Table 1.16. In 1979 mining was done primarily in the Al Hayat sector, producing as mt/yr: 26,400 petalite, 9600 lepidolite, 9600 spodumene and considerable feldspar (separated from the petalite). Hand sorting produced a comparatively high-grade product, but a large fraction of the ore was lost in the rejects and fines. Consequently, detailed (and successful) studies were made on both froth flotation and gravity separations, but they were not cost-competitive at Bikita until the mid-1980s when a 10,000 mt/yr heavy media separator was installed. It has allowed fine petalite to be recovered from the waste stockpile as a 4.4% Li2O product, and added 30 years to the mine’s life. The deposit’s lepidolite by 2002 was nearly depleted except for the reject piles, and in 1994 10,000 mt/yr of spodumene concentrates began to be produced using a flotation separation step (Harben and Edwards, 1998; Symons, 1961). North Carolina; Chemetall (Foote) Foote Mineral Co. purchased its North Carolina spodumene deposit in 1951, and limited mining and milling (327 mt/day capacity) in a somewhat improved plant began in 1952. The mine and processing plant had previously been operated by the Solvey Process Co. under a US Government contract from 1943 to 1946. Their ore averaged 15– 20% spodumene, and was mined from an outcropping pegmatite that appeared to be over 305 m long and up to 91 m wide. Strip mining was employed, with the top 1.8 –2.4 m of rock being considered as overburden, and the severely weathered ore beneath that somewhat selectively rejected. In the mill the ore was crushed to 2 7.6 cm in a jaw crusher, and then screened to a 2 1.9 cm size, with the oversize being further crushed in a cone crusher. The 2 1.9 cm ore was then fed to two wet pebble mills with granite blocks as the grinding media, working in series

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with a 40 mesh screen. The slurry from the screen was deslimed by a hydrocyclone and then sent to Humphrey spirals to remove about 5% of the ore as heavy particles, which were reprocessed and treated as a by-product. The remaining ore was then fed to flotation cells that did the unusual job of floating the waste rock (quartz, mica, feldspar and some hornblende). The silicate minerals collector was a fatty acid amine, the spodumene depressant was dextrine (which works best in an alkaline solution), caustic was the pH regulator (kept below 11.3), and pine oil was the frother. The spodumene underflow was dewatered and sent to storage, while the overflow which contained about 7% spodumene was sent to scavenger flotation cells, then to a gravity separation table, and the recovered spodumene added to the product. An expansion plan to raise the mine and mill capacity to 910– 1090 mt/day was being considered in 1953 (Goter et al., 1953). After an extensive drilling program from 1954 to 1956 a new mine was established, a new mill and ore processing plant was built, and production started on a much larger and more efficient scale. The operation was expanded in 1978 at a cost of $22 million, and an adjacent processing plant built with a capacity of 12 million lb/yr of LCE. The mine and plant were later de-bottlenecked to 16 million mt/yr by 1984. Ore was obtained from their open pit mine, which then operated on a cluster of eight pegmatites that were 3– 62 m thick, but required considerable selectivity in the mining operation. The pit was initially designed to have 10 benches and become 61 m (200 ft) deep, but that could later be expanded to a 122 m depth. The benches were 6.1 m high and 9 m wide since the 10 –20 m thick amphibolite and clay overburden in those pegmatites could not support very high vertical walls (Kesler, 1976). The first stage of mining was to excavate a 53 m slot across the ore, and then typically develop a pit 400 m in diameter and 61 m deep. In 1969 the pit had 11 benches, and usually three benches were mined simultaneously, with a new, deeper bench being opened as the lower ones became exhausted. The exact mining method for each pegmatite depended upon its width, but the barren rock was usually removed first on three sides to minimize contamination of the ore. The lowest side of the pegmatite was then mined first in order to allow the easier down-slope blasting. After the ore had been blasted the larger boulders were moved to one side and broken by 2273 – 4545 kg Ni-Hard “drop balls” before being hauled to the plant. Maintenance of the mining equipment was done in a nearby service building, and all of the machinery surfaces contacting the ore had to be frequently rebuilt or replaced because of the ore’s abrasiveness, which was similar to quartz in hardness (Bach et al., 1967; Johnson, 1958). In 1960 the maximum haul for the ore to the plant was 800 m, and 670 m for the overburden to the tailings pile. Primary gyratory crushers first reduced the ore to a 2 15 cm (2 6 in.) size, and then it proceeded to a series of size reduction and screening steps until its liberation size had been reached for a flotation separation (in 1960 there was also heavy media separation). Prior to flotation the ore was acid washed and then given an abrasive scrubbing to remove stains on the crystal surfaces

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Figure 1.84 Photograph of Foote’s Kings Mountain Spodumene Mine and Processing Plant (Anon., 1976; reprinted with permission from Chemical Week, October 20, 1976, Chemical Week Associates).

(from groundwater penetration) that would block the flotation reagents (Kesler, 1976). The flotation concentrate (now the cells’ overflow) was settled and shipped to their plants at Sunbright, Virginia and at Exton, Pennsylvania, while the flotation underflow was filtered, washed and discarded. A picture of the mine and processing plant during this period is shown in Fig. 1.84. In the processing plant in 1969 the spodumene concentrates were mixed as a slurry with finely ground limestone in the ratio of one part ore, 3.5 parts of limestone, and the slurry was fed into a 3 m diameter, 104 m long (10 £ 340 ft) coalfired rotary kiln. It was discharged at a temperature of 1030 – 10408C (9808C, Williams, 1976), with limestone first being calcined to form lime, and the spodumene changed from the a to the b form. The lime then reacted with the bspodumene to form dicalcium silicate and lithium oxide, plus various impurity by-products, with the mixture discharged from the kiln as about 2.54 cm pellets. This “clinker” was next cooled and ground, and leached with hot water in a six-stage countercurrent mixer –settler (thickener) system. The thickener underflow (the dicalcium silicate, etc.) was vacuum filtered, washed and discarded, while the overflow lithium hydroxide solution was pressure-filtered and then concentrated in a three-stage evaporator-crystallizer. A slurry of the crystallized lithium hydroxide monohydrate was continuously withdrawn, settled, centrifuged and dried as the final product, while the remaining liquor and filtrate were further evaporated. This allowed additional recovery of lithium hydroxide crystals, and the production of a saleable 50% solution of sodium hydroxide (Fig. 1.85; Bach et al., 1967). Foote Mineral also sold spodumene concentrates to LCA prior to their establishing their own mine, and to glass and ceramic producers in three grades: chemical, ceramic and low iron (Stinson, 1981; O’Neill et al., 1969). To produce the low-iron concentrates Foote first converted the concentrates from a to b spodumene

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Figure 1.85 A general flowsheet for producing lithium hydroxide from spodumene by the lime-roast process (Wilkomirsky, 1998).

by firing them in a kiln at about 11008C. This material was then cooled to 3008C and contacted with chlorine gas in a pressurized reactor, which formed both FeCl3 and AlCl3 vapor. The temperature had to be carefully controlled, since at higher temperatures LiCl would also vaporize, and at lower temperatures the reaction with iron would not be complete. The spodumene from the reactor was typically converted from a 0.67% Fe2O3 to a 0.075% Fe2O3 product (Heinrich et al., 1977). The recovery of lithium from spodumene ore was discontinued by Foote in 1984 (1986, USGS, 2000) and the mine and plant placed in a “stand-by” condition. It was officially closed in 1991 and the mine and plant dismantled in 1994 (USGS, 1997). However, the Kings Mountain conversion plant (converting lithium hydroxide to other lithium products) continued as a major processor of Clayton Valley and Salar de Atacama lithium carbonate into other lithium chemicals and lithium metal. Butyl lithium was produced at their New Johnsonville, Tennessee and Taiwan plants, and many other lithium chemicals were made at their parent company, Chemetall GmbH’s plant at Langelsheim, Germany. North Carolina; FMC (Lithium Corporation of America, LCA) LCA, who were later purchased by FMC, initially were located (as Metalloy Corp. until 1947) in St. Louis Park, near Minneapolis, Minnesota. They purchased

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hand-picked ore or concentrates of spodumene or petalite to produce their lithium products, and in the 1940s began to utilize some ore from their South Dakota mines. During this period they also performed extensive experimental work to determine the most favorable spodumene recovery method, and developed a patented sulfuric acid roast process that could also be used with other ores. In their Minneapolis plant concentrates were used as-is, and coarse ore was first crushed to a 2 2.5 to 2 5 cm size in a 23 £ 41 cm jaw crusher (Fig. 1.86). The ore was next fed to a 12.2 m long, 1.52 m outside diameter, 1.22 m inside diameter fire brick-lined rotary kiln revolving at about 1 rpm, with the ore flowing in a countercurrent manner to the flame and flue gas. The kiln heated 1 –2 t/hr of ore to 1050– 11008C, which converted most of the spodumene from the a to the b form. The changes in the ore from heating allowed the lithium in the spodumene to be more easily dissolved, since a-spodumene is almost insoluble in even strong acids. The hot ore then went to an adjacent 0.91 £ 7.9 m water-cooled rotary (7 rpm) cooler which discharged the ore at 95– 1208C. The ore had been decrepitated in the kiln to a 2 2.5 cm (average 8 mesh) size, and was next ground in an air-flow roll mill to a predominantly 2 200 mesh size. The fine ore then entered a 0.3 £ 3.1 m, 20 rpm screw conveyor along with a 35– 40% excess of 668 Be (93%) sulfuric acid, and was mixed and conveyed to a 0.91 £ 7.9 m co-current fired steel rotary kiln. The acidified ore was heated to 2508C, and then dumped into an 8000 gal (2.7 £ 7.9 m) wooden, air-agitated, batch operated dissolving tank. In this overall reaction from decrepitation to leaching the ore maintained its original structure (it even had very similar X-ray diffraction lines), and merely exchanged lithium for hydrogen: Li2O·Al2O3·4SiO2 ! H2O·Al2O3·4SiO2. Ground limestone was also added to the leach tank to neutralize the excess sulfuric acid, bring the pH to 6 –6.5 and precipitate much of the iron and aluminum that had dissolved. After about 30 min of agitation the slurry (in batches containing 10 mt of ore) was withdrawn, filtered and washed on a vacuum drum filter. The solids with 30% moisture and , 1% of the initial lithium were discarded, and the filtrate was treated with lime to a pH of 12. Then a small amount of soda ash was added to remove the remaining calcium, magnesium, iron and aluminum to produce, after filtration, a fairly pure 100 gpl Li2SO4 solution. This solution was adjusted to a pH of 7, evaporated to 200 gpl, activated carbon was added to remove organics, and the mixture filtered. Finally, the solution was maintained at 908C, soda ash was added to precipitate lithium carbonate, and the slurry was centrifuged. The lithium carbonate cake was then washed and dried to yield about 6000 lb/day of product. The filtrate was cooled to crystallize glauber salt (Na2SO4·10H2O; which was made into a salt cake [Na2SO4] by-product), and the remaining solution recycled to the evaporator (Fig. 1.86; Ellestad and Clarke, 1955; Hader et al., 1951). The advantages claimed for the process compared to an alkali or other leach were that: (1) Fine grinding of the ore was not necessary, since b spodumene is porous and even coarse particles react fairly rapidly with sulfuric acid (2 200 mesh ore is leached in 3– 15 min). (2) The kiln operations only required a short residence time

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Figure 1.86 The initial flowsheet for LCA’s production of lithium carbonate from spodumene by the acid roast process (after Hader et al., 1951).

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since both reactions occurred rapidly. (3) There was no mixing problem with the ore and sulfuric acid, and the subsequent water-leaching was rapid. (4) They felt that the process was the only one amenable to treating low-grade lithium ores directly, without first producing concentrates. In their experience with weathered (outcrop) South Dakota ore the flotation efficiency to form concentrates was only about 45– 50%, and the leaching processes about 85% efficient. By contrast, the acid-roasting process could give an 80% yield on 0.7% Li spodumene ore directly, instead of their estimated overall yield of 38 –43% with South Dakota concentrates. In 1955 they planned on using it on run-of-mine ore in their Minnesota and new Bessemer City, North Carolina plants (Ellestad and Clarke, 1955; Ellestad and Leute, 1950). The Bessemer City plant was thus designed to handle any grade of ore, including ore directly from the mine, and its initial operation in 1956 was with as-mined 1.0– 1.5% Li2O (0.46 –0.70% Li) ore. However, the process in the plant did not run as smoothly as in the pilot plant tests, and many changes had to be made. The conversion of the ore to b spodumene was not complete below about 11008C, but the impurities in the ore would fuse at temperatures above about 11608C and considerable lithium would become insoluble. This required very careful operational control of the feed rate and firing temperature. Also, the large amount of fines in the ore caused very high dust losses with the flue gas in the kiln, and the air flowing through the cooler following the kiln. Consequently, an indirect 3 £ 61 m rotary water cooled unit was used with a more modest air flow, as well as a fan and dust collector so that the cooling air could be used in the kiln’s burner. This also further complicated the kiln and cooler control, and a two-stage scrubber was needed for the kiln’s flue gas. The unprocessed ore required considerably more sulfuric acid in the acid roaster, which in turn required that the leach tank be lead and acid brick lined, and the agitator and pump be neoprene lined. In the new plant, as a process improvement a slurry delivery system was installed for the soda ash instead of the air-slide conveyors used for most of the other solid phase handling. The run-of-mine ore process was utilized for 11 months, but at the end of that time it was felt that the plant capacity could be increased by 180% when operating on concentrates, so that became the permanent feed material, with very little loss in overall recovery (Andrews, 1958). Obviously their South Dakota 45– 50% flotation yield should have been greatly increased in North Carolina, and it is very likely that the low grade ore added impurities to the product. In the mid 1960s LCA developed their Cherryville open pit spodumene mine, located about 10 km NE of the Foote Kings Mountain deposit (Figs. 1.35 and 1.36). The spodumene contained about 0.7% Li (1.5% Li2O), and was first ground to a small particle size and concentrated by flotation (initially by heavy media) to about 3% Li (6.5% Li2O; initially 3.73 or 8.0% Li2O). After being thickened and filtered the concentrate was sent to the nearby Bessemer City plant with an improved process (Fig. 1.87) and roasted in an 82 m (250 ft) long rotary kiln to 1075– 11508C. This changed the molecular structure of the spodumene from the a to the b form, and increased its surface area by 30%. The converted ore was discharged from the

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Figure 1.87 Flowsheet for the acid-roast conversion of spodumene into lithium carbonate (Wilkomirsky, 1998).

kiln, cooled and ground to a 2 100 mesh size in a ball mill, and mixed with a slight excess (over the stoichiometric amount of lithium) of 93% sulfuric acid. The acid mixture was then heated in a rotary kiln to 200 –2508C (250 –3008C, Bach et al., 1967), and the discharged material leached with water to form a fairly pure lithium sulfate solution. The slurry was neutralized with ground limestone, settled, and the underflow filtered in rotary vacuum filters, washed and discharged. The thickener overflow and filtrate were treated with small amounts of lime and then soda ash to remove the calcium and magnesium, and settled and re-filtered (Howling, 1963). This purified solution was neutralized with sulfuric acid to a pH of 7– 8, and then concentrated to 200 – 250 g/liter in a five-effect evaporator. The strong solution was treated to remove alumina, again filtered, and then a strong soda ash solution added to precipitate lithium carbonate at 90 –1008C. The solids were settled, centrifuged, washed and dried as the main product. The remaining solution (thickener overflow and centrate) still contained about 15% of the original lithium, so it was cooled to 08C to crystallize (as the decahydrate) the major impurity, sodium sulfate, and then recycled to the ore leach system. The sodium sulfate decahydrate was centrifuged and washed, and then evaporated to yield sodium sulfate as a by-product. The lithium carbonate was either sold as is, or formed into other products, such as being reacted with hydrated lime to produce lithium hydroxide (Stinson, 1981; Bach et al., 1967; see the Roast, Acid Leach section below for more details on this process). The processing of ore was discontinued in 1998, following a 3 year $18 million expansion to 36 million lb/yr LCE in 1981, but the Bessemer City plant adjacent to the mine remained as FMC’s major production facility for other lithium chemicals, including the metal and organo-lithium compounds. Some lithium chemicals are

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made at Bayport, Texas, butyl lithium and lithium metal are also produced at Bromborough, Mersyside, UK (Lithium Corp. of Europe), and they had a 100 mt/yr battery grade lithium metal operation for lithium batteries as a joint venture (Asia Lithium Corp.) at Kagawa, Japan. In 1981 they produced over 70 lithium compounds at their Bessemer City plant (Lloyd, 1981), and in 1996 $30 million was spent to modernize and further expand the plant (USGS, 1997). American Lithium Chemicals; Bikita Lepidolite In the mid-1950s the government contracted with the three lithium producers (LCA, Foote and the American Potash & Chemical Co.) for the purchase of lithium hydroxide monohydrate in order to extract most of its 6Li content. Since American Potash did not have an adequate supply from its brine processing operation (discussed above), they constructed a new plant near San Antonio, Texas, and from 1955 to 1960 imported lepidolite ore from Bikita, Southern Rhodesia (now Zimbabwe) as their raw material. One part of the ore was mixed with three parts of local limestone (one of the reasons for the plant’s location) and the mixture wet ballmilled to a 2 200 mesh size. The slurry was filtered and the wet filter cake then heated in a 3.66 m (12 ft) diameter, 99 m (325 ft) long rotary kiln to about 9118C. The hot discharge was quenched, wet ball-milled again, and leached in a countercurrent mixer –settler system. The underflow slurry was filtered, washed and the solids discarded, while the overflow liquor and filtrate were treated to remove aluminum, pressure filtered and sent to a triple-effect evaporator. In it an impure lithium hydroxide monohydrate was crystallized, which was removed, centrifuged and then dissolved and recrystallized into a pure product. The remaining liquors were further processed and evaporated to first recover more lithium, and the residual concentrated solution sent to storage tanks. After the contract period ended a new company, San Antonio Chemicals was formed to treat the end-liquors from the process by evaporating them much further to recover saleable caustic soda and an alkali crystalline mixture that contained, after reacting with soda ash: 70% K2CO3, 23% RbCO3, 3% Na2CO3, 2% Cs2CO3 and 1% Li2CO3. Plans had been made to separate the mixture into pure products, but the operation was not deemed to be profitable, so the by-product was sold as a mixture and the plant closed (Bach et al., 1967; Symons, 1961). Quebec Lithium Corp Spodumene had been produced from the Preissac-Lacorne deposit by the Quebec Lithium Corp. from 1955 to 1959, supplying ore to LCA’s North Carolina plant. Mining was conducted underground, with the initial area developed about 450 m in diameter from a 171 m deep shaft. After mining, the ore was concentrated by flotation (Kesler, 1960) in a mill near the shaft site with a capacity of 2000 mt/day of concentrates. This operation closed when LCA opened their own mine in North Carolina (Kunasz, 1994).

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However, in 1960 mining started again to supply ore for the production of lithium carbonate or hydroxide, using the Quebec Government’s Archambault –Oliver (1963) spodumene process. In it the ore was first ground to a small size and heated to about 10008C to convert the spodumene to the b form. After cooling the ore was ground to a finer size, slurried with water and a recycle solution, as well as soda ash in slight excess (such as 10 –30%) over the lithium content of the ore. The slurry (9 – 50 wt.%) was then pumped into an agitated pressure vessel at about 140 –3008C (preferably 185– 2508C; 50 –600 psi) with an ore retention time of about 1 hr. This leached the lithium and converted it to lithium carbonate, while the spodumene remained as a zeolite-type insoluble aluminosilicate. The slurry from the pressure vessel and recycle lithium carbonate precipitation filtrate then went to a carbonation tower operating at 208C, where carbon dioxide converted the insoluble lithium carbonate and lithium silicate into soluble lithium bicarbonate (0.65 – 1.8% Li greater than the concentration in the recycle solution). Sodium silicate was also converted into silica, and any soluble aluminum or iron silicates were precipitated. The slurry from the carbonation tower was next thickened and filtered, and the solids washed and discarded (they were claimed to be saleable as a zeolite). The overflow and filtrate were then heated to about 958C which converted much of the bicarbonate to carbonate, and precipitated the desired lithium carbonate product. The remaining solution and the carbon dioxide formed from the bicarbonate were both recycled to the carbonation (leach) tower. Alternately, the slurry from the bicarbonate tower could be reacted with lime to produce a lithium hydroxide solution and insoluble calcium silicate. After filtering the solids the solution could be evaporated and lithium hydroxide monohydrate crystallized (Archambault and Olivier, 1963). The process operated commercially for several years, and the production was increased fivefold since it started in 1961 to reach an annual rate of 2 million lb/yr of lithium carbonate. However, after the US Government’s purchase of lithium ended the plant could only run at reduced levels because of the intense competition, and closed in 1965 (Flanagan, 1978; Anon., 1967). Other Operations Black Hills, South Dakota In the early days of the lithium industry in the United States the most heavily developed mining area was in the Black Hills of South Dakota, where the Eta mine was the largest and most long-lived operation. However, there were also many smaller deposits such as those operated by LCA from 1941 to 1953, where they developed mining and mineral processing techniques that were useful for their subsequent North Carolina deposit. They used both open pit and underground largescale, mechanized mining methods at their Edison, Mateen and Longview-Beecher deposits, with each being a cluster of small outcropping spodumene pegmatites. To process the ore they first used a 12 t/hr heavy media separator with the results shown in Table 1.35, along with a list of the densities of the minerals involved. The yield was reduced because some of the ore was partially altered, which lowered its

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Results of LCA’s Heavy Media Separation of Spodumene Ore from the Black Hills, SD (Munson and Clarke, 1955) Product Sink Float Fines Composite Spodumene Quartz Microline Albite Muscovite Apatite Tourmaline Triphylite

Weight, Pct 7.1 66.5 26.4 100.0

Lithia, Pct

Distribution, Pct

5.36 0.16 1.19 0.80

47.4 13.4 39.2 100.0

Specific gravity, g/cc 3.1 2.65 2.56 2.60 2.76– 3.1 3.2 3.0–3.2 3.4–3.56

Courtesy of Mining Engineering Magazine; reprinted with permission of Mining Engineering Magazine.

density. Also, only 2 3.8 cm to 6 mesh ore was used because the acicular cleavage of the spodumene made the smaller particles tend to float. Flotation was next tested, using preliminary desliming, caustic washing, and an anionic fatty acid collector (Fig. 1.88). Standard equipment was used, except for changes caused by the ore’s rapid settling and abrasiveness. The results are shown in Table 1.36, and in addition to the spodumene clean mica and feldspar could also be produced by subsequent flotation steps with cationic collectors (Munson and Clarke, 1955). China’s Yichun Li – Ta –Nb mine in 1998 accounted for 90% of the country’s recoverable lithium reserves, and its lepidolite was easily obtained as concentrates from the open pit mine’s tantalum and niobium processing. However, initially there was only a small amount of lithium carbonate processed from this ore due to the high cost of the lime sintering process (see the American Potash and Chemical Co. Section, above). To reduce these costs a pressurized ammonium chloride leach process has been suggested by Xu et al. (1998). In this process lepidolite concentrates (Table 1.18) would be initially partially defluorinated by being heated to 8508C for 20 min, and then ground to a 2 74 mm (, 200 mesh) size. The roasted concentrates would be cooled and made into a 25% aqueous slurry, with 3.5 mol of ammonium chloride being present per mole of total alkaline solids in the concentrate. The slurry would be heated under pressure at 2408C for 90 min, and then cooled, filtered and washed. The process was estimated to leach about 95% of the lithium, but the filtrate would also contain most of the other alkali metals in the ore. The filtrate would consequently be evaporated to crystallize the sodium, potassium,

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Figure 1.88 LCA’s spodumene flotation process for Black Hills, South Dakota ore (Munson and Clarke, 1955, courtesy of Mining Engineering Magazine).

rubidium and cesium chlorides, and after they were separated, the high-lithium endliquor would be treated with lime and a small amount of soda ash to further remove impurities. Finally, a lithium carbonate product would be precipitated with ammonium carbonate or soda ash, and the filtrate recycled to the pressure reactor. Lithium ore has been produced in Brazil since the early 1960s, with the initial mining company being Arquena de Minerios e Metals Ltd. They mined a number of lithium minerals (spodumene, petalite and lepidolite) in Aracuai and Itinga, and also supplied amblygonite initially, and then spodumene to Cia Brasileira de Litio. The latter company had built a plant at Aguas Verelhas, Minas Gerais province, and in 1991 produced about 1000 mt/yr of lithium hydroxide and 200 mt/yr of lithium carbonate, with some government assistance (US Bur. Mines, 1992; Afgouni and Silva Sa, 1978). By 2000 Brasileira de Litio was producing its own spodumene concentrates from their underground Cachoeira mine in Aracuai (USGS, 2001), and in 2002 Metallurg was planning to produce concentrates of a lithium-bearing feldspar for domestic use (Tamlin et al., 2002). In Portugal the major lithium mining operation

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Table 1.36 Data from LCA’s Spodumene Flotation Process with Black Hills, SD Ore (Munson and Clarke, 1955) Product

Weight, Pct

Lithia, Pct

Distribution

No. 1 Partly Altered Spodumene Pegmatite Concentrate Slime Tail sand Composite Flotation efficiency,a 75.8 pct

15.9 31.8 52.3 100.0

4.92 0.68 0.46 1.26

63.5 17.2 19.3 100.0

No. 2 Altered Spodumene Pegmatite Concentrate Slime Tail sand Composite Flotation efficiency,a 76.0 pct

13.1 38.3 48.6 100.0

5.13 0.75 0.46 1.21

57.1 24.1 18.8 100.0

No. 3 Hard Rock with Altered Spodumene Concentrate Slime Tail sand Composite Flotation efficiency,a 76.0 pct

14.2 20.4 65.5 100.0

3.94 0.53 0.27 0.82

66.4 12.8 21.0 100.0

Courtesy of Mining Engineering Magazine. Lithia recovery from deslimed flotation feed.

a

in 2002 was the Soc. Minera de Pegmatite who sold unprocessed lepidolite ore. In Namibia the mining company in 1996 was Intermetmin Ltd who sold petalite and some lepidolite concentrates from their Rubicon mine (Harben and Edwards, 1998). In mid-1998 the company suddenly ceased operating after there had been a fairly rapid succession of owners and a $2.6 million government loan (USGS, 2000). North Korea has produced limited quantities of low-grade spodumene (Tamlin et al., 2002). Various Proposed Processes A detailed laboratory study has been presented by Dresler et al. (1998) on the Roast, Acid-Leach Process that exactly follows the process previously used by the LCA (now FMC) at their North Carolina deposit. It is reviewed here because of the additional details that it supplies on this very effective spodumene process. Their ore was from Wekusko Lake, Manitoba, which as at LCA’s deposit was an unzoned and low-grade (0.79% Li) spodumene ore, but with a very small crystal size. Electron micrographs showed that each of the minerals in the ore (Table 1.18) was present as discrete crystals, but the liberation size was quite small. Consequently, they ground the ore to a 2 212 mm (, 65 mesh) size, and made into a 23% slurry to be agitated and conditioned with 2 kg of sodium hydroxide/mt of ore for 20 min.

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The Lithium Recovery and Li2O Content with Rougher and Successive Stages of Cleaner Flotation (Dresler et al., 1998) Rougher Li2O (wt.%) Li recovery (%) Al Ca Fe K Mg Na S Si P a

4.40 96.9 Li2CO3 product impurities, ppma 24 1630 15 9.9 357 1007 1323 209 15

Cleaner stages 5.87

6.59

6.80

87.97

76.36

70.36

7.05 56.8

Starting with 6.6% Li2O, 0.9% Fe, 0.2% Na and 0.15% K spodumene concentrates (from a 76% flotation recovery), roasted at 700–12008C, 50% excess of sulfuric acid added and roasted at 2508C for 15 min (giving a 97% Li recovery after leaching). The product was precipitated with soda ash at a pH of 6.5–11 and 908C, producing 98% Li2CO3 (Wekusko Lake, Man., Canada spodumene).

This aided in the desliming of the 2 15 mm particles, and resulted in a 4 –5% Li loss. The ore was then filtered and washed to remove the NaOH, and further conditioned with 400 g/mt of oleic acid for 3 min. Next, it was again made into a 23% slurry, the pH adjusted to 6.8 with sulfuric acid, Dowfroth 250 added and sent to a rougher flotation cell. The rougher product contained 4.40% Li2O (2.04% Li) and had a 96.9% yield, but needed to be cleaned in 2 –4 successive flotation steps to reach an acceptable grade, reducing the recovery to 76% (without recycling, regrinding and scavenging flotation steps; Table 1.37). The final concentrate in these tests averaged 6.6% Li2O (3.07% Li), 0.9% Fe2O3 (originally 0.94 – 1.64%), 0.2% Na and 0.15% K. After the concentrate was filtered, washed and dried it was heated to 11008C for 15 min to form b-spodumene. The roasted ore was then cooled, mildly crushed, a 50% excess of 93% sulfuric acid added and the mixture heated to 2508C for 15 min. The reacted mass was next cooled and then leached with water at room temperature. A 97% yield of lithium in the b-spodumene was obtained, and lime was added to the leach solution to neutralize the excess sulfuric acid, raise the pH to 6.5 and precipitate most of the iron, aluminum and sulfate. The slurry was filtered and the pH raised to 12.0 by the addition of a small amount of soda ash to precipitate the calcium. This mixture was filtered, some activated carbon added to remove organics, the solution was refiltered and the pH reduced to 7.0 with sulfuric acid. The clear solution was then evaporated to 200 g/liter lithium sulfate concentration. Finally, soda ash was added to the strong, hot solution to precipitate lithium carbonate. When filtered, washed and dried analysis showed . 98% Li2CO3, with the impurities listed

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in Table 1.37. The lower purity than LCA’s product previously made by the same process may have been due to LCA’s further process optimization, or the weathered ore (from an outcrop) used in these experiments. Another of the established lithium ore processing methods was the Roast, Lime or Limestone Leach Process, combining an initial roast with limestone followed by water leaching, or roasting and then leaching with lime. There are a large number of patents and articles on the process, and it was commercially practiced by Foote, American Potash (both are discussed above) and others. Again, it is being further reviewed here to provide additional processing details. One of the early patents was by Nicholson (1946) who suggested grinding spodumene ore to 2 200 mesh, mixing it with ground limestone and roasting the mixture at 11208C. The cooled ore would then be reground to 2 100 mesh and leached with water at 1008C. The roasting temperature was claimed to be low enough to prevent much of the silica and aluminum from reacting with the lime, an excess of limestone improved the lithium leach, and the process resulted in 80% lithium yield. Alternately, the ore could be ground to 2 100 to 2 200 mesh, roasted at 1100– 11508C to form b-spodumene, and then reacted with lime under 15 – 250 psi pressure at 194 – 2048C. After filtration, a 1– 4% LiOH solution would be obtained that contained essentially no silica or aluminum. In either case the solution could be purified and lithium hydroxide crystallized, or soda ash added or the solution carbonated with carbon dioxide to form lithium carbonate, as desired. Research on the Big Whopper petalite deposit in the Separation Rapids area of Ontario, Canada has reported a patented process that can produce . 4.7% and 4.0% Li2O petalite products. It also separates 11.5 –12% K2O with $ 1.0% Rb2O K-feldspar and $ 10% Na2O Na-feldspar, as well as concentrates of mica, spodumene, tantalum, cassiterite, garnet and perhaps silica. After crushing, grinding and desliming the spodumene would first be floated with fatty acids, and then petalite with a fatty amine and proprietary reagents. K-feldspar would next be floated, followed by Na-feldspar, and the heavy metals would be recovered by gravity separation at several points in the process (Pearse and Taylor, 2001). A large number of other processes have been suggested to recover lithium from its ores (over 60 US patents on this subject were issued between 1900 and 1965). In the early days of the industry high-grade ores of lepidolite, amblygonite and zinnwaldite (a lithium – iron mica) were merely heated with sulfuric or hydrochloric acid, and the mixture then leached with water, but spodumene could not be leached in this manner. However, the leached solution contained most of the metals in the ore (particularly iron and aluminum), and this required a complex purification process (e.g., Vyas et al., 1975; Gauguin et al., 1961; Siegens and Roder, 1934). Later there were base-exchange processes such as roasting finely ground ore, or reacting slurries at high temperatures (100 –3008C) with an excess of potassium, sodium, calcium or magnesium sulfate (as pressurized solutions or melts) to form a leachable residue. Lithium sulfate

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could then be crystallized and the reactant-salt purified, crystallized and recycled (e.g., Lindblad et al., 1943). This again was an expensive and difficult process. Several other process suggestions were based on adding calcium sulfate or calcium chloride to the ore and then roasting it with limestone or lime. Lithium hydroxide with lithium sulfate or chloride would be leached from the roasted mixture and then the separate components crystallized (e.g., Vyas et al., 1975 who roasted lepidolite with limestone and calcium chloride at 9508C. A water leach gave an 80% yield of 4 –8% LiCl). In one version of this type of reaction the use of high enough temperatures to volatilize lithium chloride was suggested, and even operated on a small scale by Solvey in North Carolina before Foote purchased their deposit (Bach et al., 1967; Ellestad and Clarke, 1955). As a more novel process Goodenough and Stenger (1958) have suggested contacting very finely divided bspodumene (roasted spodumene ore) with much larger sized strong-cation ion exchange resins (in the hydrogen form) at about 1008C to leach lithium. The resin is then separated from the ore on screens, and lithium recovered by contacting (eluting) the resin with a strong acid. This also regenerates the resin for reuse. In the early days of the lithium industry considerable attention was paid to the recovery of lithium from moderately high-lithium Clay. Lien (1985) noted that in laboratory tests some clays could have as high as an 80% lithium extraction with a simple sulfuric acid leach, but that most required a more complex process. In brief tests a roast at 7508C with two parts of clay and one part limestone, followed by a leach with an excess of 20% hydrochloric acid gave a 70% lithium yield. In a second series of tests five parts of clay, three parts of gypsum and three parts of limestone were roasted at 9008C. A water leach resulted in an 80% recovery of lithium as lithium sulfate. In the later process the raw materials were first ground together to a 2 100 mesh size and then formed into 6.5 mm pellets before being roasted. The pellets reduced the dust loss and increased the particles’ contact with the flue gas. The roasted pellets were next ground to a 2 100 mesh size and leached with water in an agitated container. The leach liquor contained 2.5 –3 g/liter of lithium, considerable sodium and potassium sulfate, some gypsum (0.6 g/liter Ca) and other impurities. Soda ash (in recycle liquor) was added to the leach liquor, and the calcium carbonate that formed was removed by filtration. Then the solution was evaporated nearly to the sodium and potassium sulfate crystallization point (9 – 10 g/liter Li, 120 g/liter of both Na2SO4 and K2SO4), and the near-boiling solution reacted with soda ash to precipitate lithium carbonate. The remaining liquor still contained 4.0 – 4.5 g/liter of lithium carbonate (, 1900 –2100 ppm Li), and was then cooled to 0 to 2 38C to crystallize glauber salt and K2SO4, leaving 70 g/liter Na2SO4 and 105 g/liter K2SO4 in solution. After the crystallized salts were removed, the solution could have been further evaporated to crystallize more potassium sulfate, and the remaining solution recycled to the leach liquor evaporator (Fig. 1.89). As a second example of processing clay, Amer and Rashed (2002) processed 2 40 mm El-Fayoum (Egypt) bentonite in an autoclave at 2508C for 90 min with about 50%

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Figure 1.89 Flow sheet for the production of lithium carbonate from clay (Lien, 1985; reprinted form Lithium, Ed. R. O. Bach, Sec. 6, q 1985 by permission of John Wiley & Sons, Inc.).

sulfuric acid. They had fairly good lithium recoveries, but there was also considerable iron, aluminum and other impurities in the leach solution. The low grade of the clays (, 0.60% Li) and the complex processes would make them noncompetitive with lithium brines or ore concentrates. Lithium has also been recovered from obsolete lithium-containing batteries. In 1997 the US Navy awarded a $10.5 million 4 year contract to ToxCo, Inc. to recycle 4400 lithium batteries of 1200 kg that had been used as a back-up power source for the now-abandoned missile silos. Their subsidiary LithChem International had developed a grinding-hand sorting process in 1994 for recycling the lithium, aluminum, nickel and stainless steel, and the batteries began to be processed in a plant at Trail, British Columbia, Canada in 1998. Originally, the lithium was to be sent to Pacific Lithium Ltd in New Zealand to be converted into high-grade lithium carbonate (USGS, 1997), but instead they converted it themselves into lithium carbonate and hydroxide in a plant at Baltimore, Ohio. In 1996 they had also purchased excess lithium hydroxide (containing mostly 7Li) from the US Government, to be sold to the lithium grease market. In 1999 they purchased a fluorine products company, and began making lithium fluoride, lithium hexafluorophosphate and other battery electrolytes (USGS, 2000). Many other proposals have been made to recover lithium from various sources, such as small lithium batteries. For instance, Tanaka and Shimamune, 2003 suggest dissolving the electrode materials, filtering, adjusting the pH to 7– 10 to precipitate cobalt, and then recovering lithium. Alternately they propose electro-depositing the

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dissolved cobalt, and then forming lithium carbonate after the remaining solution has been purified. An alternative procedure was suggested by Lee and Yang (2001) that involved heating, shredding, calcining, leaching and lithium precipitation. The batteries were heated to 5008C and then shredded, followed by the metal being removed and the remaining powder heated to 8008C for 2 hr. It would next be leached with 2 M sulfuric acid (or 1 M HNO3) containing 20% excess hydrogen peroxide as a reducing agent. More than 95% of both the cobalt and lithium were claimed to be solubilized, and could then be separately precipitated. To precipitate LiCoO2 the Li/Co ratio in the solution was adjusted to 1.1, and then 1 M citric acid was added to form a lithium cobalt gel. It was removed and calcined at 9508C for 24 hr to produce a product with a surface area of 30 cm2/liter and a particle size of 20 mm. Various recovery methods have also been proposed to recover the lithium used in other batteries, polymerization reactions, lithium zeolites, molten radioactive salt wastes and mixed lithium compounds, among others. Lithium Chemicals A very wide variety of lithium chemicals are sold commercially, as illustrated by Chemetall’s “Product Line” in 2002 listing 33 inorganic lithium chemicals and 36 organic lithium compounds, with others available upon request. Their major production facility for these chemicals is in Langelsheim, Germany (Fig. 1.90), and many of the inorganic chemicals are produced from lithium carbonate as the initial

Figure 1.90 Aerial view of Chemetall’s Langelsheim, Germany Lithium Chemicals Plant (courtesy of Chemetall GmbH).

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lithium source. However, even with this wide diversity, in 1993 about 50% of the total lithium usage was as (unconverted) lithium carbonate, and about 20% as lithium hydroxide. To produce lithium hydroxide either moist lithium carbonate filter cake or the dry product was mixed with about a 5% excess of lime in sufficient recycle wash and make-up water to form a 0.3 lb/gal LiOH solution. The reaction was conducted at near the boiling point in an agitated tank, and the final slurry was then settled. The calcium carbonate solids that were formed were washed in a three-stage countercurrent decantation system, and either recalcined to form the lime used in the hydroxide reaction or discarded. The wash water was returned to the reactor, and the strong overflow liquor filtered and evaporated to a strength of about 1.39 lb/gal at 1008C to be sold as a liquid product, or evaporated further to continuously crystallize lithium hydroxide monohydrate. The concentrated liquor could also be cooled in a crystallizer to about 408C to form the lithium hydroxide monohydrate. In either case the lithium hydroxide would be centrifuged and dried if an impure grade were desired, or redissolved at near the boiling point and recrystallized to form a purified product. In the latter case small amounts of lime and soda ash, followed by activated carbon would be added to remove impurities, and the slurry then filtered. The purified solution would be recrystallized, centrifuged and then dried (Hader et al., 1951). To produce anhydrous lithium hydroxide the monohydrate could be dehydrated in an inert gas (because of its strong tendency to react with carbon dioxide) dryer, possibly operating under vacuum, and often this product would be pelletized (Fig. 1.91). It has also been suggested by Bruhn et al. (1998) that lithium

Figure 1.91 A general flowsheet for converting lithium carbonate into lithium hydroxide monohydrate (Wilkomirsky, 1998).

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Part 1 Lithium

hydroxide could be made by the electrolysis of lithium chloride solution in a membrane cell, similar to the production of sodium hydroxide. Lithium chloride is made in a similar manner, starting with either lithium carbonate as wet filter cake or the dry product. It is reacted with a slight excess of hydrochloric acid (usually 31% HCl) in a rubber lined reactor with a rubber lined agitator, and sufficient reactants are added to bring the solution to a density of 1.180 – 1.195, or about 3 lb of lithium chloride/gal. Carbon dioxide that is formed is vented from the top of the reactor, perhaps aided by an air stream to control the vigor of its evolution. A small amount of barium chloride is then added to precipitate any sulfate that is present, sulfuric acid is added to remove the excess barium, and the solution is neutralized with lithium carbonate. This slurry is filtered, and the clear solution then evaporated to about 40% LiCl to make a saleable liquid product. To produce a dry product, since the solution’s boiling point and solubility are so high, it must be further concentrated by direct contact with flue gas or in a directly fired pot. In the former case, a ceramic lined packed tower can be used with the solution flowing downward, and the flue gas rising to concentrate the solution to near its solidification point. It can then flow to a direct fired ceramic lined dryer to be solidified, with the exit flue gas going to the tower, or it can more easily be solidified on a chilled roll and then sent to a dryer. In either case the lithium chloride that is formed must finally be ground in a hammer mill and screened to the proper size (usually in sizes below 8 mesh), and packaged in air-tight containers (Fig. 1.92). Since lithium chloride is acidic and very corrosive, its solutions are usually neutralized with lithium hydroxide before being shipped or solidified. In both forms it is also very hygroscopic (Hader et al., 1951). Also, as previously discussed lithium chloride may be produced directly from either lithium ores or brines (Brown and Beckerman, 1990; Stenger, 1950). Lithium bromide is made in a very similar manner to lithium chloride, except it is primarily sold as a 54– 55% solution. Lithium carbonate is reacted with hydrobromic acid (usually as 45% HBr), with the pH of this solution adjusted to . 7. It can also be produced by reacting lithium carbonate or hydroxide with bromine and a reducing agent. With either reaction, lithium hydroxide is then added until the solution has a 0.01 N hydroxide content. This solution is next evaporated to the 54 –55% LiBr concentration, and then shipped in 55 gal steel drums. Most of the product is used in air conditioning systems, but a small amount of lithium bromide is also solidified in a manner similar to lithium chloride to be used in organic chemical reactions. Lithium hydride is produced from lithium metal, with the metal first placed under high vacuum (, 1 mm pressure) to remove most of the oxygen, nitrogen and moisture. It is then melted, and hydrogen slowly admitted until a pressure of about 5 psi is reached. The reaction is highly exothermic, so no additional heat is required, and the reaction must be carefully controlled. When the reaction is complete, lithium hydride is removed (sometimes dumped from a tapered wall reactor) and ground to size. Lithium amide is made by taking the crushed lithium hydride and heating it in an oven containing only ammonia. As the amide is formed hydrogen is evolved, and it is burned as it leaves a vent in the reactor. The amide is widely used in the

Processing

177

Figure 1.92 A typical flowsheet for the production of lithium chloride and lithium metal (after Hader et al., 1951).

pharmaceutical industry (Hader et al., 1951). Lithium aluminum hydride (LiAlH4) is produced by the reaction of lithium hydride with anhydrous aluminum chloride in dry diethyl ether. The solvent can be boiled-off to produce a dry powder, and the powder can be pelletized. Alternately, it can be sold as a solution, such as 20% in diethyl ether, 10% in tetrahydrofuran or 15% in tetrahydrofuran/toluene. It is used as a versatile reducing agent in many organic reactions (Deberitz, 1993). Lithium alkyls such as n-butyl lithium can be prepared by several methods, with one being the reaction of the desired alkyl (or aryl) halide (usually the chloride) with finely dispersed lithium metal in a hydrocarbon solvent (Deberitz, 1993). There is some demand for purified lithium carbonate (99.999%), and in 1999 it was produced by various companies, including Lithium Metals Technologies, Inc. (Limtech) at the rate of 300 mt/yr. In 2000 they announced a proposed expansion to 1000 mt/yr (USGS, 2000).

178

Part 1 Lithium

Figure 1.93 An example of an electrolytic cell used to produce metallic lithium (Averill and Olson, 1977; reprinted from Energy, Vol. 3, No. 3,q1978 with permission of Elsevier).

Lithium metal can be produced by the electrolysis of a molten lithium chloride – potassium chloride mixture, such as 45% LiCl/55% KCl (the range is 40 – 60% LiCl). At LCA the reduction has been conducted at 4608C, although 4208C was noted for Chemetall, by Deberitz (1993), and up to 5008C has been stated by others. In the simplest form of the cells a steel shell can act as the cathode for the electrolysis (Fig. 1.93). The cell may have exterior ceramic or other insulation, and steel rod supports on the bottom can also be used for cathodes. The anode is constructed of graphite which slowly sloughs-off, and since this corrosion increases the spacing between the anode and cathode, and thus the voltage, usually some means of adjusting this spacing is built into the equipment. The vessel may be heated by gas firing between the outer fire brick and the vessel’s inner steel walls. Lithium metal accumulates at the cell’s surface where it automatically flows from the cell or is skimmed-off with a ladle. The metal is poured into ingots and allowed to cool under an inert atmosphere. Lithium chloride used in the cell must be quite pure and dry, and the chlorine gas that is generated (about 5 kg/kg Li) is carefully routed away from the molten lithium (LCA, 1968). Foote’s original cell design used 2.54 cm thick steel plate in the form of a 1.2 £ 1.8 m box 0.91 m deep. It was placed in a fire box with the flame impinging on the bottom, and had four 20.3 cm diameter graphite anodes 1.8 m long, supported from the top and lowered into the cell as the lower section was corroded-off and the voltage increased. The cathode was the steel box, with a fan on top to exhaust the chlorine under a slight vacuum. The cell ran at about 6 –6.5 V and 8600 A (a theoretical 80% energy efficiency) to produce 41 kg/day of lithium. The metal was

Toxicology

179

periodically withdrawn from the surface and cooled into ingots which were later remelted at 1868C and formed into the desired products. In case of fires lithium chloride was used to smother the flame (Hader et al., 1951). Initially the typical energy consumption for the electrolysis was about 46 kW hr/kg Li metal (not including the fuel requirement), or about 4000 A hr per kg of lithium at theoretically 3.76 V, with the actual range 6– 12 V. The cell voltage is reduced at higher temperatures, but the graphite corrosion rate also increases. In more modern cells the terminal voltage is more typically 6.7 –7.5 V, the current 30– 60 kA, the current density 6 –7 kA/m2, the power consumption 30 – 35 kW hr/kg Li and the LiCl consumption 6.2 –6.4 kg/kg Li. The average energy efficiency of the cells was initially about 20 – 40%, but now it is considerably higher. In the modern cells wire gauze diaphragms may also separate the electrodes and help channel the reaction products (Deberitz, 1993; Cooper et al., 1979; Averill and Olson, 1977). To produce high-purity lithium metal the lithium and potassium chlorides in the melt must be exceptionally pure, and then the carbon, nitrogen and oxygen that inadvertently enter the metal must be reduced. The liquid metal can be first filtered at 2008C through sintered stainless steel with 5 m pores to remove the calcium, carbides, nitrides and oxides of various metals. Also, the liquid lithium can be de-gassed under a high vacuum at temperatures of 20–6708C to remove gaseous impurities. “Getters” such as titanium, zirconium and aluminum can also be added before filtration to form various insoluble compounds. Liquid lithium with 20 ppm nitrogen and 150 ppm oxygen have been achieved (Averill and Olson, 1977; LCA, 1968). A process utilizing the direct electrolysis of lithium carbonate has been developed by the Lithos Corp. (LithChem International), with claims that it could reduce the cost of lithium metal by 20%. In 1997 they announced the planned construction of a pilot plant to produce 20–30 mt/yr of lithium metal (USGS, 1997). Raymor Industries Inc. also were considering commercializing a process developed by McGill University for the direct electrolysis of spodumene to form lithium metal. They claimed a 25% cost savings for the metal (USGS, 2000). The electrolysis of lithium solutions to form a lithium amalgam with mercury, and then converting the amalgam in a fused salt cell to the metal has also been proposed (Cooper et al., 1979). Lithium metal is sold in the form of either ingots, rod, wire, shot, sheet, special shapes or dispersions in both a high-sodium and low-sodium grade (LCA, 1968). In addition to the companies previously noted producing lithium metal, it has also been made by duPont in the USA; Yahagi Iron Co in Nagoya, Japan; Metaux Speciaux SA in Plombieres St. Marcel, Savoie, France; China; the former USSR, and various other companies (US Bur. Mines, 1992).

TOXICOLOGY Most of the simple inorganic lithium compounds are only moderately toxic, with values such as: (1) lithium carbonate; lowest reported lethal dose, oral LDLO

180

Part 1 Lithium

(rat) ¼ 710 mg/kg; lowest reported toxic dose, oral TDLO (human) ¼ 7 mg/kg; (2) lithium chloride; lethal for 50% of the test animals, oral LD50 (rat) ¼ 751 mg/kg; oral LD50 (rabbit) ¼ 850 mg/kg; intraperitoneal, ip LD50 (mouse) ¼ 604 mg/kg (Sax, 1979). Concentrated or solid lithium hydroxide can cause caustic burns, and skin contact with lithium halides can result in skin dehydration. Organolithium compounds are often pyrophoric and require special handling (Kamienski et al., 1993). Lithium carbonate and citrate also have some very important medical uses within a very narrow range of concentrations, but there are toxic effects beyond that range. They are very effective in the treatment of manic-depressive illness, bipolar disorder, depression, suicide prevention and for a variety of other psychiatric and medical conditions. However, it often causes minor-to-serious side effects, with weight gain and impaired coordination being the most common reasons for patients not taking the drug. Mild hand tremor is the most common side effect, with fatigue and muscle weakness second most common. There can also be lithium intoxication, constant thirst, frequent urination, blurred thinking, short-term memory deficits and the more serious renal, thyroid (endocrine) and cardiovascular complications. To be most effective, lithium should be taken at monthly intervals throughout the patients life and closely monitored for side effects, which usually immediately (or soon) disappear when treatment is stopped. Doses less than 0.6 mmol/liter of blood serum are usually not effective, and more than 1.5– 2 mmol/liter can cause lifethreatening reactions. However, the toxic effects usually wear-off fast (lithium is 50% excreted within 24 –48 hr), or in severe cases can be treated by vomiting, emesis and close monitoring of the body’s fluid electrolyte disturbances. Doses of 0.9 –1.4 meq of Li/liter (, 0.5 g/day lithium carbonate or citrate) are thus usually prescribed to alleviate the acute manic or other symptoms (Ezzell, 2003; Fieve and Peselow, 1985). Lithium metal is classified as UN 1415 or “Dangerous when wet” and will react with water to form flammable hydrogen and lithium hydroxide that can be corrosive to the skin and eyes. Lithium metal is easily ignited in the air and once burning, requires special techniques to extinguish (Kamienski et al., 1993).

USES Lithium has a wide variety of uses, and the ones that have consumed the largest volume of lithium have varied widely over time. In the early days of the industry, batteries were the largest purchaser of lithium (as metal or chemicals), then atom bombs, followed for many years by grease as the dominant customer. Then glass and ceramics followed for 10– 15 years by aluminum reduction furnaces utilizing the most lithium, and from 1990 onwards the market has become very diversified, with

Uses

181

glass and ceramics again being the largest purchaser of lithium chemicals or ore concentrates. Table 1.38 lists examples of the percentage of total US lithium sales by various categories for the period 1951 – 2000, and the tonnage sold by category from 1975– 1985. In the following sections each of the major categories of lithium sales will be separately reviewed. Glass In the glass industry lithium helps to make many types of products, such as borosilicate glass, containers and bottles, fiberglass, flaconnage, internally nucleated glass ceramics, pharmaceutical glass, photochromic glass, soda lime glass, television tubes, thermal shock-resistant cookware (for freezer-to-oven use) and sealed-beam headlights. In preparing glass lithium has many benefits, such as listed in Table 1.39. It increases the melting rate by lowering the viscosity of the glass and reducing the melting temperature. Tests have indicated that as little as 0.1 –0.2% Li can increase the productivity of the glass furnace 6 – 17% (Kingsnorth, 1988) without changing the batch cycle or reducing the glass quality (the density, refractive index, luster and transparency remain the same). The increase in the plants’ capacity and production rate with lowered temperatures also increases the life of the furnace lining. Lithium reduces the seed (bubble) count (content) in the glass, lowers its thermal expansion coefficient and provides higher chemical durability to the finished product. Another important benefit is the potential reduction of calcium fluoride (CaF2) used as a flux, and the partial or total reaction with any fluorine that may be present in the glass (as well as some of the SO2 and NOx), thus reducing corrosion and the toxic emissions from the kiln. The lower viscosity and temperature of the glass also increases the speed of the glass-forming equipment as the glass leaves the furnace (Harben and Edwards, 1998). The use of a 0.1– 0.5% Li2O addition enables container and bottle glass to produce lighter weight, thinner walled products (U.S. Bur. Mines, 1992). As an example of the reduction in melting point of glass batches, glass containing either 15% Li2O, Na2O or K2O had melting temperatures of 500, 700 or 8508C, respectively. The viscosity reduction is exemplified by a lime – soda – silica glass where the viscosity was 1012 poise at 5668C, and the same viscosity was obtained at 500, 544 and 5338C with glass containing an equal amount of lithium, sodium or potassium, respectively. An example of the improvement in the glass furnace capacity is illustrated by the replacement of 1% Na2O with 0.48% Li2O in zinc alabaster or opal glass causing the reduction of 8.4– 10.5% in the melting time, and 18.2% and 23.0% in fining time (removing bubbles), respectively (Fishwick, 1974). Some of the typical amounts of lithium added to various types of glass are shown in Table 1.39. Either lithium carbonate (or other lithium compounds) or lithium mineral concentrates may be used as the lithium source in many types of glass, with several of the pure ore melting points being: spodumene 14238C, eucryptite 13978C and

182

Table 1.38 Table A. Usage Pattern for Lithium in the United States (% of sales)a 1997

1993

1992

18 32

18 20 2 11 7 13 13 4 12

34 18 11 7 13 13 4

38 14 11 7 13 13 4

18 9 9

14

1989

46.2 16.4 20.1 4.8 1.9 9.0 1.6

1985 24 14 28 17 1.5 2.5 7.5 3.0 2.5

1980

32 33 20 0.5 1.5 6 4.5 2.5

1969

— 14 20 — — — 7 —

1953

1951

39 — 47 2 2 — 2 8

31 — 40 10 4 — 5 10

Part 1 Lithium

Glass and ceramics Aluminum Lubricants Batteries Organics Chemicals Air conditioners Other

2000

B. Useage Pattern (mt of Li) (USGS, 1986)

Glass and ceramics Aluminum Lubricants Batteries Other Total

1985

1984

1983

1982

1981

1980

1979

1978

1977

1976

1975

730 640 550 90 270

910 1000 730 55 220

550 730 410 45 270

410 550 360 27 470

820 910 450 0 730

770 820 500 0 640

820 910 450 0 730

820 1090 450 0 730

1000 1360 450 0 910

550 1320 360 0 320

640 1270 320 0 400

2270

2910

2000

1820

2910

2730

2910

3090

3730

2550 2630 (continues)

Table 1.38 (continued) C. Sales of various lithium chemicals, million pounds of lithium carbonate equivalenta 1992 Li2CO3 LiOH·H2O LiCl LiBr Other salts Metal n-Butyl lithium Ore concentrates Miscellaneous Total

48 11 — — — 5 3 — — 68

1989

1974

47.7 20.1 — — 3.0 4.8 1.9 14.9 1.5

35.0 11.4 1.9 2.2 — 1.2 1.2 — 0.1

79

53 D. Sales for aluminuma

1974 1973 1972 1971 1970 1969 1968 b

Various sources. Percent of total lithium sold.

1550 mt Li 1270 960 860 640 270 180 Uses

a

38%b 36 32 30 26 14 9

183

184

Part 1 Lithium Table 1.39 (Kingsnorth, 1988) The benefits of spodumene to glass making

A. Related to lithia content Reduced melting temperature giving Reduced energy consumption Increased furnace refractory life Reduced glass viscosity in the molten and semi-molten states, leading to Increased “pull” (production) rates (8–15%) Better glass forming characteristics, higher “pack” Improved thermal shock resistance of finished product Improved strength of glass product Low cost B. Mineral form Reduced rejection rate (0–3%) Improved glass quality with respect to fewer “seeds” and better thermal shock resistance The GGS alkali content (equivalent to 100 kg of soda per tonne on a molar basis) The GGS alumina content (170 kg of alumina per tonne of GGS) Established lithia additions in the glass industry Application

Method of addition

Typical % Li2O (% Li)

TV tubes Reduces melting temperatures Improves forming properties Good finish to glass

Spodumene concentrate Petalite and Li2CO3

0.1– 1.0 (0.05– 0.4674)

Pyro-ceramic ware Zero coefficient of expansion Improves forming characteristics

Spodumene concentrate Petalite and Li2 CO3

0.4– 4.0 (0.19– 1.86)

Fibreglass Reduced viscosity, improves continuity of fibre production

Various minerals including spodumene concentrate

0.1– 1.0 (0.05– 0.46)

Various minerals including 0.1 to 1.0 spodumene concentrate

(0.05– 0.46)

Various minerals including spodumene concentrate

0.1– 0.8 (0.05– 0.37)

Safety glasses Improved Strength

Vacuum flasks, perfume bottles Ease of forming Good finish and strength

This table appeared in Industrial Minerals No. 244, February 1988, p. 24. Published by Industrial Minerals Information, a division of Metal Bulletin plc, UK. qMetal Bulletin plc, 2003.

Uses

185

petalite 13568C. However, since the commercial products only contain 50 –95% of the pure mineral, their actual melting points are somewhat lower than this because of forming eutectic compositions with their feldspar, quartz and mica impurities. Kaolin is often added with the lithium ores, and within limits the aluminum and silica in the ores can be beneficial to the glass. In some cases the ore concentrates further improve the glass by making it slightly more dense than when using lithium carbonate, and the surface hardness may be improved by as much as 20% (Fishwick, 1974). Different minerals may also have advantages over other types of ore for certain glasses, such as having a low degree of expansion upon being melted, or other beneficial physical properties. Many of the lithium ore concentrates have a low enough iron (or other harmful impurity) content to be used directly in some glass formulations, while other glasses require a higher purity lithium source. In addition, some concentrates, such as highiron spodumene have a more restricted use, or require that the iron content be lowered. Some ores can even be used without being formed into higher-purity concentrates if the lithium content is high enough, and the iron is sufficiently low. An example of this is Glass Grade Spodumene from Australia with a minimum of 4.8% Li2O (2.23% Li) and a maximum 0.2% Fe2O3 (usually 0.1%; Table 1.40). In 1988 the ore concentrates only cost about 40% of an equal amount of lithium in lithium carbonate and were quite suitable for uses such as container glass and pyroceram (Kingsnorth, 1988). However, since the lithium carbonate price reduction in 1998, the cost for the lithium in most ores became roughly the same as lithium carbonate, or only slightly lower (Table 1.41). About 50% of the total consumption of lithium in the glass, ceramics and aluminum industries prior to 1998 had been from ores, but since that time the reduced price of lithium carbonate has allowed it to replace some of these ore uses. Ores constituted about 15% of the total lithium market in 1993 (Flemming, 1993a,b), and in 2002 the estimated 3010 mt of contained lithium sold as ore (assuming an average grade of 4.0% Li2O) was about 19% of the total (Tamlin et al., 2002). Ceramics Lithium is used in ceramics to make frits and glazes, porcelain enamels (for kitchenware and bathroom fixtures), sanitaryware, shock-resistant ceramics and porcelain tiles. Either alone or combined with other compatible materials such as feldspar and nepheline syenite it produces lower melting temperature mixtures with increased fluxing power. This improves the product quality, plant efficiency and productivity by lowering the firing (vitrification) temperature, reducing the firing cycle time and reducing the “soak” period. It also forms products with lower thermal expansion coefficients (and thus greater shock resistance), lower pyroplastic deformation, more brilliant body and glaze colors, greater glaze adherence and gloss, and more stain resistance. Again, both ores and lithium compounds can be used for this application, and with ores, petalite is usually preferred over spodumene because there is no volume,

186

Table 1.40 Analyses of Various Spodumene Ore Concentrates, wt.% A. Chemical analysis Greenbushesa

Tanco

Glass grade concentrates

Li2O Fe2O3 K2 O Na2O SiO2 Al2O3 MgO CaO MnO2 P2O5 F LOI

Glass grade

Ceramic grade

Low iron

Typicalc

Specificationsd

Typical

Spec.

Specifications

4.8–6.5 1.7 0.5 0.3 63.0 24.7 Trace Trace — — — —

7.2 0.9p p 0.27 0.30 64.1 26.5 Trace Trace — — — —

7.1p 0.1p p 0.14 0.35 64.8 26.3 Trace Trace — — — —

7.25–7.30 0.04–0.05 0.10–0.20 0.15–0.20 — 25 –27 — — 0.02–0.04 0.15–0.25 0.01–0.02 —

7.25 ^ 0.1 0.06 ^ 0.01 0.30p p 0.36p p — 24.0p — — 0.04p p 0.27p p — —

5.01 0.12 0.17 0.09 75.91 17.88 — — — 0.06 — 0.29

4.8p 0.13p p

7.5p 0.10p p

B. Particle size Typical

b

Greenbushesa

Tanco

Glass grade concentrates Glass grade þ 20 mesh (841 mm) þ 28 mesh (600 mm) þ 48 mesh (300 mm) þ 65 mesh (212 mm) þ 200 mesh (75 mm) p

minimum; p p maximum.

a Flemming (1993b). b Dresler et al. (1998). c Burt et al. (1988). d

Vanstone et al. (2002).

Ceramic grade — — — — —

Low iron — — — — —

Typical — — — — —

c

d

Specifications 0

p p

trace 1.0p p — 50.0p

Typical

Spec.

Specifications

— — — — —

nil — — 95.0p (105 mm)

— — — 5.0p p 60.0p

Part 1 Lithium

Typical

b

Uses

187

Table 1.41 Examples of Lithium Ore Concentrates’ Prices, $/mt (Saller and O’Driscoll, 2000; Tamlin et al., 2002) Li2O (%) Spodumene concentratesc Glass grade spodumenec Petalited Lithium carbonatee

6.9–7.5 4.8–5.0 3.5–4.6 40.4

1998–2000

1992a

1988b

$/kg Li2O, Dec. 2000

330–395 215–230 180–270 2068–2600 f

385 175 230 4320

340 160 175 3410

5.12–6.53 4.48–4.60 4.18–6.28 5.12–6.53

a

Kunasz (1994). Kingsnorth (1988). c Basis, seller’s US warehouse. d Basis, f.o.b. Durban, Australia. e Basis, bags or drums, delivered in the US. f During January 1999 $1760–2200 mt21; for year 1999 $1760– 2030 mt21; McCracken and Sheth (1999). b

structure or phase change as it is heated, as there is with spodumene. Since lepidolite is the only ore that contains both fluorine and rubidium (which are also good fluxes), it has been preferred in some ceramics and glass applications, but it is no longer as plentiful as the other ores (Anon., 2001; LCA, 1968). Glazes and enamels utilize lithium to the same extent as ceramics, and for the same reasons. An enamel is a glass-like coating bonded to a metal (steel, cast iron, aluminum, etc.) by fusion at temperatures above about 5508C, causing the metal and the enamel to permanently combine. This gives a product the hardness of glass and the strength of the base metal. Lithium imparts desirable properties to both enamels and glazes, and can allow their use without first forming frits, to make aventurine, corrosion-resistant (“glass”) coatings for steel tanks, in high-voltage porcelain because of its low coefficient of expansion, leadless glazes for dinner ware, opaque and crystalline glazes, and to produce whiteware. Lithium (0.5 – 4% Li2O) is also one of the ingredients that has allowed the production of glass-ceramics, in which glass is forced to crystallize into very fine crystals that form a dense, strong, heat-resistant ceramic material. Examples of this are Pyroceram cookware, stove tops and the nose-cone tiles on space vehicles. The increased surface hardness, and the very low thermal expansion between 2 73 and about 6008C (e.g., 0 ^ 0.05 £ 1026/8C from 0 to 508C; Deberitz, 1993) has been beneficial for optical glass and large telescope mirrors (Fig. 1.94). Lithium is also used in various applications where improved resistance to sudden temperature change, and a lower coefficient of expansion is important such as some optical glass ceramics or refractories (e.g., specialty brick for furnace linings). There are also many other ceramic-type applications for lithium, such as lithium ferrites (Li0.5Fe2.5O4), which maintain their magnetic properties up to 6808C. Lithium borosilicates can be used as

188

Part 1 Lithium

Figure 1.94 Examples of glass ceramics for telescopes; left, support; right, 8.6 m diameter (45 mt) mirror (Deberitz, 1993, courtesy of Schott Glass, Mainz/Germany and Chemetall GmbH).

the binder for high-temperature fuzed alumina grinding wheels or refractories, and lithium can act as an accelerator for cement, in mullite formation, to densify magnesium oxide, to make piezoelectric materials or self-curing paints, and many other widely differing uses (Fishwick, 1974). In some cases the lithium is first reacted with other desired metals to form compounds such as cobaltites, manganites, etc., before being added to the ceramic or glaze mixture. Aluminum Lithium is employed in the aluminum industry in amounts such as 1– 3% LiF in the bath, or ratios such as 2 kg lithium carbonate/mt of aluminum (the range is usually 1.5 –4 kg; Nicholson, 1977). It lowers the electric reduction cells’ (Fig. 1.95) temperature (, 9708C; the alumina –cryolite melting point), raises the electrical conductivity of the cell (thus lowering the required overvoltage, which reduces the power requirement), and it reduces the fluorine emissions from the electrolytic cells by 25 – 50% (Table 1.42). Lithium carbonate reacts with the cryolite (Na3AlF6) “solvent” in the cell to form lithium fluoride, which has a very high fluxing ability, electrical conductivity and low volatility. Lithium is most beneficial for older plants, where it can reduce the energy costs by 5– 10%, and bring the cell efficiency up to 90 –95%. However, most of the lithium was purchased to initially charge the cells, and the replacement amount is now comparatively small. Also, with the more modern cells the energy efficiency is already at 90 – 93%, and lithium can only add 1 –3% greater efficiency. Thus, the usage of lithium is now less cost effective, and has slowly declined to the point that in the US in 2000 only about 50% of the aluminum mills employed lithium in their cells (Jarvis, 2000). Lithium alloys with aluminum also find some use, since lithium can impart some very useful properties to the aluminum. For example, one lithium alloy retains a high strength to 2048C in contrast to conventional alloys only being effective to 121 –1778C. The alloy also increases the modulus of elasticity by about 8%, which

Uses

189

Figure 1.95 Aluminum electrolytic cells (Deberitz, 1993, courtesy of Hamberger Aluminum Werke (HAW) and Chemetall GmbH).

could facilitate an aircraft flying at speeds of 1300– 1600 mph (Kesler, 1960). At the maximum of 2 – 3%, lithium could reduce the density of an aluminum alloy by 7– 10% and raise the modulus of elasticity by 10 –15% (Deberitz, 1993). However, lithium also increases the cost of the aluminum alloy by 3– 6 fold, and since about 70% of an airplane’s weight was aluminum (before the advent of composite fiber and plastic materials) this price increase was prohibitive. There were also some problems with corrosion and fatigue cracking. In addition, lithium is extremely difficult to introduce into alloys, since lithium floats on the molten aluminum, immediately burns, and very little enters the aluminum. Either expensive and hard to manage fluxes must cover the molten aluminum or the lithium must be added to the molten aluminum as it leaves the furnace. In both cases an inert gas must cover the metal ingots as they cool, and water-quenching can not be used. Finally,

190

Part 1 Lithium Table 1.42 Advantages and Disadvantages of Using Lithium in the Aluminum Industry (Nicholson, 1977)

Advantages of Lithium Lowers the melting temperature of the molten electrolyte. Increases the cell’s electrical conductivity. The above items lower the electricity consumption and other costs, or increase the production for the same unit costs. Reduces the fluorine emissions. Reduces the consumption of anode carbons. Reduces the consumption of cryolite. Disadvantages The operational control is more difficult and more dependent upon instruments. Lithium decreases the purity of the aluminum because of the presence of some lithium and perhaps iron in the product. It sometimes causes difficulties in casting intricate shaped products due to the formation of heavier oxide layers. Scrap aluminum with lithium is undesirable.

the reprocessing of lithium alloy scrap is quite difficult and hazardous, which adds to the cost of fabrication (Anon., 1998). A considerable amount of research has been done on lithium – aluminum alloys because of their favorable properties, and in satellite and space applications, plus some of the more demanding aircraft parts, various uses have developed. An example of this is NASA’s “Super Light Weight Tanks” for the liquid hydrogen fuel and liquid oxygen used in space shuttle launches. Twenty-five of the tanks were ordered in 1996 to be 47 m long, have an 8.4 m diameter and weigh 26.3 mt (Fig. 1.96). They were made from 5.9 mt ingots that were then rolled into various metal thicknesses for different sections of the tanks. The lithium alloy contained from 0.3– 0.5% silver, and had excellent cryogenic properties, strength up to 100,000 psi, and good fatigue, fracture toughness and corrosion resistance. Each 1% Li in the alloy lowered the density of the metal by 3%, and raised the elastic modulus (stiffness) by almost 5%. Another alloy without silver (that was cheaper and 4– 5% lighter) was used in the F-16 airplanes for a critical aft bulkhead because of its superior fatigue resistance (Anon., 1998). Batteries Lithium is the most electropositive of all metals, with a standard electrode potential of 3.045 V compared with 2.71 V for sodium and 0.76 for zinc. It thus can generate the greatest electrical power per unit weight or volume of any metal, but it is also extremely reactive and thus potentially dangerous. Special designs and applications

Uses

191

Figure 1.96 Super light weight aluminum– lithium alloy fuel tank for space shuttle launches (Anon., 1998; reprinted by permission of Light Metal Age).

are thus required, and they have perhaps been best achieved with rechargeable batteries (Fig. 1.97). In 2001 there were four common types of rechargeable batteries: lithium-ion (52% of the market), lithium polymer (4%), nickel metal hydride (27%), and nickel –cadmium (17%). Amongst these, the lithium-ion and lithium polymer batteries could store and deliver the most energy per unit of space, with the commercial batteries producing 3.7 V, or about three times more than the nickel cadmium or nickel metal hydride batteries. They were lighter, had a longer shelf life, and did not have the “memory effect” problem (the amount of energy stored was decreased if the battery was charged before having been fully discharged) of the nickel batteries. They were thus preferred for the newer generations of highperformance applications such as in mobile phones, camcorders, laptop computers,

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Part 1 Lithium

Figure 1.97 The general structure of a rechargeable lithium battery (Abraham, 1985; reprinted from Lithium, Ed. R. O. Bach, Sec. 11, q1985 by permission of John Wiley & Sons, Inc.).

hand-held portable electronic devices, tracking systems, home repair or construction tools, and military and medical devices, even though they were more expensive than some other batteries (Cairns, 2002). Initially rechargeable lithium batteries used a non-aqueous electrolyte and lithium plates as the anode, allowing lithium to dissolve during use, and to be replated onto the anode when being recharged. This was a conventional oxidation – reduction battery with a very high voltage output and capacity, but it required expensive control circuitry to prevent fires or explosions if the batteries overheated (its cell reaction is exothermic). Since lithium metal is very reactive the cells could also be dangerous if water entered the battery, or under certain other conditions. This problem could be somewhat improved if the lithium was alloyed with copper or tin, since the anodes’ crystal structure would still not change much during the charge – recharge cycles, and the alloys’ lower reactivity could lessen the need for expensive safety switches in the batteries. Even with its problems these batteries were extensively used for large, special situations such as stand-by power sources for missile silos. It is possible that new technology such as the lithium polymer battery design of Fig. 1.98, or the use of a thin-film, polymer-ceramic composite electrolyte may re-establish the use of lithium metal anode batteries (Alper, 2002). An even earlier use of lithium was in the old Edison nickel –iron batteries where lithium salts in the electrolyte provided high conductivity, and prevented freezing down to – 408C.

Uses

193

Figure 1.98 The lithium polymer battery concept (D’Amico, 1996; reprinted with permission from Chemical Week, March 20, 1996, Chemical Week Associates).

The new rechargeable lithium ion batteries totally eliminated the above problems since they do not contain metallic lithium, and instead obtain their power from the concentration difference of lithium ions traveling between the electrodes. The anodes are ultrapure graphite impregnated with lithium ions, one lithium ion to six carbon atoms, and the cathodes are extremely porous lithium – cobalt, nickel or manganese oxides (such as LiCoO2). In 2002 polyvinylidene fluoride (PVDF) was used as a binder for both the cathode and anode in 80– 90% of the batteries, with a minimum of structural rearrangement to the electrodes during use. The lithium electrolyte has to be non-aqueous, and was usually lithium hexafluorophosphate (LiPF6) dissolved in ethylene carbonate, but some LiMnO2 or other paste compositions were also employed. They allow the lithium ions to transfer fairly freely (but unfortunately with some resistance) between the electrodes during the charge or discharge cycles. The graphite, or sheets of carbon anodes were loaded with lithium ions, which flow from the anode as the battery discharges, and form a complex with the metal oxide at the cathode. The lithium ions are forced back to the anode when the battery is being charged. The carbon in the anode (negative electrode) was subject to improvements, since it could be flammable, it was slow in taking up the lithium ions, and if lithium became plated onto the electrode it was dangerous. For the positive electrode cobalt oxide was most commonly used with lithium oxide, but cobalt had the problems of being very expensive, a toxic material, and it could overheat upon charging. Manganese oxide (MnxOy) with lithium oxide was potentially a much better cathode, but it had a relatively short life in the past because it looses its structural integrity as it is being used. The addition of some chromium extended its life considerably. Iron (or other transition metal) phosphates (such as LiFePO4)

194

Part 1 Lithium

when doped with aluminum, niobium, zirconium or magnesium had also been proposed as cathodes that were less expensive, less toxic, and had a good electrical conductivity (Anon., 2002; Cairns, 2002; Tullo, 2002; Alper, 2002; Anon., 2000; D’Amico, 1996). In the lithium polymer batteries a conductive lithium polymer replaced the inorganic-filled organic liquid electrolytes (Fig. 1.98), enabling the use of lightweight plastic cases of various shapes (which could even be flexible) in place of the conventional metal cases. The polymers that could be used in the batteries include PVDF copolymers and fluoropolymers, all of which could help hold the battery together as well as separate the positive and negative electrodes, even at high voltages. Since the lithium polymer technology would replace liquid electrolytes it was claimed to be safer than lithium ion batteries since it could not overheat or explode under some circumstances. It had been expected that these batteries would offer the greatest growth potential for the future, but up to 2002 that had not yet occurred. Lithium’s 47% fraction of the $3.61 billion rechargeable battery market in 1999s had become 52% and $3 billion by itself in 2002. Sony Corp. had about 33% of this market, and Sanyo Electric Company 23% in 2000. Sony originally developed the lithium-ion batteries, but in 2000 began converting much of its manufacturing capacity to the more profitable lithium polymer type. Sanyo Electric also produced about 32% of the nickel –cadmium, and 46% of the nickel hydride batteries in 2000 (Lerner, 2001; Jarvis, 2000). Considerable research has been conducted on rechargeable lithium batteries for automobiles, but by 2002 there were still major safety and construction problems. Non-rechargeable lithium batteries have been used for many years in large electric storage units by the military, and later small batteries (Fig. 1.99) began to be used in calculators, cameras, watches, microcomputers, electronic games, small appliances, toys and other applications where a long life and/or high current density are desired (for instance, pacemakers could last 8– 10 years compared to 1 year for conventional batteries). They provide higher energy per unit wight than any other metal, with the electrochemical equivalence of lithium being 3.86 Ahr/g, compared to 1.16 for sodium, 0.5 for silver, 0.48 for cadmium, 0.28 for zinc and 0.26 for lead (Deberitz, 1993). They are more expensive than ordinary alkaline batteries, but have a much higher performance, and in 2002 were the dominant battery type in some countries such as Japan. For certain applications lithium sulfur dioxide batteries have been made with a 10 year life with no reduction in performance, and were included in the Galileo spacecraft for it anticipated 6 year trip to explore the planet Jupiter (U.S. Bur. Mines, 1992). Some of the other cathode materials that have been used with lithium, and their practical energy density as mW/cm3 are thionyl chloride (SOCl2) 700 –800; copper oxide (CuO) 550– 650; manganese dioxide (MnO2) 500 –580; carbon fluorine (CFx) 450 – 500; sulfur dioxide (SiO2) 400– 450; bismuth oxide (Bi2O3), pyrite (FeS2), and lead bismuthate (PbBi2O5) 350 –500. They operate at voltages from 1.5 – 3.9 V (Deberitz, 1993).

Uses

195

Figure 1.99 Sketches of various types of lithium batteries (Marincic, 1985; reprinted from Lithium, Ed. R. O. Bach, Sec. 9, q1985 by permission of John Wiley & Sons, Inc.).

There is a considerable literature on lithium batteries as both technical articles and patents. This includes many on the construction of the batteries themselves, and on each of the battery components. The use of lithium in batteries is not a large market for lithium (about 7% of the total in 2001), but the consumption has grown at an annual rate of 15% from the late 1980s to 2002. Even though the batteries utilize lithium and other metal oxide cathodes, and lithium in the electrolyte, the amount per battery is very small. However, growth should continue, and if the use of batteries in automobiles developed this could greatly increase the sales of lithium in this market. It is also expected that the market will increase when these batteries are produced in standard sizes (as has occurred with the non-rechargeable batteries). In the past many of them have been custom made for each application (Tullo, 2002; Saller and O’Driscoll, 2000; Schmitt, 1999).

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Part 1 Lithium

Grease Considerable lithium hydroxide is used in making greases, and the demand grew at a steady 2% per year for the period 1980– 2000. It is used in military, industrial, automotive, aircraft and marine applications, and 55% of all industrial greases contained lithium in 1981 and 60% in 1993. Lithium hydroxide (about 1 pound per 45 –100 pounds of grease) is reacted with 12-hydroxy-stearic or other fatty acids, since lithium stearate forms a matrix or sponge-like gel lubricant where the lithium attaches to the metal, and the long-chain multi-hydroxyl end of the stearate molecule extends outward in the form of interlocking spirals to hold the petroleum lubricant and cushion the wearing surface. Mixtures containing 5– 10% of the lithium soap are an excellent lubricant for bearing surfaces, since they are almost totally water insoluble, and stable in consistency over a range of shear and temperatures from 2 55 to þ 2008C. The gel holds a high volume of oil, resists oxidation and hardening, and if liquefied will reform as a stable grease upon being cooled. Because of these qualities the grease is used over a wide variety of demanding service applications (Deberitz, 1993; Lloyd, 1981). Other Uses Lithium Metal The market for lithium metal was growing at about 5% per year in the early 2000s because of its use in making organic chemicals, batteries, alloying and other applications. It is made by the electrolysis of a molten lithium chloride –potassium chloride mixture in specially designed cells, with the molten metal collecting in the top and being periodically withdrawn and cooled as ingots. Most of the ingots are then converted into a wide variety of other shapes and forms, including thin sheets, pellets, powder, etc., for each specific use. Lithium is quite soft (about 0.6 on the Mohs scale), and it can be scratched and cut with a fingernail. Some lithium is alloyed into lithium – aluminum (containing up to 7.5% Li) and lithium –magnesium (up to 13% Li) metals because of their low density (Li weighs 33, Mg 108 and Al 162 lb/ft3), high-temperature performance, and improved elasticity, tensile strength and corrosion resistance. Many of the alloys commonly contain 2 – 3% lithium, and have been used in commercial or military aircraft where they have the potential of reducing the aircraft’s weight by as much as 10%. Their usage, however, has been limited by their high cost, the introduction of competitive high-performance fiber – plastic compositions, and the difficulty in forming the alloys because of lithium’s extreme reactivity (Jarvis, 2000; U.S. Bur. Mines 1992; Anon., 1981). Other uses for lithium metal include its ability in very small amounts to remove oxygen or other gases from many molten metals, its use as an intermediate or raw material in the production of organic compounds, or in batteries, as noted elsewhere in this chapter (Kunasz, 1994).

Uses

197

Air Conditioning In air conditioning lithium bromide or chloride are used in the dehumidification of air and other gases because of the very low vapor pressure of their solutions, their low viscosity, high stability, non-toxic properties and low corrosivity (the solutions are made neutral or basic, and corrosion inhibitors are usually added). Both lithium bromide and chloride are extremely hygroscopic, and can dry air or other gases down to a very low moisture content. As they remove water from the air the gas is also cooled (because of water’s high heat of vaporization), thus providing a refrigeration effect. Their solutions (such as 54– 55% LiBr) are used in very large building air conditioning systems (Fig. 1.100) to remove the desired amount of moisture from the air, and then heat or further cool the air to its most comfortable temperature. A slip stream of the lithium solution is continuously removed from the absorber and evaporated back to its most effective concentration. Small amounts of lithium hydroxide, and perhaps lithium chromate, nitrate or molybdate are added to the recirculating brine as corrosion inhibitors (Deberitz, 1993). The solutions can also be used for absorption – evaporation (chilling), refrigeration or heat-pump systems. Solid lithium chloride or bromide can be used to dry organic liquids, as a desiccant, and in dehumidification applications (Lloyd, 1981). Organic Compounds Many organic compounds containing lithium have found important industrial, medical and other uses. In these compounds the lithium is usually bonded directly to the carbon atom, and because of the covalent nature of these bonds many of the compounds are liquids or low-melting solids. They are soluble in many hydrocarbons, as well as often being soluble in polar organic solvents such as ethers, alcohols or related materials. Many of the organolithium compounds are reactive with oxygen or air, and they may ignite spontaneously in the pure state or concentrated solutions (Kamienski et al., 1997). The most prominent organolithium compound is normal butyl lithium, which is used as a stereospecific catalyst in the polymerization of butadiene, isoprene and styrene for the production of synthetic rubber, and for the production of other polymers or elastomers. These rubbers are especially useful since the lithium catalysts develop an unusual microstructure in the product that provides various superior physical properties. For example, it can catalyze copolymers of styrene and butadiene for automobile tires that are relatively abrasion-resistant, and thermoplastic rubbers that do not require later vulcanization. The catalysts can also form “castable” elastomers and liquid polymers with a wide range of molecular weights for solventless surface coatings and other uses. Normal butyllithium can be shipped in various forms, including frequently as a 15 – 20% solution in hexane in special containers (Fig. 1.101). Other organic lithium compounds are catalysts for polyethylene, polyethyleneterephthalate films and fibers, and various other polymers. Some lithium-organics

198

Part 1 Lithium

Figure 1.100 An example of a large lithium bromide air conditioning unit (Deberitz, 1993, courtesy of York International GmbH and Chemetall GmbH).

Uses

199

Figure 1.101 Typical container for organolithium compounds (Deberitz, 1993, courtesy of Chemetall GmbH).

(as is lithium metal) are useful in many Grignard-type reactions, and find applications in the pharmaceutical, agricultural and other fields. For example, lithium is used in the preparation of vitamin A, antihistamines, carotenes, some steroids, synthetic penicillins, tranquilizers and many other compounds. Usually lithium metal, carbonate or chloride are used as the starting materials in the synthesis of these organic lithium compounds (Jarvis, 2000; Kamienski et al., 1997; U.S. Bur. Mines, 1992; Lloyd, 1981; LCA, 1968). Miscellaneous Uses One of the main application of lithium carbonate is as the starting chemical to produce a wide variety of other lithium compounds. An example of this is the production of lithium hypochlorite, which finds fairly extensive use in bleaches, sanitizers and swimming pool conditioners. It is used in many large-scale laundries as a bleach, and in swimming pools it provides excellent sanitation (by killing bacteria) while minimizing algae growth without the problems of a calcium residue. It sold at the rate of three million pounds per year in 1981, and was one of five major markets for lithium. As a medicine, lithium carbonate or acetate has been used since 1949 as a very effective treatment for manic depression (approved by the US Food and Drug Administration in 1969). It is used in very small quantities (e.g., 600 –900 mg/day Li2CO3), since too much can be toxic and have serious side-effects, and too little will not be effective, hence it is closely monitored in the blood stream of patients (see the Toxicology section). The reason for its effectiveness has remained a mystery, as is the case with some of its modern competitive medicines.

200

Part 1 Lithium

In metallurgy lithium metal is used to degas (scavenge, or remove gas from) aluminum, copper, bronze (this results in these three purified metals having a higher electrical conductivity), germanium, lead, silicon, thorium and other metals. It may also be used as an alloying ingredient for various metals besides aluminum and magnesium. Lithium chloride is an additive (or flux) to salt baths for dip brazing and open hearth soldering, and lithium carbonate, chloride or fluoride are used as scavengers and cleaners since they form low-melting slags with many metal oxides. They are also used for flux welding powders and welding rod coatings for difficult to weld metals (i.e., steel alloys and aluminum) where they reduce the flux’s melting temperature and surface tension, and increase the metal’s wetability. In the building industry lithium carbonate is an additive for quick-setting cement, special adhesives and quick-curing floor tile. Lithium hydroxide can reduce the premature deterioration of concrete because of its stronger reactivity with silica (USGS, 1997). The electronics industry uses high purity lithium carbonate and other salts for solid ion conductors and monocrystals. Dyes and pigments employ lithium hydroxide as an additive for dyestuffs to increase their solubility, and for increasing the brilliance of specific pigments. Other lithium salts are used with acid dyes. Lithium chromate can be used as a corrosion inhibitor for aggressive aqueous solutions in absorption refrigerators, and lithium hydride as a gas source for air-sea rescue kits. Pellets of lithium hydroxide or carbonate have been used extensively as adsorbents for carbon dioxide in submarines, space vehicles, and portable life-support systems (Fig. 1.102). A mixture of lithium nitrate and potassium nitrate is useful in forming hot melts to vulcanize various plastics such as EPDM, EPT or EPM (Deberitz, 1993). At some time in the future molten lithium might also be used as a hightemperature heat transfer fluid (it melts at a low temperature [180.58C], but does not boil until 13478C). Large amounts of lithium hydroxide monohydrate were purchased from 1953– 1960 by the US Government for its 6Li content, which was converted into tritium for hydrogen bombs. About 75% of the 6Li was extracted, and the remaining 42,000 mt of lithium hydroxide has been slowly sold for industrial use. The 6Li also has a high neutron cross section, so it could be useful for reactor shielding, or perhaps much later for nuclear fusion reactors if they were ever to be found feasible. Lithium carbonate or other lithium salts might also have a potential application for molten fuel cells if they were to become popular for powering electric cars or other uses (Saller and O’Driscoll, 2000; Kamienski et al., 1993; U.S. Bur. Mines 1992; Lloyd, 1981; LCA, 1968).

INDUSTRY STATISTICS The production of lithium carbonate prior to 1966 came primarily from the processing of lithium minerals (since the 1960s primarily spodumene and petalite), but by 1998 this source became phased out except in China (using lepidolite) and

Industry Statistics

201

Figure 1.102 An example of an air purification, carbon dioxide removal unit (Deberitz, 1993, courtesy of Dornier GmbH and Chemetall GmbH).

Russia, and now most of the world’s supply is extracted from various brine deposits. This source, combined with solar evaporation is much more economical, and thus has allowed the price of lithium carbonate to be considerably lowered. Initially the two US producers, FMC (formerly LCA) and Chemetall GmbH (formerly Foote Minerals and then Cyprus Foote) purchased lithium minerals or concentrates, and later mined spodumene from their large North Carolina deposits (Johnson, 1958) and converted it into lithium carbonate. Then in 1966 Foote began to recover lithium from their Clayton Valley (Silver Peak), Nevada brine deposit, and in 1984 from brine in the Salar de Atacama (with the final processing being done near Antofgasta), Chile. In 1997 SQM (originally Sociedad Quimica y Minera de Chile, then SQM Chemicals, and now SQM S.A.) also began to produce lithium carbonate from the Salar de Atacama, and cut the selling price of lithium carbonate roughly in half to gain market share. In 1997 FMC opened a similar operation at the Salar de Hombre Muerto in Argentina, but closed it the next year and contracted to purchase the less expensive SQM product. They continued to produce some lithium chloride from this facility. The Cyprus Foote and FMC spodumene operations were both officially closed by 1998.

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Part 1 Lithium

The lithium carbonate production capacity of various companies over the years to 2002 is listed in Table 1.43. The primary producers in 2002 were: (1) SQM Chemicals, with a capacity of 22,000 mt/yr of LCE from the Salar de Atacama in Chile. (2) Chemetall GmbH (who acquired the former Cyprus Foote Minerals) with 16,000 mt/yr capacity on the Salar de Atacama, and 5700 mt/yr from Clayton Valley in Nevada. And (3) FMC with 20,000 mt/yr of idle capacity from the Salar de Hombre Muerto, Argentina (Jarvis, 2000; Saller and O’Driscoll, 2000). During the 1990s there were three major producers of lithium ore concentrates, as indicated by production capacity in Table 1.43 of various mines or countries for several years preceding 2002. Sons of Gwalia owned the largest high-grade lithium (spodumene) pegmatite deposit that was in production during this period, with sufficient capacity to supply all of the world’s needs. The other two large producers of lithium concentrates were Tanco in Canada and Bikita in Zimbabwe. The amount of production from China and Russia was unknown but probably substantial, and there was relatively small production in Brazil and Portugal (and until 1998 from Namibia). Two other deposits in Canada, and one in Finland were considering production in 2002, but they faced the problems of heavy competition in an over supplied market. The estimated Western world consumption of these lithium ore concentrates (or raw ore) in the year 2000 was 158,200 mt, with an equivalent lithium carbonate content of about 18,200 mt. The specifications for various ore concentrates is listed in Table 1.40, and the list prices for spodumene and petalite concentrates for 1992 – 2000 are given in Table 1.41 (Anon., 2001). The estimated lithium content of lithium carbonate (or chloride) or concentrates (or ore) produced from their own deposits by various countries is listed in Table 1.44. There is a wide variability in the accuracy of these numbers, as some are merely educated guesses, and different sources have estimated quite different numbers. Also, in some years an estimated grade of the lithium ore or concentrates that were sold was made to establish the tonnage of lithium. Even with these inaccuracies, the numbers should be approximately correct, and they indicate that the total market for lithium has grown at about 3– 5% per year from 1960 –2000. The estimated US import, export, consumption, production, and the list price of lithium carbonate for a number of years is listed in Table 1.45A, as tabulated by the US Geological Survey or Bureau of Mines. The US ore production from 1880 –1954 is listed in Table 1.45B. Table 1.46 lists examples of product specifications for lithium carbonate from several producers.

CHEMISTRY, PHASE DATA, PHYSICAL PROPERTIES Lithium is the third element in the periodic chart, and the lightest of all metals (0.534 g/cc at 208C). Its atomic weight varies widely with the source from

Table 1.43 Table

Li2CO3 from ores Cia Brasileria de Lithio, Brazil China Russia Chemetall (Kings Mountain)a FMC (Bessemer City)a American Potasha Quebec Lithiuma Gwaliaa

1991, 307; 1996, 188 1986, 680; 1992, 1280; 1996, 1500; 2000, 1880 1986, 1020; 1990, 1350; 1996, 2250 1978, 1020; 1984, 1360; 1991, 1540 1977, 1100; 1978, 1300; 1981, 2370; 1985, 3070; 1997, 3330 1956–1963, 1100 1961, 34; 1964, 170 1996–1997, 213 (continues)

Chemistry, Phase Data, Physical Properties

A. Estimated Lithium Production Capacity, mt/yr Li (Various Reported Values) Lithium carbonate operations from brine Salar de Atacama Chemetall (Foote) 1984–1986, 1190; 1990, 1360; 1991, 2220; 1996, 2560; 2000, 2730; 2002, 3010 SQM S.A. 1997–1998, 3380; 1999, 3410; 2000, 3760; 2002, 4130 Salar de Hombre Muerto FMC 1998, 3760 (2120a as Li2CO3; 1630 as LiCl) Clayton Valley (Silver Peak) Chemetall (Foote) 1966, 1190; 1971, 1530b; 1981, 1190; 1990, 1200; 1996– 2000, 1030 Searles Lakea 1945–1978, 169 1996, 564 Pacific lithiumc Raymor industriesd 1999, 56; 2001c, 188

203

204 B. Lithium Ore Concentrates as 1000 mt of concentrates/yr Gwalia (Greenbushes) 2002, 150 (80e); 2001, 150 (67.6e); 2000, 70e; 1999, 150 (65e); 1996, 55e; 1994, 53e; 1991, 40e; 1984, 13.2e; 1983, 5e Bikita 2002, 55 (41e); 2001, 55 (49.6e); 2000, 50 (37.5e); 1999, 50 (37.7e); 1996, 25e; 1994, 23.5e; 1993, 18e; 1992, 12e; 1991, 9.1e; 1990, 19e; 1986, 27e Tanco 2002, 21 (15e); 2001, 21 (15e); 2000 (18e); 1999, 21 (19e); 1996, 22e; 1994, 20e; 1992, 18.5e; 1990, 12e; 1991, 12e; 1986, 15e Russia 2000, 75; 1996, 63 (40e); 1992, 45e; 1991, 50e; 1990, 55e China 2000, 63; 1996, 50 (16e); 1990, 16e Brazil 2002, 6 est.(6e); 2000, 6 (6e); 1999, combined with Namibia est. 2 (2e); 1990–1996, 1.6e; 1989, 2.1e; 1986, 39 Namibiaa 1996, 2 (2.5e); 1994, 1.9e; 1993, 0.7e; 1990–1992, 1.2e; 1956, 8.4e Portugal 2002 and all other small operators est. 25 (20e); 1999, 12.5; 1996, 8e; 1992– 1995, 9e; 1991, 10e; 1990, 7.6e Quebec Lithiuma 1955–1959, 2/day South Dakotaa 1942, 5e; LCA 1953, 12 t/hr Searles Lakea (licons) 1976, 0.765; 1943, 0.522; 1938, 0.200 Conversion ratios: MM lbs/yr lce/11.7346 ¼ mt/yr Li; mt/yr lce/ 5.3228 ¼ mt/yr Li; st/yr lce/5.8673 ¼ mt/yr Li; mt/yr LiCl/ 6.1077 ¼ mt/yr Li. a No longer operating. b Possible for a 2 year period. c Proposed, but not built. d Converts commercial lithium carbonate to 99.999% purity. e Reported sales, not capacity.

Part 1 Lithium

Table 1.43 (continued)

205

Chemistry, Phase Data, Physical Properties Table 1.44

Estimated World Production of Lithium Salts or Concentrates, mt of Contained Li (USGS, 2002–1958) 1999

1998

1997

1996

1995a

200 2200 32 710 5674 2346 — 140 2000 700 700

1130 2100 32 700 5326 2440 28 160 2000 700 1000

8 2800 32 1600 4551 2909 40 180 2000 1070 700

8 3700 32 690 2700 2800 48 160 800 800 500

8 3700 32 660 2600 2800 50 180 800 800 520

15,900 15,700 16,400 15,100

15,400

15,170

11,800

12,150

2002 Argentina Australia Brazil Canada Chile China Namibia Portugal Russia USb Zimbabwe Total

Argentina Australia Brazil Canada Chile China Namibia Portugal Russia USb Zimbabwe Total

200 2000 220 700 6800 2400 — 200 2000 700 700

2001 200 2400 — 700 6757 2440 — 140 2000 700 700

2000 200 2400 30 710 6732 2440 — 140 2000 700 740

1994a

1993a

1992a

1991a

1990a

1989a

1988a

1987a

8 3570 32 630 2550 2800 36 180 800 900 470

6 3560 32 590 2550 2710 14 180 800 900 360

12 2860 32 580 2660 2710 22 180 900 900 260

6 2720 22 380 2100 2710 23 200 1000 900 180

1 2690 22 380 2230 2630 24 150 1100 900 380

2 1200 32 420 1430 260 37 400 2200 1800 300

2 900 29 420 1390 260 44 310 2200 — —

1 360 36 140 910 — 64 14 1640 — —

11,480 11,700 11,100 10,200

10,500

8081c

—

—

1986a

1985

1984

1983

1982

1981

1980

1978

Argentina Australia Brazil Canada Chile China Namibia Portugal Russia USb Zimbabwe

1 357 33 16 837 835 23 — 1350 3805 534

1 363 20 10 847 835 36 1 1350 4200 530

1 217 19 3 396 835 16 7 1350 4992 425

4 81 43 — — 555 14 4 1350 4450 357

3 3 56 — — 420 19 6 1350 3468 194

1 1 57 — — 390 23 6 1250 4922 111

2 — 58 — — 390 — 7 1250 4792 405

18 — 140 — — 270 230 720 1300 5300 900

Totalc

7791

8193

8261

6858

5519

6761

6905

8100 (continues)

206

Part 1 Lithium Table 1.44 (continued) 1975

1974

1973

1972

1971

1970

Argentina Australia Brazil Canada Mozambique Namibia Portugal Russia USb Zimbabwe

9 27 136 91 9 272 9 1814 4990 862

9 27 136 82 18 272 36 1633 4717 862

9 36 136 82 18 227 36 1633 4536 862

4.5 36 136 36 18 227 23 1361 4082 862

4.5 36 109 36 9 227 18 1361 3629 862

4.5 18 54 36 9 91 9 635 3266 862

7.2 26.1 — 3.1 — 51.8 — — — —

Totalc

6260

6169

5697

5352

4808

4627

—

—

1962

1961

1960

1958

1957

1956

1955

1954

1953

Canadad Namibiad USd Zimbabwed

10 33 — 710

63 68 — 670

42 — — 1492

59 269 — 2560

77 202 — 3278

71 169 — 3096

2 256 — 2465

— 219 937 1622

— 311 675 589

Totalc

—

—

—

2906

3656

3462

2843

3083

1911

1952

1951

1950

— 294 387 44

— 355 322 79

— 294 230 6

844

834

596

d

Canada Namibiad USd Zimbabwed Totalc

1967

1966 17.2 8.7 189 5.9 2.1 66.5 — — — 1590

a

Factored from the 1996 reports of ore tonnage and tons of lithium equivalent. Estimated based upon the assumed lithium carbonate production from Clayton Valley. c Sum of the listed entries. d Estimated from the tons of lithium ore shipped, and assuming a 2.73% Li content for the ore. b

6.94 –6.99, but was averaged as 6.941 in 1995 (Coplen, 1997). Its atomic radius is ˚ , while its hydrated ionic radius is 3.40 A ˚ , and the hydration energy 519 kJ/ 0.68 A mol. At high pressure lithium changes to a strange cubic structure not seen in any element. At about 39 GPa it begins to change from its high-pressure, face-centered cubic form to a rhombohedral form. Then at about 45 GPa it converts to a new cubic structure with a large unit cell containing 16 atoms and a reduced conductivity. This form appears to be stable to 165 GPa (Anon., 2000). A comparison of some of the

Table 1.45A Recent United States Lithium Statistics, mt of contained Li (USGS, 1900–2002) 2000 Imports for consumption Exports Consumptiona Price ($/kg)b Li2CO3 LiOH·H2O Employmenta

Imports Exports Consumption Government stock pile Productiond Price ($/kg)b Li2CO3 LiOH·H2O

4.47 5.74 100

2640 1330 2800 4.47 5.74 100

1998 2590 1400 2800 4.47 5.74 100

1997 975 2200 2800 4.47 5.74 230

1996 884 2310 2700 4.34 5.51 230

1995 1140 1900 2600 4.34 5.62 230

1994 851 1700 2500 4.41 5.62 230

1993 810 1700 2300 4.21 5.71 230

1992 770 2100 2300 4.32 5.53 —

1991 590 2400 2600 4.21 5.37 —

1990

1989

1988

1987

1986

1985

1984

1983

1982

1981

790 2600 2700 —

630 2600 2700 —

1000 2300 2700 —

820 1800 2450 —

640 1800 2360 3805

373 2270 2270 4200

82 2840 2910 4992

32 2360 2000 4450

27 2080 1920 3468

136 2360 2910 4922

3.41 4.33 230

3.30 4.25 265

3.39 3.39 275

3.29 4.25 300

3.26 4.25 300

3.10 4.05

3.10 4.05

1980

1979

1978

1977

1976

1975

1974

1973

1972

1971

82 2270 4153 0 6341

45 2180 3871 0 6006

9 1820 3900 5 5711

9 1640 3483 230 5114

9 1450 2540 149 3834

82 820 2620 55 3301

64 930 3420 391 3874

120 836 2870 142 3590

27 582 2440 — 3005

120 590 2340 — 2820

2.66 3.52

2.55 3.37

2.24 3.08

1.94 2.57

1.83 2.79

1.69 2.60

1.72 1.91

1.21 1.39

1.14 1.39

Chemistry, Phase Data, Physical Properties

Imports for consumption Exports Consumptiona Productionc Price ($/kg)b Li2CO3 LiOH·H2O Employmenta

2880 1310 2800

1999

1.12 1.34

207

(continues)

208

Table 1.45A (continued) 1969

45 890 1595 2440

36 890 1146 2000

1.14 1.30 1960

Imports Exports Consumptiona Productionc Price ($/kg)b Li2CO3 LiOH·H2O a

1600 1.26 1.59

1.01 1.21 1959

— 1.47 1.59

1968

1520a 3700 0.99 1.19 1958

3000 1.47 1.41

1967

1966

1965

1964

1963

875e

232e

281e

676e

562e

—

—

2500

—

—

0.95 1.21 1957

— 1.47 1.50

Estimated. Year end listed price (perhaps up to 50% higher than the actual price). c Anstett et al., 1990. d Stinson (1981) 1974–1980; others sum of above items. e Assuming 2.76% Li ore. b

1.01 1.21 1956

— 1.96 1.76

1.08 1.19 1955

— 2.16 1.94

1.14 1.23 1954

760 907

1.17 1.19

1962

1961

—

1500

1.17 1.19

2001

2002

1990 1480 1400

1920 1620 1100

4.47 5.74

Part 1 Lithium

Imports Exports Consumption Productionc Price ($/kg)b Li2CO3 LiOH·H2O

1970

Table 1.45B The Early Production of Lithium Ores in the United States, mt of ore/yr (Johnson, 1958 through 1920; U.S. Bur. Mines, 1925–1955, 1958) 1880–1900 Very small; 1900–1917 Avg.527; 1918–1920 Avg.5,900; (mt Lia: 1899 40, 1901 130, 1910 40, 1921 140) 1925 2849

1926 3357 —

1927 3786 —

1928 4173 420

1930 1630 120

1933 457 —

1934 652 —

1935 1047 —

1936 1126 —

1937 1231 —

1938 809 —

1940 1824 (210 mt Lia

1941 3476

1942 5811 —

1943 7398 —

1944 12,083 550

1945 2,219 —

1946 2781 —

1947 2214 210

1948 3521 —

1949 4389 —

1950 8442 330

1951 11,780 —)

1953 24,712

1954 34,319

1955 ! Not published

1952 14,162 a

O’Neill et al. (1968).

1939 1805 —)

Chemistry, Phase Data, Physical Properties

1920 10,611 (1000 mt Lia

209

210

Part 1 Lithium Table 1.46 Typical Lithium Product Specifications (Harben and Edwards, 1997) MINSAL (SQM) lithium carbonate specifications

Li2CO3 Cl Na K Ca Mg SO4 B Fe2O3 H2O LOI Insolubles

Powder

Granular

99% min. 0.02% max. 0.12% max. 3 ppm max. 0.04% max 0.011% max 0.1% max. 10 ppm max. 0.003% max. 0.20% max. 0.7% max. 0.02% max.

99% min. 0.02% max. 0.18% max. 3 ppm max. 0.068% max. 0.025% max. 0.1% max. 10 ppm. max. 0.003% max. 0.2% max. 0.8% max. 0.02% max.

Specifications and typical analyses for commercial lithium compoundsa Impurities (%) Maximum

Typical production sample

Lithium carbonate, technical Moisture (loss at 1108C) SO22 4 CaO Na2O þ K2O Fe2O2 Heavy metals Chlorides

0.50 0.50 0.05 0.30 0.005 0.002 0.01

0.01 0.35 0.04 0.18 0.003 , 0.001 , 0.005

Lithium hydroxide monohydrateb Cl2 NaOH SO22 4 Fe2O3 CaO CO2 Insolubles

0.003 0.05 0.05 0.005 0.08 0.20 0.01

0.002 0.01 0.02 0.001 0.05 0.10 0.005

Lithium, chloride, technical Moisture Alkalinity as Li2CO3 SO22 3 CaCl2 NaCl þ KCl Fe2O3 BaCl2 Insolubles

1.00 0.10 0.01 0.15 0.50 0.006 0.01 0.025

0.60 0.06 0.01 0.10 0.40 0.003 0.01 0.01 (continues)

Chemistry, Phase Data, Physical Properties

211

Table 1.46 (continued) Specifications and typical analyses for commercial lithium compounds (cont.) Lithium metal Na K Ca N Fe a b

0.6 0.01 0.02 0.06 0.001

Kingsnorth, 1988. Specified minimum assay as LiOH, 53.5%: typical production sample. 55.0%.

Table 1.47A Various Chemical and Thermodynamic Properties of Lithium: A. Lithium Metal1 (CAS No. 7439-93-2) Molecular weight: 6Li 6.015; 7Li 7.016, Average 6.941 Abundance: 7.42% (6Li), 92.58% (7Li) Atomic number 3 Electron shells: 1S2 2S1 ˚ ; Mg is 1.60 A ˚ Atomic radius 1.52 (or 1.55) A ˚ (or 0.60, Mg is 0.65); Hydrated 3.40 A ˚ Covalent 1.23 A ˚ Radius: Ionic 0.68 A Atomic volume 13.10 cm3/mol Density at 208C 0.534 (33.3lb/ft3; Mg ¼ 108; Al ¼ 162lb/ft3) Crystal structure: body centered cubic Melting point 180.548C Boiling point 13428C(1336–13478C) Hardness 0.6 on Mohs scale (“it cuts like cheese”) Oxidation potential (Standard electrode potential; Li ! Liþ) 23.045 V (3.038); 0.259 g/amp hr Specific heat 0.8 cal/g/8C; at 258C 0.849 cal/g; 3.56 J/g 8C liquid at M.P. 1.05 cal/g/8C Thermal conductivity 84.8 J/m sec 8C Electrical resistivity at 208C 9.446 £ 1026; 08C 8.55 mohm-cm Vapor pressures (8C, mm Hg): 702, 0.49; 802, 2.82; 902, 12.1; 1002, 41.0; 1052, 70.5; 1077, 91.0 Heat of fusion 103.2 cal/g; 3.00(or 2.93) kJ/mol Heat of vaporization 5024 cal/g; 145.92 (135; 148) kJ/mol Ionization energy (kJ/mol): 1st 520.2; 2nd 7394.4, 3rd 11,814.6 Hydrated energy 519 kJ/mol Coefficient of expansion 6 £ 1025 cm/cm/8C; Elastic modulus 11 GPa; Rigidity 4.2 GPa; Youngs modulus 4.9 GPa Thermonuclear reaction: 6 Li þ 1n ! 3T þ 4He 4.78 MeV 7 Li þ 1n ! 3T þ 4He 2 2.47 MeV Neutron cross-section 71 barns; 6Li 945, 7Li 0.033 barns ˚ : Orange 6103 A ˚ Characteristic spectrum lines: Red 6708 A a

Data from Anstett et al., 1990; Lloyd, 1981; LCA, 1968 and others.

212

Part 1 Lithium Table 1.47B Lithium Solutions

Activity coefficient at various total ionic concentration (mol), 258C: 0.975 at 0.001 m; 0.965 at 0.002 m; 0.948 at 0.005 m; 0.929 at 0.01 m; 0.907 at 0.02 m; 0.87 at 0.05 m; 0.835 at 0.1 m; 0.80 at 0.2 m Low-solubility salts: carbonate, phosphate, fluoride and oxalate; High-solubility of halides (except fluoride) in water and polar organic solvents; High solubility of lithium-alkyls in hydrocarbons. Its chemical behavior is often very similar to magnesium. Solubility in NH3 10.17%; Reduction potential in NH3 2 2.99

Table 1.47C Comparative Atomic, Ionic and Molecular Properties of the Alkali Metals (Lloyd, 1981) Li Atomic number Electronic configuration Atomic weight Heat of atomisation from standard (kcal/mol) Heat of formation of molecules from atoms (kcal/mol) Ionisation potential for gas (kcal) eV Electronic affinity (eV) Normal electrode potential (V) Electronegativity ˚) Ionic radius (A ˚) Covalent radius (A ˚) Internuclear distance in molecule (A

3 2,1 6.941 39.0 227.2 123.8 5.36 0.54 3.038 1.0 0.68 1.58 2.67

Na 11 8,2,1 22.990 25.9

K 19 8,8,2,1 39.098 19.8

218.4

212.6

117.9 5.18 (3s) 0.74 2.71 0.9 0.97 1.92 3.08

99.7 4.41 (4s) 0.7 2.92 0.8 1.33 2.38 3.91

Rb

Cs

37 18,8,8,2,1 85.67 18.9

55 18,18,8,8,2,1 132.905

211.3

210.4

95.9 4.16

289.4 3.96

2.92 0.8 1.47 2.53

2.93 0.7 1.67 2.72 4.55

properties of the alkali metals is listed in Table 1.47, along with a discussion of lithium’s chemistry. The solubility of lithium carbonate in pure water is listed in Table 1.48, and shown in Fig. 1.103. Lithium carbonate’s solubility with sodium chloride present at 258C is listed in Table 1.49, but the data from various authors is quite variable. Seidell (1965) indicates that the solubility of lithium carbonate in water at 258C is 5950 ppm Li, and that when both lithium carbonate and sodium chloride are saturated the lithium solubility is 5200 ppm, while Deng et al., (2002) listed the latter value at 980 ppm. There is no saturated NaCl – Li2CO3 data at 1008C, but Seidell’s 258C data would extrapolate to 2500 ppm Li, Deng et al.’s, (2002) to about 500 ppm Li, while Lien (1985) found 1900 – 2100 ppm Li in the system saturated with the three salts Li2CO3, Na2SO4 and

Chemistry, Phase Data, Physical Properties

213

Table 1.47D Properties of the Alkali Metals in the Metallic State (Lloyd, 1981)

Appearance

Lattice Hardness (Mohs’ scale) Specific gravity at 08C Melting point (8C) Boiling point (8C) Heat of fusion (kcal/g atom) Effective number of free electrons per atom

Li

Na

K

Rb

Cs

Silvery white solid

Silvery white solid and liquid, purple vapour

Silvery white solid and liquid, green vapour

Silvery white solid

Silvery white solid

0.6

0.7

0.5

0.972

179.5 1336 0.69 0.55

Body centred cubic 0.4

97.8 883 0.63 1.1

0.859

0.2 1.525

1.903

63.5 762 0.57

38.7 700 0.53

29.8 670 0.50

0.97

0.94

0.85

Table 1.47E Comparative Chemistry of Lithium (Lloyd, 1981) Lithium is the leading element of the Group 1a series, and as such, exhibits in many of its properties and the same characteristics as the common alkali metals: sodium and potassium. However, in some respects it shows similarities with the alkaline earth metals, in particular with magnesium. This is manifested by the formation of a normal oxide, rather than peroxide, on reaction with oxygen decomposition of the carbonates on heating direct formation of nitrides and carbides from the elements the very low solubility of the carbonates, fluorides and phosphates the high degree of hydration of the ions solubility of the salts in polar organic solvents, such as methanol and ethers the solubility of the metal alkyls in non-polar organic solvents The low atomic weight of lithium results in its compounds bearing a higher percentage of the anion than other comparable cations. Thus the perchlorate LiClO4 and nitrate LiNO3 generate a higher proportion of oxygen per unit weight of the compound; the peroxide Li2O2 and hydroxide LiOH will absorb more carbon dioxide; the hydride LiH yields more hydrogen per unit weight than any other; the hypochlorite LiOCl will generate more free chlorine per unit weight; Li on oxidation evolves more heat — 10.25 kcal/g (Na at 2.16 kcal/g) the ionisation of Li gives the highest emf per unit weight of all metals (continues)

214

Part 1 Lithium Table 1.47E (continued)

Also, because of this effect, the lithium salts in either the fused state or the aqueous state deviate most from ideal behaviour. They depress the freezing points of fluid systems. In the fused state they are good fluxes. They reduce surface tension and viscosity and because of this they bring reactants into contact, and enhance reaction rates. The small ionic radius of the lithium atom means that its compounds with other small atoms and cations are strongly ionic in bond form. The high ionic potential results in a high energy of hydration— the ion is strongly solvated in aqueous solutions, and these show the widest deviations from the ideal. By reason of being the most electro-positive of elements, with its small size it exhibits strong covalency in many compounds. This confers a special place on lithium in the field of organic-metallic chemistry. It also leads to the solvation of the ionic compounds in organic solvents which exhibit high solubilities for the salts. Similarly, the organo-lithium compounds, for example BuLi, are readily soluble in non-polar solvents such as hexane and cyclo-hexane. Even so, the solution has sufficient anionic strength to determine stereo-specific polymerisation of isoprene and butadiene.

K2SO4 (May (1952) reported 2800). It would appear that the correct lithium solubility for the Li2CO3 –NaCl saturated system at 1008C would be somewhat over 2000 ppm Li (as is encountered in some of the commercial operations). For the system saturated with both Li2CO3 and Na2CO3 (as an approximation of the Salar de Atacama’s precipitated brine), Seidell’s three authors found 4490, 4290 and 2740 ppm Li at 258C (and perhaps about 3500 ppm Li at 1008C). Deng et al.’s (2002) saturated NaCl – Li2CO3 data at 258C indicated that the solution had a pH of 7.16, a density of 1.2204, a viscosity of 1.9575 Mpa s, and a refractive index of 1.3808.

Table 1.48 Solubility of Lithium Carbonate in Water (Seidell, 1965) Li2CO3 (g) per 100 g T (8C) 0 10 20 25 30

Li2CO3 (g) per 100 g

Water

Solution

ppm Li

T (8C)

Water

Solution

ppm Li

1.54 1.43 1.33 1.29 1.25

1.52 1.41 1.31 1.28 1.24

7062 6551 6086 5947 5761

40 50 60 80 100

1.17 1.08 1.01 0.85 0.72

1.16 1.07 1.00 0.84 0.71

5389 4971 4646 3902 3298

Density of saturated solution at 08C ¼ 1.017; at 158C ¼ 1.014. 1008C sat. with NaCl , 2900 ppm Li (2500); 1008C sat. with Na2SO4 and K2SO4 1900–2100 ppm Li (Lein, 1985); 1008C sat. with Na2SO4 and K2SO4 2800 (May, 1952).

Chemistry, Phase Data, Physical Properties

215

Figure 1.103 The solubility of lithium carbonate in water at various temperatures (Seidell, 1965).

216

Table 1.49 Solubility Data of the Liþ, Naþ/Cl2, CO23 – H2O Systems at 298.15 K (Deng et al., 2002)

No. 1 2 3 4, E 5 6, F 7 8 9 10 11 12, G 13 14 15 16 17 18, H 19, A 20, B 21, C 22, D a

Equilibrium

Liþ

Naþ

Cl2

CO22 3

H2 O

2Liþ

2Naþ

2Cl2

CO22 3

H2 O

Solid phasea

0.20 0.16 0.0062 0.092 0.044 0.044 0.00048 0.09 0.098 0.16 1.91 7.88 8.16 7.36 7.76 7.96 8.16 6.82 0.17 0.00 0.00 6.56

12.30 12.38 12.14 13.92 13.52 13.20 12.60 10.60 10.30 10.26 9.40 0.084 0.086 0.026 0.03 0.082 0.082 0.00 10.15 9.72 11.11 0.079

0.17 0.86 3.24 8.76 8.42 10.66 8.74 15.56 15.88 16.00 23.20 37.90 37.68 36.52 34.70 38.48 39.60 30.24 0.00 8.64 11.34 33.64

16.77 16.12 13.13 11.15 10.71 8.39 9.05 1.05 0.43 0.54 0.89 2.10 3.50 0.95 4.22 1.96 1.87 3.89 13.98 5.37 4.90 0.00

70.56 70.48 71.49 66.08 67.31 67.70 69.61 72.70 73.30 73.04 64.60 52.03 50.57 55.15 53.29 51.52 50.28 59.05 75.70 76.27 72.65 59.73

5.11 4.11 0.17 2.14 1.07 1.09 0.01 2.74 3.06 4.91 40.23 99.68 99.68 99.89 99.88 99.69 99.70 100.0 5.26 0.00 0.00 99.64

94.89 95.89 99.83 97.86 98.93 98.91 99.99 97.26 96.94 95.09 59.77 0.32 0.32 0.11 0.12 0.31 0.30 0.00 94.74 100.0 100.0 0.36

0.85 4.32 17.28 39.94 39.96 51.80 44.98 92.59 96.93 96.16 95.66 93.84 90.10 97.02 87.43 94.33 94.70 86.79 0.00 57.65 66.19 100.0

99.15 95.68 82.72 60.06 60.04 48.20 55.02 7.41 3.07 3.84 4.34 6.16 9.90 2.98 12.57 5.67 5.30 13.21 100.0 42.35 33.81 0.00

1390.5 1394.6 1501.7 1186.6 1258.2 1295.9 1411.1 1704.0 1762.2 1729.2 1049.2 507.5 476.3 577.1 528.8 497.5 473.7 667.5 1805.0 2004.3 1670.4 699.4

CA þ LB CA þ LB CA þ LB CA þ CB þ LB CA þ CB CA þ LB þ NaCl CB þ NaCl LB þ NaCl LB þ NaCl LB þ NaCl LB þ NaCl LA þ LB þ NaCl LA þ LB LA þ LB LA þ LB NaCl þ LA NaCl þ LA LA þ LB CA þ LB CA þ CB NaCl þ CB NaCl þ LA

LA, LiCl·H2O; LB, Li2CO3; CA, Na2CO3·10H2O; CB, Na2CO3·7H2O; letters are invariant points.

Part 1 Lithium

Ja˜necke index/(mol/100 mol dry-salt)

Composition of liquid phase (mass %)

217

Chemistry, Phase Data, Physical Properties Table 1.50 Solubility of Lithium Chloride in Water (Seidell, 1965)

T (8C)

Density of sat. sol.

22.4 29.0 223.0 236.0 250.0 262.0 266.0 273c 275.9a 280b 263.0 260.4 258.0 257.0 265.6a 268b 254.0 248.0 231.0 219.2 215.6 220.5a 220 0 5 10 15 18.5 12.5b 19.0a 19.1c 25 30 40 50 60 70 80 90 96 93c 100.5b

1.268 — 1.279 — 1.293 — — — 1.296 — 1.303 1.308 — — 1.331 1.342 1.347 — —

LiCl (g) per 100g sat. sol.

Solid phase

4.0 8.0 14.0 18.0 21.0 24.0 24.4 24.85 25.0 25.3 26.4 28.2 29.6 30.4 — 28.7 30.5 30.8 33.4 36.4 37.2 — 36.9

Ice Ice Ice Ice Ice Ice Ice þ LiCl·5H2O Ice þ LiCl·5H2O Ice þ LiCl·5H2O Ice þ LiCl·5H2O LiCl·5H2O LiCl·5H2O LiCl·5H2O LiCl·5H2O þ LiCl·3H2O LiCl·5H2O þ LiCl·3H2O LiCl·5H2O þ LiCl·3H2O LiCl·3H2O LiCl·3H2O LiCl·3H2O LiCl·3H2O LiCl·3H2O þ LiCl·2H2O LiCl·3H2O þ LiCl·2H2O LiCl·3H2O þ LiCl·2H2O

40.9 42.0 42.7 43.8 45.35 40.5 — — 45.85 46.3 47.3 48.3 49.6 51.1 52.8 54.8 56.1 — 56.5

LiCl·2H2O LiCl·2H2O LiCl·2H2O LiCl·2H2O LiCl·2H2O þ LiCl·H2O LiCl·2H2O þ LiCl·H2O LiCl·2H2O þ LiCl·H2O LiCl·2H2O þ LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O þ LiCl LiCl·H2O þ LiCl LiCl·H2O þ LiCl (continues)

218

Part 1 Lithium Table 1.50 (continued)

T (8C)

Density of sat. sol.

LiCl (g) per 100g sat. sol.

94a 97 98 100 110 120 130 140 160

— — — 1.347 — 1.344 — 1.339 —

— 56.8p 57.4p 56.2 56.7 57.2 57.6 58.0 59.2

Solid phase LiCl·H2O þ LiCl LiCl·H2O LiCl·H2O LiCl LiCl LiCl LiCl LiCl LiCl

p

metastable. Data from different authors.

a–c

Table 1.51 The System Lithium Chloride–Sodium Chloride–Water (Seidell, 1965) Results at 258Ca Grams per 100 g sat. sol.

Grams per 100 g sat. sol.

Grams per 100 g sat. sol.

LiCl

NaCl

LiCl

NaCl

LiCl

NaCl

45.8 45.5 41.3 40.1 36.8

0.0 0.5 0.4 0.2 0.3

35.7 33.5 33.5 31.6 24.9

0.3 0.4 0.3 0.8 2.3

17.4 16.9 6.5 0.0

7.3 8.4 19.0 26.4

Results at 408Cb Grams per 100 g sat. sol.

Grams per 100 g sat. sol.

LiCl

NaCl

Solid phase

LiCl

NaCl

47.98 46.51 44.76 41.60 33.96

0.0 0.68 0.82 1.04 3.17

LiCl·H2O LiCl·H2O þ NaCl NaCl þ , 4% LiCl NaCl þ , 4% LiCl NaCl þ , 4% LiCl

25.48 17.52 6.14 0.0

5.26 10.13 19.05 26.65

a, b

Data from different authors; , 4% LiCl is an unknown double salt.

Solid phase NaCl þ NaCl þ NaCl þ NaCl þ

, 4% LiCl , 4% LiCl , 4% LiCl , 4% LiCl

219

Chemistry, Phase Data, Physical Properties Table 1.52 The System Lithium Chloride–Magnesium Chloride–Water (Seidell, 1965) Grams LiCl per 100 g sat. sol.

Grams MgCl2 per 100 g sat. sol.

Solid phase

Results at 2508C 31.0 29.5 28.4 25.4 23.6 22.6 15.6 17.6 18.6 20.4 21.0

0.0 2.0 4.0 6.6 8.6 8.4 7.8 6.6 4.8 1.4 0.0

LiCl·3H2O LiCl·3H2O LiCl·3H2O þ DS DS DS þ MgCl2·12H2O MgCl2·12H2O MgCl2·12H2O þ Ice Ice Ice Ice Ice

Results at 2308C 34.4 33.0 32.0 28.6 27.0 21.0 16.4 4.4 3.4 0.0 0.0 7.0 11.0 16.0

0.0 2.2 4.0 7.4 10.4 12.8 15.4 22.2 22.8 21.6 19.4 11.6 6.8 0.0

LiCl·3H2O LiCl·3H2O LiCl·3H2O þ DS DS DS þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O þ MgCl2·8H2O MgCl2·8H2O MgCl2·8H2O þ MgCl2·12H2O MgCl2·12H2O Ice Ice Ice Ice

Results at 2108C 8.8 6.4 4.0 0.0

0.0 4.0 7.0 11.6

Ice Ice Ice Ice

Results at 08C 38.8 38.2 38.0

0.0 2.6 4.8

LiCl·2H2O LiCl·2H2O LiCl·2H2O þ 1:1:7 (continues)

220

Part 1 Lithium Table 1.52 (continued)

Grams LiCl per 100 g sat. sol.

Grams MgCl2 per 100 g sat. sol.

Solid phase

Results at 08 (cont.) 32.5 29.6 28.0 23.0 18.6 13.2 5.4 0.0

8.6 11.0 12.4 14.2 17.0 21.6 28.6 35.0

1:1:7 1:1:7 1:1:7 þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

Results at 258C 45.65 43.1 40.0 40.2 38.0 37.1 35.4 33.9 29.3 p 28.8 28.0 22.0 18.9 8.4 0.0

0.0 2.68 5.74 5.68 6.99 7.42 8.32 9.34 13.7 14.2 14.3 18.0 20.0 28.3 35.4

LiCl·H2O LiCl·H2O LiCl·H2O þ 1:1:7 LiCl·H2O þ 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

Results at 308C 46.2 39.9 39.0 35.4 35.3 34.1 33.2 31.7 29.3 26.6

0.0 6.27 6.81 8.84 9.14 9.93 10.6 10.7 13.5 15.7

LiCl·H2O LiCl·H2O þ 1:1:7 1:1:17 1:1:17 1:1:17 1:1:17 1:1:17 1:1:17 1:1:17 1:1:17 (continues)

Chemistry, Phase Data, Physical Properties

221

Table 1.52 (continued) Grams LiCl per 100 g sat. sol.

Grams MgCl2 per 100 g sat. sol.

Solid phase

Results at 308C (cont.) p

26.6 25.7 24.8 17.1 10.1 0.0

16.6 16.5 16.8 21.7 27.3 35.6

1:1:17 þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

Results at 708C 51.2 46.5 39.4 38.2 37.3 36.1 26.4 22.3 21.6 20.5 p 20.0 19.0 18.8 15.3 10.2 0.0

0.0 4.58 11.2 12.3 13.9 14.5 20.4 23.5 24.1 24.4 26.4 26.1 26.3 27.9 31.2 38.7

LiCl·H2O LiCl·H2O LiCl·H2O LiCl·H2O þ 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 1:1:7 þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

Results at 1028C 56.9 30.4 29.2 22.5 19.2 17.4 13 9

0.0 23.4 23.9 29.7 34.1 35.6 39 41

LiCl LiCl LiCl LiCl LiCl LiCl þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

1:1:7 ¼ LiCl·MgCl2·7H2O; DS, Double salt of undetermined composition; p extrapolated value.

222

Part 1 Lithium Table 1.53 Solubility of Lithium Sulphate in Water (Seidell, 1965)

T (8C) 21.735 23.30 25.11 27.04 29.67 214.65 218.45 221.4 223.0 216.0 213.0 211.5 26.5 0.0 0.0 0.6 12.5 14.0 16.7 19.6 20 25 25 25 27

Li2SO4 (g) per 100 g sat. sol.

Solid phase

T (8C)

Li2SO4 (g) per 100 g sat. sol.

Solid phase

4.072 7.791 11.30 14.33 17.67 21.95 24.85 27.1 27.9 27.32 27.24 27.18 26.73 25.431 26.332 26.51 25.986 26.07 25.96 25.85 25.205 25.502 25.793 25.697 25.4111

Ice Ice Ice Ice Ice Ice Ice Ice Ice þ Li2SO4·H2O Li2SO4·2H2O Li2SO4·2H2O Li2SO4·2H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O

30 30 31.8 35.0 38.0 43.7 45.6 50 51.6 52.4 55 65.7 71.8 77.0 94.8 94.9 95.2 100.1 103.0 104.0 142.5 186 214 232.8

25.258 25.1012 25.47 24.763 25.28 25.00 24.8811 24.39 24.82 24.71 24.622 24.34 24.210 24.05 23.76 23.4 (1.182)4 24.2110 23.5 (1.179)4 23.72 23.55 (1.176)4 22.65 22.7 23.0 ,23

Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O Li2SO4·H2O þ Li2SO4

1 to 12 are data from different authors.

The solubility of lithium chloride in water is given in Table 1.50, and lithium chloride with sodium chloride in Table 1.51. The phase system lithium chloride with magnesium chloride is listed in Table 1.52, and lithium sulfate in water in Table 1.53. The more complex reciprocal salt pair of LiCl – Li2SO4 – MgSO4 – MgCl2 is plotted in Fig. 1.104. In the evaporation of more complex brines such as from the Salar de Atacama, the solubility data initially follows the seawater evaporation system (see Garrett, 1996) until carnallite crystallizes, and then more closely follows the lithium systems, with or without sulfate and/or magnesium, depending upon the solar pond process being employed. Campbell and Kartzmark (1956 –1961) have reported on the five component system: lithium, sodium, potassium, sulfate, chloride and water, and each of its subsystems.

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Figure 1.104 Phase diagram at 0, 35 and 508C of Li, Mg, Cl, SO4 –H2O (Garrett and Laborde, 1983). Reprinted by permission of the Salt Institute.

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Deng, T. L., Yin, H. A., and Tang, M. L. (2002). “Experimental and Predictive Phase Equilibrium of the Li, Na, Cl, CO3 System at 298.158K.” J. Chem., Eng. Data 47(1), 26– 29. Garrett, D. E. (1996). “Potash; Geology, Processing, Uses and Phase Data”, 734 pp. Chapman & Hall, London. Garrett, D. E., and Laborde, M. (1983). “Phase Diagrams at 08C, 358C and 508C of Li, Mg, Cl, SO4, H2O.” Sixth Int. Salt Symp. 2, 424 pp. LCA (1968). “Lithium.” Lithium Corporation of America Brochure, 16 pp. Lien, R. H. (1985). “Recovery of Lithium from Clay.” “Lithium”. (R. O. Bach ed.), pp. 61–71. Lloyd, J. E. (1981). “Lithium Chemicals.” Speciality Inorg. Chem. 40, 98–122. May, F. H. (1952). “Process for Producing Lithium Sulfate from Lithium Phosphate.” U.S. Patent 2,608, 465, 7 pp. (Aug. 26). Seidell, A. (1965). “Solubilities”, Lithium, Vol. 2, pp. 362–440. American Chemical Society, Washington, DC.

Part 2

Calcium Chloride

GEOLOGY Calcium is one of the most common elements found in nature, occurring most frequently in gypsum or anhydrite, and also in many other rocks and minerals and as one of the common ions in water and many brines. However, calcium in the form of calcium chloride is relatively uncommon. The only calcium chloride mineral that is found in massive quantities is the double salt tachyhydrite (CaCl2·2MgCl2·12H2O), occurring in three large potash deposits and as isolated crystals in a few other locations. Only one other mineral is known, antarcticite (CaCl2·6H2O), which is exceedingly rare and only crystallizes in relatively small quantities in one Antarctic pond and in one desert lake (Dunning and Cooper, 1969; Torii and Ossaka, 1965). It has been speculated to be present as isolated crystals in a few other locations such as in the occlusions of some minerals, and small amounts of calcium chloride tetrahydrate (CaCl2·4H2O; sergipite?) may possibly occur in the Sergipe tachyhydrite deposit (Sonnenfeld and Kuehn, 1993). Calcium chloride brines on the earth’s surface are almost as rare, only being found in one pond and two dry lakes as a strong CaCl2 brine, but there are several lakes with dilute CaCl2 brines, several with CaCl2 – MgCl2 brines, and a modest number of dilute CaCl2 springs (Stankevich et al., 1992; Shearer, 1978). In contrast to this scarcity on the earth’s surface, calcium chloride brines are frequently found in several types of underground, and one type of undersea formations. Most potash and some halite formations are associated with strong calcium chloride brines. They are found in porous strata either below, adjacent to, or above the deposits, gradually becoming more dilute if they approach the surface, and modified somewhat in proportion to their distance from the potash or salt deposit. They were formed by a “dolomitization reaction” with calcite and the high-magnesium chloride end liquor (or carnallite [KCl·MgCl2·6H2O] decomposition brine) from the potash (or other highly evaporated brine) deposits (Garrett, 1996). 237

238

Part 2 Calcium Chloride

Dilute calcium chloride brines are also occasionally found in coastal aquifers, and some oil or gas formation waters that have been formed from seawater by the same dolomitization reaction supplemented by the leaching of certain types of rocks. In addition, many deep sea geothermal vents contain calcium chloride that appears to have been formed by both the dolomitization process and extremely hot seawater or brine reacting with basalt or other rocks. Because of the acidic nature of the calcium chloride brines, they usually have been modified by the products of various other brine – rock reactions, and occasionally from Na/K-for-Ca ion exchange with clays or zeolites (Sanford et al., 1992; Lebedev, 1969). The occlusion or fracture waters in some hard-rock mineral deposits (and some other geologic formations) also contain a calcium chloride solution that may have been formed by the same reactions, or have been an original magmatic fluid. These various calcium chloride occurrences will be discussed in more detail in the following sections. Calcium Chloride Dolomitization Brine As seawater (and many non-marine waters) evaporate, calcium carbonate is the first major mineral deposited, and it thus is present in large quantities under or adjacent to any halite and potash deposit. In addition, since fresh seawater periodically flooded into the evaporating halite or potash basins, some calcite is present in, and usually over the deposit. Gypsum and then salt crystallize next in the seawater evaporation sequence, and finally potash, which theoretically crystallizes first as leonite (K2SO4·MgSO4·4H2O), then schoenite (K2SO4 · MgSO 4 ·6H2 O), kainite (KCl·MgSO4 ·2.75H 2O) and finally as carnallite (KCl·MgCl2·6H2O). The first two potash salts are either minor occurrences, or do not form at all, and the first three are usually converted by stronger brines (from continued evaporation) into carnallite. It, in turn is usually converted to sylvite (KCl) by later being leached with a slightly more dilute brine (H2O þ KCl·MgCl2·6H2O[carnallite] ! KCl[sylvite] þ [MgCl2 þ 7H2O][brine]). Both the evaporation and carnallite conversion produce a strong magnesium chloride end liquor that could be flushed away, but usually slowly seeped through the deposit. During this seepage, high-calcium chloride brines were then formed by the partial-to-complete dolomitization reaction of the end liquors with the basins’ abundant calcite: MgCl2 ðin brineÞ þ 2CaCO3 ðcalciteÞ ! CaCl2 ðbrineÞ þ CaCO3 ·MgCO3 ðdolomiteÞ: If there was a large excess of magnesium chloride, or the brine was very hot, the final product could be magnesite (MgCO3). Much of the impurities in the calcite (if any) such as strontium, barium and manganese would also be reacted (Ayora et al., 2001; Vlasova and Valyashko, 1980). To a much lesser extent, calcium chloride may also be formed by reactions to produce other seawater

Geology

239

salts, polyhalite or glauberite: ðMgCl2 þ 2KCl þ 2H2 OÞðbrineÞ þ 4CaSO4 ðanhydriteÞ ! K2 SO4 ·MgSO4 ·2CaSO4 ·2H2 OðpolyhaliteÞ þ 2CaCl2 ðbrineÞ; and ðNa2 SO4 þ 2NaClÞðbrineÞ þ 3CaSO4 ·2H2 OðgypsumÞ ! 2Na2 SO4 ·CaSO4 ðglauberiteÞ þ ðCaCl2 þ 4H2 OÞðbrineÞ: The extent of each reaction gives the converted brine a variable calcium chloride concentration, depending upon the amount of calcite (or gypsum) contacted. In addition, both the high-MgCl2 and high-CaCl2 brines are acidic, and they usually react to some extent with the surrounding rocks. They may also be diluted by groundwater, and if they encounter a high-sulfate water gypsum (or anhydrite) would precipitate (Poroshin, 1981). Seawater contains 4.4 times as much MgCl2 as KCl, and most of the MgCl2 does not remain with the potash or halite deposits. Epsomite (MgSO4·7H2O) is the first of the magnesium salts to crystallize, (it deposits with the late-stage halite) and it is almost always leached, or later converted to gypsum or anhydrite. Kainite (KCl·MgSO4·2.75H2O) is only a major component in very few potash deposits, but carnallite (KCl·MgCl2·6H2O) is much more common. However, in both cases the amount of magnesium that is present in most deposits is far less than the seawater ratio of magnesium to potassium or halite. It is likely that only at the end of deposition could much of the magnesium chloride be flushed away from the depositing salts, so consequently in almost all of the deposits most of the magnesium must have been removed from the formation as an end liquor which seeped away, and thus formed these massive dolomitization brine aquifers. As noted above, this type of brine is by far the most common calcium chloride occurrence, and it accompanies most of the world’s potash, and some of the latestage halite deposits that contain up to , 200 ppm bromine (indicating that the seawater evaporation extended to the point of forming a strong [. 16% MgCl2] end-liquor and almost producing potash; Table 2.1). This end liquor often accumulated in porous strata below or adjacent to the deposit, and perhaps was later forced into strata in and above the potash or halite as the sediments were compacted. Typical analyses of seawater at various stages of evaporation are listed in Tables 2.1 and 2.2, and compared with the formation fluids and occlusions actually found in a halite deposit adjacent to potash. The formation fluids are usually somewhat of an average of the end liquor formed between halite’s initial crystallization and the first kainite that deposited. However, as with most formation waters, there has been some dilution, and in some cases, such as the Carlsbad example of Table 2.2, it apparently came from a geothermal source supplying additional boron and iodine. The adjacent Carlsbad potash deposit

240

The Composition of Sea Water End Liquor at Various Stages of Evaporation (wt.%) (Garrett, 1996) Data from small experimental ponds Carnallite

Na K Mg Cl SO4 Brc Bc Ic Cac Src TDS Density pH Concentration ratiose

Brine in Carlsbad, New Mexico Halitea,b

First NaCl

First MgSO4

First kainite

First

Last

Average

Range

8.68 0.31 1.04 15.60 1.53 535 38 0.49 338 6.7d 27.26 1.219 — 8.1

6.10 0.79 2.64 15.10 3.70 1350 96 1.3 166 3.3d 28.56 1.245 — 20.6

1.65 2.00 5.90 15.20 6.20 3460 244 3.2 — — 32.90 1.313 — 52.5

0.59 1.70 6.90 17.70 4.70 4580 326 4.2 — — 33.10 1.326 — 70

0.24 0.52 7.80 20.50 3.60 5520 392 5.1 — — 33.30 1.340 — 84

8.10 1.76 2.25 18.98 1.94 1547 1716 16.7 310 2.35 30.57 1.223 6.04 23.7

5.08–9.87 1.51–2.27 1.62–4.43 17.30 –20.20 1.52–3.31 1280–2410 1420–1888 12.2–21.0 260 –381 0.87–6.90 27.30 –31.90 1.21–1.25 5.6 –6.4 20 –37 (continues)

Part 2 Calcium Chloride

Table 2.1

Table 2.1 (continued ) Carnallite First NaCl

First MgSO4

First kainite

First

Last

Bromine in the crystals of:c Halite Sylvite Carnallite

29 — —

Rubidium in the crystals of: Kainite Sylvite Carnallite

In actual deposits 85 — —

190 2500 —

250 3300 2400

300 4000 2900

35 –100b —

20 –450 f 250 –4000 340 –3460

— — —

— — —

— — —

20 50 430

10 –60 10 –480 60 –1900

a

Abitz et al. (1990). Deposits with no potash. c In ppm. d Estimate. e Concentration factor based upon bromine (multiple of the original sea water concentration). f Deposits with potash. b

Geology

241

242

Table 2.2 Examples of Potash Deposit Dolomitization Brine (Garrett, 1996)

Depth (m)

SO4

Cl

HCO3

Ca

Mg

Salinity (g/liter)

pH

170.8a –203 220.5a –236 260a –335 344.4a –386 412a 428.5b –517.5 542.2b,c – 781

65 616 3900 7130 9590 5720 860

88 562 3180 7890 83,500 178,000 257,000

164 101 91 251 110 0 455

60 156 482 740 1430 2100 95,300

14 86 284 469 966 1110 17,600

0.45 1.88 10.9 23.5 151.5 301 408

7.0 7.0 6.8 7.2 7.1 6.0 —

(B) Brine analyses in clays within and near the Stebnik deposit (g/kg) (Valyashko et al., 1973)

Br

Mg

Na

K

Total salts

Density (g/ml)

(1) Brines within the Deposit (Sea water end liquor) Tr. 3.37 192.1 Tr. 14.23 181.7 0 25.78 200.8

2.86 2.50 2.82

47.1 46.8 65.0

21.8 22.2 17.9

26.2 23.8 10.1

293.5 291.2 322.3

1.25 1.25 1.277

(2) Brines near the Deposit (dolomitization brine) 51.3 0.40 192.6 62.0 0.10 189.0

3.83 4.40

20.1 15.90

18.7 17.2

18.5 9.3

305.7 298.1

1.273 1.258

Ca

Cl

SO4

(C) Brine analyses in clays within the Prypat and lrkutsk deposits (g/kg) (Azizov, 1974) Deposit

SO4

Cl

Br

Ca

Mg

Na

I (in ppm)

K

Sr

Total salts

Density (g/ml)

pH

Pripyat lrkutsk

0.08 0.71

171.2 264.7

2.32 5.73

49.5 118.1

5.3 11.4

39.6 7.9

46 2

66.7 15.2

1.48 4.24

284 425

1.24 1.41

6.4 4.6 (continues)

Part 2 Calcium Chloride

(A) Average brine analyses in the Starobin Deposit’s overlying formations (mg/liter) (Zatenatskaya et al., 1968)

(D) Analyses of a brine pocket in Zechstein A-1 anhydrite (adjacent to Ca2 dolomite; 408 m depth; Norwegian– Danish Basin) (wt.%) (Fabricus, 1983) NaCl

KCl

CaCl2

MgCl2

Salinity

pH

3.95

3.55

15.15

8.10

31

3.18

(E) Brine analyses near the Paradox Basin potash deposit (ppm) (Mayhew and Heylmun, 1966) Mississippian Formation Ca Average Range

4560 240 –12000

Na

Mg

Cl

SO4

Br

Fe

TDS

pH

120,100 110,000–140,000

2870 270–7500

191,600 153,000–220,000

5090 1800–7400

240 —

550 90 –1000

319,300 252,000 –374,000

5.5 4.6– 6.7

Paradox Formation Low average High average

Ca

Na

Mg

K

Li

Cl

Br

SO4

HCO3

B

TDS

pH

5030 42,700

22,200 99,900

4300 32,900

21,400 52,500

66 173

120,900 217,200

1320 2660

255 4390

208 1260

702 1630

274,668 38,2651

4.9 6.2

Also present in some samples Sr PO4 1300

1010

NH4

Fe

F

I

Mn

Rb

Al

Zn

Cs

Cu

Pb

Ba

850

750

280

264

260

95

66

60

16

6

6

0

(F) Salt Range, Dhariala, Pakistan (Gill and Akhtar, 1982). Well at 1200 m depth, flowing at 1,500 bbls/hr at 721 psi. Similar brine at 1347 m and 2439 m depth

a

MgCl2

KCl

NaCl

4.6

16.5

6.7

5.7

Siltstone. Argillaceous marl. c Halite with seams of marl and sylvite from 525 –1262 m (bottom of drill hole). b

Geology

Average (wt.%)

CaCl2

243

244

Part 2 Calcium Chloride

Figure 2.1 Map of the Carlsbad, New Mexico potash deposit (Austin and Barker, 1990; reprinted by permission of the Society of Economic Geologists).

(Fig. 2.1) is known to have had a geothermal input because of its hightemperature minerals (such as langbeinite, K2SO4·2MgSO4), and various heavy metal impurities. Similar formation and occlusion brine analyses in Carlsbad halite have been given by Stein and Krumhansl (1988) and others, in halite near the Stebnik (Table 2.2) and in the Pripyat Deep, Russian potash deposits by Valyashko et al. (1973) and Poroshin (1981), respectively, and in other deposits by many authors. The diluted dolomitization brine may also have later dissolved varying amounts of halite, potash or other soluble salts that it encountered. Several examples of the dolomitization brines that have been formed from these end liquors are listed in Tables 2.2 and 2.3. In some areas, such as the Michigan (Figs. 2.2– 2.4) and Paradox Basins (Fig. 2.5) in the United States, most of this brine appears to still be very near the potash/halite deposits, while in others its former presence is well documented (e.g., in Saskatchewan; Fig. 2.6; Table 2.3) but much of it has migrated some distance away or has escaped.

Table 2.3 Average Brine Analyses Encountered in Saskatchewan Potash Mines, (mg/liter)a (Garrett, 1996) Mine, sample location

Depth (m)

dD

d 18O

61 –73

2 147

219.1

888

1300

84

28

10

250 2 128 (122 –378)

216.1

2430

3950

44

145

41

220

2 138

217.6

1660

2630

64

137

Rocanville Shaft Cory Shaft

486 –572

2 126

214.8

29,300

16,800

256

524 –673

2 127

215.4

30,900

19,600

Allan Shaft

610 –774

2 138

217.5

41,900

Average

630

2 130

215.9

Rocanville Shaft Allan Shaft Cory Shaft

604 –890

2 70

823 –860 647 –975

Average

Recent glacial drift Upper section of shafts Average

Na

K

Ca

Mg

Br

SO4b

HCO3

pH

Density

18.5

1460

311

7.4

0.997

54

59

5350

151

7.65

1.001

26

30

39

3410

231

7.5

0.999

1089

647

110

23

—

100

7.0

1.032

958

2177

681

16

—

151

7.1

1.036

27,800

1406

2200

448

175 (1 at 12.900) 223

—

100

7.4

1.050

34,000

21,400

873

1820

592

169

18 (1 at 155) 19

6220

117

7.2

1.039

24.1

159,000

108,000

1830

4320

1530

1250

74

—

64

6.7

1.166

2 116 2 103

213.5 29.0

171,000 182,000

122,000 88,600

— (4000)

2 96

28.7

171,000

(18) 160 (1 at 1464) 81

6.2 6.7 (1 at 5.7) 6.8

1.194 1.158

790

112 108 (1 at 508) 98

930 1200

2 98 2 94

28.8 26.5

128 63

5.8 5.2

1.229 1.232

1065

2 96

27.7

96

5.5

1.231

5210 (21,930)

106,000

1990 1590 (1 at 16,300) 1800

218,000 222,000

55,500 84,600

6700 48,800

220,000

70,000

27,800

59,600 7090 (1 at 18,900) 33,300

4770

6.5

I

2420 1670 4680 1820 (1 at 28,400) (1 at 21,700) 2880 1580 11,000 9770

4080 4650

302 82

10,400

4370

192

—

250

1.173

Geology

Cory Shaft Rocanville Shaft Average

Cl

(continues)

245

246

Mine, sample location

Depth (m)

dD

d 18O

Cl

Na

K

2 7.8

274,000

11,100

12,300

Mg

Br

I

SO4b

HCO3

pH

Density

6110

85,600

8920

65

(1200)

(73)

4.6

1.257

(750) (73)

4.3 4.45

1.280 1.269

Ca

Lanigan Mine Allan Mine Average

1200

2 88

1200 1200

2 60.5 2 3.15 274 2 5.5

293,000 284,000

11,900 11,500

31,300 21,800

37,000 21,600

113,000 99,300

4070 6500

— 65

Lanigan Mine Cory Mine Average

1200

2 62

2 3.9

325,000

2750

7100

141,000

782

20,500

782

—

—

4.4

1.286

1200 1200

2 62 2 62

2 1.85 2 2.9

338,000 332,000

4620 3690

15,700 11,400

152,100 147,000

21,400 11,100

19,400 20,000

406 594

— (300)

138 138

4.4 4.4

1.351 1.319

Values in parenthesis imply limited number of analyses or data; not considered in the average. Reproduced from Potash; Deposits, Processing, Properties and Uses, Table 1–7, pages 51–52, q1996 with kind permission of Kluwer Academic Publishers. a Wittrup and Kyser (1990). b Calculated by ion balance difference.

Part 2 Calcium Chloride

Table 2.3 (continued )

Geology

247

Figure 2.2 Estimated area and thickness of the Michigan Basin potash deposit (Matthews and Eagleson, 1974, courtesy of the Northern Ohio Geological Society).

The dolomitization reaction had been observed to take place relatively rapidly and completely in modern evaporating seawater tidal flats when calcite sand or mud is present, the brine had evaporated to a fairly high strength, and the brine pool was quite warm (Levy, 1977; Kinsman, 1966; others). Normally, however, for seawater or similar low-concentration brines, the dolomitization reaction appears to take place only at very high temperatures (e.g., deep sea vents), or with exceedingly long residence times and some temperature elevation (e.g., in some coastal aquifers or oil field brines). For the more concentrated calcium chloride brines associated with potash or late-stage halite deposits, there is also usually a very direct correlation between dolomite formations and the calcium chloride brine. An example of this is with

248

Part 2 Calcium Chloride

Figure 2.3 Area of the Detroit River brine aquifers, and depth to the top of the formation (Sorensen and Segall, 1974, courtesy of the Northern Ohio Geological Society).

the very large Paradox Basin halite and potash deposits in the south-central section of the United States (Fig. 2.5; Utah, Colorado and New Mexico), where most of the original calcite layers in the deposit have been converted to dolomite. In addition, the brine occlusions in secondary halite is of the calcium chloride type (Hite, 1983). This very large deposit generated massive amounts of calcium chloride brine, found primarily in various porous strata under or near the deposit, but some of the high-CaCl2 aquifers are scattered over a great distance. The brine’s concentration and composition vary widely, depending upon the extent of dolomitization, its mixing with other waters, and the leaching of other minerals. As with most other calcium chloride brines, there are a large number of minor elements present (Table 2.2). It is also not uncommon for the dolomitization brine to have been forced into porous formations in or above the potash or halite deposits, as indicated by

Geology

249

Figure 2.4 Area of the Sylvania and Filer Sandstone aquifers, and their thickness (Pavlick, 1984).

the Michigan Basin in the United States (Table 2.4), the Starobin deposit in Russia (Table 2.2), and the gigantic Saskatchewan potash deposit in Canada (Table 2.3). In these cases, there is almost always a progressive dilution of the brine with meteoric water as it is found in shallower aquifers, and the sulfate content in near-surface water is greatly diminished because of increased reaction with the brine’s calcium ions. Also, the isotopic signature of the water and various ions gradually changes from being predominately of a marine to a meteoric source. When the dolomitization brine is found in halite or sylvite occlusions, these minerals are always secondary, and have been recrystallized after the dolomitization brine was forced back up through the deposit. In a similar manner, when it is found within the deposit’s void spaces, there also always is a progressively more dilute dolomitization brine at the higher elevations.

250

Part 2 Calcium Chloride

Figure 2.5 Map of the Potash and Halite Deposits in the Paradox Basin (Hite, 1961; reprinted by permission of the United Nations Economic and Social Commission for Asia and the Pacific).

Michigan, USA Brines This very large halite/potash deposit has many aquifers that contain a highcalcium chloride brine (Table 2.4; Figs. 2.3 and 2.4) both above and below the Silurian Salina Group’s halite (NaCl) and 33,700 km2 (13,000 mi2) potash (sylvinite, a mixture of NaCl and KCl) deposits. The major aquifers are the overlying Devonian carbonate and sandstone beds, but there are also many lesser aquifers (Fig. 2.7). There is an extensive area of strong calcium chloride brine directly over the potash deposit and extending to the south southeast (Fig. 2.4) in the first porous bed above the potash (the Sylvania Sandstone Formation) that is at

Geology

251

Figure 2.6 Map of the Prairie Evaporite Formation and the Saskatchewan Potash Deposit (After Fuzesy, 1982; reprinted courtesy of the Saskatchewan Department of Energy and Mines).

nearly the same concentration as potash end liquor. The small Filer Formation of sandstone to the northwest contains a similar, but slightly more dilute brine. Several thinner and less abundant aquifers also occur under the potash beds with equally strong, or stronger calcium chloride brines (Table 2.4). Each of the aquifers have roughly the same ratio of salts, but as they approach the surface become progressively more dilute. It would thus appear that in this basin, the potash end liquor that originally seeped through and under the potash deposit (and reacted with calcite) was much later forced from its original sediments as they were compressed by deep burial (perhaps aided by the thick glacial ice that formed over this basin) into the overlying porous strata. Their variable mineral content, as seen in Table 2.4, resulted from their considerably different migration history and variable dilution by meteoric or other groundwater (as is strongly indicated by the brine’s deuterium and 18O analyses), precipitation (such as gypsum), and their different contact with rocks that they could partially leach or react with. However, in this basin these reactions were limited, since the porous (average , 20%) carbonate strata contains fairly pure carbonates, and the sandstone strata fairly pure silica (quartz arenites) cemented by dolomite

252

Table 2.4 Brine Analyses in Various Michigan Aquifer Formations

Formation (increasing depth ! )

NaCl CaCl2 MgCl2 Br I (ppm)

Parma

Marshall

15 0.5 0.5 0.05 nil

13 10 3.5 0.15 nil

Berea

Traverse

Dundee

Sylvania

Filer

20 4 2 0.1 3

10 15 3 0.2 20

13 9 2.6 0.1 5

5 19 3.6 0.26 30

2 17 10 0.25 nil

(B) Range of concentrations, wt.% (McKetta and Cunningham, 1975) CaCl2

NaCl

2 – 23

2– 20

MgCl2

KCl

Br2

2 –4

1–2

0.1–0.3

(C) Sylvania Formation (wt.%) (1) Egleson and Querio (1969) ppm

Average Range

CaCl2

NaCl

MgCl2

KCl

SrCl2

Br

NH3

LiCl

B

SO4

I

Rb

TDS

Density

20.09 19.88 –20.75

5.33 4.82–5.65

3.61 3.40– 3.83

1.52 1.08– 1.80

0.56 0.52– 0.64

0.281 0.259– 0.293

398 250– 488

351 220–440

273 236– 410

44 36– 54

34 18– 46

15 10 –19

31.50 30.6– 33.1

1.2914 1.2788–1.2959

(2) Sylvania and Filer Formations (Wilson and Long, 1992) ppm

Average 25 samples

Ca

Na

6.10

2.50

Mg (in ppm) 9440

K (in ppm) 7820

Cl

Br

HCO3

SO4

17.20

2070

73

48

TDS 27.75

(continues)

Part 2 Calcium Chloride

(A) The major calcium chloride formations (wt.%) (Pavlick, 1984)

As compounds

CaCl2

NaCl

MgCl2

16.90

6.36

3.69

KCl 1.49

Br

Density

0.207

1.25 (assumed)

(D) Various minor formations (wt.%) (ppm; Martini, 1997)a

Berea Sandstone Niagara Limestone St. Peter Sandstone

Ca

Na

Mg

K (in ppm)

Sr (in ppm)

Ba (in ppm)

Cl

Br (in ppm)

pH

4.78 9.75 7.43

7.14 2.49 3.35

1.08 1.13 0.62

1790 4030 3070

1860 4030 3150

1910 901 16,140

21.31 26.55 21.34

1770 2850 2180

4.55 4.53 5.3

(E) Some of the Devonian Aquifers (Wilson and Long, 1993)b (ppm) Formation

Ca

Na

Mg

K

Sr

NH4

Li

Rb

Cs

Cl

Br

Berea Traverse Dundee DetroitR. Richfield

37,200 26,500 22,600 63,800 51,800

57,800 53,900 62,950 28,270 32,400

5780 4980 4770 9270 6810

528 1356 1410 9790 5570

1640 981 757 1820 1930

66 130 103 453 202

6.6 28 24 65 38

3.0 4.4 3.5 30 9.6

2.4 2.4 1.8 2.7 4.2

160,000 143,600 146,100 167,000 162,200

1200 1030 901 2470 1570

Berea Traverse Dundee DetroitR. Richfield

HCO3

I

B

Si

TDS

pH

Depth

8C

No.

dD

d 18O

54 143 187 0.55 69

16 33 34 140 138

19 17 12 30 17

– 46 27 207 111

2.8 3.7 2.5 1.6 2.3

259,290 232,070 236,530 282,530 245,390

4.9 5.3 5.0 4.5 4.7

755 805 960 1057 1277

26.4 27.8 31.4 33.8 38.7

3 11 28 3 10

217 237 239 234 240

0.7 2 2.1 2 2.5 3.5 2.9

87

Sr/ 86Sr

0.70912 0.70898 0.70845 0.70782 0.70814

The analyses appear to be in error since the anions far exceed the cations. The Berea Sandstone is gas-bearing. The Niagara Limestone is below the halitepotash deposit; the St. Peter Sandstone is much deeper. b No. is the number of samples averaged together; depth is in meters; DetroitR. is Detroit River.

Geology

a

SO4

253

254

Part 2 Calcium Chloride

Figure 2.7 Stratigraphy of the Michigan Basin (After Wilson and Long, 1993, 1992; reprinted from Water–Rock Interaction by permission from Swets & Zeitlinger publishers; reprinted from Applied Geochemistry, Vol. 8, p. 83, Fig. 2, q1993, with permission from Elsevier).

or quartz (Martini, 1997). The variability in dolomitization brine composition seen here is also typical of the world’s other dolomitization brine formations. There is a general synclinal structure to the strata under the Michigan Basin, and

Geology

255

examples of the specific stratigraphy to the southeast of the center of the basin at Midland are shown in Fig. 2.7 (Egleson and Querio, 1969). The Detroit River Group consists of 0– 350 m of variable porosity carbonates, and at its base there is 0 – 90 m of porous sandstone called the Sylvania Formation. Each of these formations cover about 40% or more of the Michigan Basin, and contain strong calcium chloride brines at depths of 300 – 1400 m. Their brines have been commercially recovered in the past, and were generally only considered to be economic below about 880 m. The brines’ total dissolved solids (TDS) and the amount of CaCl2 increases fairly consistently with depth from 3 to 23% CaCl2, and the NaCl and MgCl2 concentrations vary inversely with the CaCl2. In the Sylvania Formation, the CaCl2 usually ranges from 14 to 22%. Additional information on the brine in other aquifers and the various reactions and changes that have occurred with them has been given by Martini and Wilson (1997) and Wilson and Long (1993, 1992). The Michigan Basin brines’ very low pH helps to explain their ability to leach and react with other rocks, as is indicated by their high contents of strontium, barium and other metals, although much of the Sr and Ba probably came from the reaction with calcite. Geothermal water also probably mixed with some of the formations, as indicated by the variable presence of iodine, boron, lithium, cesium, rubidium and other rare metals. With most of the brines, the calcium concentration is somewhat higher than its magnesium equivalent in seawater end liquor from a potash deposit, and the potassium a little lower. Wilson and Long (1993) speculated that this occurred by the conversion of the clays kaolinite and smectite to illite: 2Kþ þ CaCO3 þ 3Al2 Si2 O5 ðOHÞ4 ðkaoliniteÞ ! 2KAl2 ðAlSi3 ÞO10 ðOHÞ2 ðilliteÞ þ Ca2þ þ 4H2 O þ CO2 ; or : 3½Ca0:33 ðMg0:33 Al1:66 ÞSi4 O10 ðOHÞ2 ·nH2 OðsmectiteÞ þ 2Kþ þ Al3þ þ CaCO3 ! 2KAl2 ðAlSi3 ÞO10 ðOHÞ2 ðilliteÞ þ 2Ca2þ þ Mg2þ þ ðn þ 1ÞH2 O þ CO2 þ 6SiO2 : Small amounts of glauberite (CaSO4·Na2SO4) and polyhalite (2CaSO4·K2SO4·MgSO4; both of which can also produce some calcium chloride) have also been found in the basin. Finally, some of the calcium chloride aquifers have a slightly elevated ratio of 87Sr/86Sr (they vary from 0.7080 to 0.7105; seawater is 0.70919, further indicating that there was some rock leaching [Martini, 1997]). Various Other Dolomitization Brines Even though most of the world’s many potash deposits, as well as some of the even more common halite deposits, have formed dolomitization brine, in

256

Part 2 Calcium Chloride

many cases the brine has migrated away, or has not been examined and reported. Consequently, only a few examples of the more well-known brines will be discussed. In Canada, studies have been made on the dolomitization brine found in the potash mines of south-central Saskatchewan (Fig. 2.6), and at higher elevations in their mine shafts (Table 2.3). Later Bernatsky (1998) reviewed the sampling of various aquifers in southern Saskatchewan, and found that most of them contained a seawater, calcium chloride-type brine that correlated well with the Prairie Evaporate Formation brine of the potash deposits. In both the mines and aquifers, the shallower brines were progressively diluted with meteoric water, and some of the calcium was precipitated by the surface water’s sulfate content. All of the brine contained ions from the leaching of other rocks, some of the deeper brine showed signs of having leached halite, and occasionally it appeared to have increased in strength by shale adsorption or other means. In what is perhaps an unrelated study, Bottomley et al. (1999) noted what appears to be strong seawater dolomitization brine in the Canadian Shield (at a gold mine) far to the north in the Northwest Territories toward the presumed seawater entrance of the potash deposits. The Canadian Shield brine had similar calcium chloride, bromine, strontium and lithium concentrations, as well as d34S, d6Li and deuterium values as in the potash deposits. Since very little of the Saskatchewan potash deposit end liquor dolomitization brine is now directly under, above or adjacent to the deposits, it would appear that most of it became diluted in the area’s aquifers, or traveled long distances from its origin. The Zechstein Formation is a series of massive halite – potash deposits extending from England, through Norway, France, Germany, and ending in Poland. It had 4 major periods of potash deposition, each of which have supported a number of potash mines (Garrett, 1996). In the last (youngest), or Zechstein 4 period in North Yorkshire (Boulby), England, a high-calcium chloride dolomitization brine oozes from cuts in the Boulby halite that underlies the potash deposit (that is being mined; Talbot et al., 1982). In this area of England, there are also high-calcium chloride brines (up to 16.1% salts) in several aquifers such as the upper Carboniferous Coal Measures, often in the lower Carboniferous, and occasionally in the Devonian Old Red Sandstone (Anderson, 1945). In the German Zechstein potash deposits (Fig. 2.8), Herrmann (1980) has examined over 190 brine samples from mine water or aquifers in 6 mining districts. Of these samples, 14 contained 5 – 25%, and 2 contained 40 – 50% CaCl2 (of the total salts present), each with a typical dolomitization composition, but differing degrees of conversion and dilution. The other samples had the composition of seawater end liquor or of having leached various minerals in the formations. Many authors have also noted dolomitization brine in the occlusions of secondary halite in the Zechstein (Wolfgramm and Schmidt, 2000; Zwart and Touret, 1994 [7 – 15% CaCl2]; Fabricus, 1983). The Zechstein dolomitization brine has found its way into some of the country’s deeper aquifers such as in southern Bavaria (Udluft, 1976), and also formed a number of calcium chloride springs.

Geology

257

Figure 2.8 Map of the German Potash Deposits (Garrett, 1996; reproduced from Potash; Deposits, Processing, Properties and Uses, Fig. 2.16, p. 112, q1996 with kind permission of Kluwer Academic Publishers).

At Stecklenberg in the Harz mountains, a dilute calcium chloride brine rises from underlying aquifers through a network of fissures to the surface (Haller and Mestwerdt, 1938; Harrassowitz, 1935). Two other probable Zechstein formation dolomitization brines are the strong calcium chloride subsurface waters of the Permian and Triassic Sudeten Foreland monocline in Czechoslovakia (Depowski and Llaszcz, 1968), and the similar brines that occur in occlusions in recrystallized Upper Triassic halite of the Lorraine Basin in France (Fanlo and Ayora, 1998). In the Paris Basin, the Triassic Keuper halite contains small amounts of trapped seawater dolomitization brine that has been enriched in Li, B and Sr (Fontes and Matray, 1993). In India, at the eastern end of the Salt Range at Dhariala in the Punjab region, typical low-conversion dolomitization brine (Table 2.2) has been found at a depth of 1370 m (Gill and Akhtar, 1982). The Salt Range is a narrow plateau with an

258

Part 2 Calcium Chloride

average elevation of 610– 910 m, extending about 240 km from Jhelum to the Indus river. There is a massive halite deposit under the Salt Range, containing numerous beds of potash (sylvinite and mixed sulfates; Alam and Asrarullah, 1973). Levy (1977) found a similar brine in the coastal sabkhas of the northern Sinai of Israel. It had been formed by the reaction of evaporated seawater with the fine grained calcite in the playa muds. The former Russian states have a large number of potash and halite deposits (Fig. 2.9), and calcium chloride brines have been associated with many of them. Sturua (1974) has presented a map of various Russian calcium chloride groundwater occurrences, and the majority of them are closely grouped near major potash deposits. In the Caspian Depression, Moskovskiy and Anisimov (1991) have reported such brines, and the Carpathian group’s Stebnik potash deposit in south-central Russia (which contains potassium sulfate minerals, and not the normal sylvinite), still has dolomitization brine near the deposit (Valyashko et al., 1973). Similar brines have been found with 80– 169 g/liter Ca in the Ukraine’s Dnieper– Donets Basin, grading to 50, then 6 and finally 1.5 g/liter CaCl2 as the

Figure 2.9 Map of several of the Russian potash deposits (Garrett, 1996; reproduced from Potash; Deposits, Processing, Properties and Uses, Fig. 2.20, p. 123, q1996 with kind permission of Kluwer Academic Publishers).

Geology

259

brine approaches the surface (Petrichenko and Shaydetska, 1999, 1998; Lyalko and Tereshchenko, 1973). The adjacent Belarus Pripyat Trough (or Deep; Azizov, 1975, 1972) also contains calcium chloride brines with a total salts content of up to 300 g/kg beneath the area’s halite and potash deposits. A detailed study of inclusions in the halite has shown that only typical end liquor is found in originally deposited salt, but that recrystallized salt contains high-calcium chloride brine with a very reduced sulfate content. For the latter salt, there was also evidence of adjacent dolomite formations and an excess of gypsum. They suggested that the reaction: MgSO4 ðbrineÞ þ 2CaCO3 ðcalciteÞ ! MgCO3 ·CaCO3 ðdolomiteÞ þ CaSO4 ðgypsumÞ accompanied the dolomitization reaction, since some MgSO4 is always present in seawater end liquors (Poroshin, 1981). In the Kama River area near Perm, the aquifers’ salinity increases with depth, and their analyses become more typical of dolomitization brine (Baldina and Sverdlov, 1959). In the South Siberian Platform, the groundwaters appear to be dolomitization brines, but in one occurrence the deuterium content was nearly that of meteoric water (Pinneker et al., 1968). Near numerous intruding basaltic sills the halite has been reported to contain inclusions with the solids CaCl2, CaCl2·KCl and nCaCl2·mMgCl2 (Grishina et al., 1992). In the Tungussky Basin of the Siberian Platform, dolomitization brines were found with an additional content of leached rock minerals (Shvartsev and Bukaty, 1996). In the Angara-Lena artesian basin of the Irkutsk oblast in Siberia, the deepest waters are only saturated with salt, but at higher elevations they change to the dolomitization type with up to 25.4% CaCl2 and 37.3% TDS, and then gradually become more dilute as they approach the surface. All of this basin’s calcium chloride waters are unusual in containing up to 600– 2000 ppm of hydrogen sulfide and 1600 –1900 ppm Li (Pinneker, 1967; Ryabtsev et al., 2002). Very strong calcium chloride brines were also found directly under the Irkutsk halite –potash deposit (Azizov, 1972). Under the cis-Ural region’s halite – potash deposits, both primary seawater end liquor and dolomitization brine (100 – 150 g/liter CaCl2; 200– 280 g/liter total salts) have been found. The amount of dolomite in the region roughly corresponds to the amount of CaCl2 formed, and the reaction has created enough porosity in the dolomite to provide space for a large petroleum reservoir (Popov, 1988). In the United States, the seawater end liquor in the large Carlsbad potash deposit (Fig. 2.1) has been previously noted (Table 2.1), and the brines in the Salado Formation below and adjacent to this deposit have been studied by Jones and Anderholm (1995). Similar studies have been made by Graf et al. (1966) on the calcium chloride brines in the Illinois Basin, and its adjacent Michigan Basin (discussed above). The two basins have a sequence of fresh to calcium chloride brine varying with depth, with several indicators showing that the deep brine had

260

Table 2.5 Analyses of a Few Oil and Gas Field Brines (wt.% or ppm)a

Field

Ca

England, Apedale well 3.51 Michigan, Freda sandstone 1.60 England, Renishaw well 1.36 1.03 Texas, Spur; Cisco Fm.b Oklahoma, Tulsa Co. (2)b 0.95 Texas, Mid-Continent Field 0.74 Ventura, Shiells’ Canyon (2)b 0.70 Ventura, CA So. Mtn. Fld.b 0.55 Ventura, CA So. Mtn. Fld. 0.56 Ventura, CA Wiley Canyonb 0.53 San Joaquin, CA Midway Fld. 0.29 Appalachian, Penn., Ohio, W. Va(3)b 0.26 Fresno, CA Seabord Fld.c 0.22 Ka251, Baa42, NH4a31, Bra108, NO3a44, SiO2a36, Ia21, Ba8.3, Fa0.6, pH 6.4 Texas, Mid-Continent Field 0.18 Israel (46)d 0.166 Bra138, Sra63, Sia10.4, Ia12, Lia20, d6Li 2 23.5, 1549 m depth

Mga

Na

Cl

SO4a

120 30 3000 879 1350 2900 — 43 190 40 7 180 832

4.23 0.76 3.74 2.10 3.39 3.70 0.24 0.40 0.55 0.46 1.05 0.50 1.48

13.16 4.04 9.06 4.41 7.37 7.88 1.56 1.61 1.89 1.60 2.10 1.31 2.90

270 — 450 155 440 tr. 144 49 0 315 31 42 0

16 — — tr. 653 156 0 346 293 — 193

21.02 6.52 14.50 9.45 11.95 12.63 2.57 2.61 3.01 2.66 3.48 2.13 4.70

464 536

1.11 1.936

2100 186

313 257

3.58 5.66

Cl

2.00 3.413

HCO3a 0

TDS

(B) Land and Mcpherson (1992) Gulf of Mexico Sedimentary Basin (1) Cenozoic Formation Waters Ca

Na

Ka

Mga

Lia

Sra

Baa

Zna

Fea

Mna

High-Ca; 3 wells Average; 40 wells Low-Ca; 6 wells

2.48 3.02 3.10

1580 911 505

149 376 443

33 32 10

1770 1210 564

629 492 70

7.8 25 1.2

33 57 23

33 14 3.5

3.75 1.78 0.84

10.67 8.03 6.64 (continues)

Part 2 Calcium Chloride

(A) Lane (1927)

Br

I

B

Org.e

469 245 123

86 22 13

73 73 65

240 287 240

(2) Same Area, Shallower, NaCl-type Waters Ca Na Ka

Mga

Lia

Sra

Three Formationsh

819

(ppm)

SO4

High-Ca Average Low-Ca

4.7 26 9.8

0.21

Si 33 33 21

3.52

4.1

153

I

B

Org.

92

30

38

855

(3) Cretaceous Reservoirs, Texas Ca Na

Ka

Mga

Lia

Sra

Sixteen wellsi

887

899

132

1852

SO4 4.1

1.14 SO4a 65

24

3.90 a

Si 49

a

Br

I

a

a

B

426

65

197

d13C

d18O

17.14 13.21 10.79

— 26.5 27.2

4.8 5.6 4.0

Baa

Zna

61 e

Br

(ppm)

Si

180

TDS%

0.2 13

Fea 10

d C

d O

9.83

23.3

2.2

Zna

137

69

TDS%

18

d O

11

d B

14.01

29.5

17

Srf

8C

mg

25 25 27

0.7088 0.7086 0.7080

127 132 104

3572 3565 2881

Mna

Cl

0.7 18

TDS%

Baa

d11B

95 Sr

Sr2

8C

mg

36

0.7083

80

2435

Mna

Cl

45

8.48

8C

mg

d B

Fea

f

0.7087

5.87

11

146

3689

(C) Oil field brines with a high bromine content, wt.% (evidently mixed with potash or halite end liquor; Collins, 1970) Mg

Na

Ka

Sra

Lia

Cl

Bra

SO4a

Ba

Ia

TDS

Location in Figs. 2–10

7.0 6.0 4.0 3.5 3.0 3.0

1.5 1.0 0.5 0.4 0.9 0.6

1.4 2.8 6.6 7.2 5.5 7.3

8000 40 1.0% 2000 2500 600

1500 3000 2000 — — —

60 10 100 90 40 15

20.0 17.9 19.8 18.6 16.6 18.4

2500 3200 700 1800 1200 200

1200 1200 300 12 1200 180

300 40 — — 90 60

40 40 20 20 25 20

31.15 28.37 32.29 30.09 26.42 29.39

10 1 4 13 3 8

261

(continues)

Geology

Michigan Utah N.Dakota W.Virginia Texas Louisiana

Ca

262

Table 2.5 (continued ) (C) Oil field brines with a high bromine content, wt.% (evidently mixed with potash or halite end liquor; Collins, 1970) Mg

Na

Ka

Sra

Lia

Cl

Bra

SO4a

B

Ia

TDS

Location in Figs. 2–10

3.0 3.0 2.0 2.0 1.4 1.0 3.23

0.5 0.5 1.1 0.5 1.1 600a 568

7.4 6.8 0.8 5.8 1.6 5.1 4.16

700 4000 — 3000 9000 100 3700

900 — — 1000 800 1000 1280

25 100 — 10 — 5 263

18.2 17.5 7.9 14.3 9.0 9.8 12.71

2000 5000 — 1300 1500 600 —

1800 180 1050 180 60 90 100

— — — — — 10 91

40 10 — 30 40 1000 10

32.74 28.72 11.84 23.14 14.24 16.23 20.71

9 7 5 12 11 6 —

Values in parenthesis imply number of wells. a In ppm. b Hudson and Taliaferro (1925); So. Mtn. Fld. K 110 (in ppm), Br 67 (in ppm), I 13 (in ppm); Wiley Canyon B 101 (in ppm), Br, I 0; 408C; 1515 m depth. c White et al (1963). d Bentor (1969), Chan et al (2002) (Heletz field). e Organic acids. f 87 Sr/86Sr. g Depth, meters. h Oligocene Frio, Miocene Frio, Pliocene –Pleistocene (the latter also contained 300 ppm HCO3). i Also 296 ppm HCO3. j An average of 11 wells in the Smackover Formation, ppm: Ca 29,070, Na 73,400, K 5060, Li 365, Rb 11.2, Cs 6.1, F 4.6, density 1.229 g/cc (Collins, 1976). k Holdorf et al (1993). r , 1.20. When evaporated, as wt.%: Ca 12.12, Sr 0.495, Na 0.37, Mg 0.295, K 0.15, Li 0.098.

Part 2 Calcium Chloride

Kentucky Arkansasj Kansas Pennsylva. Ohio Oklahoma Germany, Altmark

Ca

Geology

263

a seawater source. In the Paradox Basin’s large halite– potash deposit (Fig. 2.5; with 33 NaCl –potash cycles) the correlation of the amount of dolomite near the deposit with the amount of CaCl2 in the underlying aquifer is also in approximate agreement (Williams-Stroud, 1994).

Calcium Chloride in Oil and Gas Field Brines Calcium chloride occurs occasionally in oil and gas fields brines since the oil and gas are also of marine origin. The brines vary from being more dilute than seawater to usually not more than 6 times more concentrated. Normally, the calcium content is far greater than the magnesium, and the sulfate content is very low, as seen by some typical examples in Table 2.5. Also, as shown in Fig. 2.10, calcium chloride oil and gas field brines may be found throughout the United States, as well as throughout the world. These waters appear to be typical dolomitization brines, originating from the seawater that was trapped with the petroleum deposit as it formed. Possibly, in the coastal shallows near to the area accumulating the organic debris that later formed the oil and/or gas the seawater would have become somewhat concentrated, and then because of its increased

Figure 2.10 Location of some of the United States’ oil and gas field calcium chloride brines (Collins, 1970; excerpted by special permission from Chemical Engineering September 21, 1970, copyright q1970, by Chemical Week Associates, New York, NY 10038). (See Table 2.5 for brine analyses).

264

Part 2 Calcium Chloride

density, flowed downward and covered the debris. Its higher concentration may have even helped to preserve the organics from bacterial decomposition. Later, as the organics were converted to oil or gas and the mixture migrated into its final “trapped” formation, the accompanying seawater would have occasionally contacted enough calcite, and been warm enough to nearly complete the dolomitization reaction. Much of its sulfate content would have also precipitated as gypsum (as governed by the Caþ þ SO22 4 X CaSO4 solubility product). Very long residence times would have usually been involved, allowing even seawater that was diluted during its travels to have completed the dolomitization reaction. In the migration, the now-acidic brine would have also reacted with or leached other minerals, adding many minor constituents to it. Most observers believe that the brine’s occasionally high-iodine content would have come from the marine organic debris that formed the petroleum deposit. Collins (1975) has reviewed the general geochemistry of these oil field waters, and called attention to a few (in 1970) that definitely appear to be potash or halite end liquor dolomitization brines (Table 2.5(C)). Both bromine (in Arkansas and Michigan) and iodine (in California and Michigan) are or have been commercially recovered from such brines. A few other oil or gasfield brines appear to have contacted geothermal or magmatic brines because of their high boron and/or lithium (and perhaps heavy metal) content. Land and Macpherson (1992) have made a detailed study of the calcium chloride-containing oil field brines in Texas and the entire U.S. Gulf Coast area (Table 2.5). All of these waters appear to be dolomitization brine that was, on average 2.4 –5.5 times more concentrated than seawater, based upon the very similar Na, Cl, K and Br ratios to seawater for each oil or gas field. Some brines had a fairly high iodine content, and comparatively high B, Ba, Fe, Li, Mn, Si, Sr and Zn values. The boron isotope d11B numbers were in a range only found in geothermal brines (Garrett, 1998), and were very similar to the boron found in the underlying or nearby Louann Salt mass, which was also known to have had a geothermal component. These geothermal brines probably also contributed the lithium and silica (as silicic acid) to the oilfield waters, as well as some of the other metals. Land and Macpherson (1992) also postulated that the depth of burial heated these calcium chloride brines sufficiently so that they could leach some components from various rocks that they contacted. When magmatic rocks such as feldspars were present and the brine warm enough, it could do some selective leaching and convert much of the plagioclase (including its anorthite component) into the less soluble albite. The Gulf Coast Basin has sandstone units that are feldspar-rich (e.g., the Frio sandstone contains up to 30% feldspars, of which 70% are plagioclase, having an average content of 20% anorthite), and widespread albitization in these aquifers is well documented. When such reactions take place, they can produce additional calcium chloride and liberate strontium, barium

Geology

265

(in addition to that from the calcite) and other metals from the plagioclase: CaAl2 Si2 O8 ðanorthiteÞ þ 4H4 SiO4 ðsilicic acidÞ þ 2NaCl ! 2NaAlSi3 O8 ðalbiteÞ þ CaCl2 þ H2 O (where other metals sometimes substitute for the Ca. The silicic acid is perhaps hypothetical, but dissolved silica is present in many of these waters). This type of reaction would help explain the higher amounts of equivalent magnesium (now as calcium and other divalent metals) in the oilfield brine to that accounted for by the dolomitization reaction with calcite (and other carbonates) and seawater. The 87 Sr/86Sr ratios in the brines are indicative of both seawater and volcanic rocks, while the d18O and d13C values are similar to seawater that has undergone some dilution with meteoric or geothermal water. Many other articles have been written on the calcium chloride in oil and gas field brines, such as the general articles on calcium chloride brine – rock leaching chemistry by Martini (1997), Rankin (1990), Liu (1982), Collins (1975) and Lane (1927), on the fluorine and boron content of such brines (Tageeva, 1942), and the mechanism of sulfate removal from ground or seawater by calcium chloride brines (Gavriell et al., 1995; Hite, 1983; Rotkin et al., 1973; Valyashko et al., 1973). Other articles have discussed calcium chloride brines in specific oil and/or gas fields. In the Presqu’ile dolomite barrier reef gas fields of British Columbia, Canada, there is a close correlation between the reservoir’s dolomitic cement and the calcium content of the gas field brines. In the shallow fields, the dolomitization reaction is speculated to have occurred at about 508C, and in deeper fields at 1508C (Morrow et al., 2002). The Heletz –Kokhav oilfield in Israel is an example where the brine in the Heletz field only averages 2.15 times the concentration of seawater (based upon the very consistent average of Na, Cl and Br), and dolomitization has only slightly taken place. The concentration of Ca (0.15%) is about twice the seawater ratio even after reacting with about 97% of the sulfate, and the Mg (.06%) is about half of the seawater ratio. The adjacent Kokhav field’s brine has only concentrated 1.38-fold, but the Ca enhancement (0.09% Ca) and Mg and SO4 depletion are about the same. Both fields have about 5 times more lithium than would be found in seawater (2.0 and 1.5 ppm, respectively; seawater is 0.17%), and d6Li values of 2 23.5 and 2 19.7 compared to seawater’s 2 32.3. This would appear to indicate some dilution of the oilfield brine with geothermal water (or perhaps some unusual rock-leaching), but other indicators such as temperature, boron, fluorine and heavy metals were not analyzed to help make this determination (Chan et al., 2002). Holdorf et al. (1993) reported on high-calcium (38.8 g/liter), and unusually high-lithium (0.315 g/liter) brines in the Altmark gas field in Germany (also containing 110 mg/liter B and 12 mg/liter I). Aren and Depowski (1965) noted the calcium chloride brines in the Podlasie Basin gas field of Poland. Evans (1991) reviewed the hydrocarbon formation and Ca-brine migration in the Appalachian

266

Part 2 Calcium Chloride

Basin, USA, while Hudson and Taliaferro (1925) considered the calcium chloride brines of the Ventura, California oil fields. With these brines, there was only a very rough correlation of the calcium content with depth, and the brines varied widely in their calcium content and degree of dolomitization over relatively short distances within the field (see two examples of this with the South Mountain Field in Table 2.5). A strong calcium chloride and calcium fluoride brine has been reported in the Mobile Bay gasfield (Schutz et al., 2000). Klosterman (1981) studied the strong calcium chloride brine inclusions in calcite of the Smackover Formation in the Walker Creek and Mount Vernon oil fields of southern Arkansas, USA. Dodonov et al. (1948) considered the order of crystallization from calcium chloride brines in the Carboniferous gas-bearing formation of the Saratov District in Russia, while Romanyuk et al. (1973) discussed the calcium chloride brine anomalies in the Paleogene gas condensate deposits in Rassolnaya and Kosmach, Russia. Chochia (1972) noted the calcium chloride formation waters and speculated on the amount of oil and gas occurring in the Zyryanka Basin, Russia. Shtogrin (1971) discussed the chemistry of the calcium chloride brines in both the gas-bearing and non-productive structures of the eastern part of the North Crimean Trough (in Russia), while Kolodiy and Dobrov (1969) noted that the Turkmenian calcium chloride oil and gas field brines contained less deuterium than the near-surface waters. Krotova (1959) discussed the relationship of deep, strong calcium chloride brines with the amount of oil and gas in Siberia.

Calcium Chloride in Geothermal Brines Deep Sea Vents A third well-documented occurrence of calcium chloride is in the brine being emitted from a number of deep sea vents in rift or subduction zones. There are also a few inland geothermal springs containing calcium chloride, and in one location, the Salton Sea Trough in California and Mexico (also a rift zone) there is a massive deep, very hot, concentrated calcium chloride brine formation. The composition of these brines is quite complex, although all have either a known seawater origin (the deep sea vents) or a possible seawater connection. The exact reactions between the seawater and the rocks that they contacted is not known, but as with the oil field brines, in most cases it may be assumed that there was an initial contact with calcite. The resulting acidic and very hot brine probably then went through reactions such as the plagioclase (including its anorthite component) conversion into the less soluble albite, and the leaching of feldspar and other rocks (Shemiakin and Korotkov, 1979). Among the specific reactions that might occur, Von Damm et al. (1985) have suggested the following for the liberation or removal of several ions.

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(1) Liberating calcium: (a) (as noted above) CaAl2 Si2 O8 ðanorthiteÞ þ Naþ þ SiðOHÞ4 ! NaAlSi3 O8 ðalbiteÞ þ Ca2þ þ AlðOHÞ2 4 : (b) Ca2 Al2 Si3 O12 ðOHÞðepidoteÞ þ 3Mg2þ þ 2Fe2þ þ 9H2 O þ ! Mg2 Fe3 Al2 Si3 O10 ðOHÞ8 ðchloriteÞ þ 2Ca2þ þ AlðOHÞ2 4 þ 7H

(2) Freeing Na: NaAlSi3 O8 ðalbiteÞ þ 3Mg2þ þ 2Fe2þ þ AlðOHÞ2 4 þ 6H2 O ! Mg3 Fe2 Al2 Si3 O10 ðOHÞ8 ðchloriteÞ þ Naþ þ 8Hþ (3) Removing Ca: Ca2þ þ CaAl2 Si2 O8 ðanorthiteÞ þ AlðOHÞ2 4 þ SiðOHÞ4 ! Ca2 Al2 Si3 O12 ðOHÞðepidoteÞ þ Hþ þ 3H2 O

(each of the above minerals are members of the greenschist group). Pastushenko (1967) has proposed a more elaborate series of reactions that might produce calcium chloride: 1. Clinozoisite and epidote can replace plagioclase in vitreous tuffs and volcanic ash, forming sericite scales and calcite grains: 6CaAl2 Si2 O8 ðanorthiteÞ þ KCl þ H2 O þ CO2 ! 2Ca2 Al3 Si3 O12 ðOHÞ þ 2KH2 Al3 Si3 O12 þ CaCl2 þ CaCO3 2. Plagioclase (anorthite) is replaced by sericite, albite and kaolin: 3CaAl2 Si2 O8 ðanorthiteÞ þ 4NaCl þ H2 O þ CO2 þ 8SiO2 ! 4NaAlSi3 O8 þ H4 Al2 Si2 O9 þ CaCl2 þ CaCO3 ; or : 5CaAl2 Si2 O8 ðanorthiteÞ þ 2KCl þ 4NaCl þ 4H2 O þ 2CO2 þ 6SiO2 ! 2NaAlSi3 O8 þ 2KAl3 Si3 O10 þ 2NaAlSi2 O6 ·H2 O þ 3CaCl2 þ 2CaCO3

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3. Reactions with analcite to form hydrous glass: CaAl2 Si2 O8 ðanorthiteÞ þ NaAlSi3 O8 þ SiO2 þ 2NaCl þ 3H2 O ! 3NaAlSi2 O6 ·H2 O þ CaCl2 4. The conversion of amesite and chlorite into pennine and clinochlore: CaAl2 Si2 O8 ðanorthiteÞ þ NaAlSi3 O8 þ H4 Mg2 Al2 SiO9 þ H4 Fe2 Al2 SiO9 þ 4NaCl þ 4H2 O ! 3Na2 Al2 Si3 O10 ·2H2 O þ H4 ðMgFeÞAl2 SiO9 þ CaCl2 þ MgCl2 þ FeO 5. Stronger calcium chloride solutions to form silica: 2NaAlSi3 O8 þ 2H2 O ! Na2 Al2 Si3 O10 ·2H2 O þ 3SiO2 ; or : 4CaAl2 Si2 O8 ðanorthiteÞ þ 6NaAlSi3 O8 þ 8NaCl þ 14H2 O ! 7Na2 Al2 Si3 O10 ·2H2 O þ 4CaCl2 þ 5SiO2 6. Reactions accompanied by the adsorption of silica: CaAl2 Si2 O8 ðanorthiteÞ þ 2NaCl þ 2H2 O þ SiO2 ! Na2 Al2 Si3 O10 ·2H2 O þ CaCl2 ; or : 6CaAl2 Si2 O8 ðanorthiteÞ þ NaAlSi3 O8 þ 2SiO2 þ 6NaCl þ 15H2 O ! 3Na2 Al2 Si3 O10 ·2H2 O þ NaCa2 Al5 Si5 O20 ·6H2 O þ CaAl2 Si3 O10 ·3H2 O þ 3CaCl2 :

Many authors have also proposed different reactions for the contact of very hot seawater in the rocks of deep sea vents, including Hardie (1990) who considered the primary reaction to be with basalt to form spilitic greenstone. With vents and various other hot calcium chloride brines Shvartsev and Bukaty (1996), Stura (1977, 1974), Azizov (1975), Kissin and Pakhomov (1969), and Pastushenko (1967) have proposed reactions with a wide variety of other rocks. Several authors have also discussed rock leaching, adsorption, precipitation or ion exchange to add or remove components in the dolomitization brine.

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Figure 2.11 Lava flows and vents in a portion of the East Pacific Rise (Von Damm et al., 1985; reprinted from Geochimica et Cosmochimica Acta, vol. 49, q1985, with permission from Elsevier).

Deep sea vents have been found in many locations where the earth’s crustal plates are rifting apart, or one is being subducted under the other, as shown in Fig. 2.11. Table 2.6 lists the analyses of several of these vents, including deep wells in Iceland that are in a continuation of East Pacific Rise rift system (Fig. 2.12). In all cases, it is quite certain that the vents emit seawater that has seeped deep into the earth’s interior through cracks in the active zone, been heated to very high temperatures, and then forced back up through other fissures to form the vents. For instance, the deep temperatures in an Indonesian vent’s cores varied from 220 to 3488C, and the NaCl þ CaCl2 content of the deep brine from 2 to 25 g/liter (Moore et al., 2000). The vent’s composition is that of a seawater dolomitization brine, and based upon the analysis of some of the crustal material (e.g., in Iceland or from drill cores), probably most of the seawater has contacted calcite in its travel path. However, the deep basalt and fissured rocks in the vents have also been considerably altered, thus indicating that there was considerable leaching and reaction with the rock in the conduit by the very hot seawater, and probably its acidic dolomitization brine. This would also be consistent with the brines containing more calcium than indicated by the seawater’s original magnesium content, and the presence of many other constituents (including metals) in the brine. This latter factor is so significant that many of the vents are actively forming metal deposits in the surrounding areas, and it has been speculated that many of the world’s lead –zinc and other metal deposits may have formed in this manner. The analyses of deep cores have shown that the rocks are leached of their calcium, lithium and some other minerals at depth, but that there is considerable re-precipitation of various minerals in the upper sections of the vents. The vent brines usually have a total salt content nearly that of seawater, but may also have been diluted by other waters, or concentrated by evaporation.

270

Table 2.6 Composition of some Ocean-Vent and Geothermal Brines (wt.%) (A) Deep Sea Hydrothermal Vents

Sea watera Galapagosb East Pacific Risec Mid Ocean Ridgea Reykjanes (Wells36)d Red Seae

413 427 648 1611 1630 5150

Cl a

Sea water Galapagosb East Pacific Risec Mid Ocean Ridgea Reykjanes, Icelandd Red Seae

Na 1.08 0.602 1.048 1.12 1.01 9.29 SO4 (in ppm)

1.935 2711 1.210 2687 1.826 0 2.110 214 1.965 59 15.550 840

K (in ppm) 393 400 939 731 1340 1870

Mg (in ppm) 1294 1261 0 0 18 764

SiO2 (in ppm)

HCO3 (in ppm)

— 13.2 1050 21.9 510 59.1

144 — — — 2290f 142

Mn (in ppm) — — 49 63 — 82 Br (in ppm) 65.4 — — — — 128

Fe (in ppm) — — 80 — 0.49 81

Sr (in ppm) 8.2 — 7.1 — — 48

B (in ppm)

TDS

4.7 — — — 12 7.8

3.419 2.438 3.16 3.35 3.31 25.80

Ba (in ppm)

Zn (in ppm)

— 86 1.4 6 — 0.9

— — 5.5 — — 5.4

8C

pH

— 8 357 — 246 56

,7.6 — 3.5 — 5.7 —

Pb (in ppm) — — 0.54 — — 0.63

Li (in ppm) 0.17 0.30 7.2 — — 0.30

(B) Geothermal Springs

Dead Sea Springsg Asal,Spg.NE Cor.h Qarhan Spg.Chinai Tiberias Springj Reykjanes, Icelandd Utah Hot Springl

Ca (in ppm)

Na

9450 5920 5200 3930 2260 1140

2.11 2.11 8.40 0.691 1.43 0.703

K (in ppm) 2650 1440 510 77k 1670 904

Mg (in ppm) 17,000 1520 4700 825 123 70

Fe (in ppm) — — — 7.8 0.192 0.4

Mn (in ppm)

Ba (in ppm)

— — — — — 1.9

— — — — — —

Li (in ppm) — — — — 7.4 —

Cl 10.34 5.63 15.20 1.92 2.91 1.33

SO4 (in ppm) 1140 204 1450 891 206 189

(continues)

Part 2 Calcium Chloride

Ca (in ppm)

Table 2.6 (continued ) (B) Geothermal Springs

Stinking Springl Guilietti Springsm

g

Dead Sea Springs Asal,Spg.NE Cor.h Qarhan Spg.Chinai Tiberias “Springj Reykjanes, Icelandd Utah Hot Springl Stinking “Springl Guilietti Springsm

Ca (in ppm)

Na

K (in ppm)

946 386

1.26 803 (in ppm)

571 61

SiO2 (in ppm)

HCO3 (in ppm)

Br (in ppm) 2250 — — 1.4k 98 8.2 15 –

— — — 23 544 38 48 79

118 38 — 366 5 192 324 63

Mg (in ppm)

Fe (in ppm)

Mn (in ppm)

Ba (in ppm)

Li (in ppm)

Cl

— —

— —

4.1 —

— —

2.16 1880 (in ppm)

H2S (in ppm)

B (in ppm)

I (in ppm)

F (in ppm)

TDS

42 — — 77 0.2 — 60 –

— — — — 12 — — 1

— — — — 7.4 — — —

— — — — 0.5 — — —

16.64 8.755 24.80 3.23 5.216 2.28 3.65 3400 (in ppm)

297 35

SO4 (in ppm)

8C 42 60 — 61.9 99 57 48 44.6

111 173

pH — 6.8 — 7.3 6.2 6.7 — 7.5

a

Garrett (1996). Galapagos Spreading Center, Edmond et al. (1979); also as ppm F 1.39, Rb 0.14. c Von Damm et al. (1985); 4 vents; also as ppm: H2S 202, Rb 2.54, Cu 1.41, NH4 and As 0.18, Al 0.12, Cd 0.10, Co 0.08, Ag 0.03, 87Sr/86Sr 0.70314. d Bjornsson et al. (1972) (36 well analyses; average 1027 m deep; also as ppb Mo 8.5, Ge 5.8, Ti 5.7, Ga 5.1, V 1.6). e Also 0.1–0.5 ppm As, Co, Cu, Ni, Se, U; Manheim (1974). f Reported as CO2. g Vengosh and Rosenthal (1994); Average of Zohar and Yeha springs, west edge of the Dead Sea. h Lake Asal, southern end of the Danakil Depression; Valette (1975). i Lowenstein et al. (1989). j Northern Israel. k Vengosh and Rosenthal (1994) list K at 341, Br 244 ppm. l Near the Great Salt Lake, Utah. m Edge of Guilietti Lake in the Danakil Depression, Ethiopia; 10 samples; Martini (1969). b

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Figure 2.12 Location Map for the Reykjanes Thermal Brine Area, Iceland (Bjornsson et al., 1972; reprinted from the AAPG Bulletin. AAPGq 1972, reprinted by permission of the American Association of Petroleum Geologists, whose permission is required for further use).

In some deep-sea vents, the brines have not undergone the dolomitization reaction, and/or have such massive circulation rates near the ocean floor that they have only moderately increased calcium and fairly normal seawater magnesium contents, augmented by other constituents. Examples of this are the vents at the Galapagos spreading center (Edmond et al., 1979), and many wells and thermal springs in Iceland (Bjornsson et al., 1972). The fairly large number of seawater vents in the world has given rise to the rather surprising speculation by some authors that their calcium content could precipitate enough of the sulfate in the world’s oceans to have periodically made them essentially sulfate free (Hardie, 2000). This conclusion was reached despite the tremendous amount of highsulfate river water that constantly enters the oceans, and the relatively small volume of the thermal vents. Also, studies such as by Ayora et al. (2001) and many others have demonstrated that the sulfate content in occlusions within the same formation in individual marine basins varies widely, thus indicating the presence of either dolomitization brine or seawater, and not a change in seawater’s composition. As seen in Tables 2.6 and 2.7 the Red Sea vent brine of Egypt and Saudi Arabia is somewhat different from the brine in most other deep sea vents. It

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appears to be slightly concentrated seawater that has dissolved additional halite (the Br, B and Li concentration ratios are similar to each other, and average 1.73 times seawater; the Na and Cl ratios are similar to each other, but average 8.34 times seawater). Presumably, as the seawater in the bottom of the Red Sea sank through the rift zone sediments it dissolved additional halite that underlies much of the Red Sea. It then was heated, underwent the dolomitization reaction as it contacted calcite, and finally leached some of the fissure’s rocks as it was forced back to the sea floor by convection flow or steam eduction. Massive halite outcrops from under the Red Sea to the south near Jizan, Saudi Arabia, and to the north on the coastal plain at Safga, Egypt. Limited potash (sylvinite) beds have also been found in some locations with the very thick halite deposits (Notholt, 1983), and elsewhere langbeinite has been indicated by gamma ray logs (Hite and Wassef, 1985). The Red Sea vent’s metal content is similar to many other deep ocean vents, with quite high strontium values. Pushkina et al. (1982) also commented on the amounts of Ba, Cu, Fe, Mn, Ni, Pb and Zn in the brine, while Manheim (1974) in addition noted the presence of elevated amounts of As, Co, Se, U and perhaps F. Older vents in the same rift area appear to have formed Fe – Mn – Ba, Pb – Zn, and Pb– Zn –Cu – Ba deposits along the Red Sea far to the north and south of the current vents (Figs. 2.13 and 2.14). The sea floor Late Miocene shale-anhydrite breccia just south of the vents contains about 5% Zn. At the vents, the rocks contain 14 – 21% Fe, 0.8% Zn, 0.6 –0.8% Cu and 0.5– 0.7% Pb. Based upon marine sediments and the presence of several deposits of this type, it appears that the Red Sea once filled the adjacent Afar (Danakil) Depression clear to the Ethiopian Plateau (Fig. 2.13; Manheim, 1974; Bonatti et al., 1972). In addition to the dilute vent brine, there are also indications that hot potash end-liquor and its dolomitization brine have been formed under the Red Sea in other areas. It has been speculated that both of these brine types have migrated separately and formed springs and groundwater along the Red Sea fault line many hundreds of kilometers to the north in Israel (Tables 2.6 and 2.8; Fig. 2.16). Inland Geothermal Brines There are also a very limited number of calcium chloride-containing geothermal springs in inland locations (Table 2.6B), since most geothermal springs are of a sodium carbonate-type and/or low in calcium (Garrett, 1992). Most of the CaCl2 thermal springs appear to be seawater dolomitization brines because of their very similar concentration and composition, but their connection to the ocean or potash/halite deposits is uncertain. As previously noted, the Reykjanes, Iceland spring is near the ocean and in a highly fractured volcanic belt (containing considerable calcite) that is a continuation of the Mid-Atlantic Ridge (Bjornsson et al., 1972; Fig. 2.12). The Danakil Depression in the Northern Afar Rift zone (Fig. 2.14) has springs that are in an active volcanic area near the Red

274 Part 2 Calcium Chloride

Table 2.7 Geothermal Brines from the Salton See, Red Sea and Cerro Prieto, Mexico, ppm except as noted (Garrett, 1996) Cerro Prieto Brine Power plant discharge Cl (ppm) NaCl (wt.%) KCl (wt.%) CaCl2 (wt.%) MgCl2 (wt.%) LiCl (ppm) Br Rb Cs Sr B Ba U As Pb Te

— 2.20 0.32 0.12 — 100 20.5c 9.4c 3.5c 15.7c 12c 9.73c — 1.20c — —

From pre-conc. pond — 22.2 3.23 1.22 0.06–0.19 1100 — — — — — 12.4d 7.0d 7.1d 3d 2.2d

From salt pond

Typical Salton Sea brinea

— 17.83 6.87 2.59 — 2400 — — — — — 17.0d 31.0d 6.5d 6.7d 4.0d

142,000 –209,000 12.7–17.8 2.48–6.53 6.26–10.81 0.27–2.23 1340–1950 200 25– 100 24 540–2000 400–500 200 — 312 90– 210 520

Red Sea geothermal brineb 155,500 23.6 0.357 1.426 0.292 1.6 128 — — 48 7.8 0.9 — — 0.63 — (continues)

Table 2.7 (continued ) Cerro Prieto Brine

Fe Si Sb Mn Zn Cu Ag NH4 S¼ Density pH

Power plant discharge

From pre-conc. pond

From salt pond

Typical Salton Sea brinea

Red Sea geothermal brineb

0.94d — — 0.64c 0.2c 0.12c — — — — —

3.3d 2.4d 1.8d 0.6d — 0.4d 0.6d — — 1.207d —

0.4d 2.1d 3.3d 1.8d 1.7d 0.6d 0.8d — — 1.250d —

1200– 3700 40 — 1000– 2000 500 –700 6–20 1–2 650 15–30 1.18–1.26 4.6 –5.5

81 27.6 — 82 5.4 0.26 — NO3 0.044 SO4 840 — —

Reproduced from Potash; Deposits, Processing, Properties and Uses, Table 2-12, pages 184 –185, q1996 with kind permission of Kluwer Academic Publishers. a Various sources. b Craig (1969, 1966). c Average of original well brine (Mercado, 1976); Other Cerro Prieto analyses (Galinzoga, 1981). d Vazquez (1981). Geology

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Figure 2.13 Location of the Red Sea Thermal Brines, Metal Deposit, and the Danakil Depression (Bonatti et al., 1972; reprinted with permission from Economic Geology, vol. 67:6, p. 718, Fig. 1, Bonatti, E., Fisher, D. E.; Joensu, O., Rydell, H.S. and Beyth, M., 1972).

Sea coastal plain. They are below sea level and they are also adjacent to a large inland potash deposit. Also as noted above, the Israeli high-calcium chloride springs are presumed to have originated in the Red Sea fault system, with its large halite –potash deposits. Both the Icelandic and Ethiopian springs have a composition very similar to seawater, and they are in areas with recent volcanic activity. They also contain many metals that appear to be derived from rockdolomitization brine reactions. There are also CaCl2-thermal springs in a thermal zone in Utah, USA near the Great Salt Lake, which has a brine composition very similar to seawater. Some geothermal calcium chloride waters have also been reported in the Cascade Range in Washington (Mariner et al., 1994). Several hot springs in Japan contain calcium chloride, such as the Tsurumaki Spa (Oyama et al., 1987), the Togo-Matsuzaki hot springs in Totturi prefecture (Umemoto et al., 1958) and the Komumi and Ishibu-Iwachi springs on the Izu

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Figure 2.14 Metal deposits from ancient or current Red Sea thermal vents (Bonatti et al., 1972; reprinted with permission from Economic Geology, vol. 67:6, p. 727, Fig. 2.8, Bonatti, E., Fisher, D. E.; Joensu, O., Rydell, H. S. and Beyth, M., 1972).

Peninsula. The latter springs appear to be of the seawater-type, containing 3.1– 15.8 g/liter salts with a high bromine and low K, Mg, Na and SO4 content (Kanroji and Tanaka, 1980). Also, in the Matsushiro area, short-lived calcium chloride springs were formed as a result of an earthquake swarm (Yoshioka et al., 1970). All of the Japanese springs are near the ocean and rift zones, and have a seawater dolomitization-type brine. The numerous calcium chloride geothermal springs in Israel have previously been mentioned (Fig. 2.16). In Turkmenia, Lebedev (1972) has described a few calcium chloride geothermal springs in Cheleken, near Ashkhabad (and also near a potash deposit). Miscellaneous Springs Various other cool-to-warm calcium chloride springs have been reported. Low volume and dilute calcium chloride springs feed into one section of the Qarhan Lake in China (Fig. 2.15), and some other dilute CaCl2 thermal springs have been reported in China (Xu, 2000). In England, there are extensive areas of highcalcium chloride ground water and springs that may have been derived from England’s Zechstein Formation halite and potash deposits (Anderson, 1945). There are calcium chloride springs in the Harz Mountains at Stecklenberg, Germany

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Figure 2.15 Structure of Qaidam Basin, China (Sun and Lock, 1990). Legend: 1. mountains; 2. Playa surface; 3. major potash deposits; 4. smaller potash deposits; – – –. outline of Qarhan Playa; Q. Qarhan Salt Lake; D. Dabuxun Saline Lake. Reprinted by permission of Science Press (China). (See Fig. 1.19.)

where the brine probably originated in the Zechstein beds and then rose through a network of fissures (Haller and Mestwerdt, 1938; Harrassowitz, 1935). In the United States in the Trans-Pecos region of Texas, the deeper Permian and lower Cretaceous aquifers contain calcium chloride brines, and their mixing with shallower aquifers results in some quite dilute calcium chloride springs (i.e., the San Solomon, Griffin and Phantom Lake springs; Hart, 1992). Similar waters have been observed in the Texas Panhandle (Bein and Land, 1982), and both appear to have originated from the large Carlsbad potash deposit or perhaps from oil field brines. Salton Sea Geothermal Brine A very large pool of extremely hot high-calcium chloride brine occurs in the Salton Sea trough, south of the Salton Sea, California and extending further south to Cerro Prieto and the Gulf of California in Mexico (Table 2.7). The highest brine concentrations and temperatures are found near the Salton Sea, and the temperatures rise again at Cerro Prieto, but with much more dilute brine (about 1/10th the concentration). The northern geothermal area covers about 50 km2, and the southern area is somewhat smaller. In the north, 3008C brine is found at about a 900 m depth, and 3608C brine at 2100 m (White, 1968). At Cerro Prieto, there are power-production wells drawing 250– 3448C brine from a depth of 780– 1450 m (Dominguez and Vital, 1976). The brines are very complex, variable and have

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an unusual heavy metal content. The basin containing these brines is an extension of the Gulf of California rift-fault system, and filled with at least 4300 m of deltaic sediments from the Colorado River emptying into the Gulf of California. The basement rocks are at a depth of about 6100 m. Various major faults of the San Andreas system extend from the north into the trough, and no data have been published on the source of the heat anomaly in the area. The northern zone is estimated to contain over 300 million tons of calcium chloride. Originally the brine, as with the somewhat similar brine under the Red Sea, was thought to be a pure magmatic ore fluid with a deep thermal source. However, as numerous isotopic studies were made on the water in the brine it was indicated that much of it was of local meteoric origin. White (1968) suggests that rain falling on the Chocolate Mountains several kilometers to the east of the Salton Sea seeped through fault zones to depths of 3000 m or more. As it descended, it was heated, and perhaps it dissolved small potash (mostly carnallite [KCl·MgCl2·6H2O] with some kainite [KCl·MgSO4·2.75H2O]) deposits that had been formed in this gulf area’s ancient tidal basins (as are now present near the Gulf). It also may have mixed with some magmatic brine, and most of the magnesium ions from the potash (and magmatic) brine then reacted with calcite in the formation to form calcium chloride. As the resultant high-strength brine circulated through the upper porous strata, its heat and acidic nature (because of its high-CaCl2 and MgCl2 content) allowed it to dissolve a wide range of other elements. The thermal currents that were formed caused some circulation throughout the basin, but the brine still maintained a variable temperature, dilution and mineral content. A model for the convective mixing in the brine, as driven by thermal and concentration (density) differences has been proposed by Oldenburg and Pruess (1997). They suggest that this mixing is still occurring. The deuterium content of the geothermal brine, thermal springs in the area, and the springs at the base of the Chocolate Mountains are all very similar, and far different from ocean, Salton Sea or Colorado River water. The brine also has a very large 18O “shift” (increase), which is proportional to its temperature, and typical of very hot water reacting with rocks much higher in 18O. It is assumed that the deuterium content of a magmatic water would not be as high as the geothermal brine (although no examples of magmatic fluids are known; Craig, 1966), and that at best 25% magmatic water could be present based upon the isotopic and deuterium data (White, 1968). However, the 34S values are very similar to that assumed to be present in magmatic fluid. The ratios of NaCl and KCl to Br, and KCl to Rb and Cs are similar to what would be obtained by leaching a carnallite deposit, and the low bromine concentrations would appear to rule out the presence of normal seawater dolomitization brine. Also, there are no large potash or halite deposits in the area to have supplied such a brine. A probably related dilute calcium chloride brine source occurs in the flow from an artesian well on Mullet Island in the Salton Sea near Niland, California. The brine only contains 3.01% salts, distributed as: 0.56% Ca, 0.54% Na, 0.02%

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Mg and 1.89% Cl. For a short period, this brine was the feed for a salt works, and the by-product calcium chloride end liquor was sold as a dust-control road salt (Ver Planck, 1957). Calcium Chloride Lakes The Dead Sea, Israel and Jordan As with calcium chloride geothermal springs, there are also only a few calcium chloride lakes in the world. Most of them contain considerable (or even predominantly) MgCl2 as well as CaCl2, are fairly dilute and relatively small, with the exception of one gigantic concentrated lake, the Dead Sea (Tables 2.8 and 2.12). It is located in both Israel and Jordan, being split in half in a north– south direction between the two countries (Fig. 2.16). It is extremely large, containing an estimated 6.8 billion tons of CaCl2 in the complex brine that is commercially processed for potassium chloride, magnesia, bromine and salt. The Dead Sea is the low point in a long, narrow valley, which extends northward from the Gulf of Aqaba to beyond the Sea of Galilee (also called Lake Tiberias). It is the lowest natural lake on earth, about 66 km long and 10– 20 km wide, with a surface area of about 930 km2, and a mean depth of 396 m. It is within the Jordan Rift Valley, an active slip strike fault with high, steep-sided escarpments on both its east and west boundaries. This trough is a branch off of the Red Sea, and an extension of the massive Rift Valley system to the south in Africa. A thick mass of evaporites, shales and conglomerates of the Dead Sea Group of Plio-Pleistocene to Recent Age fill the rift trough. The lower, or Sedom Formation consists predominately of marine halite, while the upper, or Dead Sea sediments and the Lisan Formation are of terrestrial origin. The bromine content of the lower halite averages a little over 100 ppm, while that of the upper two formations are quite different, averaging 209 in the Lisan, and 292 in the Dead Sea halite. This indicates a much higher bromine content in the terrestrial basins’ waters than seawater. The Lisan Formation contains abundant diatoms, fish and plant remains, typical of hyper saline inland lakes. It would thus appear that the last seawater incursion into the Jordan Valley occurred during the Amara period, and that during the latter three depositional periods, the basin has been closed to the sea, concentrating inland (and in the last phase, geothermal) waters of a quite different composition than seawater. Following the marine period a freshwater lake (Samara Lake) was first formed, then the brackish-to-saline Lisan Lake, and finally the Dead Sea (which may have started to form before the Lisan Lake totally evaporated). The oldest sediments within the Dead Sea, based upon carbon dating are about 25,000 years old (Zak, 1974, 12,000 years; Bentor, 1961). The source of the Dead Sea’s mineral content, and especially the calcium chloride, is somewhat uncertain. Its composition (Table 2.8) is very similar to a seawater dolomitization brine (from potash end liquor) that has had only limited

Table 2.8 Analyses of the Dead Sea and Adjacent or Related Springs and Waters (wt.% or ppm)a Dead Sea Kenat (1966) Ca Mg Na K Cl Br SO4b HCO3b SiO2b TDS pH 8C Density b

Mt. Sdom Springs

Bentor (1969)

Bentor (1969)

Bentor (1969)

Jordan Rivera,b

Tiberias Hot Springs

Tannur Springs

1.58 4.20 3.49 0.76 20.80 0.59 540 240 — 31.50 — — —

3.99 4.27 2.22 1.57 24.97 0.29 95 127 — 37.33 — — —

8.29 1.98 3.49 2.87 26.52 0.30 108 0 — 42.01 — — —

129 95 224 32 762 9 97 181 — 1529 — — —

3930 825 6910 77 19,200 1.4 891 366 0.2 32,300 — — —

364 105 978 27 2230 20.5 173 240 — 41,390 — — —

1.2 –1.6 2.8 –4.2 2.7 –3.5 0.5 –0.8 15 –21 0.3 –0.6 550 –660 200 –230 8–10 25 –32 6.2 –6.4 19 –40 1.18–1.22

Amira (1992), Hardie (1990). In ppm.

Geology

a

Near Lake Tiberiasa,b

Sdom Well

281

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Figure 2.16 Dead Sea basin drainage area and general geology (Zak, 1974, courtesy of the Northern Ohio Geological Society).

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contact with calcite, but there are no potash deposits in the area. The underlying marine halite contains only minor amounts of potash, and its average bromine content is only half that required for the beginning of potash crystallization (see Table 2.1). There are, however, potash deposits along the fault line under the Red Sea to the south and west, but that area’s present brine (the geothermal springs in the Red Sea, which are perhaps related to the similar ones far to the north at Lake Tiberias in Israel) is low in bromine. However, the ground water in the Dead Sea area, as represented by the nearby Sdom well and the Mt. Sdom, Zohar, Yeha and several other hot springs do contain a high bromine content, but are indicative of a higher dolomitization conversion (i.e., less magnesium) than found in the Dead Sea. If they, or brines like them, were the primary mineral source for the Dead Sea, the blending of Jordan River water (which contains the Lake Tiberias spring water, and has more magnesium) over the years could easily account for the Dead Sea brine’s composition. Bentor (1961) calculated a onethird Jordan River, two-third groundwater mixture as the Dead Sea source. This would assume that the strong calcium chloride Dead Sea geothermal wells and springs contain brine from a different aquifer than the Red Sea – Lake Tiberias source, and that it was a more conventional potash end-liquor dolomitization brine. It could have come from the Red Sea potash deposits, or even been residual from much older, now extinct deposits. A similar differentiation of brines with different compositions following separate fault lines is occurring at Lake Assal in Ethiopia at the present time. In the past, before the present extensive irrigation and other uses, the Jordan River carried up to one million tons of salts per year to the Dead Sea, and brought about 75% of its yearly water supply. It had a fairly high calcium, magnesium, sulfate and bicarbonate content, and in the Dead Sea all of the calcium soon precipitated as calcite (and in some periods gypsum). Despite the Dead Sea’s high magnesium chloride content, this calcite probably was not later converted to dolomite, since it would have been deposited with suspended sediments in the dilute end of the Dead Sea. The resultant excess of river water’s magnesium would then change the Dead Sea’s Mg/Ca ratio, and perhaps resulted in the present high magnesium chloride content compared to that in the Dead Sea wells’ or springs’ brine. At present, it is assumed that the Sdom and other springs near or under the Dead Sea supply 70% of its yearly mineral input (Zak, 1980), and if all of this input had a composition similar to the Sdom well and spring it would help to support this theory. Bristol and Cadiz Dry Lakes, California The only strong-CaCl2, low-MgCl2 lakes in the world are the very small Don Juan Pond, Antarctica and Bristol Lake. Cadiz Lake (adjacent to Bristol Lake) has a somewhat more dilute, higher magnesium brine. Bristol and Cadiz Lakes are relatively small late Quaternary dry playas in the Bristol – Danby Trough of the

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south-central area of California’s Mojave Desert, which is also in the western United States’ Basin and Range province. They are located about 360 km east of Los Angeles and 135 km west of Needles, California. Bristol Lake lies just south of the main line of the Santa Fe Railroad and the small town and rail station of Amboy (Figs. 2.17 and 2.19). Its elevation is 180 m, it has an area of about 130 km2 (155 km2; Rosen, 1992), and its lacustrian sediments are very deep (. 500 m, of which 260 m are halite; . 1000 m, Jachens and Howard, 1992). The lake is in a broad closed basin with a 2000 km2 drainage area (before a fairly recent lava flow it had been 4000 km2) formed by the Bristol and Marble Mountains to the north and northeast. The mountains are composed of Jurassic granite, metamorphosed Paleozoic sedimentary rocks and some Precambrian granite. On the south and southeast of Bristol Lake are the Bullion, Sheephole and Coxcomb Mountains

Figure 2.17 Location map for Bristol and Cadiz Dry Lakes, showing elevations and adjacent lakes (Bassett et al., 1959).

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with relatively steep alluvial slopes, and to the east southeast there are low ridges that are mostly buried by drifting sand. They separate the adjacent, smaller and 31.7 m lower Cadiz Lake to the southeast by only a 15.8 m ridge, although it has been theorized that Cadiz Lake may still drain into Bristol Lake at considerable depth through a subterranean channel. On the west and northwest edges of Bristol Lake there is a 62 km2, , 2 m thick, roughly circular and level black basaltic lava flow on top of the playa surface that originated at the same time as the Amboy Crater (also called the Bagdad Crater or Bagdad Cinder Cone; Fig. 2.18) in its northeast corner. There are no lake sediments or shoreline markings on the lava, indicating that it was formed since the last time the lake was full, presumably since the last ice age ended , 10,000 years ago (Recent to Late Pleistocene). Most of the basins in the Mojave Desert were filled with water at times during the ice age periods from 50,000 to 10,000 years ago, and were actively depositing lacustrine sediments (Smith, 1976, 1979). Bristol Lake could have also been filled during the milder Tioga period 6000 years ago, but few of the Mojave basins were, so the lava’s age more likely could be as much as 10,000 years. This still makes it the youngest lava formation in Southern California. The lava field is so uniformly thin and flat that it appears to have been formed rapidly (such as within a few weeks) by a relatively hot and thin fluid. It has a low silica and high-potassium content, which would reduce the lava’s viscosity, and its lack of large crystals indicates that it was very hot when erupted, and then cooled rapidly. The 76 m high and 472 m diameter cinder cone 3 miles southwest of Amboy appeared to have formed at the same time and from the same lava by gas escaping from at least three closely connected vents, making it an unusually broad, low, nearly circular and flat-topped crater of cinders and larger ejecta (Hazlett, 1992).

Figure 2.18 The Amboy Crater and lava field viewed across the desert northwest of Bristol Lake.

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Figure 2.19 Stratigraphy of Bristol and Cadiz Lake Sediments (Bassett et al., 1959).

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The lava divided the original Bristol Lake into two sections, forming the smaller Alkali Dry Lake to the northwest. The area is one of the most arid in the United States with a mean annual rainfall of , 100 mm, and there are 2 – 3 year periods with no rain at all. The evaporation rate in the area is very high, and even though the lakes are the terminus for fairly large drainage basins, they only occasionally flood to a shallow depth after heavy storms, and then evaporate rapidly (Rosen, 1992; Bassett et al., 1959; Gale, 1951). Both Bristol and Cadiz Lakes have three types of surfaces, with the principal one being “puffy (and) efflorescent, called self-rising ground, caused by the evaporation of (fairly shallow) capillary(-flow) brine, (and the crystallization of the brine’s halite [primarily] and calcite content).” There are also some areas of hard-packed clay where the water table is lower, and other areas (generally at the lakes’ lowest levels) where there is salt on, or just below the surface (Bassett et al., 1959). These areas become smooth and sparkling white after the rare lake flooding. The outer periphery of the playa is filled with alluvial deposits of rock, sand and clay washed down from the surrounding mountains, with the average particle size being considerably graded as it approaches the lake. In many areas, these porous sediments are filled with a dilute calcium chloride brine containing roughly the same ratio of salts as in the lake. A ring of gypsum-containing sediments surrounds the lake where the sulfate ions of entering groundwater have precipitated with the lake’s calcium chloride brine. The Bristol Lake Basin was initially drilled with two holes to a depth of 307 m (1007 ft), finding salt, silt and clay in both mixtures and alternating beds (averaging about 40% halite; Fig. 2.19) throughout the depth of the drill holes. Some of the more than 40 massive and relatively pure halite beds were in excess of 3.6 m thick (totaling 160 m of halite). Later four other very similar-appearing holes were drilled up to 537 m in depth, with 260 m of halite, and the sediments near the bottom of the holes indicated an age of 3.65 mybp (million years before the present; Rosen, 1992). In the southwest area of the lake there is a 13 km2 zone called the Salt Lake, which began to be mined for salt in 1909. In 1951, it contained an average 1.5 m (0.9 –2.1 m) thick salt bed under a roughly equal thickness of surface clay and salts, and the salt was pure enough to have been sold without further processing. Fairly large amounts of this salt were mined, such as 220,000 st in 1942 and 120,000 st in 1951, but more recently salt production is obtained from some of the calcium chloride solar ponds, and it is then drained and washed to an even higher purity. Beginning in 1982, a second salt area formed to the south of the lake, growing by 1988 to about a 0.4 m thickness. Gypsum that was used to produce plaster was also recovered from the lake from 1907 to 1924. In one area near the lava flow, small grains of gypsum had been deposited as dunes that could be excavated to a depth of 1.8 –2.4 m. The cleanest gypsum consisted of a “spongy white mass, granular to coarsely crystalline, that was easily broken into blocks or crushed.” A final mineral in the

288

Part 2 Calcium Chloride Table 2.9 Examples of the Calcium Chloride Brine in Bristol and Cadiz Lakes Bristol Lake Sample 1

Cadiz Lake Sample 2

Sample 3

(A) as ppm (Smith, 1966) Na Ca K Mg Sr Cl SO4 B4O7 (B) TDS CaCl2

46,070 17,190 1479 598 — 104,600 1048 88(25) 171,000 4.14

57,370 43,000 3303 1074 — 172,900 210 30(8.4) 279,000 12.02

22,600 4500 1040 410 — 44,760 280 — 73,600 1.22

(B) % of total salts (Ver Planck, 1957) Na Ca K Mg Sr Cl SO4 B4O7(B) TDS CaCl2

26.86 10.02 0.86 0.35 0.23 60.99 0.61 0.05 171,000 24.2

20.55 15.52 1.18 0.38 0.34 61.95 0.08 0.01 279,000 43.1

30.71 6.12 1.41 0.56 — 60.82 0.38 — 73,600 16.53

(C) Maximum solar evaporation, Bristol Lake (wt.%) (Gale, 1951, 1940 data) Original % of TDS Na Ca Mg

23.6 25.3 1.1

Cl SO4 HCO3

49.91 0.04 0.05

Concentrated CaCl2 Na (and K)Cl MgCl2 Total

48.70 0.49 0.30 49.49

(D) Lithium and chloride in various brines (Vine, 1980) Bristol Lake Cadiz Lake

Li 68 –108 ppm, Li 20–67 ppm,

18–21.6% Cl 8–18% Cl (continues)

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Table 2.9 (continued ) (E) Typical Bristol Lake product, 40.68Be (1.39 g/cc) wt.% CaCl2 NaCl KCl MgCl2 pH

36.0 1.0 0.4 0.9 4.1

ppm Sr B Ba HCO3 SO4 F

2400 48 9 150 120 2

Samples (1) Drainage canal brine from National Chloride in 1953. (2) Solar concentrated brine in 1953. (3) Brine from 11 m depth in 1920.

Lake that once had a very limited commercial value was celestite (strontium sulfate). It is found on or near the surface in muds, particularly along the southern border. It occurs as nodules with a potato-like shape composed of milky white, minute crystals, and “two cars” of these nodules were gathered in 1942 for sale (Gundry, 1992; Gale, 1951). Bristol Lake is unique in the world in containing a fairly strong calcium chloride, low-magnesium chloride brine in its near-surface sediments. Only a few other lakes in the world are of the calcium chloride type, but most of them are much more dilute and/or contain a higher proportion of magnesium. Drilling has indicated that the calcium chloride brine (Table 2.9) is found throughout the lake, but only in the upper 9.75 m (32 ft) of its sediments. Its source is unknown, but the fact that it only occurs at shallow depths might indicate that during the recent lava flow and formation of the Bagdad Crater (over a dry playa) the brine near the surface was heated to temperatures high enough to react with calcite in the lake’s sediments. Based upon the brine-type in many nearby and adjacent basins (i.e., Dale Lake, Danby Lake, etc.) the residual brine from the crystallization of Bristol Lake’s massive halite deposits would have almost certainly been of the strong magnesium chloride type, and the lake sediments do contain large amounts of calcite. The heat from the lava flow would have been sufficient to initiate the dolomitization reaction, and then both thermal and brine concentration (density) effects would have caused a circulation of the lake’s brine. This would have allowed all of the lake’s brine to be reacted, and the MgCl2 to be almost totally converted to CaCl2. The heat and the acidic nature of the brine would have also resulted in some rock leaching. After a little overflow to Cadiz Lake,

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the remaining brine would have then been trapped in the near-surface sediments by the lake’s frequent layers of low-permeability clay. Alternately, the calcium chloride might have originated from waters that accompanied the lava, but no similar lava (or magma)-flow calcium chloride brines are known. Similar basalt flows in the West African Rift Valley are accompanied by high-carbonate brine, in the Andes by high-boron and lithium brines, and in some pegmatites by high-metal content brines. Calcium chloride appears to only accompany lava flows when seawater (or seawater-type lacustrine brine) is present to have its magnesium converted by the dolomitization reaction, or to assist in very hot rock leaching. Possibly an ion exchange reaction may have been initiated by the lava’s heat between clays and the lake’s brine to liberate calcium, but most clays are naturally in the sodium form (because of the predominance of salt in the clay-forming basins), calcium is more strongly adsorbed than sodium, and heat is usually not a significant factor in this exchange. Possibly the heat may have caused gypsum to react with salt to form glauberite and calcium chloride, but this has never been known to occur at any temperature, glauberite is much more soluble than gypsum, and no glauberite is present in the lake’s sediments. Finally, the heat from the lava may have resulted in sufficient rock leaching to form the calcium chloride. However, if this were the case the lake’s original MgCl2 should still be present, there should be many other metals in the brine, and there should be massive evidence of altered minerals. Many articles have been written about Bristol Lake, including one noting that most of the gypsum in the lake is near the edges, and thus apparently formed by groundwater seepage into the lake (Rosen and Warren, 1990) Handford (1982) examined the lake’s sediments, and the general geology of the region was noted by Miller et al. (1982). A number of other articles have been written on more specific aspects of the area’s geology. Cadiz Lake is located just to the southeast of Bristol Lake at an elevation of 165 m, and also contains a calcium chloride brine in its near-surface sediments. The brine is similar but more dilute than at Bristol Lake, and has a higher ratio of sodium to calcium chloride. However, in other features the two lakes are quite different, and do not appear to have ever been joined as a single lake in a deeper basin. A 152 m core hole (Fig. 2.19) showed that its sediments contain no halite except for one 0.3 m thick bed 2.7 m below the surface, and that its strata are quite different from that at Bristol Lake. The Cadiz basin contains much more gypsum, sand, silt, and even many fresh and saline water fossils. Near the center of the lake, there is a low area with the above-noted halite close to the surface, surrounded by large areas with an efflorescent surface. However, the northeastern and southern parts of the playa have a fresh water aquifer, and there are some large expanses of dry clay flats which are now utilized for agricultural production. The differences in the lake’s sediments would indicate that after the calcium chloride was formed at Bristol Lake, there was a limited period of overflow of sodium and calcium chloride brine into Cadiz Lake, but that during

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their prior long period of basin filling there was no connection between the basins. If there was some occasional overflow into Cadiz Lake, then the salts must have been quickly diluted (there are frequent fresh-water fossils) and rapidly flushed from the basin. Alternately, Danby and Cadiz Lakes may have been presettling and pre-evaporation basins for a then-lower Bristol Lake, but the current relative elevations make this unlikely, and there should have been occasional zones of similar sediments (or marker beds) in each lake, which there is not. Finally, Danby and Cadiz Lakes contain some fairly old sediments from the Gulf of California, which Bristol Lake does not have (Brown, 1992). Among other articles on Cadiz Lake there have been gravity and magnetic studies (Mickus et al., 1988), and a geologic map prepared (Howard et al., 1989). Danby Lake is southeast of Cadiz Lake, and is the last member in the Bristol– Danby Trough. It is 3.2 –4.8 km wide by 22.5 km long, at an elevation of 189 m, and the divide between it and Bristol Lake is 152 m high (Fig. 2.17). Its sediments are somewhat similar to Cadiz Lake, but it does not contain any calcium chloride. Its brine and salts are predominately of the sodium chloride – sulfate type (Calzia, 1992), which would have precipitated any calcium chloride that may have overflowed from the other lakes. Danby Lake has had limited commercial production of sodium sulfate, and considerable production of salt (Garrett, 2001). Lake Giulietti, Dallol Salt Pan, Lake Asal, Ethiopia The Danakil (Afar) Depression is a former branch off the Red Sea that extends to the northwest from the Gulf of Tadjoura and closed about 40,000 years ago. Metal deposits such as may now be forming from the deep Red Sea vents are located in its northern section, and there are coral ridges along its borders (Fig. 2.14). The Depression is roughly parallel and about 45 km from the Red Sea, and below sea level, with an elevation down to 2 120 m. It lies between the Danakil Alps near the Red Sea coast and the Ethiopian Plateau in the interior (Figs. 2.13, 2.20 and 2.21; Table 2.10). Lake Asal is located at the southern end of the Danakil Depression, while Lake Giulietti is near its middle-west in a mountain valley that drains into the Depression. The Dallol Salt Pan is in its northern section. Also in its northern end in the Masli area there is a large Pleistocene –Quaternary potash deposit. Lake Giuletti is 70 m below sea level, has an area of about 70 km2 and a depth of over 100 m. It is in an actively rifting area in a valley formed by volcanic peaks in the center of the Danakil Depression. The active volcano Alid lies southeast of the lake in the Danakil Alps (Erta’Ale) range, and there are other active volcanoes in the central range further north and to the south. A number of calcium chloride-containing hot springs flow into the southwest edge of Lake Giulietti, and the lake appears to occasionally drain into the nearby Dallol Salt Pan (playa). The dilute calcium chloride spring water is similar to slightly diluted

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Figure 2.20 Map of the Northern Afar Rift in relation to the Red Sea. 1. Assumed vent metals deposit. 2. Lake Giulietti. 3. Salt playa. (Bonatti et al., 1972; reprinted with permission from Economic Geology, vol. 67:6, p. 719, Fig. 2, Bonatti, E., Fisher, D. E., Joensu, O., Rydell, H. S. and Beyth, M., 1972).

dolomitized seawater (Table 2.6; Fig. 2.22; Barberi et al., 1970). There are marine formations containing calcite and coral between the lake and the Red Sea, and since the lake is below sea level in a rift valley connected to the Red Sea, it would appear that seawater flows into the hot, volcanic strata to form the hot springs. Some intermittent streams also flow into the lake, but they are quite dilute and do not contain calcium chloride. The lake’s mineral content is similar to the calcium chloride springs (Bonatti et al., 1972; Martini, 1969). Lake Asal is at the southern end of the Afar Depression and even closer to seawater, being only 9 km from the Goubet el Kharab bay, which connects to the Gulf of Tadjoura off of the Red Sea (Fig. 2.23). The lake is 155 m below sea

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Figure 2.21 Location map of the Danakil Depression, showing Lake Giuliette, Lake Asal and the Dallol Salt Pan. (Reprinted with permission from Economic Geology, vol. 63:2, p. 125, Fig. 1, Holwerda, J. G. and Hutchinson, R. W. 1968).

level, about 8 km wide and 14 km long, up to 35 m deep, and has a brine-covered area of 54 km2 and a salt – gypsum – clay area of 61 km2. The lake receives about 200 mm/yr of rain, and the average air temperature is 308C. There are a large number of springs surrounding the lake, with many in the east and northeast being hot (40 – 808C), having a low flow rate and being high in calcium chloride. Those to the south and southeast are warm (30 – 408C) with a high flow rate (. 100 liter/s) and a composition similar to seawater. The springs to the southwest are similar to those in the northeast but hotter (808C) and with a lower flow rate (1 –100 liter/s). There are only a few cold springs (308C) flowing into the lake in the north and west. The average analyses of seven high-calcium chloride springs to the northeast are listed in Tables 2.6 and 2.10, with the composition of the lake’s brine in Table 2.10 (which remains fairly constant throughout the year). It is seen that the seawater-type brine entering from the south results in the precipitation of

294

Table 2.10

Lakesa Don Juan Pondb Don Juan Pondc Deep ground waterb Kzakhstan ad Namib Deserte Dallol Salt Panf Vandac Guiliettif Guilietti Springf Red Lake, Crimeag Asal, Spg. NE Cor.h Ushtagand Kzakhstan bd Asal, Averageh

Don Juan Pondb Don Juan Pondc Deep ground waterb Kzakhstan ad Namib Deserte Dallol Salt Panf Vandac

Ca

Na

121,800 83,700 64,200 39,700 27,500 23,290 21,500 13,600 386 6030 5920 3380 556 2000

3470 1180 6190 45,870 75,200 46,020 6410 39,300 803 59,700 21,090 55,090 2930 103,600

HCO3

CO3

— 36 — 690 340 — 375

— — — — 100 — —

K 157 100 84 1340 8000 8860 685 2450 61 — 1440 73 — 5000 NO3 — — — — 22,500 — —

Mg 1850 — 902 11,190 15,900 16,570 — 805 35 33,400 1520 5330 504 11,900

Cl 136,100 148,900 128,000 173,800 200,000 168,600 68,040 83,400 1880 200,900 56,290 103,880 5135 200,000

B

F

— 4.7 — — 157 — 4.2

— 223 — — 55 — 84

SO4

SiO2

7.4 63 90 797 568 0 630 1020 173 — 204 3930 578 2500

— — — — 46 — — 66 79 — — — — —

TDS 363,400 385,100 199,500 273,500 350,000 263,300 165,900

Part 2 Calcium Chloride

Composition of Some Calcium Chloride Lakes (ppm)

pH

8C

4.1 5.57 6.55 — 8.4 — 5.67

0– 8 3.0 216 21.5 — — 21.0

Density 1.355 1.276 1.164 — — — 1.092 (continues)

Table 2.10 (continued )

f

Guilietti Guilietti Spg.f Red Lake, Crimeag Asal, Spg. NE Cor.h Ushtagand Kzakhstan bd Asal, Averageh

HCO3

CO3

115 63 — 38 256 211 131

— — — — — 8 —

Sr Don Juan Pondc Vandac

862 73

Li 235 27

NO3 — — — 250 — — 307 Mn 0.90 2.24

B

F

20 1 — — — — —

5 6 — — — — — Ba

1.18 .58

Fei 42 256

pH

8C

140,600 3400 300,100 87,550 171,140 9922 325,730

6.6 7.5 — 6.8 — — 7.0

— 44.6 — 60 25 — —

Zni

Cui

Nii

118 2.7

1.7 .71

2.7 .53

TDS

Density 1.13est. — — — — — — Pbi 102 .38

Coi

H2 S

Eh/V

.34 .19

,0.01 10.3

0.63 21.00

a

Also, a small section of Qarhan Lake, China has 0.6–8.0% Ca in the near-surface sediments (Bihao et al., 1986). Harris and Cartwright (1981). c Wright Valley, Antarctica; the Lake Vanda data are at 70 m depth; also ,0.045 ppm I; the Don Juan Pond is late in the summer evaporation period (Goguel and Webster, 1990; Webster and Goguel, 1988). d Central Kazakhstan lakes, Russia; b ¼ average of four other lakes; Posokhov (1949). e A coastal salt pan in Namibia; also 0.22 ppm Fe (Cagle and Cruft, 1970). f Danakil Depression, Ethiopia; 10 samples; Martini (1969). g Russia; Hudson and Taliaferro (1925). h Afars–Issas area, Ethopia; Valette (1975). i In ppb. b

Geology

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Part 2 Calcium Chloride

Figure 2.22 Location of Lake Giulietti, its thermal springs and adjacent volcanoes (Martini, 1969; Barberi et al., 1970 [Numbers 1–10 are hot springs; other numbers are sampling locations.]; reprinted from their Philosophical Transactions, 1970, vol. 267, p. 67 by permission from the Royal Society of London).

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Figure 2.23 Location map of Lake Asal (Valette, 1975, courtesy of the International Union of Geodesy and Geodynamics).

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much of the lake’s calcium, leaving the average lake water only slightly enriched in calcium chloride. The calcium chloride springs from the northeast contain a modest concentration of manganese and lead, while those from the southwest contain SiO2, Sr, Fe, Ni and Co. The seawater-type springs have a slightly elevated fluorine content. The sediments in the lake bottom are primarily gypsum, while in the northeast and southwest there are also traces of many metals, and elsewhere there is some calcite, dolomite, fluorite, plagioclase, pyroxenes, halyte, kaolinite and montmorillonite (Valette, 1975). There is considerable limestone in the sediments between the lake and sea, and many fault lines. Since the lake is below sea level, it appears that dolomitization brine is the source of the calcium chloride. The hotter springs with the greater content of calcium chloride appear to contain brine that traveled through a somewhat different path than the seawater brine, and through some of the recent hot volcanic rocks of the Afar rift zone’s mountains. Lopoukhine (1974) has also reported on the chemistry of the lake’s hot springs. Qarhan Lake, China; Kazakhstan Lakes Several springs and a small area of the Qarhan Lake, China (into which the springs flow by means of the Golmud River; Fig. 2.15) contain a modest content of calcium chloride (Table 2.6 for the springs; some lake brine has 0.6 – 8% CaCl2). The remainder of the lake contains a seawater-type brine, and gypsum precipitates in the zone between the high-calcium and high-sulfate sections. The lake is in the Qaidam Basin, and in 2001 was being prepared for large-scale potash production. There is no indication of the source of the calcium chloride, although the majority of the water entering the lake evaporates to form a strong magnesium chloride brine. As this heavy liquor seeps into the underlying calcitecontaining sediments, it should form a dolomitization brine. Some of this brine that has penetrated semi-impermeable zones might then be forced to the surface as springs by the slowly building compaction pressure of the lakes sediments. In the lake, some tachyhydrite has been reported in small isolated pools in the limited calcium chloride zone (Spencer et al., 1991). In Kazakhstan, Russia there are a group of six somewhat similar lakes (listed as “a” in Table 2.10), with Lake Ushtagan containing a fairly strong calcium chloride brine, and the others having only a modest calcium chloride content. They were all formed by dilute calcium chloride springs (Posokhov, 1949), which almost certainly originated from dolomitization brine formed by the area’s extensive potash deposits. Lake Vanda and the Don Juan Pond, Antarctica One of the most unusual of the world’s calcium chloride lakes is Lake Vanda and the nearby Don Juan Pond in Antarctica. The lake is covered with ice, there is nearly fresh water under the ice, and below that a stratified strong calcium chloride

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299

brine. It is located in the Wright Valley of Southern Victoria Land on Antarctica’s southeastern side, the “Dry Valley Area” that includes the Soya Coast and Vestfold Hills (Fig. 2.24). In this area, there are more than 90 saline-to-fresh water lakes, and all but Lake Vanda (and the nearby Lake Bonney) contain a seawater-type brine and/or snow or glacial melt water. Lake Vanda is totally different from the others in containing a high-calcium chloride, medium-magnesium, almost nobromine (a seawater indicator) brine. It, along with the small nearby Don Juan Pond with an even stronger calcium chloride brine, are also the only highly concentrated lakes in Antarctica, except for Lake Bonney in the adjacent Taylor Valley. Lake Bonney (Fig. 2.24) contains an ice-covered, stratified strong sodium – magnesium chloride brine with medium-low calcium and an unknown bromine content (Table 2.11). Tachyhydrite has been reported in soils near Lake Bonney (Pastor and Bockheim, 1980), but this is unlikely, since tachyhydrite only forms at temperatures above 22.88C, and it is acidic, while the soil was stated to be very alkaline. Lake Vanda has no outlet, and is fed by the Onyx River for brief periods in the summer. There is a permanent 3.6 m layer of ice covering it, then fairly dilute water, and below about 50 m depth the brine rapidly becomes quite strong, warmer and contains many heavy metals (Tables 2.10 and 2.11; Fig. 2.25). There is some suspended halite (NaCl) in the lower brine, and the floor of the lake consists primarily of calcite and gypsum (the latter with þ 18.9– 20.9 d34S values, or the same as seawater’s , þ 20; Matsubaya et al., 1979). Lake Vanda is about 8.5 km long, 2.5 km wide, 70– 74 m deep, and its surface elevation is at about 89 m. The Don Juan Pond’s elevation is 116 m, 8 km to the west southwest in the same valley. The Onyx River flowing into Lake Vanda contains almost pure water, which is not of the calcium chloride type. Sulfate-reducing bacteria appear to be present in the bottom of the lake, forming hydrogen sulfide that partially precipitates some of the dissolved metals. In the shallow waters at the margin of the lake there are thick microbial mats, dominated by cyanobacteria (Webster et al., 1996). If the lowest, most concentrated brine in Lake Vanda was its original mineral source, and it slowly became diluted following an incursion of fresh water, then diffusion calculations would indicate an age of about 12,000 years since that event (Wilson, 1964; Webster and Goguel, 1988). This later flooding might be further indicated by the essentially constant values for both dD and d18O between 5 and 60 m depth, indicating a uniform original water supply, since it was not further isotopically fractionated (by evaporation or freezing) with depth, even though the concentration of salts greatly increased. However, if the data of Matsubaya et al. (1979) and Webster and Goguel (1988) are both correct, then it would appear that there has been about an 18% additional dilution of the lower brine during the 15-year period between their two samplings (January 1973 –1988; average of the 70 m Cl, Ca, Na and K ratios; compare the two investigator’s Cl versus depth plots in Fig. 2.25). These data

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Figure 2.24 Location map of Lake Vanda, the Don Juan Pond and Lake Bonney in the Dry Valley Area of Southern Victoria Land, Antarctica (Matsubaya et al., 1979; reprinted from Geochimica et Cosmochimica Acta, vol. 43, q1979, with permission from Elsevier).

Table 2.11 Analyses of the Brines of Lake Vanda and the Don Juan Pond, Antarctica (g/kg) (Matsubaya et al., 1979; data from January 1973) Depth (m) Lake Vanda 4 ma 8a 14a 16a 20a 30a 40a 44a 45a 48a 50a 55 56 60 Average

71.3b 72.2 76.2 80.2

Ca

Mg

Na

K

2dD

2d18O

Density

240 200 300 500 500 600 600 900 1100 1600 5300 28.4 33.3 53.8

52 61 96 165 180 174 195 276 314 496 702 10.7 14.8 19.7

14 15 22 34 42 45 43 66 88 132 472 3.02 4.19 5.56

39 44 59 90 97 98 150 127 151 192 554 2.39 3.46 4.33

8.3 9.2 11 19 19 19 19 19 26 29 67 0.247 0.351 0.385

244 245 246 245 245 248 247 252 246 256 248 246 248 244 247

31.2 31.3 31.5 31.8 32.0 32.0 31.7 31.9 31.9 31.7 31.8 31.5 33.3 30.5 31.5

— — — — — — — — — — — 1.039 1.046 1.069

60.9 70.0 74.1 73.5

19.2 24.5 33.1 25.6

5.81 7.05 8.70 —

5.12 6.34 7.96 6.98

0.402 0.551 0.604 0.586

237 246 248 244 246

29.7 29.5 29.2 29.4 29.5

1.075 1.094 1.096 1.095

73.7 93.2 109 114

— 28.9 33.2 —

— 9.03 10.5 —

— 7.82 9.75 —

— 0.980 1.09 —

239 236 240 242

29.3 28.2 27.0 26.8

1.093 1.122 1.142 1.150

Bra

SO4a

25 28

463 284

301

(continues)

Geology

64 65 65.5 68 Average

Cl

302

Table 2.11 (continued ) Ca

Mg

Na

K

2dD

2d18O

Density

Bra

SO4a

Don Juan Pond c Streama,d 4–74 7–74 12–74 12–74 1–75 87–88e 1–71 1–71 12–71 11–69 11–73 1–76d 12–68 12–73 12–63 1–65

140 158 148 163 176 182 149 151 197 201 209 216 236 236 244 247 251

28 81.1 74.1 — — — 83.7 76.4 99.3 102.5 107.2 112.7 122 127.1 — 132.2 137.1

22 1.1 1.1 — — — — 0.45 1.5 1.27 1.6 1.6 1.85 1.8 — 2.6 1.8

42 — — — — — 1.18 19.6 16.0 — — — 3.47 3.51 — 2.16 1.63

2 — — — — — 0.10 — 0.14 — — — 0.157 0.20 — 0.23 0.26

— 206 206 183 180 170 — 156 162 195 214 209 — 183 197 193 186

— 19.5 20.2 13.6 12.4 10.3 — 9.4 8.3 14.4 17.4 16.6 — 10.8 11.8 13.5 13.9

1.000 1.224 1.208 1.233 1.255 1.265 1.276 1.216 1.283 1.288 1.298 1.324 1.355 1.361 1.370 1.380 1.386

— — — — — — 106 — — — — — — 111 — 120 123

60 0 30 30 10 30 63 20 0 0 0 0 7.4 0 — 0 0

3.8 mb 6.1 9.4 15d 58d

139 87.3 91.0 120 134

— 44.4 46.7 61.9 65.8

— 0.44 0.46 0.86 0.93

— — — 5.93 6.35

— — — 78 89

211 179 176 — —

21.0 16.1 16.3 — —

1.203 1.119 1.121 1.164 1.182

— — — — —

2.74 2.30

251 252

25.0 25.2

1.181 1.203

— —

Lake Bonney, East Lobe, Taylor Valley (a NaCl–MgCl2 lake) (g/kg) 26 m 143 1.35 27.3 43.9 32.5 162 1.22 21.0 56.9

— 367 353 86 85 2.75 2.94 (continues)

Part 2 Calcium Chloride

Cl

Date or Depth (m)

Depth (m)

Cl

Na

Ka

Onyx River 5m 15 25 35 40 45 50 55 60 65 70f

7.0a 265a 321a 476a 568a 513a 575a 1.34 10.0 31.9 55.9 48.0

2.8a 50a 58a 63a 86a 83a 63a 142a 889a 2.54 4.66 6.41

1.5 13 15 17 23 23 16 34 143 361 557 685

148.9

1.18

Lia

SO4a

Fa

HCO3a

Webster and Goguel (1988); Lake Vanda (g/kg) 0.001 4.0 23 0.19 0.10 18 43 0.29 0.12 18 50 0.50 0.15 20 42 0.67 0.22 28 63 0.86 0.21 25 60 0.89 0.16 26 52 0.91 0.51 33 67 2.14 4.08 138 99 15.0 13.8 263 166 43.1 22.9 519 278 75.9 27.0 630 375 84.2 Don Juan Pond (1987–1988 season)g 235 63 36

100

223

Bra

Density

0.02 0.06 0.10 0.14 0.16 0.16 0.18 0.47 3.80 9.19 15.0 19.6

1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.01 1.04 1.07 1.092

7.2 8.3 7.9 8.1 8.1 8.0 7.5 7.4 7.0 6.4 5.9 5.7

3.2 2.5 2.5 3.0 4.0 7.0 7.0 8.5 12.5 17.0 20.5 21.0

1.276

5.6

3.3

106

pH

8C

h

Sr Onyx River Lake Vanda Don Juan Pond

0.023 73 862

B 0.005 4.2 4.7

Other elements (ppm ) (1987–1988 season) Ba Mn 0.002 0.58 1.18

0.019 2.24 0.90

Zn 4 £ 1023 3 £ 1023 0.118

Pb 0.3 £ 1023 0.4 £ 1023 0.102

H2S ,0.01 10.3 ,0.01

a

Geology

303

in ppm. Analyses below the line are of groundwater under the Lake (11-1973) or Pond (12-1973), except the first line in Don Juan Pond represents a continuous series of dates. c In the first column, the numbers are dates, not depth, except for the groundwater. d Harris and Cartwright (1981) (1-20-1975 data). e Webster and Goguel (1988) (1987–1988 season). f Also Ca ¼ 21.5 g/kg. g Also Ca ¼ 83.7 g/kg. h Goguel and Webster (1990) (1987– 1988 season). b

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Part 2 Calcium Chloride

Figure 2.25 Stratification in Lake Vanda, Antarctica (Goguel and Webster, 1990, courtesy of the New Zealand Antarctic Record).

Geology

305

would also indicate that during the same period the lake gained 2 m in depth (from 68 to 70 m). Finally, the temperatures at the base of the lake are far higher than the area’s mean-average temperature, and perhaps could be explained by a much more recent “salt gradient effect” (the heating of a lower [light-absorbing] brine by solar radiation passing through a thin and transparent cover layer of water). This effect has been shown to be capable of heating a brine to 40– 608C even in the cold winters of the high Andes, but it is hard to believe that enough radiation could pass through the ice and upper water layer to still heat the lower more concentrated and opaque brine. If this effect is not the cause of the increased temperatures with depth, then the area must posses a much higher than normal thermal gradient. The Don Juan Pond (Figs. 2.26 and 2.27) is a very small basin in the same valley and up slope from Lake Vanda, containing such a concentrated calcium chloride brine that it remains at least partly unfrozen all year, even at 2 508C temperatures. During the 1975– 1976 season the pond varied in size from 178 £ 595 m2 (59,500 m2) to 121 £ 406 m2 (40,500 m2 area; 7780 –3080 m3 volume; 7.6 –13 cm depth) in its fairly flat 300 £ 700 m basin. The valley’s sides are quite steep (1000 m rise in less than 1 km; Fig. 2.27), and the valley rises 200 m east of the pond before it slowly drops to Lake Vanda. In some summers, the pond evaporates to dryness, forming crusts of halite, gypsum and the rare mineral antarcticite (CaCl2·6H2O). In late summer a small river flows into the pond containing a low-salt, non-calcium chloride-type water with a pH of 7.7, indicating a considerable HCO2 3 content. Under the pond is 14 m of “silty sand” containing gypsum and calcite that is frozen elsewhere in the valley, but it is not known how deep it remains unfrozen beneath the pond. Under this surface “soil” is 30 –40 m of fractured Ferrar dolerite (also called diabase, an intrusive

Figure 2.26 Location map of Lake Vanda and the Don Juan Pond, Wright Valley, Antarctica (Harris and Cartwright, 1981; q1991, American Geophysical Union, reproduced by their permission).

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Part 2 Calcium Chloride

Figure 2.27 Oblique view of the Don Juan Pond (Harris and Cartwright, 1981; q1991, American Geophysical Union, reproduced by their permission).

mixture of labradorite [feldspar of the plagioclase series containing equal amounts of Na and Ca] and pyroxene). The fractures in the dolerite contain a high-calcium chloride brine that appears to be sealed to form an artesian pressure, and beneath it is the granite basement complex (Fig. 2.28). The average precipitation (as snow) in the valley is 50– 100 mm/yr, and the evaporation rate of the strong brine at the pond in November, 1975 averaged 1.1 mm/day (it increased over 10-fold when there were heavy, cold and dry [, 20% humidity] winds). Brine samples taken at random locations and at different periods before the pond’s complete evaporation are listed in Table 2.11, and are notable for their high strontium and various trace metal content. Throughout the valley, there are seawater aerosol-formed mirabilite – thenardite crusts and small deposits, as are common in all of the coastal regions of Antarctica (Garrett, 2001; Harris and Cartwright, 1981; Goguel and Webster, 1990). Gypsum in the Don Juan Pond has a d34S value of 30.5, while at the edges of Lake Vanda the values were 39.1– 48.8. Both are estimated to have been precipitated from seawater sulfate, and then subjected to considerable sulfate-reducing bacterial action. Glacier melt water gypsum in the Wright Valley has a d34S of 9.3 –9.6, and snow melt water gypsum 12.6 –13.4 (both of the terrestrial [d34S , 10] sulfatetype; Tomiyama and Kitano, 1985).

Geology

307

Figure 2.28 Geologic structure of the Upper Wright Valley and under the Don Juan Pond (Harris and Cartwright, 1981; q1991, American Geophysical Union, reproduced by their permission).

308

Part 2 Calcium Chloride

The source of the calcium chloride in these deposits is unknown, but based upon there being no CaCl2 in the analyses of the river, glacial and other surface waters it is very doubtful that local run-off alone could be the source. This is further indicated by Tomiyama and Kitano (1985) who sampled many small pools of water and salts in the valley from the Wright Upper Glacier to the Don Juan Pond. The water compositions varied from 283 mg/liter of a calcium sulfate –sodium sulfate-type to 3.5 –104 g/liter of a sodium – magnesium-sulfate – chloride-(seawater) type brine (without considering the dissolved gypsum or calcium bicarbonate, which they did not analyze). The salts found on the ground

Figure 2.29 Lake Vanda plot of depth versus the brine’s chloride concentration; and plots of chloride versus calcium and density.

Geology

309

and near brine pools were predominantly gypsum, sodium sulfate and calcite, with some sodium chloride. Calcium chloride-type water was not found in any sample, although the salts near one small brine pool a short distance from the toe of the glacier, in its central area may have had some bishoffite (MgCl2·6H2O) in one sample and antarcticite in another, although since neither carbonate nor bicarbonate were analyzed these salts were probably merely dolomite and calcite. By contrast to the river and surface samples, as noted above, Harris and Cartwright (1981) found that under the valley’s 14 m of frozen surface soil, the groundwater in the fractured dolerite (resting on granite bedrock) contained an artesian high-calcium chloride brine. This brine appeared to be trapped and to have no pathway for entering the pond or lake, making it possible that each of these three occurrences were older calcium chloride brines that possibly originated from the same source (since their Cl versus Ca [analyses] and density plot consistently, indicating primarily a difference of dilution [Fig. 2.29]). There are some slight differences in the ratio of ions, but these can probably be attributed to their differing flow paths and basins. In Lake Vanda, there is more magnesium and sodium, and less bromine than the Don Juan Pond. The dolerite brine is similar to, but more dilute than the Don Juan Pond, with perhaps a sightly higher sodium and an unknown bromine content. Water balances based upon the river flows to the ponds and lake, and evaporation rates (after the necessary corrections for small pan-to-pond data were made) indicate that no brine leaves the lake or pond, and only a very small input of salts, with no calcium chloride are added to them each year. Thus, because of the low values for bromine and other indicators none of the brines (Lake Vanda, the Don Juan Pond and the brine in the underneath fractured dolerite) can be considered as concentrated seawater dolomitization brine. In Lake Vanda, for example, the concentration ratios of key ions to seawater at 70– 76 m depth are: Br , 0.44, Cl 3.5 –5.6, Na 0.7 – 0.9, K 1.5 –2.8 and Mg (including its equivalence in Ca) 15– 19. However, seawater aerosol also appears to have a quite different ratio of salts than seawater itself. All of the coastal areas of Antarctica receive large amounts of seawater spray each year, with sodium sulfate (as the decahydrate) being the principal mineral crystallized from it. Much less sodium chloride crystallizes, perhaps because of its lower freezing temperature, but more likely there is less sodium chloride in the aerosol. There are no analyses of this spray, but based upon the limited and partial analyses of river and surface water in the area, the sulfate and magnesium contents are proportionally high, and the bromine is very low compared to seawater. Most of the aerosol’s sodium sulfate crystallizes, thus perhaps leaving the high-MgCl2, low-Na, K and Br brine found in the surface water (and similar to that found in Lake Bonney in the next valley). It has also been predicted that there were at least four periods of glacial intrusion into the entire valley, scouring it, and almost certainly leaving large amounts of feldspar and plagioclase (from the dolerite and granite) in the layer of glacial till soil (called

310

Part 2 Calcium Chloride

sands and silty sands) under the lake and pond. The till is also said to contain considerable calcite. Finally, there are bench marks 20 – 30 m above the Don Juan Pond indicating its depth as a lake in some former wetter period. To combine these observations, the deposited modified-seawater aerosol end liquor (after Na2SO4·10H2O crystallization) would have slowly accumulated in both the Don Juan Pond and Lake Vanda basins, flushed by the glacial and snow melt waters that are currently observed. During dry periods (such as when the ocean ice shelves were greatly reduced 2500 –4000 years ago; Perkins, 2001) this water would evaporate and concentrate, eventually crystallizing salt and mirabilite (some of which dehydrates to thenardite), and perhaps forming playas as at present in the Don Juan Pond. Some of the salts on the surface would have slowly been blown away (as now happens with the very fine particles of thenardite and some salt elsewhere in Antarctica), or at times been dissolved and flushed from the basin. The strong magnesium chloride residual brine that was formed would stratify to the bottom of the lakes, and at times been strong enough to melt the ice in the underneath till (snow melting is a current use for MgCl2 and CaCl2), allowing it to react (at its slightly elevated temperatures) with the calcite to form dolomite and exchange the magnesium for calcium (the dolomitization reaction). Its acidic nature would also allow it to react with the plagioclase and feldspar to liberate considerable additional calcium and minor amounts of many other cations, as now occurs with many other CaCl2 brines. However, as the strong magnesium –calcium chloride residual brine melted the underneath ice it would become diluted, lowering its density and causing it to be displaced by stronger brines and forced to the surface. In this manner, the surface brine would slowly contain increasing amounts of calcium chloride, with even more melting power, and the circulation and calcium conversion would continue. After this process had melted all of the frozen soil, it would begin to replace water in the underneath fractured dolerite formation, where the plagioclase reaction would even more completely liberate calcium and replace magnesium with calcium. During wetter periods, both basins would become lakes (as the current Lake Vanda), and store its dilute seawater aerosol run-off waters until the next dry period. Occasionally, the lakes would overflow and flush out the surface salts. Differences in the extent of concentration, dilution, and the calcite and plagioclase content of the underneath soil and dolerite would explain the differences in the present brines. The reason why there is calcium chloride in these brines, and why it only formed in these locations might thus be due to: (1) their being in the Dry Valley Area (with a very dry, high evaporation rate climate), (2) their nearness to the ocean in order to receive seawater aerosol spray, (3) their presence in impermeable basins (the underneath granite), (4) the presence of calcite and reactive plagioclase in the basins’ surface soil, and (5) the slightly elevated rock temperatures to enhance the brine’s reactivity. Lake Bonney in the adjacent valley has collected a strong magnesium chloride brine, perhaps in the same manner as assumed here (its low sulfate and potassium contents [Br has not been

Geology

311

analyzed] probably imply that it also did not originate from trapped seawater), but apparently the floor of the Taylor Valley does not have sufficient calcite or plagioclase, or higher rock temperatures to convert much of its magnesium chloride into calcium chloride, or liberate as much additional calcium. Many other articles have been written about Lake Vanda and the Don Juan Pond, such as on the lakes’ chemistry (Nakaya et al., 1984; Yamagata et al., 1967), when the stratification occurred (Roberts and Wilson, 1965; Wilson, 1964), organic life in the Don Juan Pond (Siegel et al., 1979), and Lake Vanda’s former climate (Perkins, 2001). Antarcticite Since the Don Juan Pond forms the world’s only substantial deposit of antarcticite, its characteristics will be noted here. Its crystal habit is hexagonal, its density is 1.715 (1.700 is theoretical), and its strongest X-ray diffraction lines are 2.16, 2.80, 2.59, 2.28 and 1.978. Its optical properties are v ¼ 1:550 and e ¼ 1:495 (Torii and Ossaka, 1965). Small amounts of “divergent groups of colorless, prismatic crystals and compact aggregates (of antarcticite)” have also been found in small isolated pools at Bristol Lake in California (Dunning and Cooper, 1969; Muehle, 1971). Several authors have also reported crystals of antarcticite trapped in other minerals’ occlusions, and small amounts have been calculated to be present in the Sergipe, Brazil tachyhydrite deposit, although both sources are unconfirmed. Calcium Chloride Groundwater Throughout the world there are a number of calcium chloride groundwater occurrences that often are of an unknown origin. In many cases, they may be seawater dolomitization brines or oil field calcium chloride waters that have traveled great distances, but in some cases they appear to have originated by other means. Many authors have presented a variety of theories as to their formation, with the currently most popular of these being that they resulted from the reaction of seawater (or other water) with plagioclase to form albite. Others have suggested that the reaction was between lava or similar volcanic rocks with sea or other water (Hardie, 1990), the leaching of feldspar or mica, or the reaction of the gases that accompanied volcanoes (H2S, HCl, CO2, etc.) with calcite. Many favor the ion exchange of sodium in the original water with calcium in clays, zeolites or other rocks. All of these reactions must have occurred to some extent, but it would appear that by far the most dominant reaction to produce calcium chloride is strong-magnesium chloride brines (seawater or end liquors from extensive inland water evaporation) with calcite. Some of the various calcium chloride groundwater analyses that have been reported are listed in Table 2.12, and their descriptions will be reviewed in the following section. One of the minor occurrences of calcium chloride groundwater, but one where there has been a study on the source of the calcium chloride is in the

312

Table 2.12 Examples of Calcium Chloride Groundwater (ppm)

Namib Deserta 1 Namib Desertb 2 So. High Plainsc England, Av. No–Bad England, Av. Hi–Bad Israele 1 (5) f Israele Oil (46) f Israele 2 (8) f Israele 3 (23) f

Na

K

Mg

Cl

1823 137 4430 3500 5020 1170 1660 15,300 11,570 39,050 39,930 82,870 15,800 50,100 63,600 36,000

1620 34.3 9700 8580 18,270 6890 19,360 22,120 34,440 29,600 22,210 20,520 34,940 32,700 16,300 45,000

78 8.4 209 43 26 177 355 586 761 1160 15,670 28,730 7560 9600 — 4000

2187 86.7 3310 313 593 532 536 1350 1440 5150 42,670 19,770 41,960 7900 283 900

3850 20 32,000 20,040 39,280 12,790 34,130 64,470 77,340 130,100 249,670 265,150 208,020 170,800 138,300 142,000

HCO3

Br

I

F

TDS

476 227 143 155 193 216 257 92 164

— — — 29 39 55 138 727 620

— — — — — — 12 — 4.1

3.2 0.3 — — — — — — —

9305 305 52,260 32,900 65,570 23,850 56,630 107,785 126,920

SO4 1105 18 1470 240 64 2025 186 810 571 741 95 108 540 133 15 110 pH 6.85 7.4 6.35 — — — — — —

SiO2 50 8.8 — — — — — — — — — — — — — —

NO3 200 0.7 — — — — — — — — — — — — — —

Depth (m) — — 50 — — 1457 1549 2326 2404

— — — (continues)

Part 2 Calcium Chloride

Namib Deserta 1 Namib Desertb 2 So. High Plainsc England, Av. No–Bad England, Av. Hi–Bad Israele 1 (5) f Israele Oil (46) f Israele 2 (8) f Israele 3 (23) f Israele 4 (6) f Israele 5 (4) f Israele 6 (2) f Israele 7 (1) f Siberian Plat.g W. Caucasush Germany

Ca

Table 2.12 (continued )

e

f

Israel 4 (6) Israele 5 (4) f Israele 6 (2) f Israele 7 (1) f Siberian Platformg W. Caucasush

HCO3

Br

I

F

TDS

pH

Depth (m)

103 127 0 240 382 327

2550 2940 2990 5920 3100 3.6

— — — 0.1 — 0.5

— — — — — —

207,463 373,310 420,120 314,980 274,600 218,800

— — — — — 7.1

2085 0 0 0–400 2558 1153

— — — — r 1.264 est. 1.118 est.

a

Coastal plain in Namibia; a well in the Swartberg farm (Cagle and Cruft, 1970). As 1, except the Spes Bona farm. c Southern High Plains of Texas, Kiamichi Formation, Sanford et al (1992). d Anderson (1945); The Hi –Ba water (15 samples) also contained, as ppm: Fe 103, Li 64; no-Ba water (17 samples): Fe 13. e Bentor (1969); All are ground waters except “Oil”; No. 1 is found in very near the coast, and more diluted in the central plateau; Jurassic and Lower Cretaceous formations; Oil is oil field brines in the near-coastal Lower Cretaceous formations; No. 2 is in the Negev; Paleozoic; No. 3 is western Negev and deeper Jurassic; No. 4 is Rift Valley Arava 1 and Heimar 1; No. 5 is the Mount Sdom spring (in the Rift Valley); No. 6 is the Sdom 1 artesian well (in the Rift Valley); No. 7 for comparison is the Dead Sea (in the Rift Valley). f Values in parenthesis imply the number of wells in the average. g Bakhta megaexposure, Siberian Platform, Russia; 6 wells; Shvartsev and Bukaty (1996). h Okumi-II well; Bajocian volcanic deposits, Western Caucasus, Russia; Pastushenko (1967). i Holdorf et al. (1993). Also 290 ppm Li. b

Geology

313

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Double Lakes areas of the Kiamichi Formation in the Southern High Plains of Texas (Sanford et al., 1992). Samples were taken of the Na, Ca and Mg content of wellwater with depth, and corresponding samples were taken of the exchangeable cations in the soil. The surface water in the area has a high content of calcium/magnesium bicarbonate, while the upper formation (the Ogallala) originally contained a high-sodium/magnesium sulfate/chloride water. In areas of the High Plains where current or ancient dry lakes have formed, such as the Double Lakes, the high-sulfate water has been concentrated, and many lakes have deposits of mirabilite (sodium sulfate decahydrate), and some of them halite. In the Double Lakes basin, the Ogallala Formation contains this strong sodium chloride –sulfate brine (as wt.%: Na 2.7, Cl 4.5, SO4 1.7, Mg 0.338, K 0.337; as ppm: Ca 655, HCO3 141; pH 7.08). With increasing depth (in the underlying Kiamichi Formation) the sodium, potassium and magnesium decreases, the calcium increases, and much of the sulfate has been precipitated. This formation consists of a marine smectite clay shale that after it had been uplifted was converted to the calcium and magnesium ion exchange form (as indicated by soil samples taken some distance from the lakes). However, under the Double Lakes the soil contained considerable more exchangeable sodium (Fig. 2.30; potassium was not measured). The authors concluded that ion exchange had occurred in the soil, and that considerable sodium and potassium from the descending strong lake basin water had replaced some of the soil’s calcium, forming a CaCl2 brine (Table 2.12). This theory appears to be at least partly correct because of the low potassium in the calcium chloride aquifer, but dolomitization could have also formed this water. When the Ogallala waters evaporate to deposit mirabilite and perhaps halite (as is found in many of these lakes) a high-magnesium chloride end liquor is formed. Being heavier, it would seep into the Kiamichi Formation and have undergone the dolomitization reaction in the playa sediments to form a brine such as is now present. Later the typical Ogallala (less concentrated) water would have migrated back into the upper formation, as it now does in lakes that have commercial sodium sulfate production. In either case, gypsum would have precipitated as the change in the soil took place, forming gypsum with the low 34 S content (þ 7 – þ 10) that it now has, rather than the original marine value of , þ 20. The change observed in the ion exchange capacity in the Kiamichi Formation would have also occurred as the strong sodium – magnesium chloride/ sulfate end liquor passed through it. In northwest Queensland, Australia strong calcium chloride brines have been reported in the Proterozoic Cloncurry District, (Williams, 1994) with no indication of their source. In Yellowknife, in the North West Territories of Canada (Fig. 2.6) a very strong calcium chloride brine has been found in a gold mine and throughout the Canadian Shield area. Isotopic studies and a high bromine content in the brine strongly indicate that it is of seawater origin, and it is strong enough

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Figure 2.30 Cation exchange capacity of the Kiamichi Formation adjacent to (top figure) and under (bottom figure) Double Lakes, Texas (Sanford et al., 1992; reprinted from Water–Rock Interaction by permission from Swets & Zeitlinger Publishers).

to have been a potash deposit dolomitization brine, possibly from the Saskatchewan deposit (Bottomley et al., 1999). In the Variscan foreland of southern Belgium, calcium chloride ground waters have been reported (Muchez and Sintubin, 1997) that very likely are related the area’s Zechstein Formation halite and potash deposits. In England, there are also extensive areas of high-calcium chloride ground water and springs that may have also been derived from England’s Zechstein Formation. Anderson (1945) reported on a number of these occurrences, and especially those in the North East Coalfields. Many of the waters have a high barium, iron and lithium analysis (Table 2.12), and their total salt content varies considerably from 0.12 to 4.6 times that of seawater. In Cyprus, the Ophiolite formation contains calcium chloride brine that appears to be related to a seawater incursion, as evidenced by carbon and oxygen isotope determinations (Kerrich and Vibetti, 1985). A similar situation exists in Greece with the northern aquifer of the Filiatra limestone (Tavitian and Sabatakakis, 1994), in the soil around the Ennur coastal tract in Madras,

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Figure 2.31 Major Groundwater Basins of Israel (Vengosh and Rosenthal, 1994; Bentor, 1969; reprinted from the Journal of Hydrology, Vol. 156, p. 391, Fig. 1, q1994, with permission from Elsevier. Reprinted from Chemical Geology, Vol. 4, p. 85, Fig. 1, q1969, with permission from Elsevier).

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Tamil Nadu, India (Shepherd et al., 1994), and in the sediments of Lake Veli near the southwest coast of India (Anirudhan et al., 1991). The calcium chloride groundwater in Israel is a rather special case, since its source might have been any or all of: the Dead Sea or brine in the fault system that formed the Dead Sea; other (perhaps Red Sea-derived) calcium chloride springs in the country; from oil and gas formations; and seawater intruding into the coastal plains. However, even with such well-known potential sources, the exact nature and origin of the calcium chloride aquifers in various parts of the country are still in doubt, and many articles have been written about them. General discussions of the subject have been presented by Vengosh and Rosenthal (1994), Rosenthal (1988, 1985) and Sass and Starinsky (1979) (strontium in the brines), and Bentor (1969). The latter author tabulated many analyses of calcium chloride groundwater in various aquifers of the country (Fig. 2.31 and Table 2.12), with all of them appearing to be different forms of dolomitization brine. Those very near the coast (the first three of Israel 1 in Table 2.12) were almost pure seawater with much of the magnesium replaced by calcium, and very little of the sulfate yet precipitated. The other two samples of Israel 1 were from the central plateau, and of the same composition but considerably diluted with meteoric water. The oil field and Negev aquifers (Israel Oil and 2, 3) are slightly more concentrated seawater dolomitization brines, but the Rift Valley aquifers, springs and the Dead Sea represent considerably altered potash deposit dolomitization brine that appears to have traveled along the Red Sea fault system to Israel (Bentor, 1969). Among other articles on Israel’s groundwater, the Bet Shean – Harod multiple calcium chloride aquifer system has been reviewed by Rosenthal (1988), and the similar Eocene aquifer of Alonim Shefara by Azmon (1993). Several articles have been written about calcium chloride in the coastal aquifer (Artzi et al., 1996; Vengosh et al., 1991), and about calcium chloride brine in the coastal sabkhas (playas; Levy, 1977). In the coastal plain of Namibia, there are both very dilute and more concentrated calcium chloride ground waters (Table 2.12), and in one basin the evaporation of similar water to produce salt also forms a quite concentrated calcium chloride brine (Table 2.12; [1]). This area is similar to northern Chile in that the morning coastal fogs allow the capturing of catalytically formed nitric oxides, which in turn produce nitrates (Garrett, 1985), giving the groundwater and brines a high nitrate content (Cagle and Cruft, 1970). The source of the calcium chloride in the aquifer is unknown, but it likely has resulted from a seawater dolomitization brine –meteoric water mixture in this coastal area (i.e., the salt pan is only 32 km from the ocean). Russia has a number of calcium chloride ground waters, just as it has many marine halite and potash deposits (Fig. 2.9) and oil and gas fields. However, even though there are potash deposits in the general area of most of the following calcium chloride occurrences, it is not known whether this is their source, and many alternate theories for their origin have been suggested. Poroshin (1981) reviewed the strong calcium chloride brines found in various locations, and Shvartsev (1973) discussed

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several possible sources for the calcium, strontium and barium in these brines. Samarina et al. (1971) suggested ion exchange and rock leaching as the calcium chloride source in arid regions such as south central Russia. Seletskiy et al. (1973) discussed the supersaturated calcium chloride brines in the Angara-Lensk basin, while Pastushenko (1967) theorized that the alteration of volcanic rocks to form zeolites liberated calcium chloride for the brines in the middle Jurassic volcanic deposits of the western Caucasus region. In the eastern and central area of the Caucasian foreland, Nikanorov and Volobuyev (1968) discussed the calcium chloride brines in the Mesozoic –Cenozoic aquifers, and Volobuyev (1967) noted their increasing concentration with depth. Ovcharenko and Kurishko (1971) reviewed the characteristics of the calcium chloride brines found in the Melovyy Uplift in the Crimean Steppe. Gamalsky (1956) noted that in the Russian Platform calcium chloride brines are found in all of the deep formations, and that the different circulation patterns in the brine resulted in changes in the mineralization. In the Khapchagay uplift in the Vilyuysk area of Siberia, the deep ground water is of the calcium chloride type (Shabalin and Grubov, 1969), as are the aquifers in the Siberian Platform (Shvartsev and Bukaty, 1996). In the Volga region, certain areas have soils with a high-calcium chloride content, which is similar to some sediments in the bottom of the Caspian Sea (Slavnyy, 1966). In Saudi Arabia, the Umm Er Radhuma limestone in the eastern part of the country contains calcium chloride (Sen and Al-Dakheel, 1986), possibly either of oil field or direct seawater origin. Various calcium chloride groundwater observations have also been made in the United States. A late Alleghanian migration of such waters has been noted in the central Appalachian mountains (Evans and Battles, 1996), while in Massachusetts some calcium chloride has been found in the groundwater of Middlesex county (Toler and Pollock, 1974). The Waste Gate Formation in Maryland contains a CaCl2 water with about 6.5% TDS (Hansen, 1982). In the Trans-Pecos region of Texas, the deeper Permian and lower Cretaceous aquifers contain calcium chloride brines, and their mixing with shallower aquifers results in some quite dilute calcium chloride springs (i.e., the San Solomon, Griffin and Phantom Lake springs; Hart, 1992). Similar waters have been observed in the Texas Panhandle (Bein and Land, 1982), and both almost certainly originated from the large Carlsbad potash deposit (Fig. 2.1). In Zimbabwe, aquifers in the southern Zimbabwe Craton and northern margin of the Limpopo Belt contain calcium chloride waters (Schmidt Mumm, 1997). Tachyhydrite Deposits There are three massive tachyhydrite (CaCl2·2MgCl2·12H2O) deposits in the world, with two of them probably having once been joined and then split by a continentforming rift system (Fig. 2.32). The deposits in Brazil and the Congo appear to be positioned exactly next to each other before the South American and African continents drifted apart. Both areas contain potash deposits of Lower Cretaceous

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Figure 2.32 Location of the Congo salt – potash basin, and presumed pre-rift location of South America (DeRuiter, 1979; reprinted with permission from Economic Geology, Vol. 74:2, p. 426, Fig. 12, deRuiter, P. A. C., 1979).

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Aptian age, with significant and very similar amounts of at least 10 trace elements, most of which are quite different than found in other potash deposits (Wardlaw and Nicholls, 1972). The formations are not of normal marine origin, since there are no underlying or admixed layers of calcite (or dolomite) or gypsum (or anhydrite), and they were crystallized directly on top of continental conglomerates or shale, and tightly bounded by similar rocks (Fig. 2.33). Their bromine and rubidium contents, however, are fairly similar to seawater deposits (Table 2.13).

Figure 2.33 Structures in Cretaceous tachyhydrite-bearing evaporite formations: (1) tachyhydrite; (2) bischofite; (3) carnalite rock and sylvinite; (4) halite; (5) anhydrite; (6) carbonates; (7) clays and argillites; (8) sandstones, siltstones, and conglomerates; (I) Sergipe– Alagoas Basin; (II) Gabon Depression; (III) Congo Depression (Vysotskiy, 1988; reprinted with permission from International Geology Review, Vol. 30, No. 1. pp. 31–35, qvol. H. Winston & Sons, Inc., 360 South Ocean Boulevard, Palm Beach, FL 33480. All rights reserved).

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Geology Table 2.13 Bromine and Rubidium Analyses in Sergipe, Brazil Minerals (ppm) (Wardlaw, 1972) Marine depositsa Mineral

No. of Anal.

Range

Average

Range

Average

360 1140 4730 3670

20 –450 250 –3900 350 –4000 2000–4400

320 1850 3000 3500

7–480 60 –500

50 200

(A) Bromine Halite Sylvite Carnallite Tachyhydrite

63 6 30 43

93–643 830–1340 1960–8150 3050–4730 (B) Rubidium

Sylvite Carnallite a

6 29

18–55 55–126

38 74

Garrett (1996).

In considering the origin of the calcium chloride in these deposits, if seawater dolomitization brine from other potash or halite deposits were the source it could have provided the tachyhydrite’s seawater-type bromine and rubidium values. However, there is almost no possibility of there having been predecessor potash or halite deposits in the area, so this source is quite unlikely. It is much more probable that the deposits were formed by a brine similar to that found in some of the modern rift systems (i.e., the deep sea vents and various hot springs as in the Afar Depression, Iceland and Japan; Vysotskiy, 1988). They are also derived from seawater, but much of the magnesium has been replaced by calcium, additional calcium and many trace metals have been added by rock-leaching, and most of the bicarbonate and sulfate has been precipitated. Consequently, if during Cretaceous times such rift brines flowed into a closed, shallow inland playa-type lake in a subsiding basin, with a hot, arid climate, and did so for sufficient time, they would have formed these deposits without crystallizing the normal seawater companion masses of calcite and gypsum. This type of brine source for the deposits would also appear to be indicated by the analysis of brine in fluid inclusions found in halite in all three of the deposits. They contain 23.1– 24.6% Ca, 4.8– 5.9% Mg and 0.42 –0.49% K, with no detectable sulfate (Timofeeff et al., 1998). The geothermal nature of these brines would have also been a necessity in forming the deposits, since at present even in the harshest of deserts and hottest conditions (i.e., such as at the Dead sea, Bristol Lake and the Salton Sea [experimental] solar ponds) bischoffite (MgCl2·6H2O, a companion and

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predecessor mineral in the deposits) rarely forms except for small amounts in isolated locations, and tachyhydrite essentially never is crystallized except as an occasional rare mineral. It could form in low-humidity areas such as the highAndes or Antarctica, but the tachyhydrite deposits are in more tropical zones and near oceans. Thus, it would appear that only a geothermal brine could evaporate at high enough temperatures to form such massive tachyhydrite deposits. As an example of such present-day evaporation, in the nearby African Rift Valley at Lake Magadi, Kenya geothermal brine pools evaporate at temperatures of 45– 858C, and have formed a several billion ton trona deposit. Deep sea vent-type springs might well have been available in inland, near-rift basins to form the tachyhydrite deposits in Brazil and the Congo, and the similar Lake Giulietti-type brine might have been present at the non-rift basin in Thailand. The latter deposit is fairly near to the ocean, and has had magma intrusions such as have formed the Danakil Basin’s Lakes Asal and Giulietti calcium chloride hot springs. As a further indication of this, the potash deposits in both the Congo and Brazil still have abnormally high temperatures, such as 448C at 500 m depth in the Sergipe mine (Mraz et al., 1996). The phase chemistry of the rift-type seawater rock-leaching, dolomitization brines would have allowed deposits like each of these tachyhydrite occurrences to have been formed. However, in each of the three deposits it is also quite certain that they were formed in a non-marine basin with some terrestrial water and sediments in addition to the geothermal calcium chloride brine input. As an alternate source for the calcium chloride, it possibly could have been a magmatic fluid that originated with the molten rocks of the rifts, but based upon data from the present-day rifts and the deposit’s low boron and lithium content, this would appear to be quite unlikely. Stankevich et al. (1992) have discussed the mineralogy (including the major and minor minerals) of such deposits, while Vysotskiy (1992, 1988), and Valyashko (1975) have also speculated on the possible geochemical conditions that formed these deposits. Sergipe, Brazil This deposit occurs in three separate sub-basins of the Ibura member of the Muribeca formation in what appears to be fault bounded troughs, although the presence of 18 similar marker beds in all three basins implies that they were connected as the deposit formed, or that the evaporating conditions were very similar in each basin (Fig. 2.34; Borchert, 1977). The potash and other salts extend to the very edge of the basins, instead of the usual seawater deposits merging into halite, anhydrite or dolomite over a much wider area. A common sequence (from the bottom to the top) of the 10 – 800 m thickness of soluble salts is halite, carnallite and then tachyhydrite, although in some cases carnallite (or its sylvinite replacement) lies on top of the tachyhydrite (Fig. 2.35). The tachyhydrite is usually very coarse-grained, pure and with a transparent-to-slightly milky color (Wardlaw, 1972). These characteristics could be explained by the crystallization

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Figure 2.34 Map of the Sergipe Potash Deposit Sub-Basins and their Generalized Structure (Szatmari et al., 1979; reprinted with permission from Economic Geology, Vol. 74:2, p. 444, 445 Fig. 9, 11, Szatmari, P., Carvalho, R. S. and Simoes, I. A. 1979).

being primarily from evening cooling, followed by considerable daytime redissolving and then evening recrystallization to form large, clear crystals. There are sylvinite beds within the carnallite zones, or sometimes they occur alone. An example of the latter case is the commercial potash mine at Taquari – Vassouras, which covers an area of 85 km2. The deposit’s depth varies

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Figure 2.35 The Formation Stratigraphy at the Sergipe Potash Mine (Mraz et al., 1996; reprinted from Rock Mechanics by permission from Swets & Zeitlinger Publishers).

from 500 to 850 m, and in the initial mining area fine grained halite overlay the sylvinite (which contained some carnallite), and tachyhydrite was underneath. The tachyhydrite had a low mechanical strength, high creep tendencies and was very hygroscopic, all of which contributed to mining difficulties (Garrett, 1996). More recently, the mining area has been moved, the new mine has a depth of about 500 m, and the sylvinite is between two halite beds (Fig. 2.35). In this area, there are nine evaporite cycles, with the lower six having a sequence of halite, carnallite and then tachyhydrite. Above that the sequence is carnallite, sylvinite, halite, and limestone with shale (Mraz et al., 1996; Zharkova, 1985). Fernandes (1976) has prepared contour maps of the carnallite and tachyhydrite (Fig. 2.36) ore in one area of the deposit, and tabulated the total reserves of magnesium. In the current mining area there appeared to be 533 million tons of CaCl2 (and 387 MMmt of MgO) in massive and easily solution-minable tachyhydrite. Andrade (1984) analyzed various trace metals in the run-of-mine (impure) minerals of the deposit, and these values are compared with fairly pure minerals in Table 2.14. The values are very similar to those in the Congo deposit, and many of the world’s calcium chloride brines. They are particularly high in strontium, and relatively high in some of the heavy metals. In Sonnenfeld and Kuehn’s (1993)

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Figure 2.36 Thickness isopachs of tachyhydrite in the Sergipe mining area (Fernandes, 1976, courtesy of Minerals and Metallurgy [Brazil]).

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Part 2 Calcium Chloride Table 2.14 Various Analyses of the Sergipe Deposit (ppm) Halite

Carnalite

Sylvinite

Tachyhydrite

And.

W&N

And.

W&N

And.

W&Na

And.

W&N

Ca, %Pureb

—

—

Mg, %Pureb

—

Na, %Pureb

—

K, %Pureb

—

Cl, %Pureb

— — 1070 — 1600 700 — 130 48 190 — 21 100 24 14 — 0.5 0.6 —

— 1.60 — 0.04 39.35 35.08 — 0.20 60.65 59.81 — — 3.2 0.6 1.5 — — — — — — — — 5.2 — — 1.4

— 0.04 8.75 8.38 — 1.95 14.08 12.55 38.28 38.73 — — 1.9 0.6 6.2 — — 4.0 — — — — — 7.2 — — 2.4

— — 2.5 — — — — — — — 180 — 360 180 — 72 70 170 — 25 92 50 11 — 0.8 0.6 —

— 0.04 — 0.08 — 4.28 52.45 42.50 47.55 48.42 — — — — — — — — — — — — — — — — —

— 6.8 — 6.3 — — — — — — 1930 — 1300 130 125 100 71 63 — 34 27 13 6 — 1.7 1.0 —

7.74 7.37 9.39 9.46 — 0.08 — 0.15 41.00 40.44 1820 1360c 7.2 0.4 1.5 — — 2.5 50c — 7.5 — — 3.3 — — 0.4

Sr Ba Mn Zn Cr V Cu Fe Ni Pb Mo Co F Ag Be B

— —

— — 750 — 930 125 — 150 54 30 — 22 54 40 9 — 1.6 0.5 —

And. ¼ Andrade (1984) (apparently a run-of-mine sample). W&N ¼ Wardlaw and Nichols (1972) (apparently fairly pure crystals). a Sylvite. b As wt.%; Pure means pure crystals (theoretical analysis); the second line is that of actual samples. c Sonnenfeld and Kuehn (1993). They also listed for tachyhydrite 440 ppm K, and 240 ppm Na (also presumably fairly pure crystals).

chemical analysis of tachyhydrite, they indicated a very slight excess of Ca in the sample. This caused them to speculate that perhaps as much as 3.5% antarcticite (CaCl2·6H2O; or CaCl2·4H2O, sergipite?) might be present, but neither sulfate nor carbonate were analyzed, so the excess Ca could have been gypsum or calcite. Mraz et al. (1996) have given a detailed analysis of the amount of closure (creep) occurring in a previously mined sylvinite panel, while Trummel (1974) examined

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the laboratory creep characteristics of sylvinite, halite, carnallite and tachyhydrite from the mine. Many other articles have been written about the deposit, such as by Hite (1973). Congo The Lower Cretaceous salt deposits of the adjacent Gabon and Congo basins (Fig. 2.37) are very similar to those at Sergipe, Brazil. They appear to have been formed as interconnected terrestrial lakes whose depositional period ended with seawater encroachment, perhaps as the period of rifting, or continent separation and drifting became more severe. There are only terrestrial fossils under the deposits, but marine fossils begin to appear in the overlying gypsum bed that caps the deposits (which probably formed by the calcium chloride brine reacting with intruding seawater). In the Congo area, the Douala (in Cameroun), Gabon, Congo (in Congo, Zaire and Angola) and Cuanza (in Angola) basins appear to have been

Figure 2.37 Typical depositional cycle of the Congo potash deposit (DeRuiter, 1979; reprinted with permission from Economic Geology, Vol. 74:2, p. 425, Fig. 11, deRuiter, P. A. C., 1979).

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connected during the time that the potash – tachyhydrite deposit was forming. Their lower depths are filled with fresh water “rift sediments” and fossils, and the section immediately below the salts has a high organic content that has formed oil shale. The deposits began to form apparently just as the subsidence and rifting started. Even though the salts in the Congo deposit are very similar to those in Brazil, as seen in Fig. 2.33, the detailed stratigraphy is somewhat varied. The Congo area contains less sylvinite, about the same amount of carnallite (estimated to be about 15 –20% of the deposit) and more bischoffite. In the central part of the deposit,

Figure 2.38 Location map of the Khorat Plateau, Thailand (Hite, 1986; reprinted by permission of the United Nations Economic and Social Commission for Asia and the Pacific).

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the normal depositional sequence is a thin black layer of shale followed by halite, then carnallite and halite (with some of the carnallite converted to sylvinite), and finally up to 150 m of tachyhydrite and bischoffite. In the lower and upper sections of the basin sometimes there are relatively thin “inverse cycles” where the carnallite is followed by increasingly pure halite. The tachyhydrite beds are often lenticular, and sometimes occur as 1 –6 m intercalated beds of tachyhydrite or bischoffite –tachyhydrite with carnallite or carnallite –halite. There are very few insolubles in the soluble salt beds (i.e., 2 –5%; Vysotskiy, 1988). The Gabon and Congo basins’ deposits are similar to each other, and the strata correlate well, but they are totally different than the adjacent basins. The others began with similar sediments, but as the rifting started they were apparently filled with seawater, and became marine formations (Garrett, 1996). Zharkova (1985) has provided an additional review of the stratigraphy and lithology of this deposit’s Congo Basin. Some of the impurities in one sample of tachyhydrite are listed in Table 2.14. Thailand The Khorat Plateau of Thailand and part of Laos contains a large (, 170,000 km2) evaporite basin from Cretaceous to early Tertiary age (Fig. 2.38). In its Maha Sarakham Formation, there is a massive halite, carnallite (with local zones of sylvinite), tachyhydrite and bischoffite formation. The deposit is divided into two sections: the northern Udon –Sakhon Nakhon (21,000 km2) and southern Khorat –Ubol (4000 – 5000 km2) basins. As with the tachyhydrite deposits in Congo and Brazil, it lies directly on terrestrial sediments, with very little anhydrite and essentially no dolomite in or near the deposit. It differs from the other deposits in having three intrusive basaltic flows in or adjacent to it, and being interbedded and overlain with terrestrial red bed sediments (with little anhydrite and dolomite, and considerable clay). Essentially the only insoluble impurity in the salts is boracite (MgClB7O11, a high-temperature borate salt, usually of geothermal origin), which averaged 3.7% in one drill core section. The bromine content in the different salts is variable but similar to the other tachyhydrite deposits. A typical stratigraphic section of the formation is shown in Fig. 2.39. The deposit’s thickness averages 250 m, but reaches 1.1 km in some areas. The character of the salts are similar in both basins, indicating that they were connected as they formed. The small amount of anhydrite in the occasional interbeds of clay have d34S values of 6.4– 10.9, indicative of lacustrine water (the d34S in seawater is , 20). Also, some of the thin interbeds of anhydrite in the halite had a sieve-like structure and 34S values of 14.3 – 17.0 (El Tabakh et al., 1999), which are typical of non-marine sulfate that remains behind after considerable sulfate-reducing bacterial attack. Since seawater vent or hot basalt brine has such a low-sulfate content, most of the sulfate would have entered

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Figure 2.39 General stratigraphy of the Khorat Plateau, Thailand (Japakasetr and Workman, 1981; reprinted from the AAPG Studies in Geology. AAPGq 1981, reprinted by permission of the American Association of Petroleum Geologists, whose permission is required for further use).

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the basin from terrestrial sources, as with the “red clay” marker beds. The lack of massive amounts of gypsum and dolomite in the basin preclude there having been a direct seawater input. The tachyhydrite in the deposit is yellow to honey colored, and it is always found in alternating bands with carnallite, primarily in the central and western part of the basin in beds 1 –95 m thick. More than half of the beds are 20– 50 m thick, and the average is 39 m. Smaller amounts of tachyhydrite are sometimes found with carnallite as the basement halite slowly grades into the other salts. The sylvinite appears to be a replacement product of the carnallite (as is always the case) and overlying it, and the sylvinite is never accompanied by tachyhydrite (Garrett, 1996). Many other articles have been written about this deposit, such as by Japakasetr and Workman (1981). Other Tachyhydrite Occurrences Small amounts of tachyhydrite have been observed in several other potash deposits, and appear to have been formed by some form of sub-surface concentration process. It has been deposited on the collars of drill holes into the very strong dolomitization brine near the Irkutsk potash deposit. Small amounts have also been found in the German Zechstein potash deposits (along with some baeumlerite, KCl·CaCl2), and there are many Zechstein references such as by Heide and Kuehn (1965) and Marr (1959). Turkmenian potash deposits composed of mixed potassium – magnesium salts with a small amount of tachyhydrite have also been reported (Azizov, 1972). In each case, geothermal formations have been indicated near the deposits. Tachyhydrite Characteristics Tachyhydrite has a rhombohedral crystal structure, with a density of 1.673 and Z of 3. Its bonding structure can be represented by [CaCl6][MgCl2(H2O)6]2, and its interatomic distances have been estimated by Clark et al. (1980). Erd et al. (1979) have determined its X-Ray diffraction pattern. It does not crystallize at temperatures below 22.88C (Charykova et al., 1992). Sinclair et al. (1996) studied its phase transitions, while Horita (1989) determined the fractionation of deuterium and 18O in the water of hydration as tachyhydrite was being crystallized. As would be expected, it was indicated that deuterium was depleted in the water of hydration, but surprisingly 18O was enriched. Many articles or patents have also been presented claiming methods for the formation, recovery or processing of tachyhydrite. Metallic magnesium of . 99% purity has been claimed after the dehydration, fusion and electrolysis of tachyhydrite. Dehydration was claimed to be achieved to 0.5– 1.0% H2O by heating for 4 hr at 2608C and then 8 hr at 7258C with a chlorine partial (vapor) pressure of 0.1 –0.95 atm, or under vacuum in the presence of HCl, H2O or N2. It could be recovered by solution mining, or separated (dry) from carnallite by freefalling electrostatic separators. It could also be recovered from the wastewater

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of iodine –bromine production. It is said to be formed in some oil deposits in dolomite formations, and in natural gas pipe lines under certain conditions. It was also formed as one of the trace constituents when salt from an ion exchange process was dried (Niino et al., 1993). Finally, it could be used as a chloridizing agent to solubilize spodumene by heating to 11508C. Calcium Chloride Brine in Mineral Deposits Since microscopes with temperature-controlled specimen holders, and analytical instruments that can focus on small occlusions in other rocks became commonly available, calcium chloride brine has been inferred to have been present in the formation of many minerals and rocks. Occlusions with a freezing (initial crystallization) point (Tf) or melting (solubilization) temperatures [Tm] from below about 2 15 to 2 308C (even with the two temperatures far different from each other, such as by 20 –808C; Davis et al., 1990) have been assumed to be calcium chloride solutions. Also, several micro-instrumental (analytical) methods (Sampson and Walker, 2000) have frequently indicated CaCl2 in both the occlusion crystals and brine. Its presence has been explained by: (1) having been one of the reactants in the solution from which the mineral was crystallized, (2) being formed at a later time by the original brine in the occlusion reacting with the mineral, (3) being present as an inert material as the deposit was forming, or (4) high-CaCl2 brine penetrated, leached and recrystallized the rocks or deposit at a later time. However, very reproducible data can also be obtained from very large supersaturation (metastable equilibrium, or crystallization-point lowering) when crystallizing solutions of MgCl2, NaCl, potassium double salts and many others. This is exemplified by the frequent large differences obtained between Tf and then Tm on the just-crystallized solids, when the values should be the same. Such supersaturation is especially common with inclusions, as illustrated by the extensive work done on halite and potash deposits (see Fabricus (1983)). For instance each of the values for Th (homogenization of vapor and liquid), Tf (or c) and Tm with primary halite and sylvite crystals have indicated that the massive Zechstein Formation crystallized at temperatures of 97– 1578C, or near or far above seawater’s boiling point, which of course is impossible (Garrett, 1996). Instrumental methods on these small samples enclosed in other minerals can also be misleading, and unfortunately, very few positive identifications have been made on these calcium chloride occlusions by chemical analyses. As examples of the literally hundreds of reported calcium chloride occlusions, CaCl2 has been inferred in occlusions in the West Gore antimony – gold deposit in Nova Scotia (Kontak et al., 1996), and in barite deposits in the western Jebilet of Morocco. It was speculated that the latter deposit contains fluids that may have been related to central Atlantic rifting (Valenza et al., 2000). Similar inclusions have been found in the Long Lake

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calcite – fluorite vein, central Adirondacks, USA (Bird and Darling, 1996) and the calcite cement in the Smackover Formation of southern Arkansas, USA. In the latter deposit, the Th values indicated formation temperatures of 58– 1588C, with the CaCl2 inferred from Tf values from 2 51 to 2 688C, and Tms from 2 15 to 2 328C (Klosterman, 1981). The Mount Isa, Australia copper ore was also indicated to contain CaCl2 inclusions (Heinrich et al., 1989, 1991). Other copper deposits that have reported calcium chloride inclusions are the Little Stobie deposit at Sudbury, Canada (with precious metals; Molnar et al., 1999), the Granisle porphyry copper deposit in British Columbia, Canada (Quan et al., 1987), and the Aquas Claras deposit in the Carajas region, Para, Brazil (Gaia da Silva and Villas, 1998). Small amounts of a calcium chloride brine have been found in a strike-slip fault system in Belgium (Muchez and Sintubin, 1997), and Parry (1998) has noted that solutions with up to 19% CaCl2 have been found in a few other fault zones. Microscopic observations on these (and many other occlusions) indicate that they were formed at elevated temperatures and pressures, which the authors suggested initiated rock-brine reactions that formed the calcium chloride. The South Platte fluorite-REE pegmatites, Colorado, USA occlusions when crystallized were indicated to contain ice, hydrohalite and antarcticite by cryogenic laser spectroscopy (Walker, 1998). Raman spectra of occlusions in garnet, apatite and quartz in the Austroalpine Otzal Basement complex indicated that when cooled they contained hydrohalite and antarcticite (Kaindl et al., 1999). The gold and copper deposits of the Cloncurry district, Queensland, Australia was inferred to contain a calcium chloride solution that formed the deposit at 220– 3608C (Williams et al., 2001), and a similar deposit in the Mallery Lake area, Nunavut, Canada also recorded strong calcium chloride in lower temperature (90 –1508C) occlusions (Turner et al., 2001). A strong-calcium chloride brine was interpreted to be in fluid inclusions in the Leinster granite and quartz veins of SE England (Moran et al., 1997). Calcium chloride brine has also been indicated in the Soultz-sous-Forets, France granite alteration zone (Savary et al., 1997), and in a granitic pegmatite in Colorado a cryogenic Raman spectroscope indicated hydrohalite and antarcticite (with a Tm from 2 50 to 2 708C; Samson and Walker, 2000). Calcium chloride was also indicated in a shear zone of the Harare greenstone belt in Zimbabwe (Mutemeri et al., 1997), and in the greenstone of the southern Zimbabwe Craton and the northern margin of the Limpopo Belt (Schmidt Mumm, 1997). The Abitibi greenstone belt in Canada has also reported calcium chloride in its gold-bearing quartz veins, with some indications that the deposit was formed from ancient sea floor hydrothermal vents (Neumayr and Hagemann, 2002; Kerrich and Ludden, 2000). As noted in previous sections, the occurrence of strong calcium chloride dolomitization brine in secondary halite is quite common (Zimmermann, 2000; Grishina et al., 1992). The lead –zinc – barium district at Thunder Bay, Ontario, Canada appears to contain calcium chloride brine (Haynes, 1988), as does the Largentiere lead – zinc

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deposit in France (Leost, 1999). The migmatites of the Tatra Mountains, western Carpathia have occlusions with a strong calcium chloride brine (along with pure nitrogen; Hurai et al., 2000), and molybdenum – bismuth mineralization in the Cadillac deposit Quebec, Canada has indicated antarcticite in frozen occlusions by SEM –EDS analysis (Taner et al., 1998). The Ni – Cu –Pt deposit at Lindsley, Sudsbury, Canada (Molnar et al., 1997) is said to contain calcium chloride, as is the PGE mineralization at the New Rambler deposit in SE Wyoming (Nyman et al., 1990). Numerous authors have reported calcium chloride brine in quartz vein inclusions, such as in the Batum Salt Dome, N. Jutland, Denmark (Fabricus, 1984), the Sandia Pluton contact (Pletsch-Rivera et al., 1998), the Valles caldera, New Mexico (Sasada, 1988), the Strzegom pegmatites, Poland (Kozlowski, 1994), the Spanish Central System (Martin et al., 1997), the Bushveld Complex, South Africa (both CaCl2·6H2O and CaCl2·4H2O crystals were indicated in the quartz occlusions by Raman spectra, implying a . 42% CaCl2, boiling formation brine; Schiffries, 1990), boreholes in northern Switzerland (Mullis and Stalder, 1987; and the Mori geothermal field in Hokhaido, Japan [Muramatsu and Komatsu, 1999]). Calcium chloride was indicated in rare earth deposits (Shmulovich et al., 2002), and the Puebla de Lillo, Spain talc deposits (formed with 0– 25% CaCl2 at 280 –4058C; Tornos and Spiro, 2000). There were calcium chloride occlusions in the San Rafael Lode tin deposit in Peru (up to 40% CaCl2; Kontak and Clark, 2002), a tin skarn in the Yukon Territory (Layne and Spooner, 1991), and the tin –polymetallic ore field in Guangxi, China, which is said to contain daughter crystals of antarcticite (Fu et al., 1993). The Kombolgie uranium deposits in the Northern Territory, Australia (Brisset, 2000, Derome et al., 2000) and the Gays River and Jubilee, Nova Scotia, Canada zinc and zinc – lead deposits are also reported to contain fluid inclusions of CaCl2 (Kontak, 1998; Savard and Chi, 1998), as is the Mascot-Jefferson City zinc district, Tennessee, USA, using SEM and “energy dispersive” analysis of occlusion “decrepitates” (Haynes et al., 1989). The Balmat-Edwards District zinc (sphalerite) deposit, northwest Adirondacks, NY, USA is also reported to contain CaCl2 occlusions (Hill and Darling, 1997), along with the Lyonsdale area of northern New York (“a fossil shield CaCl2 brine”; Garside and Darling, 1993). Occlusions in the King Island scheelite skarn deposit may contain CaCl2·6H2O crystals as well as a CaCl2 brine, and it is felt that the calcium chloride was formed during deposition by reactions of the type: (1) Fe2þ þ CaCO3 ðcalciteÞ þ þ 3þ 2SiO2ðaqÞ þ 2H2 O ! CaFeSi2 O6 ðhedenbergiteÞ þ HCO2 þ 3 þ 3H ; or (2) 2Fe p 3CaCO3 ðcalciteÞ þ 3SiO2ðaqÞ þ 3H2 O ! Ca3 Fe2 Si3 O12 ðandraditeÞ þ 3CO2 þ 6Hþ ; and (3) 6Hþ þ 3CaCO3 ðcalciteÞ ! 3Ca2þ þ 3COp2 þ 3H2 O þ pore space (pthe CO2 would probably be in the form of CaCO3). In zones where there was no calcite there was no calcium chloride, and in other zones the CaCl2 concentration was proportional to the amount of Fe-calc-silicate crystallized (Kwak and Tan, 1981).

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Reactions of Calcium Chloride with Minerals The leaching (partial solubility) of a number of minerals in calcium chloride solutions has been discussed in previous sections, and reviewed by Korzhinskii and Shmulovich (1988) and Vakulenko and Razina (1992). The solubility of anhydrite at temperatures of 250 – 3258C was studied by Templeton and Rodgers (1967), being unexpectedly high, which they attributed to the formation of Ca2SO2þ and Ca(SO4)22 ions. The solubility of anorthite in supercritical 4 2 calcium chloride solutions was studied by Roselle (1993) at temperatures of 400– 6008C, a pressure of 2 kbar, concentrations of 0.005 – 6 M and pH values of 0.06– 6.2. The amount of calcium liberated, and the formation of albite was appreciable. In the studies it was assumed that aqueous SiO2 and Al(OH)3 would always be present, and that the calcium chloride occurred in the forms of CaCl2 3, CaClþ, Ca(OH)þ and CaCl2 under these conditions. Phase relationships in the reaction of calcium chloride with apatite to produce chlorapatite at temperatures of 600 –11508C have been studied by Morton (1961), as were the kinetics of the CaCl2 – hydroxyapatite reaction. The rate of calcium chloride solutions’ conversion of aragonite to calcite at 1008C was investigated by Berndt and Seyfried (1999), while determinations were made on calcite’s solubility in calcium chloride solutions by Nagy and Morse (1986), its precipitation rate (Tomiyama and Kitano, 1984), its speciation in supercritical CaCl2 solutions (Baumgartner, 1991), and inhibiting agents for its precipitation (Akagi and Kono, 1995). X-ray studies of CaCl2 – chabazite and dehydrated CaCl2 –natrolite were made by Fang (1961). When clinoptillite was treated with CaCl2, the cations were exchanged, and upon heating the Ca – clinoptillite inverted (Shepard and Starkey, 1966). The solubility of fluorite in cool-to-supercritical solutions of calcium chloride was determined by Malinin and Kurovskaya (1999, 1994) and Ryzhenko et al. (1999). At low temperatures, the CaCl2 appeared to be disassociated into Ca2þ and 2Cl2 ions, but at higher temperatures a wide variety of other ions and complexes were formed. The surface properties of mica in CaCl2 solutions was studied by Kjellander et al. (1990). Liquid occlusion analyses, d18O data and drill core analyses have indicated that portions of the Troodos ophiolite (basalt-gabbro) formation in Cyprus has been converted by marine dolomitization brine at high temperatures into actinolite, anorthite, epidote, albite and quartz, and at temperatures below 1008C into laumontite and calcite (Kerrich and Vibetti, 1985). Kalinin (1966) studied the reaction of plagioclase with calcium chloride solutions, producing garnet of the grossularite –andradite series along with anorthite. Dujon and Lagache (1986) considered the influence of fluid mixing on the cation exchange of CaCl2 solutions and plagioclase at 7008C. The solubilities and thermodynamic properties of solutions of calcium chloride and silica were determined by Popp and Frantz (1979), and their electrokinetic potential by Jichova and Havlica (1999). The solubility of amorphous silica in calcium chloride-type geothermal waters was investigated by

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Gallup (1989), while the change of stishovite to a denser phase by CaCl2 was examined by Shu et al. (1996). Winkler and Luettge (1999, 1997) studied the reaction of tremolite with calcite and quartz to form diopside, CO2 and H2O, with and without the presence of CaCl2. The temperatures ranged from 630 to 7208C (the equilibrium temperature was 6178C), the pressure was 5 kbar, the concentration 0 – 5.5 wt.% CaCl2, and the reaction time up to 34 days. The presence of the calcium chloride increased the reaction’s rates more than 10-fold to values that were much higher than expected. Zimmerman et al. (1996) performed the opposite type of reaction, and studied the formation of tremolite from calcium chloride solutions. Studies on the vermiculite mine at Libby, Montana have been made concerning the origin of the deposit. The ore lies in an augite pyroxenite that has been altered to biotite, hydrobiotite and vermiculite, with numerous syenite dikes in the pyroxenite. Ion exchange tests have shown that biotite alters to vermiculite at room temperatures in solutions with . 0.1 wt.% CaCl2, and it was theorized that the augite was altered to biotite by the intruding syenite dikes. Later some of the biotite was altered to hydrobiotite and vermiculite by descending calcium chloridecontaining near-surface waters (Bassett, 1959). In a different type of reaction, potassium chloride solutions can be reacted with gypsum to form the mineral syngenite and calcium chloride, and the syngenite in turn can be dissolved in water to form potassium sulfate. This reaction has been used commercially in several operations, but was only economically successful as a means of lowering the end liquor loss at a K2SO4 plant. Orlova et al. (2000) have recently noted that the presence of carbamide greatly improves the reaction. Clay Minerals Bentonite and the other clay minerals have considerable ion exchange capability, and their physical characteristics change dramatically depending upon the content of the salt solution in their pore space, and whether the clay is in the calcium or sodium form. As examples of publications considering the influence of calcium chloride on clays Sjoblom et al. (1999) noted the slowness of water to penetrate compacted bentonite (Na –montmorillonite is its major constituent), but that a calcium chloride solution could rapidly penetrate the bentonite and allow it to be washed away. Fresh water causes compacted bentonite to swell and to produce free-surface particles by exfoliation. These particles form a gel which further closes the pores to water uptake, while the CaCl2 solution causes the exfoliated material to shrink (or at least swell less). This allows more solution to enter the pores, causing differential expansion and a lower gel strength so that the clay may be more easily washed away. Di Malo (1996) also studied the osmotic and mechanical effects of calcium chloride solutions upon bentonite. Water-saturated bentonite when exposed to CaCl2 solutions became more consolidated, had a large decrease in deformability, and an equally large increase in shear strength. These effects could not be changed by water washing, and X-ray diffraction analysis showed that calcium had ion

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exchanged for the sodium in the bentonite. The osmotic effects became almost negligible. Anson and Hawkins (1998) also noted the increase in shear strength with bentonite and kaolinite when treated with solutions as dilute as 80 mg/liter of Ca2þ, and Ahmad et al. (2000) found that a 50:50% mixture of bentonite and kaolinite had about the same coefficient of consolidation and permeability as bentonite alone, and could form effective pond liners for strong acidic or alkaline solutions, but not when calcium chloride was present. Wang et al. (1996) studied the effect of CaCl2 solutions on the dehydration of Ca-exchanged montmorillonite. Haydon (1983) noted that as a calcium chloride brine passed through compressed smectite at elevated temperatures and pressure ion filtration and stable oxygen isotopic fractionation occurred. Dennis (1991) studied the compaction and swelling of Ca –smectite in calcium chloride solutions. Soils Many articles have been written about the interaction of calcium chloride with soils, as indicated by the following examples. Ernani and Barber (1993) have studied the effect of cation leaching (ion exchange) of acid soils by treatment with CaCl2 solutions, while Chen and Shao (2001) noted that the presence of calcium chloride greatly aided in the coagulation of suspended sediments in aqueous solutions. Dakshinamwite and Chandool (1966) examined the conductivity and ion exchange capacity of soils and clays. They found ion exchange capacities of 10– 105 meq/ 100 g, and a relationship between the soil conductivity in a CaCl2 electrolyte and the ion exchange capacity. Baudracco and Tardy (1988) studied the dispersion and flocculation of clays in unconsolidated sandstone reservoirs by percolating CaCl2 solutions. Murashko et al. (1970) examined the penetration coefficients of various salt solutions on clay soils, finding that CaCl2 solutions ranked much higher than NaCl and pure water. White et al. (1964a) noted the improved permeability of oil and gas formations if first flushed with a CaCl2 rather than NaCl brine, due to the reduced swelling of clays in the structure. Stern and Shackelford (1998) studied calcium chloride solutions’ permeability in sand – clay mixtures, while Aoubouazza and Baudracco (1992) did similar studies on clayey sandstones. Little (1992) determined the relationship between soil pH measurements in water and CaCl2 suspensions. Alsharari (1999) studied the reclamation of highly alkaline soil with gypsum, langbeinite and calcium chloride, while Mukhtar et al. (1974) studied the effect of CaCl2 solutions on the structural stability of two Vertisols: Gezira clay from Sudan, and Houston Black clay from Texas.

PROCESSING The processing of calcium chloride brines can be either a very simple or a complex series of operations, depending upon the purity required, the desire to produce solid as well as liquid products, and whether other by-products are

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also recovered from the brine. A general article on the subject has been presented by Kotsupalo et al. (1999), and descriptions of the current actual production processes are reviewed below. Other information on processing is included in the commodity reviews by the U.S. Bureau of Mines (1992 –1982), articles in Chemical & Engineering News (2001) and in chemical engineering encyclopedias. Michigan Dolomitization Brines Dow Chemical Company The Michigan calcium chloride brines have the distinction of playing a very important role in the early development of the United State’s chemical industry. By 1897, Dr. Herbert Dow had investigated these unique brines and raised capital to commence the recovery of some of the components in one of the fairly shallow brines at Midland, Michigan. The basic chemicals were then converted into a series of much more valuable products, and the total facility became a highly diversified chemicals complex. The operation was successful from the beginning, and soon was one of the major facilities in the country’s fledgling chemical industry. It still remains as one of Dow Chemical Company’s basic operations, and since 1914 calcium chloride has been one of the major products recovered from the Michigan brines. Figure 2.40 shows the major steps involved in the original process. The first operation was with brine from the relatively shallow and dilute Marshall formation (Table 2.4), and began by passing the brine through an electrolytic cell (Dr. Dow was an electrochemist). The current was carefully controlled in order to oxidize the bromide ions to elemental bromine, which could

Figure 2.40 Flowsheet for the original (1914) Dow Chemical Brine Recovery Process (Pavlick, 1984).

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then be blown out of the cell with air, condensed and converted into various compounds. The brine was next sent to a triple effect evaporator where the solution was concentrated and much of the sodium chloride crystallized. The sodium chloride was removed, dissolved, and sent to electrolytic cells to be converted into chlorine and caustic soda (along with small amounts of hydrogen). Evaporation was then continued to crystallize some magnesium chloride (with the salt), and it was separated and converted into magnesium oxychloride, a material popular at that time for producing floor tile. The brine was next reacted with slaked lime (Ca[OH]2) to precipitate almost all of the remaining magnesium, and the magnesium hydroxide was filtered, washed, and converted into a variety of magnesium compounds (such as magnesium sulfate heptahydrate, epsom salts). The remaining brine was again sent to evaporators to be concentrated to a 38% CaCl2 solution as the final step in the operation, and the sodium chloride that crystallized during evaporation was also removed and sent to the chlorine – caustic plant. This brine was adequately pure for most calcium chloride sales, and could be marketed directly or further evaporated to form various solid products. Later brine from the deeper and stronger Sylvania Formation (Table 2.4) was utilized, and since it also contained iodine, it and many of its derivatives were also produced (Fig. 2.41). The brine was first acidified and treated with just enough chlorine to react with the iodide ions and covert them into elemental iodine. The iodine was then blown from the brine with air in a packed tower, and re-absorbed for separate processing in a smaller tower. The residual brine was next heated, passed through a second large packed tower, and blown with steam and chlorine. The bromide ions were converted to bromine by the chlorine and carried by the steam from the tower. As much water as possible was separated from the bromine, and the bromine was further dried with sulfuric acid and re-distilled in a separate tower. Magnesium chloride and calcium chloride were then recovered from the brine (Pavlick, 1984). Figure 2.42 shows how the operation, and the chemicals derived from processing the Michigan brines has changed over the years. In 2002, brine was only extracted (Fig. 2.43) from the comparatively shallow and rich Filer Formation near Ludington at a 910 m (3000 ft) depth. This , 130 km (80 mi) long sandstone formation is about 30 m (100 ft) thick and has a 10– 15% porosity. The Filer brine contained 17.1– 17.4% CaCl2, 9.2 –9.9% MgCl2, 2.0 –4.0% NaCl, 0.8– 0.9% KCl and 0.232– 0.241% Br. Its specific gravity varied from 1.277 to 1.287 g/ cc at 258C, its viscosity was 2.1 cp, specific heat 0.69 cal/g/8C, pH 4.4 –4.8, and its boiling point 111.68C. To recover the brine, Dow utilized 17 wells, with about 10 normally in service at any given time, pumping 100– 150 gpm/well, or a total of 850 – 1200 gpm. To maintain a vigorous flow rate, it was necessary to re-inject a similar amount of wash, dissolved salt and other water near the periphery of the well field. Since the Filer brine does not contain iodine, bromine was the first product recovered from it by the “steaming” process as described above. The brine was

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Figure 2.41 Some of the products previously produced from the Sylvania Formation by the Dow Chemical Co. (Pavlick, 1984).

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Figure 2.42 Dow Chemicals process for recovering calcium chloride and other products from Michigan Brine in 2002 (after Pavlick, 1984).

next reacted with slaked dolime (CaMg[OH]2) to precipitate magnesium hydroxide. The dolime (CaO·MgO) had been purchased from Ohio and delivered to the plant by rail, and then reacted with water (slaked) in the plant’s previously used dolime roasting– slaking facilities. The precipitated magnesium hydroxide trihydrate was then settled, filtered, washed and finally re-pulped, with about 67% of it pumped to a nearby company that converted it to dead burned magnesia, and the remainder was sold for water treating or other uses. The residual brine had a pH of 8– 9, contained 24 –25% CaCl2, and was next sent to triple effect evaporators to be concentrated to 32 –45% CaCl2. The first effect (installed in 1978) was of a forced circulation type with an external heat exchanger. The second and third effects were of the calandria type (dating from 1942), with an agitator to increase the flow rate through its internal heat exchanger tubes. This greatly reduced the scaling problem, and re-tubing was only required on about a 10-year schedule. The salt that crystallized during the evaporation process was settled and then removed from the brine in solid bowl centrifuges, to be re-dissolved and added to the re-injection water into the Filer Formation. The centrifuges were rebuilt every 2 –3 years. Steam for the evaporators was purchased from an adjacent cogeneration power plant, and cooling water for the third stage barometric condenser was withdrawn from an estuary of Lake Michigan next to the plant. However, Zebra mussels tended to grow prolifically in the cooling water, and had to be periodically killed with chlorine and removed from the system. A vacuum pump was used for the second stage non-condensables. Part of the concentrated solution was sold directly, and part was further concentrated to 77 – 78 CaCl2 flakes or 90– 94% pellets (Table 2.21(A) and (B)). The 42– 45% CaCl2 brine contained about 2.5% NaCl (with the KCl expressed as equivalent NaCl) and 1% other salts. To make solid products, this solution was

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Figure 2.43 Schematic drawing of the brine recovery and reinjection wells in the Michigan Sylvania and Filer Formations (after Pavlick, 1984).

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first further concentrated in a single-stage forced circulation evaporator using high-pressure steam. For the production of 74– 78% CaCl2 flakes, some of this concentrated solution was sent to heated pans below large rotary drums. The cooled drums were dipped into the hot solution, causing some of the calcium chloride to cool and freeze on the drums’ outer surface as the drums rotated. These approximately calcium chloride dihydrate solids were then scraped off on the other side before the drum dipped into the pan again. The thin flakes that were formed (about 0.5– 2 mm thick by 6 –12 mm wide; Fig. 2.44) dropped into the top of a very large multi-tray dryer. In the top two-thirds of the dryer the flakes were contacted by hot flue gas, and in the lower third they were cooled by air. The cooled product was then bagged, put into bulk bags or shipped to the customers in bulk. To produce the 90 – 94% CaCl2 solid product, a portion of the concentrated liquid was sprayed into a dryer and directly heated by very hot flue gas. Solids were removed from the dryer, cooled and screened to the desired particle size, while the exhaust gas was scrubbed and then vented. The oversize was crushed and returned to the screen, while the fines were recycled to the 45% feed solution, or compacted into almond-shaped pellets. The 90 –94% CaCl2 pellets were also shipped in either bulk or bags.

Figure 2.44 Typical flaked calcium chloride (Dow, 2001, courtesy of The Dow Chemical Company).

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Figure 2.45 The Dow Chemical Calcium Chloride Plant at Ludington.

In 2002 the 600,000 st/yr capacity (of equivalent 100% CaCl2 product) Ludington plant (Fig. 2.45 shows the entire plant, and Fig. 2.46 the solids plant) had 225 employees. Its products were bromine, sodium bromide, calcium bromide, magnesium hydroxide, and the calcium chloride in liquid, flake, pellets or a food grade made by filtering the final 45% CaCl2 solution. The products were shipped by pipeline, truck, rail or by barge through the Great Lakes, and a large inventory was maintained for the seasonal demand. In early 2003, it was announced that within about 1 year the plant would stop manufacturing magnesium hydroxide, and close their brine-gathering wells. They would instead purchase a magnesiumdepleted 24% CaCl2 brine from Martin Marietta Magnesia Specialties Co. in Manistee, Michigan (their former primary magnesium hydroxide customer), and deliver it by pipeline to their plant. This would cause a reduction in their staff of about 30 people (Busch, 2003, 2002; Dunklow, 2002; Chemical Market Reporter, 2002; Pavlick, 1984). Other Companies The General Chemicals Group in Manistee and the Wilkinson Co. in Mayville also produced calcium chloride from Michigan brines in 2002. However, in late 2002 the General Chemical Group announced that they had closed their Manistee plant with a capacity of 450,000 st equivalent flake/yr. None of the world’s other

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Figure 2.46 Evening Picture of Dow Chemical Company’s Solid Calcium Chloride Plant at Ludington, Michigan (Van Savage, 2002; reproduced by permission of the Chemical Market Reporter).

dolomitization brines are or have been worked on a large-scale to produce calcium chloride, but the Cory potash mine near Saskatoon, Saskatchewan for a period recovered its very strong calcium chloride mine water brine (Table 2.3), and sold it, as is. However, the quantity of mine seepage gradually decreased with time until the operation no longer remained profitable. Bristol and Cadiz Lakes The high-calcium chloride brine of Bristol Lake has been continuously recovered on a commercial basis since 1910 (Jachens and Howard, 1992). The brine permeates the uppermost 9.75 m section of the playa sediments, and appears to be present throughout the entire lake. However, the permeability of the sediments

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varies widely, being the greatest in the halite layers, less in the sandy silt beds, and much less in the high-clay beds. The seepage of brine through the sediments thus varies considerably throughout the lake, and appears to be the greatest in the lowest areas (called the Salt Lake zones) where there is more salt near the surface. In the commercial operations, the brine (Table 2.9) is gathered by seepage trenches or pits, and a brackish calcium chloride water is evaporated from wells adjacent to the lake. By 1951, some of the trenches had been extended up to 8 km in length, and by 1992 there were 43.5 km of canals, pits and solar evaporation ponds (Gundry, 1992). The initial brine has a density of about 198Be (1.15 g/cc), and it is usually concentrated to about 408Be (1.38 g/cc), or about 32 –36% CaCl2 (Table 2.9). The solubility of the sodium chloride in the brine decreases during the solar evaporation process to a steady value of about 1% NaCl at 35% CaCl2. During the evaporation, salt thus crystallizes (starting at about 258Be [1.21 g/cc]) in the trenches, pits and ponds, and must be periodically removed from them. Because of the area’s very high summer temperatures and low humidity, the dilute brine from the playa can be evaporated (with pond temperatures reaching 598C) to a commercially saleable strength in as little as 2 weeks or up to 1 –2 months. However, during storms, cool evenings or in the winter (humid periods) the brine will actively absorb moisture from the air. Since most of the uses for calcium chloride do not require a very high purity, the evaporated product is normally sold without further treatment. When processed into the solid form, or if the magnesium concentration is too high the solution becomes quite corrosive, and its pH must be adjusted to 7.5– 8.5 with caustic soda. A corrosion inhibitor such as sodium chromate or dichromate (or a less toxic salt) may also be added (Gale, 1951). As an example of the production of calcium chloride on Bristol Lake, the National Chloride Company of America started large-scale production in 1950. They own some of the land on which they operate, and have placer mining claims from the U.S. Bureau of Land Management on other areas (some dating from 1908). Most of their land is in the southern part of the lake, including much of the lowest section, the Salt Lake areas. Over the years, they have excavated a large number of drainage canals, (Fig. 2.47) with the longest extending for 10 km (6 mi) and another for 4.8 km (3 mi). Many of the trenches drain to a central gathering area, and portable diesel pumps (Fig. 2.48) move the brine from others to the central canals or the ponds. The early trenches were made with a Northwestern 80 dragline, while the newer ones are excavated about 4.3 m (14 ft) deep and wide by a large backhoe (Fig. 2.49). Some deep pits (Fig. 2.50) are also utilized as brine collectors, with one dug to a 23 m (75 ft) depth, but the extra depth was not found to be beneficial. As a final method of brine collection, brackish water from aquifers just outside the lake is pumped (Fig. 2.48) to trenches to be evaporated and join the other brine. This large aquifer of a dilute calcium chloride brine (, 10,000 ppm) in some of the alluvial fan sediments next

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Figure 2.47 A typical brine seepage trench at Bristol Lake (Courtesy of National Chloride Company of America).

to the lake appears to have migrated from the lake and been diluted by entering groundwater. By the time brine in the collection trenches has reached at least 27– 288Be it is transferred to holding ponds (Fig. 2.51) to be further concentrated to the product strength of 408Be. It is then sent to deep ponds (Fig. 2.52) to be held until it is shipped in stainless steel tank trucks by their customers (one customer also re-ships some brine by rail from Amboy). The principal evaporation period for brine production is generally from April/May through September. Salt is periodically removed from the trenches by the back hoe, and from the ponds by a 613C Catapiller scraper-carrier. The seepage of brine into the trenches and pits is very much a function of the weather, with wet years greatly increasing the amount of brine gathered, and in dry periods (such as during the 1960s and in 2002) reducing the flow. During rainy periods (perhaps every 2 –3 years), some water briefly floods onto the low section of the lake, causing the salt surface to become very smooth and white. There can also be occasional flash floods in limited areas of the lake, and some of the surface water may get into the canals. To minimize the damage from such isolated rainfall and surface flooding, the 10 km trench has been dammed in the middle. As a rare 25 –50 year event, in 1983 the entire lake flooded, filling the ditches with a dilute brine, and necessitating raising the access roads across the lake (Beeghly, 2002).

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Figure 2.48 Typical brine pumps utilized at Bristol Lake. Top, brackish well at the edge of the lake. Bottom, portable transfer pump for the trenches, pits and ponds. (Courtesy of National Chloride Company of America).

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Figure 2.49 Backhoe cleaning salt from a seepage trench at Bristol Lake. (Courtesy of National Chloride Company of America).

Figure 2.50 A typical brine seepage pit at Bristol Lake. (Courtesy of National Chloride Company of America).

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Figure 2.51 A typical brine holding (evaporation) pond at Bristol Lake. (Courtesy of National Chloride Company of America).

The Tetra Technologies operation on Bristol Lake relies more on seepage pits than trenches, and utilizes considerable solar ponding of dilute brine pumped from the edge of the lake. Brine is pumped from the seepage pits into tank trucks, and then hauled to the evaporation ponds. Their operation on Cadiz Lake is entirely based upon dilute brine pumped to solar ponds. In both operations, salt is harvested from the solar ponds, washed and then drained for sale. It is shipped by truck, or rail from the siding at Saltus (near Amboy) where the Bristol Lake salt is washed (Morrow, 2002). In prior years, some of the Bristol Lake calcium chloride product was converted at Amboy to a 75– 78% CaCl2 flake by the Hill Brothers Chemical Co. The concentrated brine from the ponds was sent to open-pan evaporators where it was first heated to 1328C and then evaporated at 171– 1778C to the desired concentration (Fig. 2.53). This very viscous solution was next spread on chilled rolls to be cooled and solidified into thin sheets, which were scraped from the rolls and ground to the desired particle size. The flaked product was then sent to turbo dryers, from there cooled and packaged in plastic-lined, air-tight bags, and shipped to their customers. Some of the early operators on Bristol Lake were the Calcium Chloride Group, The Desert Properties Co., the National Chloride Co. and the California Salt Co. (Gundry, 1992). The latter company at first only mined salt from the lake, but later also produced calcium chloride as a by-product. The company was sold to the Leslie Salt Co., who continued the production of calcium chloride, and eventually

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Figure 2.52 A typical brine product shipping pond at Bristol Lake. (Courtesy of National Chloride Company of America).

recovered all of their salt from the calcium chloride ponds. However, they in turn sold their 324 ha (800 acre) of solar pond land, and 450 placer claims in 1978 to the Cargill Co., and Cargill sold to Tetra Technologies on August 2, 1998. Cargill continued to market the operation’s by-product salt, and Tetra also later purchased the similar calcium chloride producer, the Lee Chemical Co. on Cadiz Lake. Lee Chemical had started calcium chloride production in 1960, and was soon followed

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Figure 2.53 Flow chart showing method of recovering calcium chloride from Bristol Lake brine. (Ver Planck, 1957, courtesy of the California Division of Mines and Geology).

by the Delta Chemical Co. They later combined their operations as Lee Chemical. In 2002, Hill Brothers Chemical Co. was a brine sales agent for National Chloride, and ground some Tetra flake (from other locations) to a powder in their very lowhumidity Amboy plant (Chemical Market Reporter, 1999; Chemical Week, 2001). General Processing Technology Various literature articles have discussed the general subject of producing calcium chloride (Kotsupalo et al., 1999), and others have focused upon the specific operations involved. Wang (1998) and Postoronko et al. (1985) reviewed

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the problem of calcium sulfate scaling in evaporators used to concentrate calcium chloride brine, and suggested means of preventing its formation or removing the scale. The solubility of calcium sulfate in CaCl2 solutions was studied by Li and Demopoulos (2002). Bunikowska and Synowiec (2002) noted that sulfate impurities in pH 2.2 calcium chloride brine could be removed by hydrated zirconium oxide ion-exchange pellets. The sulfate could then be removed from the pellets by a 908C water wash so that they could be re-used. Alternately, the sulfate could be precipitated as barium sulfate. Babkina et al. (1981) suggested the use of titanium heat exchanger tubes in the evaporators, while Semke et al. (1975) discussed an automatic control system, and later (1977) a mathematical simulation of the evaporation process. Makabe et al. (2002) discussed automating the evaporators by means of an attenuated total-reflectance IR spectrometry instrument that continuously analyzed the concentration of all components in the brine. Kotsarenko et al. (1975) proposed an automatic condensate control and removal system for the evaporators, while Zhukov et al. (1981) noted some of the problems occurring due to the hygroscopic nature of the sodium chloride by-product crystallized in the calcium chloride evaporators. Ross and Sloyer (1953) suggested a purification process when a higher purity CaCl2 brine was required for the preparation of other calcium salts or for food use. They suggest diluting the CaCl2 to about 22%, and then acidifying with dilute phosphoric acid to a pH of 6.5. The filtered solution would contain very little magnesium and less than 0.0001% Fe and , 0:0005% heavy metals. Niino et al. (2002) also suggested the removal of heavy metals by their adsorption on activated carbon. Wheeler (1999) described the precipitation of magnesium from calcium chloride brines with lime, and the analytical monitoring methods. To produce solid calcium chloride, the starting brine is usually first evaporated to 40– 50% CaCl2 in triple effect evaporators (Hedley, 1951). Above this concentration, the large boiling point rise of the liquor, and its high viscosity make multiple-stage vacuum evaporation difficult, so after the crystallized salts (NaCl, CaSO4, etc.) have been removed the brine may be sent to single-effect evaporators to concentrate it to 55 – 65% CaCl2. In prior years, and smaller operations it could be sent to open-pan evaporators to be concentrated to 70– 72% CaCl2, with the temperature rising to about 1708C. At this point, the liquid can be cooled and solidified (such as on rolls), with the solids that are produced being in nearly the dihydrate form (the dihydrate contains 75.5% CaCl2). The solids can then have additional water removed to 76 –78%, or even 94% CaCl2 in conventional dryers or kilns. It has also been suggested that evaporators operating under 5 – 25 psig pressure can take the brine concentration to up to 85.5% CaCl2 (near the monohydrate composition), and operate at temperatures below the solidification point of about 2358C (Graves, 1958). The most common of the solidification methods is to form flaked calcium chloride by having a cooling roll dip into a tray of the hot liquid, as noted above, or spreading the hot solution onto the top of the cooling roll. The solidified sheet

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Figure 2.54 Typical Flaked Calcium Chloride Process. (Pittsburgh Plate Glass Co. [Allen] 1963).

of calcium chloride can then be scraped off of the roll and sent to a dryer to reach the desired 76– 78% CaCl2 concentration (Fig. 2.54). The flakes are screened and the oversize crushed to form a product with the desired particle size range (Table 2.21(A) and (B); Fig. 2.44). A number of articles and patents have been issued on the production of calcium chloride in the flaked form. Allen and Walton (1963; Pittsburgh Plate Glass) recommended adding low-value fines (2 20 mesh) from the product screens back to the (171 – 1828C) melt tank, so that the water-cooled rotating drum picked up an 8 –12% slurry instead of the normal 70 –76% CaCl2 solution. The product scraped from the drum was 0.76 – 2.54 mm thick, and after being broken and dried in a multi-tray dryer it was screened to the desired flaked particle size (20 mesh– 9.53 mm [3/8 in.]). The , 1228C flakes were heated in the upper hearths with , 3168C air, and cooled in the lower hearths with , 168C air. If a 74% CaCl2 solution were in the melt tank, a 78% CaCl2 product would be formed. The Asahi Glass Co. (1982) patented a method for reducing the corrosion problem when producing flake calcium chloride. They stated that it was most effective to treat a 35% CaCl2 solution from the primary evaporators with either 60– 1800 ppm of sodium silicate, 60– 900 ppm sodium silicate with 500 –1000 ppm of Ca(OH)2, or 1000– 1500 ppm Ca(OH)2 and 100– 200 ppm of gypsum. On the flaking rolls, they suggested 60 ppm SiO2 (as sodium silicate) and 1500 ppm of Ca(OH)2. This produced a white product with only 60 ppm Fe, in contrast to untreated material having a slightly brown color and up to 400 ppm Fe. Zaikin and Stankevich (1981) recommended cooling the flaked product from its normal 70 –1228C temperature to 36 – 408C before grinding by using a watercooled rotating drum augmented with a countercurrent flow of air preheated to 38 –418C to prevent the flakes from sticking together or on the drum walls. Other articles, such as by Markelov et al. (1972) discussed a means of level control in the drum feeder to provide a more uniform product thickness.

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Figure 2.55 Typical granular calcium chloride (Dow, 1980, courtesy of The Dow Chemical Company).

The preparation of granular calcium chloride (as shown in Fig. 2.55) can be accomplished by several different means, with perhaps the most common being fluidized bed dryer-granulators such as described by Schwalm et al. (1986), and mathematically analyzed by Sulg et al. (1973). In the former article, it was suggested to spray a neutralized 55– 60% CaCl2 solution into a gas-fired fluidized bed of granules, and to constantly withdraw the desired amount of product. The production rate from 8.5 t/hr of a 56.4% CaCl2 solution would be about 2.3 t/hr of a dry 76% CaCl2 product, and 2.5 t/hr of dust to be recycled by being redissolved in a more dilute (i.e., such as 30% CaCl2) feed solution. A 94% CaCl2 product could be obtained with higher gas temperatures or feed concentrations. Safrygin et al. (2002) suggested having the gas velocity in the fluid bed at 1.5– 3.5 m/s, and that the fines from the product screens be mixed with the entering feed solution. Scherzberg et al. (1985a,b) suggested that the final product could be made less dusty by treating it with 0.7– 3% of a more dilute CaCl2 solution, or 0.2– 0.7% of oil. For example, 1000 g of 92.5% CaCl2 granules with . 5 g of dust were cooled from 1708C in a fluidized bed cooler, and sprayed with 2.5 g of a 30% CaCl2 solution. The resulting product analyzed 90.7% CaCl2 and contained , 1 g of dust. Other types of granulators include prilling towers where a more concentrated CaCl2 solution is sprayed into the top of a tall, open tower with a counterflow of hot, dry flue gas (James, 1994). Hedley (1951) noted that spraying a 70 –72% CaCl2 solution at 170– 1758C into a tower maintained under a vacuum of 100– 400 mm of mercury could allow the solution to evaporate to 76% CaCl2 without any additional heat, but that a countercurrent flow of hot (1808C) air

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rising past the granules would allow the process to be more flexible. He also noted that the heat of crystallization of the dihydrate can supply much of the heat needed for the evaporation. Moore (1978) and Misumi and Asagao (1974) suggested spraying a hot (225 – 2858C) 74 –78% CaCl2 solution into a refrigerant (such as CCl3F) or solvent (such as CH3CCl3 at 25 – 678C and 25– 120 psig pressure to produce a 2 4 to þ 20 mesh, or 3 –5 mm product. Drum granulators have also been employed, such as in the procedure suggested by Bennett and Carmouche (1953). They sprayed a calcium chloride solution onto a bed of recycled ground semi-dry particles (13% moisture) in a rotary kiln being heated by 250 – 5008F flue gas. The solution contained . 50% CaCl2, was at a temperature within 258C of its boiling point, and the solids in the kiln were at a temperature between 150 and 1908C. The rate of recycled dust and small particles to the kiln was from 8 to 30 times the amount of the CaCl2 in the sprayed solution. The particles from the drum granulator were sent in succession through a low, and then high-temperature shelf dryer, with two stages of screening. The product was rounded, hard, dense (55 –65 lb/ft3), dust free 3 – 20 mesh pellets of nearly anhydrous calcium chloride (94% CaCl2). In a subsequent patent, (Wilcox and Speer, 1966) the temperature in the drum granulator was increased to 500 – 17608F, which reduced the recycle load to 2 – 8 times the sprayed amount of CaCl2 solution, and increased the production rate by 30 –90%. The hot gas entered the granulator in a concurrent manner, and was regulated in amount so as to not cause melting, puffing or to allow the solids to rise above their previously specified temperature. Lebedenko et al. (1982) also described a drum granulator where hot concentrated CaCl2 solution was sprayed into an inclined rotating drum of granules in counterflow to hot flue gas. The still moist granules that were formed were then sent to a conventional dryer. A model was made of the granule’s growth. In a similar manner with a disk granulator, as described by Novotny and Kadavy (1975), the hot, concentrated CaCl2 solution was sprayed onto a bed of granules on an inclined rotating disk. The moist granules overflowed from the bottom of the disk, and were conveyed to a conventional counterflow dryer. The disk could be operated at ambient temperatures, or heated by steam or flue gas. A horizontal disk granulator was suggested by Comstock (1954) to prepare large particles 6.4 –19.1 mm in diameter, with a density of 1.32– 1.64 g/cc. A 45 – 60% CaCl2 solution was sprayed inside a thick bed of heated granules, and the granules were then circulated to the surface. For larger particles, a pug mill could also be employed, and be able to process solutions as dilute as 60% CaCl2 and at temperatures of 120 –1508C. The hot liquid would be fed into the mill at one end with recycled fines or off-size particles, and be slowly conveyed with mixing to the discharge end. Heat could be applied through the mill’s walls, or by a hot air stream blowing over the top. The discharged calcium chloride would be of a mixed size and need to be screened, with the oversize ground to produce the final product (Hedley, 1951). Small granules may be prepared by the use of

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a high-temperature (4008C inlet flue gas) spray evaporator-dryer feeding 30 –40% CaCl2, as suggested by Gleason and Sui (1982). Other forms of solid calcium chloride can also be produced for special uses, such as pelletizing grains or powders, the crystallization of CaCl2·2H2O (Trypuc and Buczkowski, 1991), pouring , 72% CaCl2 into small molds (Tokuyama Soda Co., 1982), or forming solid calcium chloride tubes (Korolev, 1974). Articles on the drying of calcium chloride flake or granules have also been presented, including the patent on a multiple hearth dryer by Iwamoto et al. (1979). With a 70% CaCl2 solid feed material in a three-hearth unit the bed thickness in the first hearth was kept at 40 mm, the second hearth had 92 mm of product, and the third hearth 151 mm. With a 1608C flue gas, the feed material dried to 98% CaCl2, and it did not agglomerate or form a large amount of dust. Other Sources of Calcium Chloride As a quite different source of calcium chloride to the natural brines discussed above, it is also produced on a large scale from the waste liquors of Solvey soda ash plants. However, this process is beyond the scope of this book, and its production is quite different from that required in most natural calcium chloride processes. The Solvey end liquor from producing soda ash is normally a fairly dilute solution, such as containing about 9 –10.5% CaCl2, 5.5– 7% NaCl, 0.3– 0.7% SO4 and , 3% other salts. It is usually first carbonated to remove any free ammonia or Ca(OH)2, treated with an oxidizing agent to remove some of the organics, then with CaSO4 to precipitate some of the sodium sulfate (as glauberite), and finally evaporated to crystallize much of the NaCl and other salts (Novotny, 1968; Shitov et al., 1981). Commercial calcium chloride is also manufactured by reacting by-product or waste hydrochloric acid with limestone. The limestone and HCl are agitated to produce an approximately 37% CaCl2 solution with a pH of about 9. It is then pressure filtered and the pH adjusted to 8 before the solution is evaporated to its desired concentration (Gomes, 1997). Calcium chloride can also be produced from some waste solutions, and various of them have been discussed in the literature. Coal or other flue gas limescrubbing liquors remove SO2 and lesser amounts of volatile chlorides, and calcium chloride will accumulate in the scrubbing liquor. After gypsum is precipitated and removed from a slip stream, the liquor can be evaporated and spray dried. Riedel et al. (1999) have studied the corrosion problems encountered with the evaporator heat exchanger tubes in this application. In laboratory tests, they found that in the first effect with , 100 g/liter CaCl2 and , 858C that Alloy 31 (X1NiCrMoCu32-28-7) was satisfactory, but that in the last effect with , 300 g/liter CaCl2 and 858C Alloy 59 (NiCr23Mo16Al) or Alloy C-276 (NiMo16Cr15SW) were required. These results were confirmed in a 9 month plant test, where Alloy 59 had the lowest crevice corrosion in both effects. Alloy tubes were recommended over the competitive graphite tubes because of graphite’s susceptibility to mechanical damage.

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Other proposed sources of calcium chloride include the thermal decomposition (at 270 –2808C) of scrap PVC, with the flue gas being absorbed in a lime or limestone scrubber (Aoki et al., 2002). Others have suggested the washing of incinerator or fly ashes to recover calcium chloride, or the leaching of blast furnace slag with acids. Dust from scrap steel shredders could also yield calcium chloride in the absorbed incinerator flue gas.

USES OF CALCIUM CHLORIDE Calcium Chloride has a wide range of uses, but the major ones are for deicing, dust control, road stabilization, concrete curing, oil well drilling, tire ballasting and various industrial and miscellaneous uses. The percentages of the total sales as divided into these categories are shown in Table 2.15. There are many excellent discussions of these uses, such as publications by the major manufactures that give detailed instructions on each of their applications, product handling and environmental factors. There are also a few general articles in the literature on calcium chloride uses, such as by Kotsupalo et al. (1999). Deicing Calcium chloride’s usefulness in reducing or eliminating ice and snow on roads, or in preventing bulk commodities (such as coal in open rail cars) from freezing solid (or to thaw them when frozen) is because pure solutions at a concentration of 30.22% CaCl2 have a freezing point as low as 2 49.88C (2 57.68F). When ice is contacted by CaCl2, it is rapidly melted by both the freezing point-lowering effect and the heat of dilution, and it also appears to weaken the inter-crystalline bonds of ice to make it more friable and easily broken or removed. In actual practice, calcium chloride is considered to be most useful when the temperatures are below 2 178C (1.48F), and to be quite efficient from 2 29.4 to 2 31.68C (2 21 to 2 258F). This compares with salt being most effective from above 2 2 to 2 48C (25 – 288F, although it can melt ice to 2 21.28C [2 6.28F] at a concentration of 23.2% NaCl), while magnesium chloride is most effective to 2 158C (58F), although at a concentration of 21.38% MgCl2 it can melt ice to 2 33.68C (2 28.58F). Thus, calcium chloride is the only material inexpensively available for deicing in very cold climates, or for very cold storms. Dow Chemical recommends an application rate of 15– 30 gal per lane mile of roadway for their 30% CaCl2 corrosion-inhibited Liquidow q Armor q solution (Dow, 2002). They also note that the solution meets the PNS bulk storage test of less than 1% crystallizing after 7 days at 2 298C (2 208F). Magnesium chloride deicers can only pass this test at 2 188C (08F). When calcium chloride is applied with salt, the freezing point can be lowered to any desired intermediate temperature between their two limits or practical

Table 2.15 Calcium Chloride Use Pattern in the United States (Per Cent of Total Consumption) Deicing

Dust controla

Industrial processing

Concrete additiveb

Oil and gas drilling

Tire ballast

Other

2001 1999 1998 1998 1997 1996 1995 1994 1993 1992 1991 1990 1989 1988 1987 1986 1985

22 21 35 30 38

20 25 20 20 18

20 13 20 20 20

12 11 10 12 12

17 22 10 10 4

5 — — 4 4

4 8 5 4 4

40 35

20 20

20 20

5 5

10 12

— 4

5 4

40

20

20

5

5

3

7

35

22

21

5

10

—

7

40

22

20

5

5

4

4

(continues)

Uses of Calcium Chloride

Year

359

360

Table 2.15 (continued )

1984 1983 1982 1981 1980 1974 1972 1969 1966 1965 1963 1957 Average

Deicing

Dust controla

Industrial processing

Concrete additiveb

Oil and gas drilling

Tire ballast

Other

35

20

20

5

12

4

4

28 25

23 30

20 20 20 14 10 — 10 17 18

5 10 10 11 13 13 13 13 9

15 7 — — — — — — 11

3 3 2c 5c 5c 5c 5c — 4

6 5 8 10 17 — 17 25 5

55d 63d 55d 30

25 55d

15 31

30 22

C&EN (2000), Chemical Market Reporter, Chemical Week; US Bureau of Mines. a And road stabilization. b Cement manufacture and concrete accelerating. c Refrigeration brine. d Combined deicing and dust control.

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361

ranges, and thus the use of salt “pre-wetted” with calcium chloride (Dow recommends 8– 10 gal of their Liquidow q Armor q per ton of salt; Dow, 2001; Fig. 2.56) is more economical for less severe temperatures. For temperatures from 0 to 2 28C (16 – 328F), the ratio of flake calcium chloride to salt in CaCl2 – NaCl mixtures should be from about 1 to 3, from 2 9 to 2 188C (0 –158F) 1 to 2, and from 2 18 to 2 238C (2 1 to 2 108F) the ratio should be 1 to 1 (Dow, 1998). Calcium chloride works much faster than salt, and is most effective if applied before the low-temperature period to prevent ice from bonding to the roadway. However, once the ice has formed the calcium chloride can still be applied in either the liquid, flake or pellet form, and will rapidly melt the ice and break the bond between the ice and road surface. Its ability to then adhere to the road

Figure 2.56 Trucks removing loose snow and ice, and spreading salt pre-mixed with calcium chloride (Allied, 1980).

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surface often reduces the frequency of plowing and repeated applications, thus compensating for being more expensive than salt. For the same reasons, abrasives such as sand or cinders scattered onto the ice-covered road can also benefit from being pre-treated with calcium chloride, or mixed with it (Tetra, 2002). There is considerable literature on deicing, with perhaps the most useful sources being the excellent brochures issued by the calcium chloride producing companies. Among the general studies on deicing are a series of articles (in Japanese) by Sugawara et al. (1986 – 1999). They considered the mechanisms involved by both the heat of solution and the concentration factor on ice melting, and have developed a numerical model for deicing. Kirchner (1992) studied the comparative effectiveness of CaCl2, KCl, urea and Ca– Mg acetate on the salts ability to undercut ice and spread at the road-ice interface. Hahn (1988) compared the advantages of deicing with CaCl2 solutions rather than solids, and Burtwell (2001) noted a large-scale test in England comparing the merits of CaCl2-pre-wetted salt compared to ordinary salt. Boley (1984) discussed coal-thawing with CaCl2, and several papers have suggested quick analytical methods for determining the quantity of ice-melting chemicals on a road surface (Della Faille d’Huysse, 1981; Salins du Midi, 1978; Martinek and Beranek, 1975). There is also a considerable literature on possible corrosion inhibitors to be added to deicing CaCl2. The suggested chemicals include Na2B4O7·10H2O (borax), 2-butyne-1,4 diol, ammonium carbamate, carboxylic acid with an amine containing alkyl or alkylbenzyl groups, Na2PO3F, sodium hexametaphosphate (NaPO3)6, hexamethylenetetramine, lanthanum or a rare earth salt plus a soluble gluconate, sodium metasilicate, organic acid salts (acetate, ascorbate, formate, lactate, saccharate or tartarate), di and mono orthophosphate, calcium phosphate or calcium hydrogen phosphate, sodium or potassium silicate, a phosphonic acid derivative (RnR1N(CH2PO3M2)22n (where R and R1 are alkyl aminoalkyl or hydroxyalkyls), sodium tripolyphosphate (Na5P3O10), urea, and zinc chloride or sulfate. The effectiveness of one of the commercial corrosion inhibitors is illustrated in Fig. 2.57. The literature on suggestions for making deicing mixtures with calcium chloride is also surprisingly large. As noted above, pre-wetting ordinary salt with CaCl2 solutions, or mixtures of salt with calcium chloride solids are commonly employed, but in addition suggestions for a suspension of fine NaCl crystals in a CaCl2 solution and forming a solid CaCl2 layer around an inner NaCl mass have also been made. There have been suggestions to combine calcium chloride with various other salts, including: the fertilizers potash or urea (perhaps coating the fertilizers with CaCl2), sodium tripolyphosphate (Na5P3O10), magnesium chloride or sulfate, ammonium chloride or phosphate, and sodium, calcium and zirconium silicates (the latter is a far-IR absorber). Suggested mixtures with minerals or inert substances include: calcite, diatomaceous earth, kieselguhr, pozzolan, pumice, sand or gravel, calcined and expanded shale, and small particles of asphalt or concrete, cinders, lava and sandstone. Many organic agents have also

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363

Figure 2.57 (Dow, 2001, courtesy of the Dow Chemical Company).

been suggested to be used with calcium chloride, often with other inert fillers. These include: polymers of methylacrylate, oxyethylene, oxypropylene, urethane, and vinyl acetate; carbohydrates (molecular weight 180 –1000), coal tar, ethyl and methyl alcohols, glycerin, glycol, 1-hexadecanol, metal lactates, 1,2-propylene, propylenediamine dioleate, polyglycol, thioacids, solvents, and surface-active agents Ossian and Steinhauser (2000). Many waste products have also been suggested to be used with calcium chloride, such as wastes from: the alcohol industry (including from molasses, sugar beets and sugar cane), ashes, biofuels, bark mulch, cheese (from milk or whey), magnesium processing, pulp and paper sludge, saccharinic acid (from the alkali treatment of bagasse, corn stalks, molasses, paper pulp, saw dust, straw, sunflower stalks or waste paper), tofu processing, whey, wine (from grapes and other fruit), and zinc. The environmental problems with calcium chloride in its deicing application, as with salt, are that when used in excess it can harm roadside vegetation and contaminate water supplies. Howard and Beck (1993) have studied the effect of road deicing chemicals upon the springs, wells, groundwater and aquifers near

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highways in southern Ontario, Canada, and found that a small, rather minimal change occurred compared to background values. Mayer et al. (1999) noted that the Canadian environmental agency was studying road salts, finding that the salts’ greatest effect on surface water was to small road-side ponds, or streams draining large urban areas. In Italy Canovi (1987) reported that there had been no visible damage to the vegetation near heavily traveled roads, nor an appreciable change in the adjacent soil composition. Rich and Murray (1990) observed that in central Connecticut roads in drainage basins with drinking water reservoirs only used calcium chloride, and not sodium chloride for deicing. Bubeck and Burton (1989) discussed the effect of deicing chemicals on underground aquifers, and Werner and Dipretoro (1998) reviewed road salts’ effect upon water-supply springs. Kjensmo (1997) found that a small thermally stratified lake near a highway in Norway was becoming less stratified due to the slow increase of road salts in its upper section. Bowers and Hesterberg (1976) studied the effect of deicing salts upon white pine near roadways in Michigan. They found that salt spray was a problem, but not CaCl2 with spray or root take-up. Similar observations on Scots pine in Finland were noted by Viskari and Karenlampi (2000) and Hautala et al. (1992), where NaCl damage only occurred within 20 –30 m of the road, and the seasonal moisture greatly effected the extent of damage. Other tree species (Sorbus aucuparia L., Acer pseudoplatanus L., Tilia platyphyllos Scop., and Platanus acerifolia W.) were studied with only CaCl2 deicing in Belgium. It was found that there was little damage, but Tilia was the most, and Platanus the least effected by the CaCl2. Platanus also became somewhat more sensitive to infection by Gnomonia venata (Paul et al., 1987). Utosawa (1995) found that deicing CaCl2 that entered a sewage system enhanced the removal of phosphates from the water. Calcium chloride can also be quite corrosive to concrete, automobiles and other structures that it contacts. However, the addition of corrosion inhibitors to the CaCl2 can reduce its corrosion problems compared to salt by more than 70% (Fig. 2.47), and each of the major producers markets an inhibited product (Dow: Liquidow Armour or Plus; General Chemical: Corguard; Tetra: Winter Thaw-DI; and Reilly Industries for magnesium chloride: Ice-Stop CI2000). Gillott (1978) has discussed the effect of deicing calcium chloride on the corrosion of concrete roadways and adjacent structures. Ludwig and Balters (1994) compared the damage to concrete from various deicers when the temperature fluctuated from 2 20 to þ 108C. The damage decreased in the order of NaCl, urea, CaCl2 and water. Tsukinaga et al. (1994) found the same order of deterioration based upon the amount of concrete scaled-off under standard test conditions. Balazs et al. (1990) reviewed the freeze –thaw corrosion problem, and Frey and Funk (1985) studied the loss of concrete’s compressive strength by deicing chemicals. Under the worst conditionsm, there could be a 70– 75% loss, but in actual practice the loss was found to be quite small (Pagliolico et al., 1997). Sayward (1984) noted

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that the mechanism of freeze – thaw deterioration of concrete by deicing chemicals was by the initial formation of micro cracks that facilitated access of salt and water for freeze-cracking, which in turn allowed CO2, water and salt to corrode the reinforcing rods, and spall the concrete. Hudec et al. (1992, 1994) suggested that potassium acetate had equal deicing properties to CaCl2, and was not a corrosive agent. Dust Control Since calcium chloride is hygroscopic and deliquescent (it absorbs and retains water) under a wide range of climatic conditions (temperature and humidity), when applied to a dusty or potentially dusty surface (i.e., such as an unpaved country road; Fig. 2.58) it wets and consolidates the dust. It then absorbs more water to help prevent additional dust from forming, and because of its solutions’ high-vapor pressure (and high-boiling point), it helps to retain this moisture. This improves safety and the ease of driving, assists in retaining all of the small particles and surfacing material (as well as the dust), aids in producing a dense surface and reduces the amount of blade work that is required for road maintenance. Without the calcium chloride it has been estimated that a car driving one mile per day on a dusty road for a year could cause the loss of 1 t of dust. As these fines are lost (causing vision and health problems), voids are created between the larger particles, and the surface begins to loosen. Then the road begins to degrade, potholes form and “washboarding” occurs, aggregate is pushed to the side, and more frequent blade work is necessary. The application of calcium chloride can reduce the aggregate loss by up to 75%, and the blade work from one-third to onefifth of that for an un-treated road (Tetra, 2002). The recommended application rate is , 1.2 liter/m2 (0.27 –0.5 gal/yd2) of 32– 40% liquid calcium chloride (or 1.5 lb of flake/yd2) in the late spring (preferably after recent blading work and a rain), followed by one-half to one-third of that amount during the summer. When sprayed onto roadways (Fig. 2.59) one company’s product is said to be effective with only one application per year, but two to three coatings are usually recommended. Various surface-active additives have been suggested to enhance CaCl2 solutions’ ability to penetrate dust, and thus help its control (Dow, 1998, 1980). One of the agents that has been recommended is sodium succinate, which when applied to various strength CaCl2 solutions in the amount of 0.2– 0.3% increases their penetration ability by 13 –18% (Zhou and Wu, 1999). Soil Compaction or Stabilization In a similar manner, calcium chloride is effective in surface consolidation, base stabilization, and compaction during road construction. For the latter use, it is added as either liquid or flake to the aggregate as it is being laid down, or to plant

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Figure 2.58 Typical dust from an unpaved road, and one type of calcium chloride spreading truck (Dow Chemical, 1980, courtesy of the Dow Chemical Company).

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Figure 2.59 Typical liquid calcium chloride spreader for dust control or road stabilization (Dow, 1999, courtesy of The Dow Chemical Company).

mixes. It absorbs water into the base material and later retards the evaporation of this moisture, contributing to the stability of the structure. Studies have shown that a moisture variation of only 1% from optimum may reduce the aggregate’s density by over 2 lbs/cuft, and increase void space by as much as 8%. Calcium chloride addition can maintain the moisture level to the desired 7 –8% during the road base compaction, which greatly aids in the interlocking process, and is especially important in the summer. Less rolling and grading are required, the load-bearing capacity and stability are improved, and the overall construction costs are reduced (Tetra, 2002). In one test where the top 6.4 cm (2.5 in.) of roadbase was graded and then compacted by rolling, with water alone the road was in good condition for 12 weeks and in poor condition for 42 weeks, even with 5 gradings during the year. When flake calcium chloride was mixed with the top 6.4 cm of road base in the amount of 2.2 lbs of flake/yd3, and then graded and compacted only two additional bladings were required over the 54 week period to keep the road’s surface in good condition for 40 weeks and poor condition for 14 weeks. With 4.2 lbs of flake/yd3 the road stayed in good condition for over 54 weeks with no additional blading (Althouse, 2001). Some of the literature on soil compaction (also called binding, hardening, increased load-bearing, stabilization or strengthening) include Alkiri et al.’s (1975) general article on the subject, and the use of additives with the CaCl2 (Angelova, 1997; Baek, 1997; Fujioka, 1997; Grott, 2000; Shimada et al., 2000; Shirasaka et al., 1998; Sugihara, 1997; Thomas, 1997; Wu, 2000; Zhou and Wu, 1999).

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Some of the suggested additives are: CaO, Ca(OH)2, gypsum, NaOH, Na2SO4, H3PO4, alum, calcium aluminate, cement, blast furnace slag, fly ash, lignosulfonic acid, sulfamic acid, flocculants, polymers and wetting agents. Materials that might be used to increase the penetration of calcium chloride solutions into soil (or dust) include: sodium succinate, sodium silicate, sodium dodecyl sulfonate and sodium dodecyl benzyl sulfonate (Zhou and Wu, 1999; Wu et al., 1998; Wu and Chen, 1998). A number of articles on the effect of calcium chloride on various clay minerals are listed in the Geology chapter. Concrete Calcium chloride accelerates the setting time of concrete by as much as 30 – 50%, with the effect being greater in richer mixtures (with a higher ratio of cement to sand and aggregate) and colder weather. It also provides increased compressive strength, especially at the early stages (i.e., it may be more than doubled after setting one day, and increased by 50% at 7 days). There is also usually an increase in flexural strength of the finished concrete, but this effect is much smaller. Calcium chloride usually reduces the amount of bleeding (water separating from the mixture), settling and subsidence of fresh concrete, and the amount of water required for good workability. Reduced water also helps to densify and waterproof the concrete to minimize later corrosion. Up to 1– 2% liquid calcium chloride (as CaCl2) is usually used in concrete formulations, although if solid calcium chloride is preferred at least 5 min of mixing is desirable, with the flaked product dissolving faster than pellets. Increasing the set time and early strength allows earlier finishing of the concrete, which reduces the labor cost and is especially important in colder weather. However, the amount of CaCl2 present is too small to have any appreciable effect on the possible freezing of the water in concrete, but it does cause an earlier liberation of the heat of hydration. Thus, the length of time during which concrete can be damaged by freezing is reduced. Special lowtemperature concrete can be prepared with a mixture of calcium chloride, various clays and other additives (Nedra, 1971; Stroitel’stvo, 1971; Go Bocan, 2001). Calcium chloride may also be used as an additive in making cement to lower its alkali content (Jardine et al., 2000; Lackey, 1992). On the negative side, the presence of CaCl2 can result in increased corrosion to the concrete’s reinforcing rods if the steel is not adequately buried (at least 5 cm [2 in.] are recommended) and it is exposed to air and moisture. Since corrosion is enhanced when the steel is stressed, pre-stressed concrete beams should not use calcium chloride. Concrete with CaCl2 is more susceptible to sulfate attack and galvanic corrosion (from stray electrical currents caused by the steel reinforcing rods contacting other buried metals, etc.), and it is more corrosive to any other buried metals (aluminum, copper, etc. pipes or conduit). Calcium chloride does not decrease the deicer freeze – thaw scaling resistance of properly designed and placed concrete, but it can increase winter deterioration with poor concrete, or

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mixtures that did not have a proper drying period (3 – 6 months are recommended before freezing). Maintaining a high pH in the concrete (it is normally , 12.5), adding some pozzolans or fly ash to densify the concrete, or even the addition of latex (which is especially advantageous for highway patching; Nagi et al., 1994) can reduce these problems (Lackey, 1992). In general, not more than 1 lb of CaCl2/100 lbs concrete should be used when the temperature is above 328C (908F), 1– 1.5 lbs at 21– 328C (70 – 908F), and 2 lbs/100 lbs concrete below 218C (708F; Dow, 1998). There is considerable literature on the use of calcium chloride in concrete, such as the general article by Anon. (1971), and Foster’s (1929) discussion of the effect of CaCl2 on the hydration of the cement minerals. Akhverdov and Batyanovskii (1986) noted that for dry molded, rapidly hardening concrete, 1% of calcium chloride also increases the strength of the concrete by 4.1%. As concrete ages, the surface areas slowly carbonate, and this reduces the harmful effect of road deicing salts in forming calcium oxychloride (Pagliolico et al., 1997). Various types of cement and other additives to the concrete (with the CaCl2) can be used for especially corrosive conditions (Epshtain et al., 2000). Many other articles have been written on the use of calcium chloride in preparing the cement for the concrete, such as with other additives, and for special concrete applications. Oil and Gas In the oil and gas industry, calcium chloride is used in well drilling muds to increase the fluid’s viscosity and density (and thus the weight [and pressure resistance] of the mud column), and as the internal phase (the small droplets in the oil) in an invert emulsion of an oil – water based system. It is also used in specially designed muds to deal with plastic-flow shale (it stiffens the shale and reduces its swelling), and in completion brines to flush-out the drilling mud prior to pumping the oil. The use of CaCl2 instead of NaCl in drilling muds and as a flushing brine can considerably improve the permeability of the formation because of calcium chloride’s ability to coagulate clays and prevent their forming into finely divided colloids that would block the flow of oil or gas (White et al., 1964b). Calcium chloride is also used as an accelerator in cementing the drill casing to the rock formation (Reddy, 2001; Reddy et al., 2000), or upon abandonment filling of the casing. Some of the literature on other applications of calcium chloride to oil and gas production include those by Reed (1968; with water or steam injection), Trebin et al. (1968; improving formation porosity), Bagci et al. (2001; the effect of CaCl2 on limestone reservoirs), Navarrete et al. (2001; with the addition of a polymer to cement casings when there are highly porous zones), Gupta and Santos (2002; a field test method for the compatibility of CaCl2 drilling fluid with the formation rocks), Blackwell et al. (1974; with a gelling agent and liquid

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hydrocarbon to form a hydrofracturing mixture), Kelkar et al. (2001; as a perforating brine for hydrofracturing), a well work-over fluid (again to clean the well of solids before maintenance), and in packing fluids (the fluid in the rubber seal between the tubing and the casing), Sabins (2001; a gravel-pack carrier fluid), Ralston and Persinski (1968; inhibiting salt deposition and crystallization), Arshinov et al. (1971; CaCl2’s prevention of gas hydrate from forming), and many others. In the production of natural gas calcium chloride is commonly used as a drying agent to remove water and prevent the formation of methane hydrates (Hodgson and Martinez, 1984; Kolodezni et al., 1977; Musaev, 1972; Arshinov et al., 1971; Hodgson, 1970; others). An example of an absorption column is shown in Fig. 2.60, where the solid pellets are loaded into the upper section of the column, and the gas introduced into the tower’s base and first passed through a phase separator to remove any liquid water or hydrocarbons. As the pellets absorb water they slowly dissolve, and this solution then flows to a modified bubble-cap section where it continues to absorb water until it reaches a concentration of about 20 – 25% CaCl2, and is then discharged. New pellets are added as needed, and the combined solid and liquid-phase dehydration provides very good efficiency and economy (Dow, 1998). Typical water absorption rates and the amount of gas processed per unit of calcium chloride are shown in Fig. 2.61, and as examples, 1 lb of 94% CaCl2 can absorb 17.3 lbs water at 95% relative humidity, but only 2.1 lb H2O at 60% humidity. If hydrate growth-inhibitor/modifying compounds are used in the system, they may sometimes be recovered from the treating solution by the salting-out effect of calcium chloride (Blytas and Kruka, 2001). When natural gas is absorbed to enrich the helium content of the residue, a calcium chloride solution is added to the absorbent to produce a dry, enriched gas (Kryukov et al., 1971). In petroleum refineries, calcium chloride is used to dry materials such as distillate, LPG, kerosene and diesel fuel, while in petrochemical plants it can also be used to dry products like chlorinated hydrocarbons. For these applications, large particles of calcium chloride are required, such as Dow Chemical’s Peladow DG, almond-shaped briquettes. Some of the specific literature on this subject include the articles by Barnett (1996) on drying distillates, Sweeney (1981) on gasohol, Hu (1995) on diesel fuel, Masini et al. (1990) on the hydrocarbons to produce chloromethanes, and Matsushita Electric Industrial Co. (1982) and Murthy (1975) on kerosene. In the latter article, the drying was done in parallel-series columns (Fig. 2.62), with a flow rate of 0.75 – 1.0 ft/min through the CaCl2 beds. The capacity was 0.93 lb H2O/lb CaCl2 to go from 95% CaCl2 pellets to the hexahydrate form, and 0.61 to the tetrahydrate. Salt was used in an initial bed to remove physically trapped water, and the calcium chloride removed most of the remaining (including dissolved) water. The beds were dumped after 18– 30 days’ operation, with a 32 –38 bbl kerosene/lb of CaCl2 capacity.

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Figure 2.60 An example of a natural gas dryer using calcium chloride briquettes (Dow, 1998, courtesy of The Dow Chemical Company).

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Figure 2.61 Factors in the drying of natural gas with calcium chloride (Dow, 2002, courtesy of the Dow Chemical Company).

Ballasting Tire weighting is widely used in order to increase the traction of farm and industrial machines under various surface conditions by increasing their weight. A standard (30 –40%) calcium chloride solution, with its high density is put into the vehicle’s tires, often increasing the machine’s weight by as much as 10% without any hindrance to its operability or performance. The axles do not carry any additional weight, and besides improving traction and drawbar pull by as much as 80%, it reduces tire wear and stress on other parts of the vehicle, reduces the vehicle’s center of gravity, and helps to prevent the lifting of the front or rear tires with the handling of heavy loads. The calcium chloride solution is pumped or pressured into the tire with its valve stem in the top position, and usually filled to 90% of the tire volume (although 75 – 100% can be used). Air is then added to the desired pressure. For tractor tires, the amount of solution can be from 16 to 213 gal per tire (Dow, 1998, 1980).

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Figure 2.62 The calcium chloride drying of kerosene (Murthy, 1975).

Food Processing Considerable calcium chloride is used in food processing, but if classified as food-grade it must meet Food and Drug Administration (FCC) standards, and be manufactured under “good manufacturing practices”. In 1990 four U.S. companies produced a special anhydrous food grade product, and by 2002 several other companies were marketing both liquid and solid food grade calcium chloride. Calcium chloride is used in canning, pickling (“briner’s grade”), water treating for soft drinks or beer, the processing of cheese (it increases the size and strength of curds, and hastens ripening; Guven and Karaca, 2001), to enhance food flavor, for food preservation and packaging, as a nutritional supplement and other food-type objectives. As a few examples of these uses, it can firm the skin of fruits and vegetables. With pasteurized jalapeno pepper rings, it gives the product a firmer, fresh-like, uniform texture (Gu et al., 1999), with macerated fruit

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or vegetables it enhances the consistency of the product (McCarthy, 1997). Diced fruit and vegetables when treated with 0.5 –1.0% CaCl2 in a fruit or vegetable juice provide maximum firmness and improved quality (Hinnergardt and Eichelberger, 1994), and a CaCl2 brine partially dehydrates the outer surface of canned potatoes, thus reducing “sloughing” (Negi and Nathson, 2002). The use of 0.2% CaCl2 in the brine for canned breadnut seeds improved their flavor (Mathews et al., 2001). Beef may also be processed (including being tenderized) with calcium chloride (Moskovskaya, 1996), and fish may be firmed and preserved in a brine with CaCl2 (Cho, 2001; Liberman, 2002). Calcium chloride as a desiccant when shipped with groundnuts reduced the nuts’ moisture content to , 0:8% (compared to 10.2% in a control sample) and prevented mold formation (Navarro et al., 1988). Industrial There are many industrial uses for calcium chloride, such as in chemical manufacturing where it is one of the common reactants used to produce various calcium compounds. It has also been suggested as a reactant in the production of nitrosyl chloride (Oesterreichische Stickstoffwerke, 1969), and it is often used as a dehydrating agent to help form other chemicals. For instance, in azeotropic distillation its presence can allow the breaking of otherwise inseparable (dilute) azeotropic compositions, such as to form absolute alcohol (Zeitsch, 1989) and acetonitrile (Bala and Botez, 1980). Similar drying of chlorodifluoromethane allows it to be used as a refrigerant (Kwon et al., 1998), and prepares supercritical carbon dioxide (CO2) for various uses (Diaz and Miller, 1985). Drying a mixture of dipropylene glycol and monopropylene also allows the two products to be separated by distillation (Popescu et al., 1983). In the preparation of ethyl acetate by an esterification reaction, drying the reaction mixture with CaCl2 greatly increases the yield (Terelak et al., 1977), and drying methyl acrylate’s raw materials allows an easier separation from the reaction mix (Firsov et al., 1979). Many products can be made with the addition of calcium chloride, such as an improved xanthan gum (Cahalan et al., 1977), printing inks for embroidery (Chen and Sun, 1990), calligraphic inks (Zhou, 1997), an acrylate flocculent (Katayama Chemical Works, 1984), deodorants (Ogura, 1995), and dry cell paste (Davassy, 1979). Calcium chloride can also be used to help produce certain chemicals, such as recovering aluminum chloride from uranium by-product solutions (Casensky et al., 1999), stabilize chlorine dioxide (Khalaf, 1996), forming cyanuric chloride (Fujimori et al., 1990), and in the low-temperature drying of ceramics made by the lost wax process (Tikhonov and Klinskii, 1974). It can be used to coagulate latex in the manufacture of rubber, and as a low temperature refrigeration brine. In paper manufacturing, it provides artificial water hardness that increases the web strength of corrugating media. In the steel industry

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it is an additive to pelletize ore for blast furnaces, and to reduce the alkali content of steel mill furnace feed mixtures. In waste water treatment it can precipitate fluorides, borates and phosphate (Tatsumi et al., 2001; Tano and Tateiwa, 2001, 2002; Kato, 2002, 2003; Koizumi, 2003; Sumita and Oyama, 2002; Kataoka et al., 2002), and help remove organics (COD) in paper industry wastewater (Palitzsch (2002)). It can break emulsions in oily wastes, and remove silicates and various solids by densifying flocs formed by coagulating agents. It has various other environmental uses, such as in reclaiming oilfield soil (Merrill et al., 1990), the remediation of lead, heavy metals and organically contaminated soil (Cline et al., 1993; Dessi et al., 2000), it can be combined with surfactants to remove tetrachloroethylene from contaminated soil (Ramsburg and Pennell, 2000; Schweninger and Falta, 1999) and its addition to other flooding solvents increase their density and penetrability (Myers et al., 2001). It has been added to an aminoalkyl acrylate flocculent to increase its effectiveness (Katayama Chemical Works, 1984). Calcium chloride has also been used in various metallurgical operations, such as flue gas scrubbing as an additive to enhance the lime – SO2 reaction (Alscher et al., 1979), in the heavy media separation of coal from inadvertently mined rock, and in the high-temperature processing of coal or other fuels (Borkowski, 1976). In other scrubbing applications, it has been used to recover heat from flue gas (Charyev and Malkovskii, 1989), or for energy conservation in vinyl chloride production (Jarosek et al., 1986). It has been used for the leaching of silver and lead from zinc roaster flue gas after the zinc was leached by dilute sulfuric acid (Zheng, 2000), and for the solubilization (chlorination) of Cu, Zn, Pb, Au and Ag by the high-temperature roasting of their sulfides or oxides. Its incorporation into a slag can cause the desulfurization of some iron alloys (Oktay, 2002), and a 40% CaCl2 solution makes a good surface quenching agent for complex steel parts (Liu, 2002). Drying Many of the uses noted both above and below result from the ability of calcium chloride to act as a humectant or drying agent, and there is an extensive literature on this subject. Air, gas and hydrocarbon drying are special cases of this ability that have been especially active. Natural gas drying has been previously discussed, but other general articles on gas (any type of gas) drying have been presented by Borkowski (1970), Hodgson (1970), Norton (1968), Popov (1967), Wilcox (1967), and many others. Thurston (1948) studied gas drying by calcium chloride solutions, and developed mass transfer coefficients for the operation of wetted wall and packed towers. Large industrial or commercial air conditioning units, as in Fig. 1.100, commonly employ calcium chloride for both moisture removal, temperature control and sometimes heat recovery (Fujioka et al., 2000; Gultekin et al., 1991; Waldenmaier, 2000).

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Air drying equipment and procedures have been described by Balobaev (1987), Drutskij et al. (1995), Koyama (1986), Vasseur (1984, 1982), and others. Ullah et al. (1988) made experimental studies on air dehumidifying with calcium chloride solutions in packed beds, and established efficiency coefficients. Gandhidasan and Al-Farayedhi (1994) and Lobo and Da Silva (1982) suggested solar evaporation of the absorbing solution. Many solid calcium chloride mixtures have been suggested for drying air, such as by Ehata et al. (1986), Mitsui Engineering and Shipbuilding Co. (1985), Suzuki and Hidaka (1977), and many others. A very large number of calcium chloride drying agents have been described in the literature. This includes suggestions for many types of containers for CaCl2 granules or flake, such as being encapsulated for controlled release, in containers with perforated caps, or being placed in air-permeable plastic bags to be used as disposable dehumidifiers in closed spaces. Even more prolific have been articles on calcium chloride mixtures and different compositions to enhance their dehydrating effectiveness for special conditions, such as reviewed by Lu (1999). This includes mixtures with other chemicals such as forming a double salt with CaO, mixed solutions with LiCl, a mixed slurry with Mg(NO3)2, and a mixed powder with MgSO4 or sodium silicate. Even more suggestions have been made for mixtures with various minerals, such as with: asbestos; a combination of calcite, talc, magnesia and cement; with activated carbon, fluorite, gypsum, powdered iron or aluminum; untreated or impregnated into foamed or swelled perlite or vermiculite; paper or pumice; sepiolite; silica or silica gel; vermiculite; and zeolite. The largest category, however, is with individual or mixed polymers, including polymers of: acrylamide; acrylic acid; acrylonitrile; alkyleneglycols; cellulose in a large number of forms (alpha starch, bagasse, carrageenan, CMcellulose, corncob meal, corn starch, Konjak powder, potato starch, Sumikagel, and wheat flour); ethylene; methylcellulose; saccharide; starch (see cellulose); vinyl acetate; vinyl acetamide; and vinyl alcohol. Because of the large number of these patents or articles they will not be listed in the reference section. Miscellaneous Other uses for calcium chloride are listed in Table 2.16, including in mining, deinking recycled newspapers, and in water treating for its addition to remove impurities and aid in solids flocculation. In agriculture it is a growth-enhancing macro-nutrient fertilizer. Increasing the Ca2þ/NHþ 4 ratio in fertilizers has been shown to result in significant growth increases, especially in crops that are roots and tubers (such as onions, beets, etc.) by increasing the rate of photosynthesis (Fenn and Feagley, 1999). It also can be used to treat high-sodium soils (that are “alkaline or salty”; Alsharari, 1999), and a calcium chloride spray has been found useful to control the rain cracking of sweet cherry trees (Fernandez and Flore, 1998; Lang et al., 1998), as well as to thin fruit blossoms (especially apples; Long, 2002). Too much calcium chloride, however, such as may occur with the

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Table 2.16 Examples of Some of the Uses for Calcium Chloride 1. 2. 3. 4. 5.

6.

7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

17. 18. 19. 20. 21. 22.

Adhesives and antifreeze in poster cements, a raw material in processing natural glue, a starch modifier (it lowers the gel temperature) or humectant. Agriculture. A fertilizer supplying calcium and chlorine, an ingredient in herbicides, liquid feed supplement for cattle, dehydrating agent, tractor tire weighting. Beverages. Precipitant or ingredient in brewing, and to control the water composition in beer and soft drinks. Canned goods. Firming agent for certain fruits and vegetables, gelling agent for low-calorie jellies and preserves, “cold peeling” tomatoes. Cement, concrete and gypsum manufacture. Chemical lowering of high-alkali content in Portland cement, ingredient in water-reducing and set-controlling admixtures for concrete, early-strength accelerator in concrete and mortar, cold weather additive, aerating aid during the calcining of hydrated calcium sulfate. Chemicals. Intermediate chemical in several reactions and various organic synthesis, catalyst, dehydrating agent, manufacture of molecular sieves, manufacture of other calcium salts such as calcium stearate, removal of water from process air. Ceramics. Porosity reducer, lubricant for ceramic bodies and refractory products. Dairy products. Aid in curd formation for making cottage cheese, additive for evaporated milk to increase its calcium content. Dyes. Precipitating agent. Fluxes. Ingredient in many flux compositions, and in some printed circuit boards1. Gas, Natural. Dehydrating agent, reduces hydrate formation. Highways. Deicing for ice and snow removal, dust control on unpaved surfaces, shoulder and base stabilizing, soil solidification, plant sterilization agent. Hydrocarbons. Desiccant. Electric lamps. Additive to cement used in making electrical insulators, intermediate chemical in producing phosphor powder for fluorescent lamps, refining tungsten ore. Lead. Flux in production of lead–calcium alloys. Mining and metals. Treatment of pelletized iron ore, removes non-ferrous impurities from iron ores, controls alkalies and “scaffolding” in iron blast furnaces, removes non-ferrous impurities from basic oxygen, open hearth and electric furnace waste dust, refining copper and nickel ores, separating copper from auto scrap, production of magnesium metal, production of sodium metal, purifying molybdenite ores, production of calcium molybdate used in alloying molybdenum to steel, production of disposable cores for aluminum die casting, separating shale from coal fines, freeze-proofing and dust-proofing coal, ore and other aggregates for shipment and storage, component in heat-treating salts, fire retardant in brattice cloth (for mine ventilation). Organic compounds. Fire retardant, dehydration agent, azeotrope breaking. Paint, varnish, and lacquer. Manufacture of calcium naphthenates used in various varnishes and lacquers, manufacture of whitewash paints. Petroleum. Drilling mud additive, weighting agent, cement additive to seal wells, completion fluids, flushing agent. In refining as a dehydrating agent, manufacture of lubricant additives. Pharmaceutical and medicinal. Manufacturing processes, hot pack compresses, cold pack dressings, dehydrating agent. Plastics and synthetic resins. Suspending agent to prevent coalescing during polymerization, control particle size development. Printing. Lithographic chemical, etching fluid for aluminum printing plates. (continues)

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23. 24. 25. 26. 27. 28. 29. 30.

Pulp and paper products. Drainage aid, bleach modifier in pulp treatment, increases web strength of corrugating media, improves dye retention. Refrigeration and air conditioning equipment. Heat exchange brine, dehydrating agent. Rubber manufacture, Reclamation. Latex coagulating agent, aid in disintegrating cord and devulcanization. Seed processing. Drying seeds. Vapor degreasing and hard surface cleaning. Desiccant when reprocessing solvents. Waste treatment. Flocculating agent for the treatment of emulsified wastes containing phosphates, oils and paint sludges, removal of fluorides, silicates. Weed killer. Fire retardant and ingredient in certain weed killers. Miscellaneous. Water antifreeze in fire extinguishers, flux for soldering and brazing, hide pickling, kelp processing.

Sources: Allied (1980); C&EN (2000); Chemical Marketing Reporter; Chemical Week; Dow Chemical (1980). 1 Iwasayama and Oba (1991).

disposal of oil industry high-density brines, can damage vegetation (Vidakovic et al., 2002). Small amounts of calcium chloride are useful in reducing the incidence of milk fever in transition cows by increasing the adsorption of calcium into the blood stream. Powdered calcium chloride has some uses in specialized cements such as metallic, non-shrink grouting cements, pre-mixed dry cement patching and waterproofing compounds, gypsum plaster (as an “aridizing agent”), and special applications of cement acceleration. It can be used in defoaming and humectant formulations, and as a dispersing agent for textiles, powdered formulations for paint pigments, “stay soft” salts (as a free-flowing agent), heat-treating salts and molten salts fluxes (Iwasayama and Oba, 1991). It has also been suggested as a heat-storage agent (Gultekin et al., 1991). In medical applications, dilute CaCl2 solutions have been used as an animal purgative, to treat animal tendons, and to soak ligaments for human transplants (Tanaka et al., 2001). It has also been used as a binding agent in many applications, such as its addition to refractory clay to reduce the firing temperature and strengthen the product (Volkova et al., 2001). Calcium chloride is said to facilitate the formation of iron ore pellets and to improve both their green and dry crushing strengths (Ivanov et al., 1985; Kawatra et al., 1998). It has even been recommended to aid in the granulation of various halogen-containing polymers (Ueno et al., 1981). TOXICOLOGY AND SAFETY Calcium chloride is not a hazardous material, and it even finds many uses in foods and medicines. However, some precautions still must be taken in its

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379

handling. Its toxicity upon ingestion, is indicated by the test on rats: oral LD50 (rat) is 1.0 –1.4 g/kg (the lethal dose for half of the test animals, in this case rats, per kg of animal weight). When applied subcutaneously, the (sc)LD50 (rat) is 2.63 g/kg, and for dogs [sc]LDLO [dog] is 274 mg/kg. When contacted intraperitoneally the (ip)LD50 (rats) is 264 mg/kg and for mice 280 mg/kg. When applied intravenously, the (iv)LD50 (mouse) is 42 mg/kg. The lowest toxic dose for oral consumption by rats over a 20 week period (oral TDLO) in one test was 1.12 g/kg. The acute toxicity when ingested is thus quite low, but still somewhat more toxic than sodium chloride [oral NaCl LD50(rat) is 3.0 g/kg], and it may irritate the gastrointestinal tract if swallowed in large doses. However, the ip and iv toxicity is quite high, so care must be taken for these low-probability exposure routes (Sax, 1979). Contact of either solid material or strong solutions of calcium chloride in the eyes is likely to produce severe irritation and/or moderate corneal injury. The effects may include a burning sensation, conjunctival irritation with edema, as well as corneal injury. Single exposures to the skin may result in some reddening, while repeated or prolonged contacts may result in an allergic reaction, appreciable irritation and possibly a mild burn. Contact with abraded skin or cuts can cause severe necrosis. Contact with dilute solutions (under 5 –10%) usually has only a slight to no harmful effect. Dust inhalation may irritate the nose, throat and lungs, but there are no threshold values (TLVs) or ACGIH dust standards for this material. However, one manufacturer recommends that the dust concentration be kept below 10 mg/ft3 (Dow, 2002). Calcium chloride’s aquatic toxicity for fish, LD50 (96 hr) is 100 mg/liter (of anhydrous calcium chloride; Allied, 1980). Safety goggles or face shields should be used when calcium chloride dust or solution might contact the eyes, and long-sleeved shirts and gloves should be used when handling calcium chloride. Dust masks should also be employed when working with solid calcium chloride in a dusty atmosphere. Contact with leather (such as shoes) should be avoided, since calcium chloride causes dehydration, and will ruin the leather. When dissolving solid calcium chloride add the solids to the water slowly, as there will be a considerable temperature rise, and boiling or splashing may occur. Never add water to the solids because of this effect, and when diluting liquid calcium chloride add it to the water (if possible), and slowly for the same reason.

PRODUCTION STATISTICS The U.S. Bureau of Mines stopped compiling statistics on U.S. calcium chloride production and consumption in 1991, and the average price and net (price) realization by the producers in 1984. Since they were the principal source for this information, the marketing data for calcium chloride after those years have become much less accurate and quite scattered. Different sources have made quite different estimates for each of the statistical numbers, so average values have often

380

Table 2.17

Total Production 1999a ,900

1998 875

1997 —

1996 ,800

1995 827

1994 823

1993 746

1992 653

1991 621

1990 682

1989 678

1988 703

1987 657

1986 770

1985 944

1984 —

1983 943

1982 853

1981 917

1980 811

1979 841

1978 1031

1977 868

1976 897

1975 828

1974 —

1973 759

1972 940

1971 995

1970 907

1969 960

1968 970

1967 870

1966 890

1965 995

1964 890

1963 806

1962 741

1961 690

1960 670

1959 740

1958 645

1957 645

1956 650

1955 610

1954 524

1953 490

1952 428

1951 514

1950 420

1949 —

1948 —

Natural Calcium Chloride Only 1984 ,838

1983 664

1982 ,617

1981 705

1980 581

1979 720

1978 773

1959–1963 437 avg. ($19/t)

1945– 1950 ,300– , 350

California only (75% basis) 1945 6.8

1946 9.9

1947 7.5

1948 10.0

1949 11.1

1950 13.3

1951 17.0

1952 18.0

1953 17.0

C&EN (2000), Chemical Market Reporter, Chemical Week, Majmundar (1985), Manville Chemical Products Corp. (1993), Smith (1966), U.S. Bureau of Mines (1992 –1982), Parker (1978), Ver Planck (1957), other sources. a CMR 3/4/02 indicates 1999 at 1108; 2000 at 1143.

Part 2 Calcium Chloride

United States Production of Calcium Chloride (1000 st of equivalent 77– 78% CaCl2)

381

Production Statistics Table 2.18 Some Values for the United States Consumption of Calcium Chloride, 1000 st/yr of equivalent 77– 78% CaCl2 1999 ,900 1238a

1998 875 1186a

1996 ,800 1160a

1995 800 1140a

1993 805 1095a

1992 827

1991 739

1983 1064

1982 916

1981 971

1980 799

1975 790

1970 925

1965 1005

1990 817

1989 818

1988 919

1987 807

1986 828

1985 915

Chemical Market Reporter, Mannville Chemical Products Corp., U.S. Bureau of Mines (1992–1982), Majmundar (1985). a CMR 3/4/02 values from 1995 to 2000; 2000 1316 Mst; Imports: 2000 252,000, 1999 211,000; Exports: 2000 79,000, 1999 81,000.

been listed in the following tables. The total production rate of calcium chloride in the U.S. since 1950 is estimated in Table 2.17. There are fewer statistics for the natural product, but scattered data from 1945 to 1984 are also listed in this table. Some of the early production rates in California (Bristol Lake) are likewise listed from 1945 to 1953. During 1945 –1985 the natural production was about 75% of the total, but in recent years because of the increasing conversion of waste hydrochloric acid with limestone to form calcium chloride, the natural product has dropped to about 50%. In the early period the California production was about 3% of the total. The United States consumption of calcium chloride is estimated in Table 2.18, indicating that there has been very little increase in the consumption rate over the period from 1965 to 1999. Table 2.19 lists an estimated average sales price for both solid and liquid calcium chloride since 1982, based upon

Table 2.19 United States Average Sales Price of Calcium Chloride, $/st 77–78% CaCl2 (Flake)a 2001 179b 1981 120.57 1967 35.0

1999 134b 1980 111.90 1966 35.0

1996 159b 1979 86.06 1965 34.0

1994 164b 1978 79.42 1964 32.5

1993 143b 1974 32.8 1963 32.5

1991 137 1973 32.8 1962 32.0

1990 137 1972 44.0 1960 32.3

1990 (215c) 1971 44.0 1959 31.0

1989 139 1970 40.8 1957 31.1

1987 125b 1969 37.4 1955 26.9

1984 125b 1968 45.0 1954 26.9

1982 120.57

1950 23.1

(continues)

382

Part 2 Calcium Chloride Table 2.19 (continued ) 40% CaCl2 (Liquid)

2001 52.8c 1965 14.0

1994 53.6c 1964 14.0

1987 38.4c 1963 12.6

1984 36.3c 1960 12.6

1982 34.64 1955 11.8

1981 34.64 1950 9.68

1980 33.22

1979 25.89

1978 31.91

1972 16.5

1970 16.2

Natural Product, Average Realization, $/st, 75% CaCl2 basis 1984 110.98

1983 107.43

1982 99.72

1981 87.54

1980 82.53

1979 72.09

1978 69.74

California Only 1945 14.6

1946 17.3

1947 14.9

1948 16.3

1949 18.0

1950 19.9

1951 22.1

1952 21.8

1953 22.6

Average Realization from all Sales, $/st, 75% CaCl2 basis 1992 ,86.6 1981 96.81

1991 122.9 1980 91.35

1990 136.5 1979 75.91

1989 129.0 1978 72.79

1988 119.4 1964 29.05

1987 133.1

1986 145.9

1985 143.9

1984 120.2

1983 118.0

1982 108.7

a

CMR 3-4-02 stated that the list price was $270/st in 2002, and steady at $250/st from 1995 to 2001. Source: U.S. Burea of Mines, Chemical Market Reporter, other sources. b List price for 100 lb bags times the 1990 ratio of average net realization to the list price ($137/215). c List price.

an early ratio of the list-to-actual prices. Usually the quoted market prices are much higher than the actual average realized prices, making these numbers very uncertain. The production capacity for calcium chloride by the various United States operators is estimated in Table 2.20, and typical product specifications for the natural calcium chloride products are listed in Tables 2.21(A) and (B).

PHASE DATA AND PHYSICAL PROPERTIES Phase Data The solubility of calcium chloride in water is listed in Table 2.22 and plotted in Fig. 2.63. Figure 2.64 shows these same data plotted on a larger scale, and Fig. 2.65 indicates the fields of the various crystalline phases in more detail. There are at least four hydrates of calcium chloride, with 1,2,4 and 6H2O (and perhaps

Phase Data and Physical Properties

383

Table 2.20 United States Production Capacity of Natural Calcium Chloride (1000 st of equivalent 77% CaCl2; Chemical Marketing Reporter) Company General Chemical, Manistee, MIa Dow Chemical, Ludington, MI. Hill Brothers, Cadiz Lake, CA Lee Chemical, Cadiz Lake, CA Magnesium Corporation of America, Rowley, Utah National Chloride, Amboy, CA Wilkinson, Mayville, MI Tetra Technologies, Amboy, CA Total

2002

1999

1993

1990

1987

1984

1981

360 735 — 15 45

300 700 3(?) 15 35

— 700b 5 15c 35d

— 700 — 15 35

— 700 — 15 10

— 991 4 30 10

— 800 — — —

20 55 25 1255

15 55 25e 1148

11 55 18 834

11 24 18 803

11 24 18 778

14 5 36 1090

14 8 30 852

a

Started by Ambar September, 1997; closed April 30, 2000; sold to General in late 2000; closed in 2003. b Also listed at 625. c Also listed at 25. d Amax sold the operation in 1993 to Renco, who changed the name to Magnesium Corporation of America. e Purchased from Cargil (Leslie Salt) in August, 1998; includes Cadiz, CA.

a fifth with 0.33H2O; Reid and Kust, 1992), with considerable debate as to whether the tetrahydrate has three forms, a, b and g. The early researchers tabulated these three forms [including Seidell (1958); see Fig. 2.64], but later studies have indicated that they are merely quite reproducible metastable phases. The minimum temperature that can be reached by freezing a pure calcium chloride solution is 2 49.88C at 30.22% CaCl2 (Seidell, 1958; Yanateva, 1946; 2 49.88C and 30.52% CaCl2; Oakes et al., 1990; 2 49.958C and 30.33% CaCl2), although this value is very difficult to determine. Earlier researchers had indicated numbers down to 2 558C and 29.8% CaCl2. The more recent data from Oakes et al. (1990) for the solubility of CaCl2 when saturated with ice (and also for the CaCl2 – NaCl –H2O system) has been obtained by a rapid, less precise procedure. It attempts to remove clear brine from ice-filled samples as the temperature is slowly falling, and estimates the liquid compositions at each withdrawal temperature by comparing their densities against standards. The method would appear to only obtain approximate data, but surprisingly, the numbers agree quite well with other investigators (Table 2.22; for the CaCl2 – NaCl –H2O system Figs. 2.66 and 2.67). Their data further indicate that the ice freezing point of NaCl and CaCl2 solutions and their mixtures, as a first approximation is only dependent upon the total wt.% of salts in the solution (Fig. 2.67).

384

Table 2.21(A) Typical Calcium Chloride Product Specificationsa (wt.%) (Maximum Limits Unless Otherwise Specified) Flakesc

Pelletsd

Pelletse (almonds)

Inhibited liquidf

28–42

77– 80

90 (min)

Alkali metals as NaCl

1.68–2.52

4.3

4.9

Magnesium as MgCl2 Other salts Ca(OH)2 CaCO3 SO4 Fe, ppm Heavy metals as Pb, ppm

0.15–0.23 — — — — — —

0.07 0.85 0.10 0.07 0.04 50 20

0.08 0.98 0.20 0.20 0.20 50 20

91 (min) Briquette: Density (g/cc): 1.86–1.88 Porosity: 15–20% Size: 0.7 in. thick, 1.1 in. long Bed voids: 45 –50% Angle of repose: 288

— Colorless to pale yellow

51– 60 White flakes

58– 66 White pellets

30 (min) ppm P 25 Ba 10 Zn 10 As 5 Se 5 Pb 1 Cr 0.5 Cd 0.2 Cu 0.2 Hg 0.05 LC50 23,452g 83.5 Brown, slight odor

— — — — —

0 0–20 — 95–100

0 0– 20 90– 100 99– 100

Chemical analysis CaCl2

Bulk density (lbs/ft3) Appearance Screen analysis .3/8 in. .4 mesh .20 mesh .30 mesh a

All Dow Chemical products (Dow, 2001). The trademark names are: Liquidow, c Dowflake, d Peladow, e Peladow DG, and f Liquidow Armor (a corrosion-inhibited liquid). g Lowest concentration that is lethal to 50% of Rainbow trout, mg/liter CaCl2. b

60–68 White .1/2 in. 85% .1/4 in. 94–100%

Part 2 Calcium Chloride

Liquidb

Phase Data and Physical Properties

385

Table 2.21(B) General Specifications for Calcium Chloride Products 28 –42% CaCl2 Liquida: As 38% CaCl2; ,0.1% NaCl (alkali chlorides), ,0.1% MgCl2, ,1% others, excluding water; pH 7 –8; density at 688F (208C) 11.48 lbs/gal (r ¼ 1.3785 g/cc); crystallization temperature 418F (58C). Recommended to melt ice to 2178F (227.28C). 32 –41% CaCl2 Liquid, food gradea: Analysis based upon 100% CaCl2: Magnesium and alkali salts ,5% (as NaCl and MgCl2), ,0.3% alkalinity (as Ca[OH]2), fluoride ,0.004%, heavy metals ,0.002 (as Pb), lead, ,5 mg/kg; alkaline. 77% CaCl2 Flakeb: ,0.2% NaCl, 0.50% MgCl2, ,1.00% others; bulk density 48– 51 lbs/cuft; particle size: 0% þ 5 mesh (3.99 mm), .24% þ 10 mesh (2.00 mm), .25% þ16 mesh (1.19 mm), ,15% 230 mesh (0.59 mm) 77 –80% CaCl2 flakec: ,4.3% NaCl, ,0.07% MgCl2, ,0.85% other, ,0.1% Ca(OH)2, ,0.04% CaCO3, ,0.04% SO4, ,0.005% Fe, ,0.002% heavy metals (Pb); 51– 60 lbs/cuft; 0% þ 3/8 in., 0–20% þ4 mesh, 95–100% .30 mesh 94 –97% CaCl2 pelletsa: $94% CaCl2, ,2% alkali chlorides (as NaCl), ,0.1% magnesium (as MgCl2), ,1% others (not including water), ,20 ppm Fe, ,500 ppm Na, ,250 ppm SO4, ,2000 ppm alkalinity (as CaCO3); pH 7–10; density ,55 lbs/cuft; screen size: 100% 24 mesh (4.76 mm), 80– 100% 28 mesh (2.38 mm), 0–5% 240 mesh (0.42 mm). 94 –97% CaCl2 food gradea: .93% CaCl2, ,1% alkali chlorides (as NaCl), ,0.1% magnesium (as MgCl2), ,0.1% others (not including water), ,20 ppm Fe, ,500 ppm Na, ,250 ppm SO4, ,2000 ppm alkalinity (as CaCO3); pH 7–10; density ,55 lbs/cuft; screen size: 100% 24 mesh (4.76 mm), 80– 100% 28 mesh (2.38 mm), 0–6% 240 mesh (0.42 mm). a

Tetra Chemicals (2002). Dow Chemical Co. (2001). c Hill Brothers Chemical Co. (2002). b

The lowest CaCl2 –H2O temperatures obtained with a commercial product and with a short residence time to reach equilibrium (allowing the potential for considerable supersaturation) are stated to vary from 2 558C at 29.6% CaCl2 to 2 618C at 29.5% CaCl2. However, the presence of small amounts of MgCl2 and NaCl (less than 2%, as is present in commercial calcium chloride) can actually decrease the freezing point to a minimum of 2 578C (Luzhnaya and Vereshchetina, 1946). More extensive data on the CaCl2 – H2O system has been published by Potter and Clynne (1978; to 978C, also with an approximate method), Brass and Thurmond (1983; low temperatures), Williams-Jones and Seward (1989; 100– 3608C), Oakes et al. (1994; supercritical, two-Phase; 0.3 –3.0 mol/kg), Doherty (1990; supercritical), and Sterner and Felmy (1995; a computer program for Pitzer equation parameters). The solubility of calcium chloride in steam is given by Martynova et al. (1966).

386

Part 2 Calcium Chloride Table 2.22 The Solubility of Calcium Chloride in Water (A. Data of Seidell, 1958) wt.% CaCl2 solution saturated with the solid phase: Ice

CaCl2·6H2O

CaCl2·4H2Oa

CaCl2·4H2Ob

CaCl2·4H2Og

CaCl2·2H2O

CaCl2·H2O

30.22(8) 28.3 25.7 21.3 14.4 0.0 — — — — — — — — — — — — — — — — — — — — — — — — — — — — —

30.22(8) 32.1 33.4 34.7 36.0 37.3 39.3 41.2 42.7 45.3 47.2 48.5 49.1 50.0(3) — — — — — — — — — — — — — — — — — — — — —

— — — — — — — 46.0a,b 47.5a 48.7a 49.3a 49.7a 49.8a 50.0(3) 50.7 51.5 52.4 52.7 53.4 53.8 56.6(6)e — — — — — — — — — — — — — —

— — — — — — — 49.0c 50.0a 51.2a 51.8a 52.1a 52.3(2) 52.5 53.1 53.9 54.7 55.1 55.9 56.3(5)d — — — — — — — — — — — — — — —

— — — — — — — — 51.1a 52.1a 52.7a 53.0(1) 53.2 53.5 54.0 54.8 55.7 56.0(4) — — — — — — — — — — — — — — — — —

— — — — — — — — — — — — — — — — — 56.0(4) 56.2 56.3(5) 56.6(6) 57.8 58.6 59.5 60.6 61.4 63.4 65.6 69.0 71.8 74.8(7) — — — —

— — — — — — — — — — — — — — — — — — — — — — — — — — — — — — 74.8(7) 75.0 75.7 76.8 77.6

t (degrees) 249.8 240 230 220 210 0 þ 10 15 20 25 27.5 29.0 29.5 30.1 32.5 35.0 37.5 38.5 40.0 41.0 45.1 60 70 80 90 100 120 140 160 170 175.5 180 200 235 200

(B) Data of Potter and Clynne (1978) CaCl2·6H2O

CaCl2·4H2O

CaCl2·2H2O

Temperature (8C)

CaCl2

Temperature (8C)

CaCl2

Temperature (8C)

CaCl2

9.13 12.75 22.28 28.16 30.08

38.96 39.87 43.12 47.15 —

30.08 33.54 35.15 38.64 42.62 44.81 45.13

— 50.37 50.94 52.23 54.20 55.47 —

45.13 49.37 61.81 68.63 85.85 97.65

— 56.03 56.85 57.28 58.69 59.75

Solid Phases: (1) CaCl2·6H2O þ CaCl2·4H2Og; (2) CaCl2·6H2O þ CaCl2·4H2Ob; (3) CaCl2·6H2O þ CaCl2·4H2Oa; (4) CaCl2·4H2Og þ CaCl2·2H2O; (5) CaCl2·4H2Ob þ CaCl2·2H2O; (6) CaCl2·4H2Oa þ CaCl2·2H2O; (7) CaCl2·2H2O þ CaCl2·H2O; (8) Ice þ CaCl2·6H2O. a Metastable. b At 148. c At 15.98. d Others reported 55.8%. e Others reported 55.9%.

Phase Data and Physical Properties

Figure 2.63 The solubility of calcium chloride in water (Seidell, 1958).

Figure 2.64 Calcium chloride solubility on a larger scale (Seidell, 1958).

387

388

Part 2 Calcium Chloride

Figure 2.65 Calcium chloride phase diagram showing the phases present (Seidell, 1958).

Tables 2.23 – 2.25 list solubility data for calcium chloride with sodium chloride, and Figs. 2.66 –2.68 plot some of these data. This system has been extensively used to guess at the composition of occlusions in various minerals (because of their observed low freezing temperature and subsequent melting point [ignoring supersaturation and frequently up to a 508C difference between these two values]), and for the deicing use of calcium chloride. Besides the extensive data of Seidell (1958); including that of Yanateva (1946), there have been publications on this system by Mun and Darar (1956) and more recently by Gibbard and Fong (1975; to 2 48C), Oakes (1992) and Oakes et al. (1990). Other articles on this system have

Phase Data and Physical Properties

389

Figure 2.66 The freezing point of ice in CaCl2 or NaCl solutions (Data of: £ ¼ Seidell, 1958; Mun and Darar, 1956; W, A ¼ Oakes et al., 1990).

been presented by Igelsrud and Thompson (1936; 08C), Vanko et al. (1988; a few points on the ice – NaCl·2H2O – solution border), Ivanov (1964; 408C), Yuan et al. (2000a; 258C; estimated data by the Pitzer equation), Williams-Jones and Samson (1990; theoretical), Naden (1996; a Microsoft Excel 5.0 computer program for CaCl2 solutions), Makarov and Shcharkova (1969) and Kirgintsev and Lukyanov (1965). There is considerable data available on the calcium chloride – magnesium chloride system, as listed in Table 2.26. Since the solubility of sodium chloride in concentrated calcium or magnesium chloride solutions is quite low, these data

390

Part 2 Calcium Chloride

Figure 2.67 (a) The freezing point of ice in mixed CaCl2 and NaCl solutions (ratio of NaCl/(NaCl þ CaCl2); Data of: Mun and Darar (1956) W ¼ 0.125, £ ¼ 0.274, D ¼ 460, A ¼ 0.687; Oakes et al., 1990 P ¼ 0.393, 8 ¼ 0.80, 2 ¼ 0.20). (b) Triangular solubility diagram for the CaCl2 – NaCl– H2O system, with the higher temperature points under 2–4.5 kb pressure (After Oakes et al., 1990; Vanko et al., 1988).

Phase Data and Physical Properties

391

Table 2.23 The System Calcium Chloride–Sodium Chloride–Water (wt.%) (Seidell, 1958) CaCl2

NaCl

Solid phase

CaCl2

Results at 258 8.40 7.60 4.40 0.0 0.0 4.80 13.80 26.40 30.80 33.80

0.0 1.20 4.60 7.40 25.60 21.20 14.80 4.60 2.42 1.40

Ice Ice Ice Ice NaCl·2H2O NaCl·2H2O þ NaCl NaCl NaCl NaCl NaCl

2.20 8.20 13.50 25.00 14.20 5.80 4.30 2.40 1.20

Ice Ice Ice NaCl·2H2O NaCl·2H2O NaCl·2H2O þ NaCl NaCl NaCl NaCl þ CaCl2·6H2O

24.20 23.80 22.00 19.60 24.70 29.40 32.00

2.60 10.50 19.00 24.40 13.80 4.80 2.36 1.40 1.30

Ice Ice Ice NaCl·2H2O NaCl·2H2O NaCl·2H2O þ NaCl NaCl NaCl NaCl þ CaCl2·6H2O

25.80 25.60 23.80 23.50 24.20 28.80 31.60

3.20 12.00 22.40 23.60 13.60 4.40

Ice Ice Ice NaCl·2H2O NaCl·2H2O NaCl·2H2O

Ice Ice Ice Ice þ NaCl·2H2O NaCl·2H2O NaCl·2H2O NaCl·2H2O þ CaCl·6H2O

1.00 1.80 4.00 4.50 4.20 2.20 1.45

Ice Ice Ice Ice þ NaCl·2H2O NaCl·2H2O NaCl·2H2O NaCl·2H2O þ CaCl2·6H2O

Results at 2408 27.20 26.80 26.20 28.60 31.00

Results at 2208 18.40 11.40 0.0 0.0 12.90 25.40

1.00 1.80 3.80 7.00 1.20 2.20 1.40

Results at 2358

Results at 2158 16.20 10.00 0.0 0.0 13.00 25.80 30.65 33.60 33.70

Solid phase

Results at 2308

Results at 2108 13.00 7.80 0.0 0.0 13.40 24.00 26.10 30.70 33.80

NaCl

1.00 2.00 3.10 2.10 1.50

Ice Ice Ice þ NaCl·2H2O NaCl·2H2O NaCl·2H2O þ CaCl2·6H2O

Results at 2458 28.60 27.80 27.40 28.30 30.30

1.20 2.20 2.50 2.20 1.60

Ice Ice Ice þ NaCl·2H2O NaCl·2H2O NaCl·2H2O þ CaCl2·6H2O (continues)

392

Part 2 Calcium Chloride Table 2.23 (continued) Results at 2208 (cont.)

29.00 30.60 33.30 33.40

3.00 2.30 1.30 1.35

NaCl·2H2O þ NaCl NaCl NaCl NaCl þ CaCl2·6H2O

Results at 2458 (cont.) 30.40

1.20

CaCl2·6H2O

Results at 2258 20.40 12.30 11.20 12.80 24.80 29.80 32.60 32.80

3.60 12.80 13.80 13.40 4.20 2.20 1.30 1.40

Ice Ice Ice þ NaCl·2H2O NaCl·2H2O NaCl·2H2O NaCl·2H2O NaCl·2H2O NaCl·2H2O þ CaCl2·6H2O Results at 08 (Separate data from two investigators)

CaCl2

NaCl

Solid phase

CaCl2

NaCl

Solid phase

37.55 37.50 35.31 27.14 20.60 16.12 9.53 5.33 0.0

0.0 0.31 0.93 3.59 7.78 11.36 17.28 21.19 26.42

CaCl2·6H2O CaCl2·6H2O þ NaCl NaCl NaCl NaCl NaCl NaCl NaCl NaCl·2H2O

0.0 13.80 26.60 31.40 34.40 37.60

26.23 14.60 4.60 2.60 1.40 1.10

NaCl NaCl NaCl NaCl NaCl NaCl þ CaCl2·6H2O

Results at 258 (Separate data from two investigators) d25 25

CaCl2

NaCl

Solid phase

CaCl2

NaCl

Solid phase

— 1.4441

84 78.49

0 1.846

45.60 0.0

0.0 26.80

CaCl2·6H2O NaCl

1.3651 1.3463 1.2831 1.2653 1.2367 1.2080 1.2030

48.58 53.47 36.80 30.08 19.53 3.92 0

1.637 1.799 7.77 10.70 18.85 32.48 35.80

CaCl2·6H2O CaCl2·6H2O þ NaCl NaCl NaCl NaCl NaCl NaCl NaCl NaCl

27.30 32.40 35.60 43.50

4.70 2.80 1.40 1.00

NaCl NaCl NaCl NaCl þ CaCl2·6H2O

(continues)

Phase Data and Physical Properties

393

Table 2.23 (continued) Results at 508 g per 100 g sat. sol.

Results at 94.58 g per 100 g sat. sol.

NaCl

Solid phase

CaCl2

NaCl

Solid phase

57.0

0.0

CaCl2·2H2O

58.1

0.8

56.3

0.9

57.4

1.1

30.9 15.1 3.3 0.0

3.6 13.2 24.4 26.8

CaCl2·2H2O þ NaCl NaCl NaCl NaCl NaCl

CaCl2·2H2O þ NaCl NaCl

45.7 32.8 15.3 11.5

1.3 4.3 15.0 18.1

CaCl2

NaCl NaCl NaCl NaCl

Table 2.24 The System Calcium Chloride– Sodium Chloride–Water at Negative Temperatures with Ice Always Present (wt.%) Temperature (8C)

CaCl2

NaCl

21a

0 0.36 0.77 1.21 1.70 2.22

1.73 1.46 1.15 0.81 0.42 0

25

0a 0c 0b 1.61a 2.45c 3.30a 4.40b 4.47c 5.07a 6.15c 6.94a 7.50c 7.60b 8.88a 8.73c 8.40b

7.82 7.73 7.40 6.44 5.60 4.94 4.60 3.80 3.38 2.32 1.74 1.07 1.20 0 0 0

Temperature (8C)

CaCl2

NaCl

2 10

0a 0c 0b 2.80a 4.36c 5.60a 7.50c 7.80b 8.42a 10.18c 11.28a 12.26c 13.00c 14.01c 14.06a 14.40b

13.96 13.90 13.50 11.18 9.59 8.40 6.37 8.20 5.62 3.84 2.82 1.75 2.20 0 0 0

2 15

0a 0b 3.69a 7.30a 10.00b 10.97a

18.71 19.00 14.75 10.96 10.50 7.10 (continues)

394

Part 2 Calcium Chloride Table 2.24 (continued )

Temperature (8C)

CaCl2

NaCl

2 15

14.43a 16.20b 17.84a 18.50b

3.49 2.60 0 0

2 20

0a 0b 4.39a 8.64a 11.40b 12.74a 16.72a 17.34a 18.40b 20.58a 21.30b

22.36 22.40 17.58 12.96 12.00 8.50 4.18 3.53 3.20 0 0

Temperature (8C)

CaCl2

NaCl

2 25

9.73a 14.46a 19.34a 20.40c 22.90a 23.80b

14.60 9.36 3.93 3.60 0 0

2 30

15.50a 20.32a 21.03a 22.00b 23.80b 24.20b 24.84a 25.70b

10.33 5.08 4.28 3.80 1.80 1.00 0 0

a

Oakes et al. (1990). Seidell (1958). c Mun and Darar (1956). b

Table 2.25 Calcium Chloride and Sodium Chloride Solutions in Equilibrium with Ice (Oakes et al., 1990) 1.000

0.593

0.195

0.000

u

wt.%

u

wt.%

u

wt.%

u

wt.%

0.00 1.84 2.03 5.42 6.98 9.22 10.27 12.49 12.86 14.91 16.91 19.17 20.64 21.48

0.00 3.08 3.40 8.45 10.48 13.11 14.25 16.46 16.78 18.57 20.18 21.80 22.80 23.34

0.00 1.70 2.06 3.01 4.05 5.19 5.21 6.02 7.08 8.10 9.10 10.03 11.02 12.19 13.24 14.08 15.14 17.04 19.06

0.00 3.13 3.77 5.32 6.92 8.50 8.54 9.58 10.88 12.05 13.10 14.03 14.96 16.00 16.89 17.55 18.36 19.69 21.02

0.00 1.55 2.03 3.37 4.40 5.05 6.00 7.02 8.09 9.02 10.04 11.13 11.97 13.96 16.09 18.07 20.36 22.01

0.00 3.18 4.09 6.34 7.87 8.75 9.95 11.13 12.27 13.18 14.12 15.06 15.75 17.22 18.65 19.82 21.08 21.94

0.00 1.42 2.08 3.68 4.06 5.21 6.02 6.03 7.06 8.08 8.42 9.08 10.98 11.62 12.03 13.50 14.21 15.30 16.02

0.00 3.13 4.42 7.09 7.64 9.16 10.15 10.20 11.35 12.39 12:64a 13.34 14.83a 15.46 15.72 16.77 17.31 18.04a 18.42 (continues)

395

Phase Data and Physical Properties Table 2.25 (continued) 0.796

u

0.00 1.71 2.01 3.05 4.01 4.55 4.94 4.98 6.12 6.13 7.03 8.84 10.12 11.00 11.91 13.32 14.19 15.19 16.05 17.03 18.01 18.99 18.99 19.00 20.05 20.96 22.13 22.48

0.393 wt.%

0.00 3.03 3.54 5.20 6.68 7.43 7.97 8.09 9.57 9.56 10.69 12.78 14.12 14.98 15.81 17.05 17.79 18.58 19.25 19.96 20.65 21.30 21.30 21.31 21.99 22.56 23.25 23.63a

0.169

u

wt.%

21.02 23.23 23.28

22.19 23.41 23.43

0.00 1.59 2.04 2.10 3.18 4.22 5.46 6.94 8.06 9.16 10.06 11.11 12.06 13.00 14.02 15.32 16.95 17.03 19.36 21.22 23.21 25.07 25.07 26.93

0.00 3.12 3.88 4.02 5.84 7.40 9.08 10.88 12.10 13.24 14.09 15.04 15.84 16.60 17.37 18.30 19.39 19.41 20.85 21.87 22.96 23.85 23.85 24.71

u

19.99 22.11 24.05 25.98 28.02 30.08 30.21 32.11 34.03

0.000 wt.%

20.86 21.94 22.85 23.70 24.54 25.34 25.39 26.09 26.75

u

wt.%

17.80 19.32 19.64 20.78 21.12 21.97 23.08 24.02 25.23 25.25 26.08 27.06 27.08 27.21 27.22 28.08 28.98 30.16 32.09 32.13 34.07 35.72 36.03 37.07 38.08 39.03 39.83 40.80 42.16 42.89 43.15 44.00 44.77 45.64 46.55 46.66 48.28 49.37 50.14 50.51 51.20

19.46 20.19 20.40 20.97 21.16 21.56 22.04 22.41 23.01 22.95 23.37 23.73 23.73 23.78 23.82 24.16 24.44 24.89 25.52 25.56 26.17 26.69 26.78 27.08 27.39 27.66 27.88 28.14 28.50 28.68 28.71 28.97 29.16 29.37 29.58 29.61 29.94 30.23 30.39 30.46 30.59a

Freezing point depressions (u, 8C) and total salinities (wt.%) for NaCl/(NaCl þ CaCl2) wt ratios shown in parentheses at the head of each pair of data columns. a Uncertain values.

396

Part 2 Calcium Chloride

Figure 2.68 Plot of the CaCl2 –NaCl –KCl–H2O phase diagram, with all points saturated with NaCl (Holdorf et al., 1993; reprinted from the Seventh Symposium on Salt (ISBN 0444891439), vol. 1, p. 573, Fig, 3, q1993, with permission form Elsevier).

can also provide a rough estimate of the CaCl2 – MgCl2 – NaCl system, where there is much less data. Some of these phase data for calcium chloride with magnesium chloride and sodium chloride in water are listed in Table 2.27, and other data are given by Luzhnaya and Vereshchetina (1946). This system would be encountered in the processing of natural brines when the magnesium is not precipitated. However, the amount of data available on this system is very limited, but eutectic temperatures down to 2 578C are indicated when the NaCl is , 2% (for example: 2 538C with 1.6% NaCl, 20.85% CaCl2, 6.95% MgCl2; Reps and Schuhmann, 1965). Charykova et al. (1992) used the Pitzer equation on the CaCl2 – MgCl2 – NaCl –H2O system to predict that tachyhydrite would only begin to be stable above 22.88C. Majima et al. (1980) found that when crystallized at 1108C, 2CaCl2·MgCl2·6H2O was formed. Some of the similar data with potassium chloride instead of sodium chloride are shown in Table 2.28 (these data are also very limited). Potassium chloride is slightly more soluble than sodium chloride, but other than that, the data should be reasonably similar. Perova (1957) presented 25 and 558C isotherms for this system with KCl instead of NaCl, and Holdorf et al. (1993) plotted

397

Phase Data and Physical Properties Table 2.26 The System Calcium Chloride–Magnesium Chloride–Water (wt.%) (Seidell, 1958) Results at 158C CaCl2

MgCl2

Solid phase

CaCl2

MgCl2

Solid phase

0.0 20.30 22.60

35.50 20.30 18.40

MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O þ CaCl2·6H2O

31.00 38.60 41.20

10.60 2.80 0.0

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Results at 258C Density

CaCl2

MgCl2

Solid phase

1.470 1.465 1.472 1.486 1.473 1.460

45.06 41.87 38.95 38.70 36.37 32.82

0.0 4.06 7.93 9.43 10.78 13.55

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O þ 1.2.12 1.2.12 1.2.12

Density

CaCl2

MgCl2

Solid phase

1.455 1.441 1.428 1.391 1.371 1.341

31.17 28.12 25.09 16.05 10.33 0.0

14.54 16.31 18.13 23.33 27.61 35.54

1.2.12 þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

Results at 258C (second author) CaCl2

MgCl2

Solid Phase

CaCl2

MgCl2

Solid Phase

45.0 44.88 44.55 42.14 41.48 39.58 38.82 38.80 33.36 31.32 22.29 7.87

0.0a 0.168a 0.612a 3.098 4.115 6.47 8.14a 9.42a 13.58a 14.60a 19.32a 29.25a

C þ Mg(OH)2 C þ Mg(OH)2 C þ 3.1.11 C C C C þ 3.1.11 C þ 1.2.12 3.1.11 þ 1.2.12 3.1.11 þ 1.2.12 þ M 3.1.11 þ M 3.1.11

7.46 48.80 46.49 44.71 43.42 43.38 43.19 41.99 40.03 38.88 38.94

29.84a 0.0 2.739 4.742 5.97 6.01 6.24 7.30 8.52 9.43 9.38

3.1.11 C4a C4a C4a C4a C4a C4a C4a C4a C4a C4a

Results at 358 g per 100 g sat. sol. CaCl2

MgCl2

Solid Phase

51.33 49.09 47.75 45.03 45.03 45.07 39.25 35.64 35.15 30.06 26.98 26.78 26.70 26.66

0.0 2.10 3.50 6.20 6.20 6.22 9.49 11.88 12.19 15.67 17.95 18.04 18.13 18.24

CaCl2·4H2Oa CaCl2·4H2Oa CaCl2·4H2Oa CaCl2·4H2Oa þ CaCl2·2MgCl2·12H2O CaCl2·4H2Oa þ CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O þ MgCl2·6H2O CaCl2·2MgCl2·12H2O þ MgCl2·6H2O

(continues)

398

Part 2 Calcium Chloride Table 2.26 (continued ) Results at 358 g per 100 g sat. sol.

CaCl2

MgCl2

Solid Phase

23.15 18.89 9.81 3.63 0.0

20.28 23.17 29.07 33.56 36.28

MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O Results at 758

58.58 55.76 52.59 52.57 45.01 34.97 29.15 19.28 13.83 12.52 10.15 9.97 8.31 8.31 8.04 3.84 0.0

0.0 2.53 5.27 5.28 8.74 14.70 18.63 25.55 29.65 30.59 32.39 32.59 33.90 33.91 34.07 36.62 39.12

CaCl2·2H2O CaCl2·2H2O CaCl2·2H2O þ CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O CaCl2·2MgCl2·12H2O þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O Results at 1108C

g per 100 g sat. sol.

g per 100 g sat. sol.

CaCl2

MgCl2

Solid phase

CaCl2

MgCl2

0.0 2.0 13.6 18.2 26.9 40.0 47.8

42.8 41.4 30.8 28.0 21.2 14.0 10.0

MgCl2·6H2O MgCl2·6H2O þ 1:1:12 1:1:12 1:1:12 1:1:12 1:1:12 1:1:12

49.0 51.0 53.5 55.7 56.8 62.3

9.5 8.0 6.2 4.7 4.0 0.0

Solid phase 1:1:2 þ 1:2:6 1:2:6 1:2:6 1:2:6 þ CaCl2·2H2O CaCl2·2H2O CaCl2·2H2O

Results at 08C g per 100 g sat. sol. CaCl2

MgCl2

g per 100 g sat. sol. Solid phase

CaCl2

(Y.) 0.0 14.00

34.80 23.50

MgCl2·6H2O MgCl2·6H2O þ CaCl2·6H2O

MgCl2

Solid phase

(P. & T.) (I. & T.) 37.44 30.22

0.0 6.82

CaCl2·6H2O CaCl2·6H2O

(continues)

Phase Data and Physical Properties

399

Table 2.26 (continued ) Results at 08C g per 100 g sat. sol. CaCl2

MgCl2

g per 100 g sat. sol. Solid phase

CaCl2

MgCl2

(Y.) 18.50 27.80 35.20 37.26

18.50 9.40 2.60 0.0

Solid phase

(P. & T.) (I. & T.) CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

27.17 25.10 22.12 14.89 13.88 13.87 5.84 5.70 4.51 0.0

9.81 11.95 14.94 22.83 23.83 23.91 29.91 30.04 30.89 34.62

Results at 2 58 (Y.)

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O Results at 2108 (Y.)

g per 100 g sat. sol.

g per 100 g sat. sol.

CaCl2

MgCl2

Solid phase

CaCl2

MgCl2

Solid phase

8.40 8.20 6.40 4.60 0.0 0.0 10.60 17.80 27.20 34.40 36.20

0.0 0.70 2.20 4.60 7.40 34.40 26.00 17.80 9.10 2.20 0.0

Ice Ice Ice Ice Ice MgCl2·6H2O MgCl2·6H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

14.80 14.20 10.80 7.00 0.0 0.0 10.40 17.60 26.60 33.60 35.50

0.0 1.00 3.60 7.00 11.60 33.60 25.40 17.60 8.80 2.20 0.0

Ice Ice Ice Ice Ice MgCl2·8H2O Ice þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Results at 2 158 g per 100 g sat. sol. CaCl2

MgCl2

g per 100 g sat. sol. Solid phase

CaCl2

MgCl2

(Y.) 18.50 17.20 13.20 8.50 0.0 0.0 10.78 17.40 26.00

0.0 1.20 4.40 8.50 14.60 32.60 24.21 17.40 8.60

Ice Ice Ice Ice Ice MgCl2·8H2O MgCl2·8H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Solid phase (P. & T.)

0.0 2.83 3.90 11.97 24.14 25.71 28.07 29.98 33.86

34.78 31.92 30.50 22.42 10.73 9.43b 7.52b 5.11b 0.0b

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O MgCl2·8H2Oa CaCl2·6H2O CaCl2·6H2O þ MgCl2·6H2O MgCl2·6H2O MgCl2·6H2O

(continues)

400

Part 2 Calcium Chloride Table 2.26 (continued ) Results at 2158

g per 100 g sat. sol. CaCl2

g per 100 g sat. sol.

MgCl2

Solid phase

CaCl2

MgCl2

(Y.) 32.80 34.50

2.10 0.0

Solid phase (P. & T.)

CaCl2·6H2O CaCl2·6H2O

25.68 26.15 27.17 28.70 31.85 0.0 14.69 10.85 4.01

Results at 2208 (Y.)

8.67 7.70 6.46 4.42 0.0 12.23 0.0 5.0 13.35

MgCl2·8H2Oa MgCl2·8H2Oa MgCl2·8H2Oa MgCl2·8H2Oa MgCl2·8H2Oa Ice Ice Ice Ice

Results at 2258 (Y.)

CaCl2

MgCl2

Solid phase

CaCl2

MgCl2

Solid phase

21.00 20.00 15.00 9.60 0.0 0.0 17.20

0.0 1.30 5.00 9.60 16.80 26.60 17.20

Ice Ice Ice Ice Ice MgCl2·12H2O CaCl2·6H2O

23.50 21.80 16.60 10.60 0.0 0.0 16.30

0.0 1.40 5.60 10.60 18.80 24.30 17.40

25.40 32.20 34.00

8.40 2.00 0.0

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

16.90 25.00 31.60 33.60

16.90 8.40 2.10 0.0

Ice Ice Ice Ice Ice MgCl2·12H2O MgCl2·12H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Results at 2308 CaCl2

MgCl2

Solid phase

CaCl2

(Y.)

MgCl2

Solid phase

(P. & T.)

25.00 23.60 18.00 11.40

0.0 1.60 6.00 11.40

Ice Ice Ice Ice

32.92 30.07 27.04 23.55

0.0 2.50 5.49 8.61

0.0 0.0 15.80

24.40 22.50 15.80

19.78b 15.23b 7.40b

12.26 16.76 25.55

18.20 24.60

15.20 8.20

Ice MgCl2·12H2O MgCl2·12H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O þ MgCl2·12H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

6.14b 19.13

27.20 9.94

CaCl2·6H2O MgCl2·12H2O

(continues)

Phase Data and Physical Properties

401

Table 2.26 (continued ) Results at 2 308 CaCl2

MgCl2

Solid phase

CaCl2

MgCl2

(Y.) 31.20 33.40

2.00 0.0

(P. & T.)

CaCl2·6H2O CaCl2·6H2O

19.44 11.39 11.27 7.63

10.20 14.70 14.65 16.99

MgCl2·12H2O MgCl2·12H2O MgCl2·12H2O MgCl2·12H2O

Ice Ice Ice Ice Ice þ MgCl2·12H2O MgCl2·12H2O MgCl2·12H2O þ CaCl2·6H2O

0.0 0.0 4.83 6.18 25.03

22.19 20.47 16.45 17.16 0.0

MgCl2·12H2O Ice Ice Ice Ice

CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

28.20 26.40 20.30 12.70 11.00

0.0 1.70 6.70 12.70 14.40

Ice Ice Ice Ice Ice þ MgCl2·12H2O

13.40 23.40 24.00 30.00 32.00

13.40 8.80 8.10 1.95 0.0

MgCl2·12H2O MgCl2·12H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Results at 2358 (Y.) 26.60 25.00 19.20 12.20 8.00 14.30 21.00

0.0 1.60 6.40 12.20 16.40 14.30 11.60

24.30 30.60 32.40

8.10 2.00 0.0

Results at 2458 (Y.) 29.50 27.60 21.10 17.10 22.60 25.00 29.40 31.80

0.0 1.60 6.00 10.40 7.50 6.50 1.90 0.0

Solid phase

Ice Ice Ice Ice þ MgCl2·12H2O MgCl2·12H2O MgCl2·12H2O þ CaCl2·6H2O CaCl2·6H2O CaCl2·6H2O

Results at 2 408 (Y.)

Other temperatures g per 100 g sat. sol. t(8C)

CaCl2

MgCl2

Solid phase

(249.88) (251.78) (255.08) (252.28) (220.78) (2 6.78) (250.28) (243.08) (233.68)

30.40 28.86 26.00 27.40 10.56 9.90 21.50 13.0 0.0

0.0 2.08 5.00 2.93 23.23 26.40 7.30 13.00 21.38

Ice þ CaCl2·6H2O Ice þ CaCl2·6H2O Ice þ CaCl2·6H2O þ MgCl2·12H2O Ice þ CaCl2·6H2O þ MgCl2·12H2O CaCl2·6H2O þ MgCl2·12H2O þ MgCl2·8H2Oa CaCl2·6H2O þ MgCl2·8H2O þ MgCl2·6H2O Ice þ MgCl2·12H2O Ice þ MgCl2·12H2O Ice þ MgCl2·12H2O

(Y.) (Y.) (Y.) (P.&T.) (P.&T.) (P.&T.) (Y.) (Y.) (Y.)

1.2.12 ¼ Tachhydrite ¼ CaCl2·2MgCl2·12H2O. C ¼ CaCl2·6H2O; 3.1.11 ¼ 3MgO·MgCl2·11H2O; M ¼ MgCl2·6H2O; C4a ¼ CaCl2·4H2Oa; 1:1:12 ¼ MgCl2·CaCl2·12H2O; 1:2:6 ¼ MgCl2·2CaCl2·6H2O; Y and P&T are two different sources; Y ¼ Second author; P&T ¼ Third author. a These solutions also contained 0.002 g MgO. b Metastable.

402

Table 2.27 The System Calcium Chloride–Magnesium Chloride–Sodium Chloride–Water (wt.%) (Seidell, 1965) Temperature (8C)

g MgCl2 per 100 g sat. sol.

g NaCl per 100 g sat. sol.

Solid phase

38.83 30.50 23.40 10.44 3.50 4.80 3.05 3.43 7.20 11.49 5.84 12.30 20.96 3.95 2.45 3.52 12.70 3.97 7.90 13.85 6.75 14.65 22.28 5.69 12.18 8.28 14.82 13.15 0.0 0.0

9.56 14.95 19.15 27.70 3.50 1.60 6.15 10.29 7.20 3.81 17.54 12.30 6.98 3.95 7.35 1.84 4.23 11.93 7.90 13.85 20.25 14.65 7.42 17.09 12.18 24.86 14.82 13.15 22.71 23.35

0.30 0.38 0.32 0.32 16.0 17.00 13.90 8.6 9.0 9.20 1.72 2.00 1.66 18.5 16.0 19.10 10.40 9.3 10.2 2.40 2.05 1.60 1.70 0.32 1.38 0.56 0.65 1.42 1.56 1.81

CaCl2·6H2O þ 2MgCl2·CaCl2·12H2O þ NaCl MgCl2·6H2O þ 2MgCl2·CaCl2·12H2O þ NaCl MgCl2·6H2O þ NaCl MgCl2·6H2O þ NaCl Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O Ice þ NaCl·2H2O NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl NaCl·2H2O þ NaCl Ice þ MgCl2·12H2O Ice þ MgCl2·12H2O MgCl2.·12H2O þ NaCl·2H2O MgCl2.·12H2O þ NaCl·2H2O MgCl2.·12H2O þ NaCl·2H2O Ice þ MgCl2·12H2O þ NaCl·2H2O MgCl2·12H2O þ NaCl·2H2O þ NaCl

System at 258C

2 21.5 2 22.5 2 24.3 2 26.5 2 26.5 2 27.0 2 39.0 2 44.0 2 52.0 2 0.5 2 1.2 2 1.3 2 6.5 2 6.5 2 8.0 2 19.5 2 23.8 2 25.0 2 27.5 2 38.3 2 42.5 2 17.5 2 27.0 2 30.0 2 35 2 26

The Quaternary Eutectic is about 2588 with a ratio of about 17:83 MgCl2:CaCl2 in solution.

Part 2 Calcium Chloride

g CaCl2 per 100 g sat. sol.

Phase Data and Physical Properties

403

Table 2.28 The System Calcium Chloride–Magnesium Chloride–Potassium Chloride– Water (wt.%) (Seidell, 1958) Results at 358 g per 100 g sat. sol.

g per 100 g sat. sol.

MgCl2

CaCl2

KCl

Solid phase

MgCl2

CaCl2

KCl

Solid phase

27.33 20.61 10.59

0.0 9.12 23.57

3.81 3.69 3.47

CþK CþK CþK

5.83 4.07 4.01

31.66 35.51 35.60

3.38 3.44 3.43

CþK CþK CþK

Results at 358

Results at 758

g per 100 g sat. sol.

Av.

Av.

Av.

Av.

Av.

g per 100 g sat. sol.

MgCl2

CaCl2

KCl

Solid phase

MgCl2

CaCl2

KCl

Solid phase

2.31 1.30 1.15 1.17 1.02 36.17 26.66 18.07 14.23 13.83 7.28 6.98 6.15 5.91 4.42 2.69 1.72 1.36 0.0 0.71 0.93 6.20 6.13 5.82 18.18 4.34a 4.22a 3.97a 3.39a 3.06a 2.05a 1.51a 0.89a

41.00 45.85 48.34 48.89 49.76 0.0 12.97 26.74 32.01 32.60 42.75 43.27 44.44 45.30 46.56 48.40 49.25 49.54 50.45 50.06 49.83 45.03 45.18 45.43 26.68 48.18 48.63 49.24 50.49 50.81 51.78 52.05 50.70

3.80 4.71 5.69 6.04 6.46 0.14 0.16 0.23 0.32 0.33 0.71 0.76 0.97 0.99 1.33 2.28 3.92 5.10 6.48 6.46 6.45 0.0 0.33 0.70 0.0 1.41 1.53 1.64 2.10 2.10 3.91 4.53 7.01

CþK CþK CþK CþK C þ K þ Ca4 C þ Mg6 C þ Mg6 C þ Mg6 þ T CþT CþT CþT CþT CþT C þ Ca4 þ T C þ Ca4 C þ Ca4 C þ Ca4 C þ Ca4 K þ Ca4 K þ Ca4 K þ Ca4 T þ Ca4 T þ Ca4 T þ Ca4 T þ Mg6 CþT CþT CþT C þ T þ Ca4 C þ Ca4 C þ Ca4 C þ Ca4 C þ K þ Ca4

29.26 23.14 16.71 10.89 6.39 4.48 3.85 3.47 38.86 36.20 33.68 26.95 16.12 7.92 6.76 5.58 0.0 1.85 5.28 4.66 4.55 4.51 4.97 5.44 6.68 0.0 2.71 33.91

0.0 8.56 18.07 27.57 36.76 41.95 44.63 46.31 0.0 4.26 8.40 16.96 32.00 44.90 46.43 46.69 57.66 55.23 52.58 52.47 52.51 52.29 51.21 50.06 46.63 51.20 47.32 8.31

5.57 5.56 5.61 5.85 6.73 7.81 8.80 9.70 0.32 0.38 0.39 0.67 1.58 3.68 5.11 6.25 3.59 3.35 0.0 2.65 3.07 3.19 3.44 3.81 5.02 10.33 9.84 0.0

CþK CþK CþK CþK CþK CþK CþK CþKþD C þ Mg6 C þ Mg6 C þ Mg6 þ T CþT CþT CþT CþTþD CþD Ca2 þ D Ca2 þ D Ca2 þ T Ca2 þ T Ca2 þ T þ D TþD TþD TþD TþD KþD KþD T þ Mg6

C ¼ KCl·MgCl2·6H2O; T ¼ CaCl2·2MgCl2·12H2O; Ca4 ¼ CaCl2·4H2O; Av. ¼ Average; K ¼ KCl; D ¼ 2KCl·CaCl2·2H2O; Ca2 ¼ CaCl2·2H2O; Mg6 ¼ MgCl2·6H2O. a Metastable.

404

Table 2.29 The Density of Calcium Chloride Solutions, rT4 (Perry and Chilton, 1978)

(%) 2 4 8 12 16 20 25 30 35 40

258C

08C

208C

308C

408C

608C

808C

1008C

1208C

1408C

1.0708 1.1083 1.1471 1.1874

1.0171 1.0346 1.0703 1.1072 1.1454 1.1853 1.2376 1.2922

1.0148 1.0316 1.0659 1.1015 1.1386 1.1775 1.2284 1.2816 1.3373 1.3957

1.0120 1.0286 1.0626 1.0978 1.1345 1.1730 1.2236 1.2764 1.3316 1.3895

1.0084 1.0249 1.0586 1.0937 1.1301 1.1684 1.2186 1.2709 1.3255 1.3826

0.9994 1.0158 1.0492 1.0840 1.1202 1.1581 1.2079 1.2597 1.3137 1.3700

0.9881 1.0046 1.0382 1.0730 1.1092 1.1471 1.1965 1.2478 1.3013 1.3571

0.9748 0.9915 1.0257 1.0610 1.0973 1.1352 1.1846 1.2359 1.2893 1.3450

0.9596 0.9765 1.0111 1.0466 1.0835 1.1219

0.9428 0.9601 0.9954 1.0317 1.0691 1.1080

(B) Potter and Clynne, 1976, r T4 Temperature (8C) wt.% Molality 125 150 175 200 225 250 275 300

1

3

0.948 0.927 0.902 0.876 0.841 0.803 0.762 0.717

0.967 0.947 0.923 0.899 0.868 0.833 0.794 0.751

5.26 0.5 0.986 0.967 0.945 0.920 0.892 0.860 0.823 0.872

7 1.001 0.982 0.960 0.936 0.909 0.878 0.843 0.803

9.99 1 1.026 1.008 0.987 0.963 0.938 0.909 0.876 0.830

14.27 1.5 1.064 1.046 1.025 1.002 0.978 0.951 0.921 0.886

18.16 2 1.101 1.083 1.062 1.038 1.015 0.990 0.962 0.931

21.72 2.5 1.137 1.118 1.097 1.073 1.051 1.027 1.001 0.973

24.98 3.0 1.173 1.154 1.132 1.107 1.085 1.062 1.038 1.013

27.98 3.5 1.208 1.188 1.166 1.141 1.118 1.096 1.074 1.052 (continues)

Part 2 Calcium Chloride

(A)

Table 2.29 (continued ) (B) Potter and Clynne, 1976, r T4 wt.% Molality 0 25 50 75 100

30.74 4.0 1.301 1.287 1.264 1.252 1.238

33.31 4.5 1.315 1.292 1.280 1.267

35.69 5.0 1.342 1.319 1.308 1.294

37.90 5.5 1.367 1.346 1.335 1.322

39.97 6.0 1.391 1.372 1.362 1.349

41.91 6.5 1.414 1.398 1.389 1.376

43.72 7.0 1.435 1.424 1.416 1.403

45.42 7.5

1.449 1.442 1.430

(C) Oakes, 1990; r TT (for r T4 divide by the r of water; 258C ¼ 0.997077, 358C ¼ 0.994061, 15.568C (608F) ¼ 0.9990412) 258C

358C

wt.%

258C

358C

41.58 39.96 38.00 36.10 35.99 33.99 31.05 29.23 28.93 28.01 26.38 22.39 21.99 21.72 21.11

1.4086 1.3905 1.3681 1.3468 1.3456 1.3232 1.2906 1.2708 — — 1.2405 1.1994 — 1.1928 1.1867

1.4021 1.3841 — — 1.3396 1.3174 — — 1.2624 1.2525 — — 1.1908 — —

17.78 16.40 14.46 14.39 12.40 10.16 9.98 6.91 6.19 5.08 3.84 3.61 1.94 1.57 1.03

1.1541 1.1410 1.1228 — 1.1038 1.0837 1.0820 1.0551 — 1.0395 1.0290 1.0270 1.0131 1.0100 —

— 1.1368 — 1.1180 — — 1.0783 — 1.0455 — — 1.0237 — — 1.0025

Phase Data and Physical Properties

wt.%

405

406

Part 2 Calcium Chloride

three isotherms of the CaCl2 –NaCl – KCl system when saturated with NaCl (Fig. 2.68). A great deal of other solubility data exist for calcium chloride with many other salts and other solvents. The latter data have been used for systems attempting to extract or salt-out calcium chloride from solution, such as the solubility data on dimethylformamide by Spiridonov et al. (1976). Physical Properties The density of calcium chloride solutions is listed in Tables 2.29 and 2.30, and plotted in Fig. 2.69. The method of reporting the density should be noted, since the common scientific notation assumes that the density at a given temperature is compared with water at 48C (where it is 1.00000 g/cc), or r 25 4 for 258C. When the data are reported as being compared with water at the same temperature (i.e., r 25 25) it must be divided by the density of water to give the true density. Additional density data for the CaCl2 – H2O and CaCl2 –NaCl – H2O systems are given by many other authors, such as Hu (2000), Zhang et al. (1997), Oakes (1992), Oakes et al. (1990), Kumar and Atkinson (1983), Kumar et al. (1982), Ivanov (1964), Klochko et al. (1959; the last two also listed viscosity and electric conductivity), Yuan et al. (2000a) and Yuan and Liu (2000; the last two made estimates based upon the Pitzer equation). Some of the density data for the ternary system CaCl2 –NaCl – H2O are listed in Table 2.31, and shown in Figs. 2.70 and 2.71. As seen by the figures, it would appear that the density of CaCl2 and NaCl mixtures are closely approximated by a linear combination of the two pure solutions’ density at that temperature. Viscosity data for calcium chloride solutions are shown in Tables 2.32 and 2.33 and Fig. 2.72. Other authors presenting viscosity data include Isono (1984; also density and electrical conductivity data), Nowlan et al. (1980) and Goncalves and Kestin (1979). The specific heat of CaCl2 solutions is shown in Table 2.34 and Fig. 2.73 (note that the values for cal/g/8C are the same as for Btu/lb/8F), and other values are given by Perron et al. (1981). Vapor pressure data for calcium chloride solutions are shown in Figs. 2.74– 2.76. Additional values are given by Lannung (1940), Mun and Darar (1956), Baker and Waite (1921; from 150 to 3508), Zarembo et al. (1980; with activity coefficients), and Tkachenko and Shmulovich (1994; from 400 to 6008C). Boiling point data are listed in Table 2.35 and Figs. 2.77 and 2.78 (the temperature when the vapor pressure is 760 mm of Hg. Above 69% CaCl2, there is no boiling point because the dihydrate decomposes). The vapor pressure over calcium chloride solutions is listed in Table 2.36, and the surface tension of CaCl2 solutions listed in Table 2.37 and plotted in Fig. 2.79. The heat evolved when solid calcium chloride is added to water is listed in Table 2.38, while the temperature increase in dissolving flake calcium chloride is listed in Table 2.39. Fig. 2.80 shows these two values graphically. The enthalpy of various calcium chloride solutions is given in Table

Phase Data and Physical Properties

407

Table 2.30 The Density, Viscosity, Index of Refraction, Freezing Point Depression, Osmolality and Specific Conductance of Calcium Chloride Solutions at 208C (Weast, 1977) A% by wt.

r D20 4

C (g/liter)

M (g mol/liter)

n

D (8C)

O (Os/kg)

h/h0

g (mmho/cm)

0.50 1.00 1.50 2.00 2.50 3.00 3.50 4.00 4.50 5.00 5.50 6.00 6.50 7.00 7.50 8.00 8.50 9.00 9.50 10.00 11.00 12.00 13.00 14.00 15.00 16.00 17.00 18.00 19.00 20.00 22.00 24.00 26.00 28.00 30.00 32.00 34.00 36.00 38.00 40.00

1.0024 1.0065 1.0106 1.0148 1.0190 1.0232 1.0274 1.0316 1.0358 1.0401 1.0443 1.0486 1.0529 1.0572 1.0615 1.0659 1.0703 1.0747 1.0791 1.0835 1.0923 1.1014 1.1105 1.1198 1.1292 1.1386 1.1482 1.1579 1.1677 1.1775 1.1976 1.2180 1.2388 1.2600 1.2816 1.3036 1.3260 1.3488 1.3720 1.3957

5.0 10.1 15.2 20.3 25.5 30.7 36.0 41.3 46.6 52.0 57.4 62.9 68.4 74.0 79.6 85.3 91.0 96.7 102.5 108.3 120.2 132.2 144.4 156.8 169.4 182.2 195.2 208.4 221.9 235.5 263.5 292.3 322.1 352.8 384.5 417.1 450.8 485.6 521.4 558.3

0.045 0.091 0.137 0.183 0.230 0.277 0.324 0.372 0.420 0.469 0.518 0.567 0.617 0.667 0.717 0.768 0.820 0.871 0.924 0.976 1.083 1.191 1.301 1.412 1.526 1.641 1.759 1.878 1.999 2.122 2.374 2.634 2.902 3.179 3.464 3.758 4.062 4.375 4.697 5.030

1.3342 1.3354 1.3366 1.3378 1.3390 1.3402 1.3414 1.3426 1.3438 1.3451 1.3463 1.3475 1.3487 1.3500 1.3512 1.3525 1.3537 1.3549 1.3562 1.3575 1.3600 1.3625 1.3651 1.3677 1.3704 1.3730 1.3757 1.3784 1.3812 1.3839 1.3895 1.3951 1.4008 1.4066 1.4124 1.4183 1.4242 1.4301 1.4361 1.4420

0.222 0.440 0.661 0.880 1.102 1.330 1.567 1.815 2.074 2.345 2.630 2.930 3.244 3.573 3.917 4.275 4.649 5.04 5.44 5.86 6.74 7.70 8.72 9.83 11.01 12.28 13.65 15.11 16.7 18.3 21.7 25.3 29.7 34.7 41.0 49.7

0.119 0.237 0.355 0.473 0.593 0.715 0.843 0.976 1.115 1.261 1.414 1.575 1.744 1.921 2.106 2.299 2.499 2.71 2.92 3.15 3.63 4.14 4.69 5.28 5.92 6.60 7.34 8.12 8.98 9.84 11.67 13.60 15.97 18.66 22.04 26.72

1.013 1.026 1.037 1.048 1.061 1.076 1.091 1.108 1.125 1.141 1.157 1.173 1.189 1.206 1.222 1.240 1.257 1.276 1.295 1.316 1.359 1.405 1.454 1.505 1.561 1.622 1.688 1.760 1.839 1.926 2.123 2.351 2.640 2.994 3.460 4.027 4.810 5.795 7.306 8.979

8.1 15.7 22.7 29.4 36.1 42.6 48.9 55.1 61.1 67.0 72.8 78.3 83.6 88.7 93.6 98.4 103.0 108.0 112.0 117.0 125.0 133.0 141.0 148.0 154.0 160.0 165.0 169.0 173.0 177.0 182.0 183.0 182.0 179.0 172.0 162.0 150.0 137.0 123.0 106.0

A% ¼ anhydrous solute weight per cent, g solute/100 g solution; r or D20 4 ¼ relative density at 208C, kg/liter; Cs ¼ anhydrous solute concentration, g/liter; M ¼ molar concentration, g mol/liter; n ¼ index of refraction at 208C relative to air for sodium yellow light; D ¼ freezing point depression, 8C; O ¼ osmolality, Os/kg water; h/h0 ¼ relative viscosity, ratio of the absolute viscosity of a solution at 208C to the absolute viscosity of water at 208C. (Water’s viscosity at 208C is 1.002); g ¼ specific conductance (electrical) at 208C, mmho/cm.

408

Part 2 Calcium Chloride

Figure 2.69 The density of calcium chloride solutions at various temperatures.

2.40. The freezing point depression, refractive index, osmolality and specific conductance are all listed in Table 2.30. An et al. (1978) list activity and osmotic coefficients. The water absorbed by solid calcium chloride is listed in Table 2.41 and Fig. 2.81. Electrical conductance and ionization constants have been presented by Frantz and Marshall (1982). Other volumetric properties of calcium chloride solutions have been given by Potter and Clynne (1973) and Tsay et al. (1988; at high temperatures and pressure). Various bulk properties are presented by Alekhin et al. (1980). The activity coefficients of calcium chloride are given by Long et al.

Phase Data and Physical Properties

409

Table 2.31 Density of Mixed CaCl2 –NaCl –H2O Solutions (Concentrations as wt.%) CaCl2 wt.% 258C Hu (2000) 1.90 0.60 1.32 4.04 1.53 2.80 5.33 3.07 5.98 8.30 4.57 7.21 9.84 12.99 5.09 8.08 11.07 13.06 17.46 15.59 13.72 11.85 8.11 6.24 20.09 16.57 13.06 8.69 22.25 17.33 14.89 10.83 8.41

Fraction

wt.% NaCl

Total wt.%

62.7 17.5 23.0 64.8 19.1 34.3 62.4 28.2 53.5 72.7 31.4 49.1 66.4 86.7 30.4 48.0 65.5 77.1 90.2 80.5 70.8 61.1 41.8 32.1 91.7 75.2 58.9 39.0 89.4 68.8 58.7 42.3 32.6

1.13 2.82 4.41 2.19 6.45 5.37 3.21 7.81 5.19 3.11 9.99 7.48 4.98 1.99 11.67 8.75 5.83 3.88 1.89 3.77 5.65 7.54 11.30 13.19 1.83 5.47 9.10 13.62 2.63 7.87 10.47 14.78 17.35

3.03 3.42 5.73 6.23 7.98 8.17 8.54 10.88 11.17 11.41 14.56 14.69 14.82 14.98 16.76 16.83 16.90 16.94 19.35 19.36 19.37 19.39 19.41 19.43 21.92 22.04 22.16 22.31 24.88 25.20 25.36 25.61 25.76

1.0208 1.0221 1.0396 1.0464 1.0567 1.0591 1.0649 1.0798 1.0852 1.0905 1.1097 1.1137 1.1186 1.1252 1.1276 1.1318 1.1367 1.1406 1.1670 1.1642 1.1609 1.1578 1.1522 1.1502 1.1925 1.1878 1.1825 1.1772 1.2210 1.2157 1.2128 1.2082 1.2059

19.91 18.00 16.63 15.64 13.91 12.09 10.36 8.80 8.25 6.38

25.02 22.62 20.89 19.65 17.48 15.19 13.02 11.06 10.03 8.02

1.1950 1.1741 1.1594 1.1490 1.1311 1.1125 1.0952 1.0799 1.0719 1.0565

258C Oakes et al. (1990) 5.11 20.41 4.62 20.41 4.26 20.41 4.01 20.41 3.57 20.41 3.10 20.41 2.66 20.41 2.26 20.41 2.05 20.41 1.64 20.41

Density

(continues)

410

Part 2 Calcium Chloride Table 2.31 (continued ) CaCl2

wt.%

Fraction

wt.% NaCl

Total wt.%

Density

1.37 1.02 0.61 0.41 0.20 0.13

20.41 20.41 20.41 20.41 20.41 20.41

5.36 3.98 2.38 1.60 0.79 0.51

6.73 5.00 2.99 2.01 0.99 0.64

1.0468 1.0338 1.0190 1.0118 1.0043 1.0018

10.23 9.77 9.52 8.67 7.92 6.30 4.73 3.17 2.28 1.57 0.79 0.40

40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67

14.92 14.25 13.90 12.65 11.56 9.20 6.90 4.62 3.34 2.29 1.16 0.58

25.15 24.02 23.42 21.32 19.48 15.50 11.63 7.79 5.62 3.86 1.95 0.98

1.2042 1.1938 1.1884 1.1695 1.1534 1.1194 1.0874 1.0568 1.0398 1.0263 1.0118 1.0044

15.96 14.70 13.78 12.38 11.14 9.90 8.42 7.42 6.21 4.95 3.77 3.01 2.50 1.86 1.21 0.62 0.34

60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68 60.68

10.35 9.53 8.92 8.02 7.21 6.42 5.45 4.81 4.03 3.20 2.44 1.95 1.62 1.20 0.78 0.40 0.22

26.31 24.23 22.70 20.40 18.35 16.32 13.87 12.23 10.24 8.15 6.21 4.96 4.12 3.06 1.99 1.02 0.56

1.2238 1.2038 1.1892 1.1678 1.1493 1.1312 1.1099 1.0959 1.0791 1.0619 1.0461 1.0361 1.0294 1.0210 1.0126 1.0050 1.0014

21.11 20.55 18.59 17.13 15.56 13.64 11.97 10.28 8.46

80.48 80.48 80.48 80.48 80.48 80.48 80.48 80.48 80.48

5.11 4.99 4.50 4.15 3.77 3.31 2.90 2.49 2.05

26.22 25.54 23.09 21.28 19.33 16.95 14.87 12.77 10.51

1.2311 1.2242 1.1999 1.1823 1.1637 1.1415 1.1226 1.1038 1.0841 (continues)

Phase Data and Physical Properties

411

Table 2.31 (continued ) CaCl2 wt.%

Fraction

wt.% NaCl

Total wt.%

7.24 5.94 5.11 4.28 3.40 2.58 1.69 0.87 0.40

80.48 80.48 80.48 80.48 80.48 80.48 80.48 80.48 80.48

1.75 1.44 1.24 1.03 0.82 0.62 0.41 0.21 0.10

8.99 7.38 6.35 5.31 4.22 3.20 2.10 1.08 0.50

1.0710 1.0574 1.0487 1.0402 1.0312 1.0228 1.0140 1.0057 1.0011

23.06 21.85 20.12 18.92 17.01 16.27 15.42

83.11 83.11 83.11 83.11 83.11 83.11 83.11

4.69 4.44 4.09 3.85 3.45 3.30 3.14

27.75 26.29 24.21 22.77 20.46 19.57 18.56

1.2478 1.2329 1.2119 1.1976 1.1752 1.1667 1.1571

26.09 26.00 24.00 23.50 21.30 20.00 18.74 17.12 16.09 16.00 15.09 12.62 12.00 10.73 9.29 8.00 6.89 4.45 4.00 2.86 2.00 1.00

— — — — — — — — — — — — — — — — — — — — — —

1.19529 1.19443 1.17726 1.17388 1.15588 1.145330 1.13545 1.12276 1.11479 1.11401 1.10713 1.08840 1.08365 1.07423 1.06364 1.05412 1.04605 1.02858 1.02530 1.01730 1.01712 1.00409

19.91 16.63 12.09 7.98 3.98

25.02 20.89 15.19 10.03 5.00

1.1899 1.1544 1.1080 1.0679 1.0302

0 0a 0a 0 0 0a 0 0 0 0a 0 0 0a 0 0 0a 0 0 0a 0 0a 0a 358C Oakes (1900) 5.11 4.26 3.10 2.05 1.02

0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 20.41 20.41 20.41 20.41 20.41

Density

(continues)

412

Part 2 Calcium Chloride Table 2.31 (continued ) CaCl2

wt.%

Fraction

wt.% NaCl

Total wt.%

0.41 0.13 10.23 7.92 7.22 4.73 2.28 1.57 0.79 0.20

20.41 20.41 40.67 40.67 40.67 40.67 40.67 40.67 40.67 40.67

1.60 0.51 14.92 11.56 10.53 6.90 3.34 2.29 1.16 0.29

2.01 0.64 25.15 19.48 17.75 11.63 5.62 3.86 1.95 0.49

1.0085 0.9986 1.1990 1.1487 1.1339 1.0833 1.0362 1.0229 1.0085 0.9977

15.96 12.37 8.42 3.77 1.21 0.62

60.67 60.67 60.67 60.67 60.67 60.67

10.35 7.67 5.45 2.44 0.78 0.40

26.31 20.04 13.87 6.21 1.99 1.02

1.2186 1.1631 1.1057 1.0425 1.0094 1.0019

21.10 17.13 11.97 7.24 3.40 0.87

80.48 80.48 80.48 80.48 80.48 80.48

5.12 4.15 2.90 1.75 0.82 0.21

26.22 21.28 14.87 8.99 4.22 1.08

1.2258 1.1776 1.1184 1.0673 1.0278 1.0026

23.08 19.99 17.00 13.06 10.03 7.03 3.76 1.51 0.51 0.49

— — — — — — — — — —

1.16508 1.14038 1.11705 1.08720 1.06490 1.04323 1.02019 1.00448 0.99759 0.99745

0 0 0 0 0 0 0 0 0 0 0 0 0

1.40211 1.38415 1.33961 1.31743 1.26240 1.25254 1.19080 1.13681 1.11799 1.07833 1.04545 1.02372 1.00248

0 0 0 0 0 0 0 0 0 0

0 0 0 0 0 0 0 0 0 0

41.58 39.96 35.99 33.99 28.93 28.01 21.99 16.40 14.39 9.98 6.19 3.61 1.03 a

100 — — — — — — — — — — — –

Perry and Chilton (1978).

0 0 0 0 0 0 0 0 0 0 0 0 0

Density

Phase Data and Physical Properties

413

Figure 2.70 The density of calcium chloride–sodium chloride solutions at 258C († ¼ Hu, 2000; D ¼ Oakes et al., 1990).

(1999) and Galleguillos et al. (1999). The activity coefficient in seawater are noted by Tishchenko and Popova (1992), while other ionic forms of calcium chloride in seawater have been suggested by Kesov and Sychkova (1983). The vapor pressure of HCl over solid phase CaCl2 as it decomposes at temperatures from 380 to 5008C are given by Bischoff et al. (1996). For CaCl2 – NaCl solutions, various physical properties have also been determined. The solutions’ density (Table 2.31), and the activity coefficients of the ions at temperatures to 2508C and 400 bars (395 psi) pressure have been determined by Oakes (1992), and the heat of dilution and mixing from 100 to 3008C and 21.5 MPa (145 psi) pressure by Oakes et al. (1998). Low temperature Raman spectroscopic investigations have been made by Samson and Walker

414

Part 2 Calcium Chloride

Figure 2.71 The density of calcium chloride–sodium chloride solutions at 358C († ¼ Oakes et al., 1990).

Table 2.32 The Density and Viscosity of Calcium Chloride Solutions at 258C, g/cm3 and Centipoise ( ¼ mPa s) (Zhang et al., 1997) wt.%

Molality

Density

Viscosity

wt.%

Molality

Density

Viscosity

0 0.244 0.265 0.518 0.651 0.755

0 0.0220 0.0239 0.0469 0.0590 0.0685

0.99708 0.9991 0.9993 1.0014 1.0025 1.0033

0.8904 0.8977 0.8982 0.9046 0.9082 0.9106

8.646 8.780 11.387 15.951 17.254 18.008

0.8528 0.8673 1.1578 1.7100 1.8788 1.9790

1.0703 1.0714 1.0947 1.1368 1.1491 1.1563

1.1386 1.1431 1.2430 1.4636 1.5409 1.5897 (continues)

415

Phase Data and Physical Properties Table 2.32 (continued) wt.%

Molality

Density

Viscosity

wt.%

Molality

Density

Viscosity

1.071 1.572 2.016 2.220 2.839 3.145 4.185 4.258 4.417 5.550 6.090 6.187 6.285 6.378 8.052

0.0975 0.1439 0.1854 0.2046 0.2633 0.2926 0.3936 0.4007 0.4164 0.5295 0.5843 0.5942 0.6043 0.6139 0.7890

1.0059 1.0101 1.0138 1.0155 1.0206 1.0232 1.0320 1.0325 1.0339 1.0435 1.0481 1.0490 1.0497 1.0506 1.0651

0.9186 0.9311 0.9425 0.9476 0.9637 0.9719 1.0001 1.0020 1.0063 1.0392 1.0554 1.0583 1.0610 1.0641 1.1182

18.849 18.947 21.769 22.647 24.406 27.655 30.259 34.579 37.299 38.409 43.198 44.868 45.290 46.648

2.0929 2.1063 2.5073 2.6380 2.9091 3.4443 3.9093 4.7625 5.3599 5.6189 6.8525 7.3330 7.4589 7.8783

1.1645 1.1654 1.1932 1.2021 1.2201 1.2540 1.2820 1.3297 1.3604 1.3729 1.4268 1.4454 1.4502 1.4650

6.6476 1.6543 1.8825 1.9666 2.1566 2.6110 3.1142 4.4138 5.7327 6.4365 11.2940 14.0283 14.8392 17.8223

m ¼ mwater ð1 þ Am0:5 þ Bm þ Cm2 þ Dm3:5 þ Em7 Þ: Table 2.33 Viscosities of Calcium Chloride Solutions, Centipoise Temperature (8C) wt.% CaCl2 0a 0 1.1 1.6 2.7 3.4 5.0a 5.0 5.3 6.8 8.7 10.0a 10.5 12.2 14.1 15.0a 16.0 17.8 18.2 19.8

225

220

215

210

0

10

18

20

25

30

40

— — — — — — — — — — — — — — — — — — — —

— — — — — — — — — — — — — — — — — — — —

— — — — — — — — — — — — — — — — — — — 6.30

— — — — — — — — — — — — — — — 4.09 4.22 4.52 — 4.89

1.77 1.79 — 1.80 — 1.85 1.84 1.92 — 2.00 2.10 2.13 2.21 2.33 2.49 2.50 2.68 2.89 — 3.12

1.29 1.31 — 1.32 — 1.36 1.35 1.41 — 1.46 1.53 1.52 1.61 1.70 1.81 1.84 1.95 2.11 2.25 2.31

— 1.06 1.09 — 1.14 —

1.02 — — — — — 1.07 — — — — 1.16 — — — 1.40 — — — —

— 0.89 0.92 — 0.96 — — — 1.03 — — 1.19 — — — — — — 1.59 —

0.79 — — — — — 0.82 — — — — 0.93 — — — 1.20 — — — —

0.67 — — — — — 0.73 — — — — 0.86 — — — 1.03 — — 1.19 —

— 1.21 — — 1.37 — — — — — — 1.84 —

(continues)

416

Part 2 Calcium Chloride

Table 2.33 (continued ) Temperature (8C) wt.% CaCl2

225

220

215

210

0

10

18

20

25

30

40

20.0a 21.8 23.7 25.0a 25.7 27.6 29.7 30.0a 30.8 31.8 34.0 35.0a 36.1 37.2 40.0a 40.0 45.0a

— — 11.45 — 12.90 14.96 — — — — — — — — — — —

— 8.52 9.33 9.94 10.51 12.01 14.27 14.27 — 17.26 — — — — — — —

— 6.82 7.48 — 8.40 9.60 11.17 — — 13.44 — — — — — — —

4.97 5.34 5.85 6.32 6.59 7.57 8.86 9.04 — 10.61 — — — — — — —

3.12 3.41 3.76 4.04 4.22 4.85 5.69 5.77 — 6.74 8.16 8.83 9.90 10.86 — — —

2.33 2.56 2.84 3.07 3.20 3.68 4.34 4.30 4.55 5.10 6.12 6.62 7.23 7.87 11.75 12.29 —

— — — — — — — — 3.74 — — — — — — 9.50 —

1.81 — — 2.38 — — — 3.33 — — — 4.99 — — 8.48 — —

— — — — — — — — 3.22 — — — — — — 7.60 —

1.54 — — 1.97 — — — 2.62 — — — 3.87 — — 6.39 — 11.50

1.22 — — 1.54 — — — 2.07 2.44 — — 3.07 — — 4.90 5.38 8.90

508C

60

70

80

90

100

0.53 0.57 0.64 0.76 0.99 1.27 1.73 2.54 4.00 6.57 11.80

0.46 0.51 0.57 0.68 0.85 1.07 1.43 2.17 3.26 5.24 9.24

0.40 0.45 0.51 0.62 0.74 0.90 1.24 1.82 2.72 4.25 7.45

0.34 0.39 0.47 0.55 0.68 0.82 1.01 1.46 2.15 3.39 5.97

0.30 0.35 0.42 0.49 0.59 0.70 0.89 1.22 1.74 2.77 4.95

0.26 0.28 0.35 0.42 0.49 0.59 0.73 1.03 1.52 2.33 4.28

0a 5a 10a 15a 20a 25a 30a 35a 40a 45a 50a

Data from Allied Chemical (1980). Data from Dow Chemical (1980).

a

(2000) and Walker (1994). Thermodynamic properties have been estimated by Galleguillos et al. (1999). A comparison of the boiling points for pure calcium chloride with that of commercial solutions of 77% CaCl2 (Fig. 2.78) indicates that the commercial

Physical Data and Physical Properties

417

Figure 2.72 The viscosity of calcium chloride solutions, centipoises (Dow, 2001, courtesy of the Dow Chemical Company).

solutions have 1– 38C higher boiling points than pure solutions. Table 2.42 lists the density of a commercial CaCl2 solution and its crystallization temperatures. The densities were measured at 608F (15.568C), and compared to water at 608F (i.e., considering it to be 1.000), while the crystallization data were taken with a very short residence time, making it quite supersaturated (i.e., indicating temperatures up to 10% lower than equilibrium solubility data). The pH of solutions of commercial calcium chloride (after caustic soda [NaOH] has been added to make it basic) are listed in Table 2.43. Calcium chloride solutions are inherently acidic, since they are the salt of a strong acid (HCl) and a week base (Ca[OH]2). However, to reduce their (acidity) corrosiveness, the solutions are usually neutralized and made slightly basic during processing. Finally, Table 2.44 lists various physical properties of CaCl2 and the calcium chloride hydrates, using the convention for thermochemical values used by the National Bureau of Standards. A negative sign indicates that the process is exothermic (gives off heat). Additional detail on thermodynamic properties is given by Sinke et al. (1985). Various other properties, such as sublimation pressures are given by Sommer (1958).

418

Table 2.34 Specific Heats of Pure Calcium Chloride Solutions

% CaCl2 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

2408C; 2408F

0.636 0.634 0.631 a

2308C; 2228F

0.656 0.649 0.642 0.639 0.635

International Critical Tables, vol. 2, p. 328, 1929.

2208C; 248F

0.695 0.686 0.678 0.670 0.663 0.656 0.649 0.644 0.638

2108C; þ 148F

08C; 328F

þ108C; 508F

208C; 688F

308C; 868F

0.768 0.756 0.745 0.734 0.723 0.713 0.704 0.694 0.686 0.678 0.670 0.663 0.656 0.649 0.643

0.882 0.867 0.853 0.839 0.825 0.812 0.799 0.787 0.775 0.763 0.752 0.741 0.731 0.721 0.711 0.702 0.693 0.685 0.677 0.669 0.662 0.655 0.649

0.887 0.872 0.858 0.844 0.831 0.818 0.805 0.793 0.781 0.770 0.759 0.748 0.738 0.728 0.718 0.709 0.700 0.692 0.683 0.676 0.669 0.662 0.655

0.892 0.877 0.863 0.849 0.836 0.823 0.811 0.799 0.787 0.775 0.764 0.754 0.744 0.733 0.724 0.715 0.706 0.698 0.690 0.682 0.675 0.668 0.661

0.897 0.882 0.868 0.855 0.842 0.829 0.817 0.805 0.792 0.780 0.769 0.759 0.749 0.739 0.729 0.721 0.712 0.704 0.696 0.689 0.682 0.675 0.668

Part 2 Calcium Chloride

Specific heat in Btu per pounda

Physical Data and Physical Properties

419

Figure 2.73 The specific heat of calcium chloride solutions, cal/g 8C (Dow, 2001, courtesy of the Dow Chemical Company).

Figure 2.74 The vapor pressure and relative humidity of calcium chloride solutions at 188C (Lannung, 1936; £ ¼ Dow, 2001).

420

Part 2 Calcium Chloride

Figure 2.75 The vapor pressure of calcium chloride solutions, lower temperature range, mm Hg (Allied, 1980).

Physical Data and Physical Properties

421

Figure 2.76 The vapor pressure of calcium chloride solutions, higher temperature range, mm Hg (Dow, 2001, courtesy of the Dow Chemical Company).

Table 2.35 Boiling Points of Calcium Chloride Solutions (Allied Chemical, 1980) wt.% CaCl2 5 10 15 20 25 30 35 40 45 50 55 60 70

Boiling Points (8C) 100.7 101.6 102.9 104.6 107.0 110.2 114.8 119.8 124.8 130.6 137.3 144.6 165

422

Part 2 Calcium Chloride

Figure 2.77 The boiling temperature of calcium chloride solutions, lower temperature range (Allied, 1980).

Figure 2.78 The boiling temperature of calcium chloride solutions to higher temperatures, and commercial solutions (Dow, 2001, courtesy of the Dow Chemical Company).

423

Physical Data and Physical Properties Table 2.36 Examples of the Vapor Pressure of Calcium Chloride Solids and Solution (Lannung, 1936) Temperature (8C)

Pressure (mm Hg)

Temperature (8C)

Relative humidity

Pressure (mm Hg)

Relative humidity

Temperature (8C)

Pressure (mm Hg)

Relative humidity

60

0.82

0.00549

18.5 47.4

0.594 4.00

0.0372 0.0492

18.0 43.4

1.33 10.24

0.0859 0.1548

17.2 22.0

2.55 3.78

0.1733 0.1906

(A) The solids CaCl2 $ CaCl2·H2Oa 0 92.3

0.0075 6.04

0.00164 0.0105

30

0.095

0.00298

(B) The solids CaCl2·H2Oa $ CaCl2·2H2O 0 28.7 54.1

0.135 1.25 5.97

0.0295 0.0423 0.0528

17.45 31.6

0.544 1.51

0.0364 0.0433

(C) The solids CaCl2·2H2O $ CaCl2·4H2Oa 0 30.0

0.244 3.65

0.0533 0.1147

17.2 31.7

1.25 4.28

0.0850 0.1221

(D) The solids CaCl2·4H2Oa $ CaCl2·6H2O 0 17.5 28.0

0.60 2.61 6.14

0.1310 0.1740 0.2166

9.7 18.0

1.36 2.73

0.1507 0.1764

(E) Two-phase equilibrium points with calcium chloride in water Vapor pressure (mm Hg)

Temperature (8C)

Solid phases

2 49.8 29.5

Ice, CaCl2·6H2O CaCl2·6H2O, CaCl2·4H2Oa CaCl2·4H2Oa, CaCl2·2H2O CaCl2·2H2O, CaCl2·H2Oa CaCl2·H2O, CaCl2

45.3 175.5 260

Relative humidity

0.0157 6.74

— 0.223

11.99

0.164

842

0.124

3-4 atm

0.08

(F) The vapor pressure of CaCl2 solutions at 188C Concentration (wt.%) 43.21 42.05 37.43 34.94 34.26 32.86 32.20 28.98

Pressure (mm Hg)

Relative humidity

Concentration (wt.%)

Pressure (mm Hg)

Relative humidity

Concentration (wt.%)

Pressure (mm Hg)

Relative humidity

4.81 5.22 6.97 7.97 8.32 8.80 9.07 10.28

0.311 0.337 0.450 0.515 0.538 0.569 0.586 0.664

27.64 26.61 25.95 23.83 22.21 17.35 14.82 14.65

10.76 10.09 11.29 12.10 12.43 13.56 14.00 14.04

0.695 0.717 0.729 0.782 0.803 0.876 0.905 0.907

12.01 6.28 5.05 4.47 3.68 3.14 2.90 2.69

14.46 15.07 15.15 15.19 15.24 15.29 15.31 15.32

0.934 0.974 0.979 0.981 0.985 0.988 0.989 0.990

424

Part 2 Calcium Chloride Table 2.37 Surface Tensions of Pure Calcium Chloride Solutions Surface tension in dynes per centimetera

% CaCl2

at 108C (508F)

at 258C (778F)

0 1.10 2.70 5.26 9.99 18.17 24.98 30.75 35.69 43.75 55.42

74.22 — — 75.74 77.42 81.12 85.22 89.17 — — —

71.97 72.32 72.75 73.49 75.17 78.87 82.97 86.92 90.37 97.07 106.97

a

International Critical Tables vol. 4, p. 465, 1929.

REFERENCES Geology Abitz, R., Myers, J., Drez, P., and Deal, D. (1990). “Geochemistry of Salado Formation Brines Recovered from the Waste Isolation Pilot Plant (WIPP) Repository.” Proc. Symp. Waste Manag. (Waste Manag. ’90), Albuquerque, NM 2, 881 –882. Ahmad, N. S., Karunaratne, G. P., Chew, S. H., and Lee, S. L. (2000). “Bentonite–Kaolinite Mix for Barrier Systems”, Geotechnical Special Publ. 105, pp. 93 –104, Am. Soc. Civil Eng. Akagi, T., and Kono, Y. (1995). “Inhibiting Effects of Lanthanum Ions on Calcite Formation from CaCl2 – NaHCO3 Solutions at 258C.” Aquatic Geochem. 1(2), 231 –239. Alam, G. S., and Asrarullah, P. (1973). “Potash Deposits in a Salt Mine, Khewra, Jehlum District, Punjab, Pakistan.” Rec. Geol. Survey, Pakistan 21(Pt. 2), 1–14. Alsharari, M. A. (1999). “Reclamation of Fine-Textured Sodic Soil Using Gypsum, Langbeinite and Calcium Chloride.” PhD dissertation, Univ. of Arizona, 179 pp. Amira, J., et al., (1992). “From Hot to Cold Crystallization in Jordan.” Phos. Potash 182, 28–32. Anderson, W. (1945). “On the (Calcium) Chloride Waters of Great Britain.” Geol. Mag. 82(6), 267–273. Andrade, L. L. (1984). “Study of Trace Metals in Sergipe Evaporites.” Rev. Bras. Eng. Quim. 7(3– 4), 55–59. Anirudhan, T. S., Sivanandan, A. V., and George, O. (1991). “Adsorption and Desorption of Inorganic Phosphates by Sediments of Lale Veli, SW Coast of India.” J. Appl. Hydrol. 4(1–2), 17–28. Anson, R. W., and Hawkins, A. B. (1998). “The Effect of CaCl2 in Pore Water upon the Shear Strength of Kaolinite and Sodium Montmorillonite.” Geotechnique 48(6), 787–800. Aoubouazza, M., and Baudracco, J. (1992). “Influence of the Ionic Forces of CaCl2 Solutions on the Permeability of Clayey Sandstones.” Reun. Ann. Sci. Terre 14, 7 pp. Aren, B., and Depowski, S. (1965). “Calcium Chloride Waters with the Podlasie Basin Gas Field, Poland.” Kwartalnik Geol. 9(9), 17 –27.

References

425

Figure 2.79 Surface tensions of pure calcium chloride solutions. International Critical Tables (1929), vol. 4. p. 465.

426

Part 2 Calcium Chloride Table 2.38 Heat Evolved in Dissolving Anhydrous Calcium Chloride in Water Heat evolved

Water per mol (111 g) CaCl2

Moles 400 200 100 50 20 10 6 a

Gram calories

Approximate temperature rise calculated from previous columna

Btu

Grams

% CaCl2 in solution by weight

Per mol CaCl2

Per gram CaCl2

Per lb CaCl2

Per lb solution

8F

8C

7200 3600 1800 900 360 180 108

1.5 3.0 5.8 11.0 23.6 38.1 50.7

17,921 17,921 17,754 17,635 17,037 15,412 13,381

161.5 161.5 159.9 158.9 153.5 138.8 120.6

290.6 290.6 287.9 286.0 276.3 249.9 217.0

4.4 8.7 16.7 31.4 65.1 95.3 110.0

4.5 9 18 37 92 160 200

0.25 0.5 1 20.5 51 89 111

Assuming no heat loss. (Allied, 1980).

Table 2.39 Temperature Increase in Dissolving Regular Flake (77– 80%) Calcium Chloride in Water at 608F (Tested in Insulated Flask, Allied, 1980) Final solution Temperature rise (approximate)

Pounds regular flake CaCl2 Specific gravity at 60/608F 1.073 1.135 1.190 1.242 1.286 1.324 1.334

Per gallon of water

Per gallon of solution

%CaCl2

8F

8C

1 2 3 4 5 6 7

0.96 1.83 2.63 3.36 4.02 4.62 5.16

8.4 15.1 20.6 25.3 29.2 32.7 35.6

18 33 48 59 68 74 77

10 18 26 33 38 41 43

References

427

Figure 2.80 Theoretical temperature increase in dissolving Dow’s 77% CaCl2 Dowflake or 94% CaCl2 Peladow (Dow Chemical, 2000, courtesy of the Dow Chemical Company).

Artzi, Y., Vengosh, A., Adar, A., and Ayalon, A. (1996). “Sources of Salinization in the Coastal Aquifer.” Israel Geol. Soc., Ann. Mtg., Abs., 2 pp. Austin, G. S., and Barker, J. M. (1990). “Economic Geology of the Carlsbad Potash District.” Econ. Geol. Guidebook Series 8, 10 –14. Ayora, C., Cendori, D. I., Taberner, C., and Pueyo, J. J. (2001). “Brine–Mineral Reactions in Evaporite Basins.” Geology 29(3), 251 –254. Azizov, A. I. (1975). “Calcium Chloride Hydrosphere in the Earth, and Hydrothermal Ore Formation.” Geol. Rudn. Mestor., 42(2), 70 –73; (1974, 1972). “Supergene Potassium as a Factor in the Formation of Potash Deposits.” Acad. Sci. USSR, Dokl. Earth Sci. Section, 209(1–6), 212 –215 (April). Azmon, B. (1993). “Geochemistry of Groundwater in the Eocene Aquifer of Alonim Shefaram in the Period 1959– 1977.” Fourth Int. TNO-BMFT Conf. Contam. Soil 1, 475 pp. Baldina, A. L., and Sverdlov, Y. G. (1959). “Calcium Chloride Waters in the Perm Area, USSR.” Geochim Sbornik 13, 286 –294.

428

Table 2.40 Enthalpy of Pure Calcium Chloride Solutions—Btu per Pound (Allied, 1980). (Reference state ¼ 08C and 1 atm where heat content of liquid water equals zero; enthalpy of CaCl2 solution equals its heat of solution)

8C

% CaCl2

8F

0

5

10

12.5

15

17.5

20

22.5

25

27.5

29.8

32

30 25 20 15 10 5 0

86 77 68 59 50 41 32

54.1 45.1 36.1 27.1 18.1 9.1 0

64.7 56.3 47.9 39.6 31.3 23.0 14.7

75.3 67.5 59.7 51.9 44.2 36.5 28.8

80.5 73.0 65.4 58.0 50.5 43.1 35.7

85.7 78.5 71.2 64.1 56.9 49.8 42.7

90.8 83.7 76.9 69.9 63.0 56.2 49.3

96.0 89.2 82.6 75.9 69.2 62.6 56.0

101.4 94.9 88.4 81.9 75.4 68.9 62.5

106.6 100.2 93.9 87.5 81.2 75.0 68.7

111.2 105.0 98.9 92.7 86.6 80.6 74.5

115.3 109.2 103.2 97.2 91.3 85.5 79.6

118.8 112.9 107.1 101.4 95.6 89.9 84.2

25

23

2 146.0

2 59.0

21.1

28.3

35.6

42.5

49.4

56.1

62.5

68.5

73.8

78.6

2 10

14

2 148.5

2 91.4

2 30.3

0.1

28.6

35.7

42.9

49.7

56.3

62.5

68.0

72.9

2 15

5

2 150.9

2 105.4

2 54.0

2 28.3

2 2.6

23.1

36.4

43.4

50.1

56.5

62.2

67.3

2 20

24

2 153.2

2 114.9

2 68.8

2 45.7

2 22.6

0.4

23.5

37.1

44.0

50.6

56.4

61.7

2 25

2 13

2 155.6

2 122.6

2 80.1

2 58.8

2 37.5

2 16.2

5.1

27.1

37.9

44.6

50.5

56.1

2 30 2 35

2 23 2 31

2 157.8 2 160.1

2 129.3 2 135.0

2 89.4 2 96.9

2 69.4 2 77.8

2 49.4 2 57.8

2 29.4 2 39.6

2 9.4 2 20.6

10.6 2 1.2

30.5 18.9

38.7 32.8

44.9 39.2

50.5 45.0

2 40 2 45 2 50 2 55

2 40 2 49 2 58 2 67

2 162.2 2 164.4 2 166.5 2 168.6

2 140.2 2 145.1 2 149.7 2 154.1

2 103.3 2 109.1 2 114.5 2 119.7

2 87.0 2 91.1 2 97.0 2 102.5

2 66.3 2 73.1 2 79.4 2 85.3

2 47.8 2 55.2 2 61.8 2 68.1

2 29.4 2 37.1 2 44.3 2 50.9

2 10.9 2 19.1 2 26.7 2 33.7

7.6 2 1.2 2 9.1 2 21.5

25.5 16.9 8.5 0.7

33.5 27.8 22.2 16.5

34.2 25.0 15.8 6.5

Below 2558C, solid mixture of ice and CaCl2·6H2O.

Part 2 Calcium Chloride

Temperature

429

References

Figure 2.81 The amount of water adsorbed by commercial calcium chloride at 258C, lbs/lb of product (Dow, 2001, courtesy of the Dow Chemical Company).

Table 2.41 Water Absorbed by Solid Calcium Chloride at Various Relative Humidities (258C)a Final Solution Relative humidity (%)

100 95 90 85 80 75 70 65 60 55 50 45

kg water absorbed/kg CaCl2

Concentration

% CaCl2

1 0 5.2 10.4 14.8 19.1 22.6 25.6 28.3 31.1 33.8 36.0 37.8

2 0 8.5 14.2 18.3 22.0 25.0 27.5 29.8 32.0 34.0 36.0 37.9

Flakeb 1 0 14.0 6.5 4.3 3.1 2.5 2.1 1.8 1.5 1.3 1.2 1.1

3 8.4 5.0 3.5 2.8 2.0 1.6

Anhydrousc 2 0 8.2 4.5 3.3 2.5 2.1 1.8 1.6 1.4 1.3 1.2 1.1

1 0 17.3 8.2 5.4 4.0 3.2 2.7 2.4 2.1 1.8 1.6 1.5

2 0 10.2 5.7 4.2 3.3 2.8 2.5 2.2 2.0 1.8 1.6 1.5

(continues)

430

Part 2 Calcium Chloride Table 2.41 (continued ) Final Solution

Relative humidity (%)

kg water absorbed/kg CaCl2

Concentration

% CaCl2

39.5 — 41.7 43.9

39.9 — 42.1 44.5

40 36 35 30

Flakeb 1.0 — 0.9 0.8

1.0

Anhydrousc 1.0 — 0.9 0.8

1.4 — 1.3 1.2

1.4 — 1.3 1.1

B. Deliquescence (lowest relative humidity and temperature at which the exposed piece of anhydrous CaCl2 will take up enough water to dissolve) Relative humidity, % Temperature, 8C 8F

20 37.8 100

30 23.3 74

40 6.7 44

43 5.6 42

1 and 2 are estimates from two producers of calcium chloride, Dow and Allied Chemical (1980), respectively. 3 is from Tetra (2002). a See Figs. 2.61 and 2.81 for similar data. b 77–80% CaCl2, average ,79.3%.

Table 2.42 Typical Densities and Freezing Points of Commercial Calcium Chloride Solutions (Dow 2001) (Courtesy of the Dow Chemical Company)

% CaCl2 0 5 6 7 8 9 10 11 12 13 14 15 16 17 18

Approximate specific gravity at 258C r 25 25

Weight (kg/liter) at 258C r 25 4

Liters per metric ton of sol. at 258C

DOWFLAKE Equiv. kg per liter of sol. at 258C

DOWFLAKE Equiv. lbs per gal of sol. at 778F

Freezing points (8C)

Freezing points (8F)

1.000 1.047 1.056 1.065 1.074 1.083 1.090 1.100 1.110 1.120 1.129 1.139 1.149 1.159 1.169

0.997 1.045 1.055 1.063 1.073 1.081 1.087 1.097 1.107 1.117 1.126 1.136 1.146 1.156 1.165

— 957 948 941 932 925 920 912 903 895 888 880 873 865 858

— 0.068 0.083 0.098 0.114 0.131 0.139 0.155 0.170 0.186 0.202 0.218 0.235 0.252 0.269

— 0.57 0.69 0.82 0.95 1.09 1.16 1.29 1.42 1.55 1.68 1.82 1.96 2.10 2.24

þ0 23 24 24 25 26 27 28 29 2 10 2 11 2 12 2 13 2 15 2 17

þ32 þ27 þ25 þ24 þ23 þ21 þ20 þ18 þ16 þ14 þ12 þ10 þ8 þ5 þ2 (continues)

References

431

DOWFLAKE DOWFLAKE Equiv. lbs Equiv. kg per gal of Freezing per liter of sol. at 778F points (8C) sol. at 258C

Freezing points (8F)

Table 2.42 (continued )

% CaCl2

Approximate specific gravity at 258C r 25 25

Weight (kg/liter) at 258C r 25 4

Liters per metric ton of sol. at 258C

1.179 1.189 1.199 1.209 1.219 1.228 1.240 1.251 1.263 1.275 1.287 1.294 1.298 1.310 1.322 1.334 1.345 1.357 1.369 1.381 1.392 1.404 1.416 1.428 1.439 1.474

1.175 1.185 1.195 1.205 1.215 1.224 1.236 1.247 1.259 1.271 1.283 1.290 1.294 1.306 1.318 1.330 1.341 1.353 1.365 1.377 1.388 1.400 1.412 1.424 1.435 1.470

851 844 837 830 823 817 809 802 794 787 779 775 773 766 759 752 746 739 733 726 720 714 708 702 697 680

19 20 21 22 23 24 25 26 27 28 29 29.6 30 31 32 33 34 35 36 37 38 39 40 41 42 45

0.286 0.304 0.322 0.340 0.358 0.377 0.396 0.416 0.436 0.456 0.477 0.490 0.498 0.519 0.541 0.563 0.585 0.607 0.630 0.653 0.676 0.700 0.724 0.748 0.773 0.848

2 18 2 20 2 22 2 24 2 27 2 29 2 32 2 35 2 39 2 43 2 47 2 51 2 47 2 37 2 27 2 20 2 12 27 21 þ4 þ9 þ13 þ16 þ18 þ21 þ26

2.39 2.53 2.68 2.83 2.99 3.14 3.30 3.47 3.63 3.81 3.97 4.08 4.15 4.33 4.51 4.69 4.87 5.06 5.25 5.45 5.64 5.84 6.04 6.24 6.44 7.07

21 24 28 2 12 2 16 2 20 2 25 2 31 2 38 2 46 2 53 2 60 2 52 2 34 2 17 24 þ10 þ20 þ30 þ39 þ48 þ55 þ61 þ65 þ69 þ78

Table 2.43 An Example of the pH of Calcium Chloride Solutions after the Addition of Sodium Hydroxide (Allied, 1980) CaCl2 (%) pH

0.25 7.1

0.5 8.5

1.0 9.4

1.5 10.0

3.0 10.3

4.0 10.3

5.9 10.3

10.3 10.3

15.6 10.1

20.5 9.9

25.1 9.7

30.5 9.3

40.0 8.6

432

Part 2 Calcium Chloride Table 2.44 Miscellaneous Properties of Calcium Chloride and Its Hydrates (Dow Chemical 1980, 2001)a

Property

CaCl2·6H2O

CaCl2·4H2O

Percent CaCl2 Molecular wt.b Melting point (8C) Boiling point (8C) Density (258C)

50.660 219.075 29.92c – 29.9 — 1.71

60.632 75.492 183.0445 147.014 45.13c – 45.3 175.5c –176 — 174 –175c 1.83 1.85

Heat of fusion (cal/g) Heat of solution, 1 dil. (cal/g) Heat of formation (kcal/mol) Free energy, 258C (kcal/mol) Heat capacity, 258C (cal/g/8C)

Coefficient of expansion Refractive findex, 208C

(1.7182 at 48, 1.68 at 178C)c 47.3c –50 40.7c 17.2 2 623.3

— 39– 40c

CaCl2·2H2O

— 21

— — 2 14.2 (0.1c) 2 71.6c,d – 72.8

CaCl2·H2O

CaCl2

86.035 128.999 187,260c 181c –183 2.24

100.00 110.983 772,782c 1935 2.16 (2.152 at 158C, 2.155 at 208C)c

— 32

— 55.0c, 54.2c 2 96.8 (0.6c) 2 176.2 (0.3c) 2 190.1 (0.6c)

—

2 480.3 (0.17, 0.4)c —

0.34 –0.40c

0.32

0.28

0.20

0.32 at 08Cc

—

—

—

—

—

—

—

—

—

—

—

265.49 (0.06c) —

61.4c –61.5

—

335.58 (0.23, 0.25, 0.31)c 179.8, 195.36c 0.16 (0.164 at 618Cc,e) 0.00062c 1.52c

a

Negative sign means that heat is evolved. The molecule CaCl2·0.33H2O has also been claimed with a heat of solution to infinite dilution of 217.06 kcal/mol. b 1995 revised atomic weights. c Other sources (Allied Chemical, 1980, Tetra Chemicals, 1992, Sinke et al., 1985, etc.) d 188C; heat of solution in 400 mol water. e Heat capacity, cal/g mol/8C ¼ 16.9 þ 0.003868K (claimed for 0–7008C). Multiply by 4.184 to convert calories to joules.

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454

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Index

Air conditioning; lithium’s use in 99, 182, 197, 198, 201 Air drying with CaCl2 374 Albite’s conversion to chlorite 266, 267 Alcohols as boron or lithium solvents 115, 116, 126–129, 136– 140 Al Hayat sector, Bikita petalite deposit 61 – 63, 65, 156, 157 Alkali Lake (near Bristol Lake), California 284, 285 Alkali metals’ properties 211– 214 Allan potash mine’s CaCl2 brine, Saskatchewan, Canada 245, 246 Altai Mountains lithium pegmatites, China, Russia 92, 95, 96 Altmark gasfield’s dolomitization brine, Germany 45, 265 Alto Ligonha’s lithium pegmatites, Mozambique 2, 91 Alumina, aluminum hydroxide adsorption of lithium 5, 132, 140– 144 Aluminum; lithium’s use in production, alloys 99, 182, 188– 190, 196 Amblygonite characteristics, deposits 4, 53, 58, 62, 67, 70, 73, 77, 78, 90 – 92, 94, 150, 154, 155, 168, 171 Amboy, Amboy Crater, Bristol Lake, California 285, 352 American Lithium Chemicals, San Antonio, Texas 99, 165 American Potash and Chemical Co., Trona, California 58, 99, 133, 165, 203

Abitibi greenstone belt lithium pegmatites, Quebec, Canada 89 Absorbant, adsorbent; lithium’s use as an 83, 99, 176, 182, 197, 198, 200, 201 Absorption of moisture by CaCl2 346, 365, 370– 372, 374– 378, 432, 433 Abundance of Li in the earth’s crust 1, 97, 98 Acid leach processes, lithium ores 59, 171, 172 Acid roast process, spodumene ore 161– 164, 169– 172 Activated carbon, use with lithium solutions 161, 170, 175 Activity coefficients, lithium, CaCl2 solutions 212, 420, 427, 430 Adsorption, reaction; lithium on clays, rocks 3, 10, 18, 21, 34, 44 Aerosols, sea water 306, 309, 310 Afar Depression’s thermal CaCl2 springs, Ethiopia 276, 277, 291– 298, 321 Afghanistan’s Hindukush Mountain lithium pegmatites 2, 56 – 57, 97 African Rift Valley 322 Africa’s lithium pegmatites (excluding Bikita) 2, 3, 91 – 92, 169, 204 Age of calcium chloride lakes 281, 285, 287, 291, 299 Age of lithium brine deposits 8, 16 – 18, 27, 28, 33 Age of lithium pegmatites 53, 60, 61, 68, 79, 83, 85, 90, 92 Agricultural uses of CaCl2 376– 378

459

460

Index

Amesite and chlorite’s conversion by CaCl2 268 Ammonium chloride – lepidolite pressure leach process 167 Analacite’s conversion by CaCl2 to hydrous glass 268 Analyses, CaCl2 products 384, 385, 430 Analyses, lithium concentrates, products 170, 184, 187, 210 Angara-Lena Basin’s high-lithium, CaCl2 brine, Russia 45, 142, 259, 318 Anhydrite’s solubility in CaCl2 335 Anorthite’s conversion to albite, epidote, etc. 265, 267, 268 Anorthite’s solubility in CaCl2 335 Antarctica 294, 295, 298– 311, 322, 383 Antarcticite 237, 298, 300, 305, 311, 326 Antimony– gold minerals with CaCl2 occlusions 332 Antofagasta, Chile 11, 109, 112, 114, 126, 132 Apatite in lithium pegmatites 60, 73, 81, 96 Apatite with CaCl2 occlusions 333 Apatite’s reaction with CaCl2 335 Appalachian oil or gasfield CaCl2 brines, groundwater 260, 265, 318 Aracuai – Itinga lithium pegmatites, Brazil 2, 90 Aragonite’s reaction with CaCl2 335 Archambault – Oliver’s spodumene CO2 – pressure leach process 166 Arcs: Aleutian, Central American, Kurile, Sunda; lithium in 45, 46 Argentina’s lithium pegmatites 2, 91 Arkansas oil or gas field CaCl2 brines 2, 45, 47, 136, 142, 262, 264, 267 Arquena de Minerios e Metais Ltd, Brazil 91, 168 Asal, CaCl2 Lake, thermal springs, Ethiopia 270, 271, 291– 295, 297, 298, 322 Australia’s lithium pegmatites 59 – 61 Australia’s Queensland CaCl2 groundwater 314 Austria’s Koralpe spodumene deposit 2, 3, 92 Avalon Ventures, Ontario, Canada 53, 87

d11B 261, 264 Bagdad Crater (Cinder Cone; Amboy Crater), Bristol Lake 285 Ballasting tires with CaCl2 brine 272 Bangkog Cuo Lake, China, high-lithium brine 38, 39, 41 Barite with CaCl2 occlusions 332 Barroso – Alvao and Guarda lithium pegmatites, Portugal 2, 3, 94, 95 Basalt; leaching, adsorbing lithium; forming CaCl2 by sea water 238, 268, 269, 273, 290, 311, 335 Base exchange processes for lithium ores 161– 164, 166, 169– 171, 173, 174 Batteries, lithium’s use in 82, 99, 182, 190– 195 Beecher Lode spodumene pegmatite, South Dakota 84, 85, 166 Belgium’s Variscan Foreland CaCl2 groundwater 315 Bentonite, effect of CaCl2 on 336, 337 Bentonite, leaching lithium from 172 Berea Formation CaCl2 brine, Michigan Basin 252– 254 Bernic Lake lithium pegmatite, Manitoba, Canada 2, 3, 55, 56, 68 –78, 150– 155 Beryl recovery from lithium pegmatites 53, 61, 62, 83, 90, 155 Bessemer City, North Carolina 2, 81, 99, 163, 164 Best Bet spodumene mine, Northwest Territories, Canada 90 Big Mack petalite deposit, Ontario, Canada 54, 85 – 89 Big Whopper petalite deposit, Ontario, Canada 53, 56, 85– 88, 171 Bihar, Rajasthan’s lepidolite pegmatite, India 56, 94 Bikita, Zimbabwe lithium pegmatite 2, 3, 52, 53, 56, 57, 61 – 68, 155– 157, 165, 204 Bikitaite 4, 62, 87 Biotite, effect of CaCl2 solutions on 336 Bischoffite from high-lithium brines 110, 113, 124, 125

Index

Bischoffite deposits 321, 320, 327– 329 Bismuth minerals with CaCl2 occlusions 334 Black Hills lithium pegmatite district, South Dakota 56, 82 – 85, 166– 169, 204 Blasting lithium ore 146, 150, 155, 158 Bleaches, sanitizers, lithium hypochlorite’s use in 199 Bob Ingersol lepidolite mine, South Dakota 83, 84 Boiling points of CaCl2 solutions 339, 350, 353– 356, 424, 425, 430 Bolivia’s high-lithium salars and springs 27 – 29 Bolivia’s lithium pegmatites 91 Bonneville Salt Flats high-lithium end liquor, Utah 28, 31, 32, 36, 37, 45 Border zones in lithium pegmatites 52, 65, 66, 71 Boric acid; lithium co- or by-product 110, 112, 118, 119, 124, 126 Boron in calcium chloride brines, deposits 264, 265, 329 Boron in lithium deposits 10, 18, 20, 21, 25, 27 – 29, 38, 52 Boron extraction from high-lithium brines 115, 116, 126– 130 Bougouni area’s lithium pegmatites, Mali 2, 92 Brackish water containing CaCl2 287, 346, 348, 350 Brasileira de Litio 91, 168, 203 Brazil’s lithium pegmatites 2, 3, 56, 90, 91, 168, 204, 206 Briner’s grade CaCl2 344, 373, 385 Briquetting lithium carbonate 108, 116, 129 Bristol Lake, California 33, 45, 283– 290, 311, 341, 345– 352 Bromine in evaporating sea water, CaCl2 brines 239, 241, 246, 252, 253, 260, 264, 274, 281, 321 Bromine in tachyhydrite deposits 321, 329 Bromine recovery from CaCl2 brines 45, 264, 338– 341 Buck and Pegli lithium pegmatite, Manitoba, Canada 89

461

Burial depth of lithium pegmatites as they formed 50, 52, 61, 83 Butyl (or n-Butyl) lithium 99, 131, 177, 180, 183, 197

d13C 261, 265, 315 Cadiz Lake, California 33, 45, 283, 284, 286, 288, 290, 291, 345, 351 Calcite with CaCl2 occlusions 333 Calcite reactions to form CaCl2 259, 334 Calcite’s solubility in CaCl2 334, 335 Calcium chloride groundwater 256, 257, 272, 279, 287, 311– 318, 346, 347 Calcium Chloride Group, Bristol Lake, California 350 Calcium chloride in mineral occlusions 238, 244, 259, 311, 321, 332– 334 Calcium chloride lakes 345–352 Calcium chloride precipitation of sulfate from groundwater 113, 239, 249, 264, 265, 293, 321, 349 Calcium chloride product specifications 288, 384, 385 Calcium chloride’s reactions with minerals 335– 337 California’s oil or gas field CaCl2 brines 260, 262 California Salt Co., Bristol Lake, California 350 Canada’s calcium chloride brines 244– 246, 251, 256, 314, 345 Canada’s lithium pegmatites 2, 3, 55, 56, 68 – 78, 85 – 90, 150– 155, 165, 166, 169– 171, 205, 206 Canadian Shield’s dolomitization brine 256, 314 Canals for CaCl2 brine 246, 347 Carbon dioxide with lithium or CaCl2 deposits 49, 83, 96, 311 Carbon dioxide– pressure leach process for spodumene 166 Cargil Co., Bristol Lake, California 351 Carlsbad, New Mexico potash deposits, brine 239– 241, 244, 259 Carnallite decomposition brine 237, 238 Carnallite deposits 237– 239, 320– 331

462

Index

Carnallite from high-lithium brines 110, 113, 124 Carpathian potash deposit’s dolomitization brine, Russia 136, 258 Cascade Range’s thermal CaCl2 springs, Washington 276 Caspian Depression’s dolomitization brine, Russia 258 Cassiterite in lithium pegmatites 53, 74, 76, 79, 83, 90, 91, 148, 155, 171 Catal Lagoon, Salar de Hombre Muerto, Argentina 22, 24,25 Catalyst, lithium’s use as a 99, 197, 199 Catamarca Province’s lithium pegmatites, Argentina 2, 91 Cat Lake-Winnipeg River lithium pegmatite area, Manitoba, Canada 68, 85, 89 Caucasus region’s CaCl2 groundwater, Russia 312, 313, 318 Caustic soda by-product, lithium ore limeroasting process 159, 160, 165, 171 Celestite production at Bristol Lake 287, 289 Cementing well casings with CaCl2 cement 369 Ceramics, lithium’s use in 99, 182, 185, 187, 188 Cerro Prieto geothermal, solar pond brine, Mexico 14, 15, 37, 274, 275, 279 Cesano’s medium-lithium geothermal brine, Italy 42, 44, 141 Cesium or pollucite recovery from lithium pegmatites 69, 83, 92, 150, 152, 154, 155, 165, 168 Chabazite’s solubility in CaCl2 335 Cheleken, Turkmenia CaCl2 hot springs 277 Chemetall GmbH 75, 99, 101, 111, 112, 114, 115, 157, 160, 174, 201, 203 Chemicals from CaCl2 374, 377 Chemicals, lithium 174– 179 Chepica de1 Salar, Salar de Atacama, Chile 111, 112, 114 Cherryville spodumene mine, North Carolina 163

Chile’s high-lithium playas, salars 10 – 21, 108– 130 Chilled rolls for flaked CaCl2 350, 352 China’s high-lithium brines 38 – 43, 135 China’s lithium pegmatites 2, 3, 56, 92, 93, 203– 205 Chita Oblast’s lithium pegmatites, Urals, Russia 94 Cia Brasileira de Litio 91, 203, 204 Clay, high-lithium 97, 98, 172, 173 Clay, impervious for solar ponds 100, 101, 106, 110, 111, 337 Clay minerals, effect of CaCl2 on 336, 337, 369 Clayton Valley, Nevada 2, 3, 5 – 10, 21, 99, 101– 108, 201, 203 Clinoptillite, effect of CaCl2 solutions on 335 Columbite in lithium pegmatites 79, 83, 89, 91 Coagulation of sediments, products with CaCl2 337, 369, 373– 376, 378 Coastal CaCl2 aquifers 238, 247, 273, 276, 312, 313, 315– 317 Coastal sabkhas, brine pools’ dolomitization brine 238, 247, 258 Commercial CaCl2’s properties 289, 384, 385 Companhia Brasileria do Litio 91, 168, 203 Concrete, CaCl2’s use in 359, 360, 368, 369, 377 Concrete’s corrosion by CaCl2, 362, 364, 365, 368, 369 Congo’s tachyhydrite deposit 318– 320, 327– 329 Consumption statistics for CaCl2, U.S. 381 Copper minerals with CaCl2 occlusions 333, 334 Cordillera de Domeyko, Salar de Atacama, Chile 10, 13, 17, 20 Cordillera de la Sal, Salar de Atacama, Chile 10, 13, 16, 17, 20, 21 Core zones in lithium pegmatites 52, 65 CORFO, Corporacion de Fomento de la Production, Chile 108, 110, 112, 118

Index

Corrosion inhibitors for CaCl2 brines or solids 346, 354, 358, 362– 365, 384 Cory potash mine’s CaCl2 brine, Saskatoon, Canada 245, 246, 345 Costa Rica Rift vents 44, 46 Cost of CaCl2 381, 382 Covas de Barroso spodumene pegmatites, Portugal 56, 57, 94, 95 Crimean Steppe’s CaCl2 groundwater, Russia 318 Cronembourg’s high-lithium geothermal brine, France 42, 44, 141 Cyprus Foote Minerals 75, 102, 201 Cyprus Ophiolite Formation’s CaCl2 groundwater 315 Czechoslovkia’s lithium pegmatites 2, 97

dD 15, 37, 245, 246, 251, 253, 256, 280, 299, 301– 303, 331 Dabuxun Saline Lake, China 135, 279 Dale Lake (near Bristol Lake), California 284, 289 Dalol Salt Pan, Ethiopia 291– 295 Danakil Depression, Ethiopia 273, 276, 277, 291– 298 Danby Lake, California 33, 283, 284, 289, 291 Da Qaidam Lake, China 38, 40, 41, 135, 279 Dead Sea, Israel, Jordan 2, 31, 32, 37, 38, 45, 142, 143, 145, 278, 281– 283, 312, 313 Dead Sea’s thermal CaCl2 springs 270, 271, 278 Deep ocean vent brine’s d6Li, d7Li 46 Deicing with CaCl2 358– 365 Delta Chemical Co., Cadiz Lake, California 352 Density of CaCl2 solutions 308, 346, 347, 383, 393, 408– 411, 430, 432–434 Density of CaCl2 – NaCl solutions 339, 412– 418 Density of lithium, its minerals and solutions 43, 54, 58, 59, 167, 211, 214, 217, 218 Density vs temperature, Salar de Atacama brine 130

463

Depth of lithium pegmatites when formed 50 – 52, 61, 83 Desert Properties Co., Bristol Lake, California 350 Detroit River Formation’s CaCl2 brine, Michigan 248, 253– 255, 342 Deuterium 37, 245, 246, 251, 253, 256, 280, 299, 301–303, 331 Disk granulators for CaCl2 356 Dnieper– Donets Basin’s CaCl2 brine, Russia 258 Dolerite, fractured, Antarctica 305– 310 Dolime, slaked (magnesium precipitant) 341 Dolomitization reaction, brines 37, 45, 238– 264, 266, 269, 272, 273, 289, 310, 311 Don Juan Pond, Antarctica 45, 283, 294, 295, 298, 300– 303, 305– 311, 322, 333 Double Lakes area CaCl2 brine, Texas 312, 314 Dow Chemical Company 338– 345, 383 Dredging lithium solar ponds 104, 106 Drilling mud, CaCl2’s use in 269 Drugs, lithium’s use in 180, 199 Drum granulator for CaCl2 356 Dryers for CaCl2 solids 350, 353– 357 Drying agent, CaCl2 as a 370– 378 Dundee Formation’s CaCl2 brine, Michigan Basin 252–254, 342 Dust control with CaCl2 365– 367 Dust reduction in solid CaCl2 products 355, 356 East Pacific Rise vents 42, 269, 270, 272 Edison spodumene mine, South Dakota 84, 166 Electric conductance, CaCl2 solutions 420, 427 El Tatio geyser field, Chile 10, 14, 17, 18, 20, 44 Emerald Fields Resources Corp., Canada 54, 87 England’s CaCl2 groundwater 45, 256, 277, 312, 315 England’s oil or gas field CaCl2 brines 260

464

Index

English River lithium pegmatites, Ontario, Canada 85 – 87 English Zeichstein Formation 45, 256, 277, 315 Enthalpy, CaCl2 solutions 427, 431 Entrained brine in solar pond salts l00– 102, 120, 124, 125 Environmental factors, deicing CaCl2 361– 365 Environmental uses of CaCl2 375 Epidote conversion to chlorite 267 Etta spodumene mine, Black Hills, South Dakota 55, 83, 84 Eucryptite 4, 50, 51, 53, 56, 57, 59, 62, 67, 69, 73, 78, 87, 94, 157 Europe’s lithium pegmatites 2, 92, 93 Evaporation plants for CaCl2 brine 339– 341, 350– 353, 355, 357 Evaporation plants for lithium brines 5, 34, 36, 133, 134, 146, 159– 162, 165, 167, 170–172 Evaporation rates at CaCl2 lakes 287, 306, 321, 322, 347 Evaporation rates at high-lithium playas 8, 25, 29, 100, 102, 108, 110, 113 Fault zones with CaCl2 occlusions 333 Feldspar conversion to albite 264, 266– 268, 310 Feldspar recovery from lithium pegmatites 53, 83, 149, 153, 157, 167, 171 Ferrar dolerite, Antarctica 305– 310 Ferric chloride –lithium chloride complex 36, 139, 160 Filer Formation CaCl2 brine, Michigan Basin 249, 251, 252, 339, 341, 342 Finland’s lithium pegmatites 92 Fire clay, high-lithium, alumina 97 Flaked calcium chloride 343, 344, 350, 352– 354, 384, 385 Flint clay, high-lithium 98 Flocculation of clay with CaCl2 336, 337, 365, 367, 369, 378 Flotation of lithium ores 53, 83, 91, 121, 133, 134, 148, 149, 157– 159, 163, 165, 167–171

Fluidized bed dryer-granulator for CaCl2 355 Fluorine-containing lithium pegmatites 92, 96, 167 Fluorite with CaCl2 occlusions 332, 333 Fluorite solubility in CaCl2 335 FMC Corp. 81, 99, 131, 160, 164, 201, 203 Food processing with CaCl2 344, 373, 374, 385 Foote Minerals Co. 34, 75, 79, 99, 101, 108, 112, 160, 163, 164, 201, 203 Formation pressure, temperature, lithium pegmatites 50 – 52, 61, 83, 96 Fractional crystallization, high-lithium magma 49 – 53 France’s lithium pegmatites 2, 92 Freeze – thaw deterioration of concrete by CaCl2 364, 365 Freezing points of CaCl2 solutions 358, 361, 383, 385, 386, 388– 398, 402– 406, 410, 434 French medium-lithium geothermal brines 42, 44, 141 Fresh water fossils, CaCl2 deposits 291, 320, 327, 328 Fresno, California oil or gas field CaCl2 brines 260 Galapagos Spreading Center vents 270, 272 Garnet as a lithium by-product 171 Garnet with CaCl2 occlusions 335 Gas drying with CaCl2 370– 372, 374, 375 General Chemicals Group 344, 383 German CaCl2 groundwater, thermal springs 256, 257, 277, 279 German oil or gas field CaCl2 brines 45, 260, 265 German Zechstein Formation potash deposits 256, 257 Glaserite from high-lithium brines 104 Glass inclusions in Olivine’s d7Li 46 Glass, lithium’s use in 99, 181– 185 Glauberite, formed with CaCl2 239, 255, 290

Index

Glazes and enamels, lithium’s use in 185, 187, 188 Gold deposits with CaCl2 occlusions 333 Golmud River (with CaCl2), Qarhan Salt Lake, China 279, 298 Goubet el Kharab, Ethiopia 292, 297 Govindpal area’s lithium pegmatites, India 2, 94 Granite with CaCl2 occlusions 333 Granular CaCl2 343, 355– 357, 384, 385 Granular lithium carbonate 108, 116, 129 Grease, lithium’s use in 83, 99, 182, 196 Great Basin lakes’ d7Li 46 Great Salt Lake’s CaCl2 thermal springs, Utah 271, 276 Great Salt Lake’s high-lithium brine, Utah 2, 21, 31, 32, 35 – 37, 45, 46, 112, 137, 139, 140, 146 Greece’s Filiatra CaCl2 groundwater 315 Greenbushes’ spodumene pegmatite, Australia 59 – 61, 146– 150, 203 Greenstone with CaCl2 occlusions 333 Groundwater with calcium chloride 256, 257, 272, 279, 287, 311– 318, 346, 347 Groundwater with medium-to-high lithium 45, 142, 259, 312, 313, 318 Guarda area’s lithium pegmatites, Portugal 2, 3, 56, 94, 95, 168, 169, 204 Guemes, Salta province, Argentina 132 Guilietti’s Lake, thermal springs, Ethiopia 271, 291– 296 Gulf of California rift-fault system 280 Gulf of Mexico oil or gas field CaCl2 brines 260, 261, 264 Gulf of Tadjoura, Ethiopia 291, 292, 297 Guogaling Cuo Lake, China 39 Gwalia Consolidated Ltd. 59, 146, 203 Gypsum with CaCl2 or lithium lakes, brine 287, 290, 298, 299 Gypsum solubility in CaCl2 brine 353 Halite with CaCl2 occlusions 244, 321, 332, 333 Hallman-Beam spodumene mine, North Carolina 81

465

Hand sorting lithium ores 53, 66, 83, 155– 157 Harding and Pidlite lithium mine, New Mexico 55, 85 Hardness of lithium minerals 54, 58, 59, 150, 158 Harvesting salts 100, 104, 107, 113, 120, 121, 123–125, 347, 349 Harz Mountains CaCl2 springs, Germany 277 Hatchobaru and Kyshu’s geothermal brine, Japan 42 Heat evolved in dissolving solid CaCl2 358, 362, 427, 429, 430 Heat of dilution, strong CaCl2 solutions 358, 427, 429, 430 Heavy media separation 152, 153, 157, 158, 163, 166, 167 Hectorite 4, 10, 97, 98 Heletz – Kokhav oilfield dolomitization brine, Israel 265 Hiddenite 54 High-CaCl2 groundwater 311– 318 High-CaCl2 lakes 281– 311, 345–352 High-intensity magnetic separators 148, 149, 153, 154 High-iron spodumene 54, 55, 75, 159, 160 High-lithium clay 97, 98, 172, 173 High-lithium groundwater 45, 142, 259, 318 High-sulfate Salar de Atacama brine 20, 21, 114, 118 Hill Brothers Chemical Co. 350, 352, 383 Hindukush Mountain Range’s lithium pegmatites, Afganistan 2, 56, 57, 97 History of the lithium industry 98, 99 Holding ponds for CaCl2 brine 347, 350 Homogenization temperature of occlusions 332– 334 Humates (organics) in brine, lithium salts 134, 135 Humectant, CaCl2’s use as a 378 Humidity at CaCl2 lakes 287, 306, 321, 346, 352, 422 Humidity at high-lithium lakes 100, 108, 113

466

Index

Hydraulic seal for solar ponds 100 Hydrochloric acid’s use to form CaCl2 311, 357 Hydrogen bombs, 6Li conversion to tritium 99, 165 Hydrogen sulfide 42, 259, 271, 299, 303, 304, 311 Iceland’s CaCl2 thermal springs 270, 272, 321 Illinois Basin dolomitization brine 259 Inclusions containing CaCl2 238, 244, 259, 311, 321, 332– 334 Inclusions in lithium pegmatites 52, 83, 96 Index of refraction, CaCl2 solutions 410 India’s Ennur CaCl2 groundwater 315, 317 India’s lithium pegmatites 2, 56, 94 Indonesian thermal CaCl2 vents 269 Industrial uses of CaCl2 359, 374, 375 Industrial uses of lithium 196, 197, 200 Intermediate zones in lithium pegmatites 52, 65, 66, 71 Intermetmin Ltd., Namibia 169 Iodine in brines 239, 252, 253, 260, 264, 272 Iodine recovery 264, 339 Ion exchange extraction of lithium 138, 141, 172 Ion exchange of soils with CaCl2 336, 337 Ion exchange to form CaCl2 238, 268, 290, 311, 314, 315, 318 Ionization constants of CaCl2 solutions 427 Irkutsk’s potash dolomitization brine, Russia 242, 259 Ishibu-Iwachi CaCl2 hot springs, Japan 276, 277 Isotopes of lithium 34, 44 – 46, 98, 99, 165, 200, 256, 265 Israel’s CaCl2 groundwater 312, 316, 317 Israel’s oil or gas field CaCl2 brines 46, 260, 265, 312, 313, 317 Israel’s thermal CaCl2 springs 270, 277, 278, 282, 283, 317 Italy’s high-lithium geothermal brines 42, 44, 141 Japan’s thermal CaCl2 springs 276, 277, 321

Jordan Rift Valley 381 Jordan River 278, 282, 283 Juan de Fuca Ridge vents 42 Kainite from high-lithium brine 110 Kama River area’s dolomitization brine, Russia 259 Kansas oil or gas field CaCl2 brines 262 Kaolin in weathered lithium pegmatites 54, 59, 91 Kaolinite conversion to illite 255 Kaolinite, effect of CaCl2 solutions on 337 Karibib-Omaruru districts’ lithium pegmatites, Namibia 2, 91, 169 Kazakhstan’s CaCl2 lakes 294, 295, 298 Kentucky’s oil or gas field CaCl2 brines 262 Kerosene drying with CaCl2 370, 373 Keystone district lithium pegmatites, Black Hills, South Dakota 82 K-Feldspar recovery from lithium ores 53, 83, 149, 153, 167, 171, 177 Kiao Qaidam Lake, China 40 Kings Mountain lithium pegmatites, North Carolina 2, 3, 55, 79, 81, 99, 159, 160, 163 Kirgiztan’s lithium pegmatites 97 Kokhav oil field, Israel 265 Koktokay lithium pegmatite, Xinjiang – Uygur province, China 2, 92 Kola Peninsula’s lithium pegmatites, Russia 2, 96 Komuni CaCl2 hot springs, Japan 276 Koralpe lithium pegmatite, Austria 2, 92 Kunzite 54, 82

d6Li, d7Li 34, 44 – 46, 98, 99, 165, 200, 256, 260, 265 Laguo Co Lake, China 38, 41 Lake Adijdata, Ethiopia 47 Lake Asal, Ethiopia 270, 271, 291– 295, 297, 298, 322 Lake Bonney, Antarctica 299, 300, 302, 309, 310 Lake Giulietti, Ethiopia 271, 291– 296, 322 Lakes, high-CaCl2 281–311, 345– 352 Lakes, high-lithium 5– 43, 101– 135, 288

Index

Lake Tiberias thermal CaCl2 springs, Israel 270, 271, 278, 281– 283 Lake Vanda, Antarctica 45, 294, 298– 311 La Motte spodumene pegmatite, Quebec 89, 179, 203 La Negra, Salar de Carmen, Chile 112, 114, 115 Lanigan potash mine CaCl2 brine 246 Lava flow, Bristol Lake, California 285 Leaching CaCl2 from rocks 238, 239, 251, 255, 264– 269, 289, 318, 321, 334 Leaching lithium from rocks 1, 3, 5, 10, 21, 42, 43, 52 Lead– zinc deposits formed by thermal vents 269, 273, 277 Lead zinc minerals with CaCl2 occlusions 333, 334 Leakage in solar ponds l00– 102, 106, 110, 113 Lee Chemical Co., Cadiz Lake, California 351, 352, 383 Lepidolite deposits 53, 61 – 75, 82, 83, 89, 91 – 94, 96 Lepidolite processing 53, 155– 157, 165, 167, 168 Lepidolite’s properties 4, 56, 58, 67, 78 Leslie Salt Co., Bristol Lake, California 350 Licons (dilithium phosphate) 56, 133– 135, 203, 204 Lime or limestone roast of lithium ores 157– 160, 165, 167, 168, 171 Lime precipitation, CaCl2 brine 338– 341, 344 Lime precipitation, lithium brine 102– 104, 106, 107, 116, 117, 127, 129, 161, 162, 164, 170 Limestone – HCl reaction to form CaCl2 311, 357 Liners (membranes) for solar ponds 101, 106, 113, 118, 123 Liquid CaCl2 specifications 384, 385 Liquid extraction of lithium brines 115, 116, 126– 130, 136– 140 LithChem International 173, 179 Lithian muscovite, mica 71 – 73, 75, 78, 89, 93

467

Lithiophilite 4, 78, 79, 154 Lithium adsorption, reaction with rocks 3, 10, 18, 21, 34, 44 Lithium alkyls 99, 177, 196, 197, 199 Lithium aluminum hydride 177 Lithium amide 176, 177 Lithium, average content in rocks 1, 97, 98 Lithium batteries 82, 99, 182, 190– 195 Lithium bromide 176, 183, 197, 198 Lithium carbonate production 34, 107, 108, 110, 114–117, 126– 131, 134, 135, 149, 159–173, 177, 182 Lithium carbonate specifications, analyses 170, 210 Lithium carbonate’s solubility 108, 134, 135, 212–216 Lithium carnallite from high-lithium brines 110, 113 Lithium chloride 116, 131, 132, 171, 176, 183, 197, 210 Lithium Corporation of America (LCA or Lithco) 81, 99, 160, 162, 163, 201 Lithium hydride 83, 176, 200 Lithium hydroxide production 99, 108, 129, 159, 160, 165, 171, 175, 176, 182 Lithium ion rechargeable batteries 191, 193– 195 Lithium isotopes 34, 44 – 46, 98, 99, 165, 200, 256, 260, 265 Lithium leaching from rocks 1, 3, 5, 10, 42, 43, 52 Lithium loss during boron extraction 115, 116, 129, 130 Lithium metal 98, 160, 177– 179, 180, 183, 188– 190, 192– 196, 211, 212 Lithium Metals Technologies, Inc. (Limtech) 177 Lithium phosphate precipitation 34, 133– 135, 143, 145 Lithium polymer batteries 191, 193– 195 Lithium potassium sulfate from high-lithium brine 104, 110 Lithium recovery from batteries 173, 174 Lithium’s average content in rocks 1, 79, 78 Lithium sulfate 110, 126, 134, 135, 161– 165, 171– 173

468

Index

Lithium’s use as a catalyst 99, 197, 199 Lithium’s uses 83, 99, 180– 201 Longview-Beecher lithium pegmatite, South Dakota 84, 166 Lorraine Basin’s dolomitization brine, France 257 Los Patos River, Salar de Hombre Muerto, Argentina 22, 24 Louisiana oil or gas field CaCl2 brines 261 Low-sulfate Salar de Atacama brine 20, 113, 118, 124 Ludington, Michigan 339, 344, 345 Magma flow to form lithium pegmatites 1, 47 – 54, 60, 61, 68, 69, 82, 96 Magmatic fluid 1, 37, 50 – 52, 238, 280, 322, 389 Magnesium chloride brine, solid 110, 124, 239– 250, 289, 299, 310, 338– 341, 344 Magnesium Corporation of America 383 Magnesium precipitation from CaCl2 brine 338– 341, 344 Magnesium precipitation from high-lithium brine 102– 104, 106, 107, 116, 117, 127, 129, 161, 162, 164, 170 Magnesium sulfate from lithium brines 110, 126, 145 Magnetic separation 148, 149, 152– 154, 156 Mahai Lake, China 38, 40 Mali’s lithium pegmatites 2, 3, 92 Manganese oxide, lithium adsorbent 144, 145 Manic depression, lithium’s use in treating 180, 199 Manitoba, Canada’s lithium pegmatites 56, 68 – 78, 85, 89, 90 Manono and Kitotolo lithium pegmatite districts, Zaire 91 Manteen lithium pegmatite, South Dakota 84, 166 Map, world; location of some lithium deposits 2 Marine carbonates’ d6Li 46 Marine clastic sediments’ d6Li, d7Li 46 Marine fossils; CaCl2 deposits 290, 291, 322

Marshall Formation CaCl2 brine, Michigan Basin 252, 254, 338 Martin Marietta Magnesia Specialties Co. 344 Massif Central’s lithium pegmatites, France 2, 92 Matsushiro area’s CaCl2 thermal springs, Japan 277 Maywood Chemical Co. 83, 99 Medicine, lithium’s use in 177, 180, 199 Melting temperature of lithium and alkali metals 211 Melting temperatures of lithium minerals 50 – 52, 149, 181, 185 Melting temperature of occlusions 332– 334 Membranes for lithium extraction 138, 145 Membrane liners for solar ponds 101, 106, 113, 118, 123 Metabasalt’s d6Li 46 Metal deposits formed by CaCl2 thermal vents 269, 273, 276, 277, 291, 298 Metalloy Corp. 160 Metallurgical uses of CaCl2 375, 377 Metallurgy, lithium’s use in 179, 188– 190, 196, 200 Meteoric water, reaction with CaCl2 brine 113, 239, 249, 264, 265, 293, 321, 349 Meteorites’ d6Li, d7Li 46 Methane with CaCl2 or lithium deposits 42, 49 Mica recovery from lithium ore 53, 83, 153, 154, 167, 171 Mica’s reaction with CaCl2 335 Michigan Basin 45, 141, 247– 255, 259, 338– 345 Michigan’s oil or gas field CaCl2 brines 260, 261, 264 Microlite recovered from lithium pegmatites 53, 61, 83 Mid Ocean (Atlantic) Ridge thermal vents 46, 270, 273 Mid Ocean Ridge basalt’s d6Li, d7Li 45, 46 Migmatites with CaCl2 occlusions 334 Minas Gerais province’s lithium pegmatites, Brazil 2, 90 Minera de Pegmatite, Portugal 169

Index

Mineral deposits with CaCl2 occlusions 238, 244, 259, 311, 321, 332– 334 Minerals, lithium 1, 4, 54 – 59, 67, 73, 74, 76 – 78 Minerex, Austria 92 Minsal (now SQM S.A.) 118, 210 Mirabilite, thenardite near CaCl2 lakes 306, 309, 310, 314 Molybdenite in lithium pegmatites 89 Molybdenum minerals with CaCl2 occlusions 334 Mono Lake, rivers’ lithium iostopes, California 33, 34, 46 Montmorillonite (Na – ), effect of CaCl2 solutions on 336, 337 Moose No. 2 spodumene mine, Northwest Territories, Canada 90 Montebrasite 4, 56, 58, 78, 94, 154 Mount Sdom CaCl2 Springs, Dead Sea, Israel 278, 312, 313, 383 Mozambique’s lithium pegmatites 91, 206 Mullet Island CaCl2 brine, Salton Sea, California 280, 281 Namib Desert’s CaCl2 groundwater, Namibia 312, 317 Namibia’s CaCl2 lakes 294, 312, 317 Namibia’s lithium pegmatites 2, 3, 91, 169, 204– 206 Nakai Trough 44 National Chloride Company of America 346, 350, 352, 383 Natrolite’s solubility in CaCl2 335 n-Butyl lithium 99, 131, 177, 180, 183, 197 Niagra limestone CaCl2 brine, Michigan 252, 254 Nickel – copper– platinum minerals with CaCl2 occlusions 334 Niobium in lithium pegmatites 59, 92, 167 Non-rechargeable lithium batteries 194, 195 Non-zoned lithium pegmatites 50, 53, 55, 75, 80, 81, 89, 94 North Afar Rift thermal CaCl2 springs, Ethiopia 273, 276, 277, 292, 293, 297, 298, 321

469

North Carolina’s lithium pegmatites 53, 75, 79 – 82, 157– 165, 203 North Dakota’s oil or gas field CaCl2 brines 45, 261 North Fiji Basin vents 42 North Korea’s lithium pegmatites 169 Northwest Territories’ CaCl2 groundwater, Canada 314, 315 Northwest Territories’ lithium pegmatites, Canada 2, 90 Norwegian – Danish Basin’s Zechstein brine 243 Noumas, Norrabees areas’ lithium pegmatites, South Africa 2, 92 Nuristan, Afghanistan’s spodumene pegmatite 2, 56, 57, 97

dl8O 15, 37, 245, 246, 251, 253, 261, 265, 280, 299, 301, 302, 315, 331 Ocean rift or subduction zones 1, 42, 44, 45, 238, 247, 266– 277 Occlusions containing CaCl2 238, 244, 259, 311, 321, 332– 334 Occlusions in lithium minerals 52, 83, 96 Ohio’s oil or gas field CaCl2 brines 262 Oil or gas field’s lithium, CaCl2 brines 45, 46, 238, 247, 260– 266 Oil and gas industries’ use of CaCl2 359, 369– 373 Oklahoma’s oil or gas field CaCl2 brines 260, 262 Ontario and Quebec, Canada’s lithium pegmatites 23, 89, 90, 99 Onyx River, Lake Vanda, Antarctica 299, 303 Open pan evaporators for CaCl2 brine 350, 352, 353 Open pit mining 90, 92, 106, 115, 155– 159, 163, 166 Ophiolite’s reaction with CaCl2 335 Organic lithium compounds (other than butyl lithium) 177, 196, 197, 199 Organics in lithium, CaCl2 brines 115, 116, 134, 135, 161, 170, 175, 261, 311 Origin of high-lithium brines 8, 9, 17 – 21, 29 – 34, 37, 38

470

Index

Origin of lithium pegmatites 49 –52, 60, 61, 68, 69 Origin of tachyhydrite 321, 322 Orinoco River sediments’ d6Li 46 Osmolality, CaCl2 solutions 410 Othake, Kyushu, Japan’s geothermal brines 42 Owens River, Lake, California 29, 33, 34 Pacific Lithium Ltd. 135, 173, 203 Pala District’s lithium pegmatites, California 55, 82 Pakeagama Lake spodumene pegmatite, Ontario, Canada 89 Pan de Azucar rail station, Chile 109, 114 Paradox Basin, Colorado 14, 45, 243, 244, 248, 250, 263 Paris Basin’s dolomitization brine, France 257 Parma Formation CaCl2 brine, Michigan Basin 252, 254 Pastos Grandes Salar, Bolivia 23 Pelletized calcium chloride 343, 384, 385 Pelletized lithium carbonate 108, 116, 129 Pennsylvania’s oil or gas field CaCl2 brines 262 Permeability of soils, effect of CaCl2 on 337 Petalite deposits 53, 61 – 68, 85 – 89, 150– 157, 204 Petalite characteristics 4, 56, 58, 67, 68, 78 Petalite conversion to spodumene or eucryptite and quartz 50 – 52, 58, 69, 94 Petalite processing 53, 152–157, 168, 169, 171 PGE mineralization with CaCl2 occlusions 334 Phase data for CaCl2 322, 382, 383, 385– 407 Phase data for lithium compounds 134– 136, 143, 172, 202, 206, 212, 214– 223 pH of commercial CaCl2 solutions 341, 346, 432, 434 Phosphate precipitation of lithium 133– 135, 143, 145 Physical properties of CaCl2 solutions 408– 435

Physical properties of CaCl2 – NaCl solutions 412– 417, 425, 429, 430, 433, 434 Physical properties of lithium and its compounds 34, 44, 46, 50 – 52, 181, 185, 200, 202, 206, 211, 212, 214 Physical properties of solid CaCl2 426, 435 Pits, trenches for CaCl2 brine seepage, Bristol Lake 346, 347, 349 Plagioclase conversion to albite 264, 265, 267, 268, 310, 311 Plagioclase’s reaction with CaCl2 266, 268, 335 Platinum minerals with CaCl2 occlusions 334 Podlasie Basin gasfield’s dolomitization brine, Poland 265 Pollucite recovery from lithium pegmatites 62, 67, 68, 72, 83, 152, 154, 155 Polyhalite reaction to produce calcium chloride 239, 255 Polylithionite 4, 58 Portugal’s lithium pegmatites 2, 3, 56, 94, 95, 168, 169, 204– 206 Potash (potassium chloride) deposits 320– 331, 339 Potash end liquor lithium, CaCl2 brine 35 – 37, 45, 46, 112, 124, 137– 139, 141– 143, 145, 237– 259, 263, 311– 318 Potash from lithium brines 104, 110, 112– 114, 118– 124, 135 Potassium carbonate, lithium by-product 165, 166 Potassium lithium sulfate from high-lithium brines 110 Potassium schoenite from high-lithium brines 112, 124–126 Potassium sulfate, lithium by-product 110, 112, 118, 119, 124– 126 Prairie Evaporate Formation, Canada 251, 256 Precipitation of metals from ocean vents 269, 273, 276, 277, 291 Preissac-Lacorne district’s lithium pegmatites, Ontario, Canada 2, 89

Index

Presquile gas field dolomitization brine, BC, Canada 265 Pressure, temperature of initial lithium pegmatites 50 – 52, 61, 83, 96 Pre-wetted salt with CaCl2 361, 362 Prilling tower for granular CaCl2 355 Primary spodumene 50 – 52, 54, 55, 59, 61, 75, 81, 83, 89 – 97 Pripyat Trough (or Deep’s) CaCl2 brine, Russia 242, 244, 259 Production capacity; lithium carbonate, ore concentrates 203, 204 Production capacity of natural CaCl2 383 Production rate of lithium (by country) 205, 206 Production statistics for CaCl2 379– 382 Product ponds for CaCl2 brine 350, 351 Product specifications for CaCl2 384, 385 Product specifications for lithium carbonate, ore concentrates 149, 210, 211 Properties of lithium metal, compounds, alkali metals 54 – 59, 211, 212 Protolithionite 58 Puga Valley’s high-lithium geothermal brine, India 44 Pug mill granulator for CaCl2 356 Pumping CaCl2 brine 339, 342, 344, 346, 348, 350 Pumping high-lithium brine 16, 102–107, 113, 118, 123 Purification of CaCl2 brine 353, 357 Purification of high-lithium brine 107, 115, 116, 126– 129, 134, 161, 164–166, 168, 170– 173 Purified lithium carbonate 177 Qaidam Basin, Da and Kiao Qaidam, China 38, 40, 41, 135, 279 Qarhan Lake, thermal CaCl2 springs, China 135, 270, 277, 279, 294, 295, 298 Qinghai Lake’s medium-lithium brine, China 2, 38, 135 Quartz (silica) as a lithium by-product 83, 153, 171 Quartz with CaCl2 occlusions 333, 334

471

Quebec, Quebec Lithium Corp., Canada 89, 90, 99, 165, 166, 203, 204 Rain or snowfall at CaCl2 lakes 287, 293, 306, 346, 347 Rain or snowfall at high-lithium lakes 8, 21, 25, 29, 100, 102, 108, 113 Raman spectra 333, 334, 430 Rare earth minerals with CaCl2 occlusions 333, 334 Rare metals in lithium pegmatites 49 –53 Raymor Resources Co. 89, 179, 203 Rechargeable batteries, lithium’s use in 191– 195 Reclamation of alkali soils with CaCl2 337 Red Lake’s CaCl2 brine, Crimea, Russia 294, 295 Red Sea’s lithium, CaCl2 thermal vent brine 38, 270, 272– 277, 282, 283, 291– 293, 297 Refrigeration brine, CaCl2’s use as a 374, 378 Relative humidity of CaCl2 solutions 422, 426, 433 Relative humidy at CaCl2 lakes 278, 306, 321, 346, 352, 422 Relative humidity at high-lithium lakes 21, 100, 108, 113 Reserves of lithium 2, 3, 10, 17, 25, 29, 36, 37, 61, 68, 75, 79, 81, 87, 89, 91, 92, 98 Reykjanes’s lithium, CaCl2 thermal vents, Iceland 42, 270– 272 Rhodesia, Southern (Zimbabwe’s) lithium pegmatites 2, 3, 52, 53, 56, 57, 61 – 68, 99, 155– 157, 165, 204 Richfield Formation’s CaCl2 brine, Michigan 253, 254 Rift basins 37, 38, 42, 276, 280, 281, 291– 293, 297, 298, 312, 313, 321, 322, 332 Rio Grande de Lipez medium-lithium water, Bolivia 28 Rio Salado’s medium-lithium water, Chile 10

472

Index

Rio San Pedro’s medium-lithium water, Chile 11, 13, 20 River water’s d6Li, d7Li 46 Roast-acid leach process for spodumene 161– 164, 169– 172 Roast-limestone or lime leach process for lithium ores 157– 160, 165, 167, 168, 171, 172 Rocanville potash mine’s CaCl2 brine 245 Rock leaching to produce CaCl2 brines 238, 239, 251, 255, 264, 265– 269, 289, 318, 321, 334 Rock leaching to produce lithium brines 1, 3, 5, 10, 42, 43, 52 Roof scaling, bolting 150, 151 Room and pillar mining 150 Rozna area’s lithium pegmatites, Czechslovakia 2, 97 Rubican Mines, Namibia 169 Rubidium in evaporating sea water, tachyhydrite 241, 321 Rubidium in lithium pegmatites 60, 61, 75, 83, 87, 89, 92, 165, 168, 171 Russian CaCl2 groundwater 317, 318 Russian lithium pegmatites 2, 3, 94 – 96, 203– 205 Russian oil or gas field CaCl2 brines 266 Russian potash deposits 242–244, 258, 259 Rwanda’s lithium pegmatites 91

d34S 256, 280, 299, 306, 314, 329 Salar de Aguas Calientes, Chile 19 Salar de Ascoton, Chile 19 Salar de Atacama, Chile 2, 3, 5, 9 –21, 99, 108– 131, 201 Salar de Atacama’s rivers and streams 11, 13 Salar de Bellavista, Chile 19 Salar de Carmen, Chile 109, 112, 114, 118, 126, 127 Salar de Coipasa, Bolivia 11, 23, 27, 28 Salar de Empexa, Bolivia 11, 23, 27, 28 Salar de Hombre Muerto, Argentina 2, 3, 5, 11, 21, 22, 24 – 27, 99, 131– 133, 144, 202, 203 Salar de Huasco, Chile 19 Salar de Lagunas, Chile 19

Salar Salar Salar Salar Salar

de Pintatos, Chile 19 de Pujsa, Chile 19 de San Martin, Chile 19 de Surrie, Chile 19, 21 de Uyuni, Bolivia 2, 3, 9, 11, 23, 27 – 29 Safety considerations when handling CaCl2 378, 379 Sales price of CaCl2 381, 382 Sales price of lithium carbonate, ore concentrates 187 Salta, Argentina 11, 24, 132 Salt cake (sodium sulfate), lithium by-product 161– 165, 170 Salt crusts, Mono Lake’s d7Li 46 Salt produced from CaCl2 brine 287, 338– 342, 350, 351 Salt porosity in high-lithium playas 13, 16, 25, 28, 29, 104, 108, 123, 131 Salt gradient solar ponds 305 Salting-out process, boric acid, lithium recovery 126, 145 Salt Lake (within Bristol Lake), California 287, 346 Salt mixtures with CaCl2 for deicing 361– 363 Salton Sea, California, Mexico 2, 14, 37, 52, 141, 142, 274, 275, 279–281 Salt Range, India, Pakistan 242, 257, 258 San Antonio Chemicals, Texas 165 San Joaquin, California oil or gas field CaCl2 brines 260 San Louis Province lithium pegmatites, Argentina 2, 91 Saskatchewan, Canada potash deposits, CaCl2 brine 45, 244–246, 251, 256, 345 Saudi Arabia’s Umm Er Radhuma CaCl2 groundwater 318 Scheelite minerals with CaCl2 occlusions 334 Schoenite from lithium brine 112, 124– 126 Sdom well, springs, Dead Sea, Israel 278, 279, 283, 312, 313 Searles Lake, California 2, 3, 5, 31 – 33, 58, 59, 133, 165, 203

Index

Sea of Galilee thermal springs, Israel 278, 279, 283 Sea water aerosol 309, 310 Sea water evaporation, cooling 49, 238– 241, 321 Sea water’s lithium content, d6Li, d7Li 45, 46, 49, 144, 145, 270 Sea water’s reaction with rocks 238, 247, 263– 269, 272, 273 Sebka El Adhibate, Tunisia 47 Secondary spodumene 50 –52, 54, 55, 68 Seepage trenches, pits for CaCl2 brine 346– 351 SEM analytical method 334 Separation Rapids pegmatites, Manitoba, Canada 57, 69, 85 –89, 171 Sergipe tachyhydrite deposit, Brazil 320– 327 Sergipite 237, 326 Setting time of concrete with CaCl2 addition 368, 369 Siberian Platform’s CaCl2 groundwater, Russia 312, 313, 318 Siberian Platform’s dolomitization brine, Russia 259, 266 Silica (quartz) as a lithium by-product 83, 153, 171 Silica’s solubility in CaCl2 335 Silver Peak, Nevada 2, 5, 6, 9, 10, 101, 103, 107, 201, 203 Single effect evaporators for CaCl2 brine 343, 353 Smackover Formation oil or gas fields 2, 45, 47, 137, 142, 262, 266 Smectite, effect of CaCl2, conversion to illite 255, 337 Smectite, lithium 4, 97 Sociedad Chilena de1 Litio Ltd. 112 Sociedad Quimica y Minera de Chile (SQM) 118, 201 Sodium hydroxide by-product, lithium limeroast process 159, 160, 165, 171 Sodium sulfate by-product, spodumene acid roast process 161, 162, 164, 170 Soil, effect of CaCl2 solutions on 318, 336, 337

473

Soil stabilization or compaction with CaCl2 365– 368 Solar evaporation of CaCl2 brine 346– 352 Solar evaporation of lithium brine 5, 36, 37, 42, 43, 47, 48, 100– 106, 108– 126 Solar ponds design and operation 100– 102, 106, 113, 118, 120, 123 Solar radiation 21, 100, 113 Solonopole and Quixeramobim lithium pegmatites, Brazil 91 Solubility data for CaCl2 322, 382, 383, 385– 407 Solubility of lithium compounds 134– 136, 143, 172, 202, 206, 212, 214– 223 Solubility (leaching), lithium from rocks 1, 3, 5, 10, 21, 42, 43, 52 Solvent extraction of boron 115, 116, 126– 130 Solvent extraction of lithium 115, 116, 129, 130, 136–140 Solvent loss during boron extraction 116 Solvey process CaCl2 357 Solvey Process Co. 99, 157, 172 South Africa’s lithium pegmatites 2, 92 South Dakota’s lithium pegmatites 56, 82 – 85, 166– 169, 204 Southern and Nigel sectors, Bikita pegmatite 61, 62 Southern High Plains’ CaCl2 groundwater, Texas 312, 314, 315 Soviet Union’s lithium pegmatites 2, 3, 94 – 96, 203– 206 Specific conductance, CaCl2 solutions 410 Specific gravity, CaCl2 solutions 339, 383, 393, 408–411 Specific heat, CaCl2 solutions 339, 420– 422 Specific gravity, heat of lithium metal 202, 211, 212 Spodumene characteristics 4, 50 – 56, 67, 78, 80, 167 Spodumene concentrates, analyses, price 186, 187 Spodumene concentration process 146– 159, 163, 165– 170, 203, 204 a-Spodumene conversion to b-spodumene 50 – 52, 159, 160– 167, 170, 171

474

Index

Spodumene deposits 53, 54, 59 – 83, 89 –92, 94, 95, 97, 146– 171 Spodumene, primary 50 – 52, 54, 55, 59, 61, 75, 81, 83, 89 – 97 Spodumene, conversion to Li2CO3 157– 166, 169– 171 Spodumene, secondary (psuedomorphs after petalite) 50 – 52, 54, 55, 68, 78 Spray evaporator – dryer– granulator for CaCl2 343, 357 Springs, inland thermal CaCl2, lithium 46, 270, 271, 276, 279 Springs, Mono Lake fresh water’s d7Li 46 Sqi (spodumene, quartz, feldspar intergrowth) 67, 70, 78 SQM S.A. 109, 118– 131, 201, 203 87 Sr/86Sr 253, 255, 261, 265 Starobin potash deposit’s dolomitization brine, Russia 242, 249 Steam Boat Springs, Colorado, Nevada 14 Stebnik potash deposit’s dolomitization brine, Russia 242, 244 Stecklenberg’s CaCl2 springs, Harz Mountains, Germany 257 Stevensite, lithium 97 Stewart lepidolite mine, San Diego county, California 82 Stinking Spring, thermal CaCl2 spring, Utah 271 Stishovite’s reaction with CaCl2 336 St. Peter Formation’s CaCl2 brine, Michigan Basin 253, 254 Stratification in Lake Vanda 299, 301– 304 Strickland Quarry’s lithium pegmatite, Connecticut 55 Strip mining 90, 92, 106, 155, 156– 159, 163, 166 Strontium isotopic ratios 253, 255, 261, 265 Sudeten Foreland’s dolomitization brine, Czechoslovakia 257 Sua Pan’s high-lithium brine, Botswana 31, 45, 48 Sulfate precipitation by CaCl2 brines 113, 239, 249, 264, 265, 293, 321, 349

Sulfate precipitation from lithium brines 107, 113, 116, 117, 127, 129, 161, 162, 164, 170 Sulfate reducing bacteria 299, 306 Sulfuric acid roast of lithium ores 161– 164, 170 Super-critical water, methane, CO2 in forming lithium pegmatites 49 – 52, 83, 96 Surface active agents, aid to CaCl2 penetration 365, 368 Surface tension, CaCl2 solutions 427 Surrie, Salar de, high-lithium hot springs, river 19, 21 Sweden’s Varutrask and Uto lithium pegmatites 92, 98 Swimming pool conditioner, lithium hypochlorite 199 Sylvania Formation’s CaCl2 brine, Michigan Basin 45, 249, 250, 252, 254, 255, 339, 340, 342 Sylvinite from lithium brines 104, 110, 112– 114, 118– 124, 135 Sylvinite deposits 320– 331 Syngenite formation to produce CaCl2 336 Synthetic CaCl2 357 Synthetic rubber, organo-lithium catalysts for 99, 197 Tachyhydrite 237, 298, 299, 318– 332 Talc with CaCl2 occlusions 334 Tanco, Tantalum Mining Corporation of Canada Limited 68, 150, 152 Tanco spodumene deposit, Manitoba, Canada 52, 55, 56, 68 – 78, 150– 155, 202, 204 Tannur CaCl2 Springs, Lake Tiberias, Israel 278 Tantalum, tantalite recovery from lithium pegmatites 53, 59, 61, 68, 75, 79, 83, 87, 89, 90, 92, 147– 160 Taylor Valley, Antarctica 299, 300, 311 Temperature and pressure of depositing lithium pegmatites 50 –53, 61, 83, 96 Temperature, air at high-lithium playas 25, 100, 108

Index

Temperature rise in dissolving CaCl2 358, 362, 427, 429, 430 Terrestrial fossils near CaCl2 deposits 281, 290, 291, 327 Tetra Technologies Co. 350– 352, 383 Texas CaCl2 groundwater 312, 314, 315, 318 Texas oil or gas field CaCl2 brines 45, 47, 260, 261, 263, 264 Texas thermal CaCl2 springs 279 Thailand’s tachyhydrite deposit 328– 331 Thermal CaCl2 springs (inland) 270, 271, 273, 276– 279, 282, 283, 291–293, 296, 298, 317, 321, 322 Thermal gradients 50, 61, 305, 310, 311 Thermal vents, deep-ocean medium-lithium brine 42, 44 –46, 238, 266– 273 Thermodynamic properties of CaCl2 solutions 430, 431, 427, 435 Thor lithium pegmatite, Northwest Territories, Canada 2 Tiberias CaCl2 thermal springs, Israel 270, 271, 278, 281– 283 Tibet Lithium New Technology Co., Zabuye Lake, China 135 Tibet’s (Xizang Plateau’s) high-lithium brines 38– 43, 135 Tidal pool dolomitization brine 247, 317 Tin, cassiterite recovered from lithium pegmatites 53, 59, 61, 74, 87, 89 – 92, 155, 171 Tin minerals with CaCl2 occlusions 334 Tin Mountain lithium pegmatite, South Dakota 83 Tin-spodumene Belt, North Carolina 53, 75, 79 – 82 Tire ballasting with CaCl2 solutions 359, 372 Togo Matsuzaki CaCl2 hot springs, Japan 276 Topaz in lithium pegmatites 66, 77, 92 Toxicity of lithium compounds 179, 180 Toxicology of CaCl2 378, 379 Tourmaline in lithium pegmatites 52, 60, 62, 70, 73, 149, 153, 154 Trans– Pecos region’s CaCl2 groundwater, Texas 318

475

Traverse Formation’s CaCl2 brine, Michigan Basin 252– 254 Tremolite’s reaction with CaCl2 336 Trenches, seepage for CaCl2, Bristol Lake 346– 351 Triphylite 4, 83, 167 Triple effect evaporators for CaCl2 brine 339– 341, 353 Tritium, 6Li conversion to 99, 200 Tsuramaki Spa’s CaCl2 hot springs, Japan 276 Tungussky Basin’s dolomitization brine, Russia 259 Turkmenia’s oil or gas field CaCl2 brines, thermal springs 266, 277 Ulexite in high-lithium lakes or salars 18, 20, 21, 24, 25, 27, 28 Ultraviolet light detection of eucryptite 157 Underground mining of lithium pegmatites 92, 150, 151, 155, 165, 166, 168 United States’ CaCl2 production capacity 383 United States’ CaCl2 production, consumption and price 380, 382 United States’ lithium deposits 2, 5 – 10, 29 – 37, 75, 79 – 85, 101– 108, 133– 135, 137– 142, 157–170 United States’ lithium production, consumption, exports and price 203– 209 Unzoned lithium pegmatites 50, 53 – 55, 75, 80, 81, 89, 94 Urals potash deposit’s dolomitization brine, Russia 259 Uranium minerals with CaCl2 occlusions 334 Uses of CaCl2 358– 378 Uses of lithium 82, 83, 99, 180– 200 Ushtagan (CaCl2) Lake, Kazakhstan 294, 295, 298 Utah Hot (CaCl2) Spring 271 Utah’s CaCl2 thermal springs 271, 276 Utah’s oil or gas field CaCl2 brines 261 Uto lithium pegmatite, Sweden 98

476

Index

Uzbekistan’s Tien – Shan lithium pegmatites 97 Vapor pressure of CaCl2 solutions 420, 422– 424, 426 Vapor pressure of solid CaCl2 426, 427 Varutrak lithium pegmatite, Sweden 92, 94 Vegetation’s reaction to CaCl2 363, 364 Vents, ocean, with lithium, CaCl2 brine 42, 44 – 46, 238, 266–273 Ventura oil or gas field’s CaCl2 brines, California 260, 266, 321 Vermiculite’s reaction with CaCl2 336 Viscosity of CaCl2 solutions 339, 410, 417– 420 Viscosity vs temperature, Salar de Atacama brine 130 Volcanic rocks’ d7Li 46 Volga region’s CaCl2 groundwater, Russia 318 Wairakei’s geothermal brine, New Zealand 42, 44, 141 Walker Lake, Nevada’s d7Li 46 Wall zones in lithium pegmatites 65, 66, 71 Waste gas or other sources of CaCl2 357 Waste Gate Formation’s CaCl2 groundwater, Maryland 318 Water absorbed by solid CaCl2 372, 427, 432, 433 Water content of depositing lithium pegmatites 49 – 52, 56, 57, 67, 78, 82 – 84, 89 Water treating with CaCl2 357, 376 Wekusko Lake spodumene pegmatite, Manitoba, Canada 56, 169– 171 Wells for CaCl2 brine 339, 342, 346, 348, 350 Wells for high-lithium brine 102– 106, 113, 118, 123, 131

Western Caucasus’ CaCl2 groundwater, Russia 313, 318 West Virginia’s oil or gas field CaCl2 brines 261 White Picacho lithium pegmatite district, Arizona 84, 85 Wilkinson Co., Mayville, Michigan 344, 383 World map of some lithium deposits 2 World production of lithium (by country) 205, 206 Wright Valley, Southern Victoria Land, Antarctica 298, 300 Xizang (Tibet’s) high-lithium lakes, China 38 – 43, 135 X-ray diffraction lines, antarctite 311, 331 Yashan Batholith, China 92, 93 Yellowknife and Thor lithium pegmatites, Northwest Territories, Canada 2, 90 Yellowstone, Norris springs, Wyoming 14 Yichun lepidolite granite, China 56, 92, 93 Zabuye (or Zabuye Caka) high-lithium Lake, China 2, 38, 41 – 43, 135 Zaire’s lithium pegmatites 2, 3, 91 Zavitaya area’s lithium pegmatites, Russia 2, 94 Zeichstein Potash Formation, dolomitization brine 45, 243, 256, 257, 277, 315, 332, 333 Zimbabwe Craton’s CaCl2 groundwater 318 Zimbabwe’s (Southern Rhodesia’s) lithium pegmatites 2, 3, 52, 53, 56, 57, 61 – 68, 155– 157, 165, 204– 206 Zinc minerals with CaCl2 occlusions 333, 334 Zinnwaldite 4, 55, 58, 94 Zoned lithium pegmatites 50, 52 – 55, 61, 65, 66, 68, 71 – 73, 82, 85, 90 – 97