Inorganic Fluorine Chemistry - ACS Publications - American Chemical

Chapter 14 Alkali M e t a l

Polyhydrogen

Fluorides

Useful Halogen-Exchange Media 1

2

Downloaded by UNIV OF NEW SOUTH WALES on September 10, 2015 | http://pubs.acs.org Publication Date: April 29, 1994 | doi: 10.1021/bk-1994-0555.ch014

Richard E. Fernandez and Joseph S. Thrasher

1DuPont Specialty Chemicals, Jackson Laboratory, DuPont, Deepwater, NJ 08023 2Departmentof Chemistry, The University of Alabama, P.O. Box 870336, Tuscaloosa, AL 35487-0336 Alkali metal polyhydrogen fluoride systems, MF·nHF where M = K, Rb, Cs or mixtures thereof and n = 0.5 - 3 are versatile fluorination media for halogen exchange reactions. The chemical and physical properties of these systems will be discussed as well as their applicability towardsfluorocarbonsynthesis. Alkali metal fluorides are well known as traditional halogen exchange or nucleophilic substitution reagents (7-5). Typically polar aprotic solvents such as N,Ndimethylformamide, iV,A^dimethylacetamide, acetonitrile, tetrahydrofuran, glymes, etc. are required. In addition, crown ethers are often added to increase the solubility of the alkali metal fluorides and thereby the halogen exchange reaction rate. The order of reactivity is Cs > Rb > Κ > Na for the alkali metals and 1° > 2° for the electrophilic carbon center. From an industrial point of view these traditional halogen exchange media present a number of severe limitations. First, alkali metal fluorides, especially cesium fluoride, are expensive and are generally not recovered. Crown ethers are also expensive and difficult to recover in addition to being toxic. Recovery and disposal problems are also of concern with any of the aforementioned aprotic solvents. And perhaps the most severe limitation of all is the fact that these halogen exchange media can only be used in a batch mode of operation. The Concept While it is known that alkali metal bifluorides are poor traditional halogen exchange reagents they should be good "solvents" for alkali metal fluorides. If we consider the simplest case where the cations of the solvent and solute are the same then the "n" value, the molar ratio of HF to alkali metal cation, describes the system. Hence a value of "n" < 1 describes a basic or nucleophilic system, a value of "n" = 1 describes a "neutral" system, and a value of "n" > 1 describes an acidic system. The "free ^responding author 0097-6156/94/0555-0237$08.00/0 © 1994 American Chemical Society In Inorganic Fluorine Chemistry; Thrasher, J., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1994.

238

INORGANIC FLUORINE CHEMISTRY: TOWARD THE 21ST CENTURY

fluoride" concentration or the melt basicity can be readily controlled by HF addition/removal. And since polyhydrogen fluorides are always in thermal equilibrium with the metal fluoride and HF, temperature can be used as a control as well.


MF-nHF MF + η HF

[1]

Thus the alkali metal polyhydrogen fluorides are "participatory solvents," able to generate, in situ, the required reagents for traditional alkali metal fluoride halogen exchange. Further the solubilized alkali metal chloride product from the halogen exchange reaction should be easily regenerated to the alkali metal fluoride by addition of HF and removal of HC1. It was of interest to apply this concept to the synthesis of CFC alternatives, and as such the reaction sequence can be written as a repetitive loop between equations 2 and 3, the organic reaction and the inorganic or regeneration reaction, respectively. Before moving to experimental results it is informative to examine both R-Cl + MF R-F + MCI

[2]

MCI + HF MF + HC1

[3]

the thermodynamics and literature precedents for these two reactions. Thermodynamics and Early References The utility of alkali metal fluorides in halogen exchange reactions has already been documented. It is also a well-established fact that cesium is most effective and lithium least effective for the transformation of an R-Cl bond to an R-F bond. This is largely due to the minimization of the difference in lattice energies between MF and MCI coupled with the increased R-X bond strength (3,6). A pertinent example is the reaction of HCFC-133a to HFC-134a as shown in Table I (7-11) where the l o g K values clearly show the advantage of the larger cation. 10

Table I. Organic Reaction CF CH Cl(g) + MF(s) CF CH F(g) + MCl(s) 3

2

3

(HCFC-133a)

2

(HFC-134a)

MF

Δ Η \ , kJ/mole

AG° , kJ/mole

HF LiF NaF KF RbF CsF

24.8 54.7 10.3 -22.1 -31.7 -42.2

24.5 51.2 7.7 -23.2 -32.7 -42.4

r

LogioKp -4.3 -9.0 -1.3 4.1 5.7 7.4

In Inorganic Fluorine Chemistry; Thrasher, J., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1994.

p

14. FERNANDEZ AND THRASHER

Alkali Metal Polyhydrogen Fluorides

239

For the inorganic regeneration reaction (eq 3) the situation is not so clear. Nowhere in the literature is it documented which alkali metal chloride can be most easily or efficiently converted to the corresponding alkali metal fluoride. A similar thermodynamic analysis for this reaction as shown in Table II (8,11) indicates a preference for the smaller cation. This is in contrast to the requirement of the organic Table II. Inorganic Reaction - Alkali Metal Fluoride Regeneration


MCl(s) + HF(g) MF(s) + HCl(g) MCI LiCl NaCl KC1 RbCl CsCl

AH° , kJ/mole r

-28.4 16.0 48.4 57.9 68.4

AG° , kJ/mole r

LogioKp 4.4 -3.2 -8.6 -10.3 -12.0

-25.3 18.3 49.2 58.6 68.4

reaction for a larger cation. However, one must also consider the formation of polyhydrogen fluorides during the regeneration reaction if more than one equivalent of HF is used. As shown in Table III (8,9,11) for the alkali metal bifluorides, the regeneration reaction is then more favored in all cases, which can be rationalized by the well-known exothermicity of the reaction between alkali metal fluorides and hydrogen fluoride (bifluoride formation). While these data suggest larger cations could be used, the trend of favoring the smaller cation still predominates. One should bear in mind that these very approximate thermodynamic calculations are for room temperature gas-solid reactions and as such should only be used as a guide. To obtain more quantitative values one must include corrections for temperature, phases, heats of solution, etc. Unfortunately, the data required to make these corrections are not readily available. Table ΙΠ. Inorganic Reaction - Alkali Metal Bifluoride Regeneration MCl(s) + 2 HF(g) MHF (s) + HCl(g) 2

MCI

AH\,kJ/mole

LiCl NaCl KC1 RbCl CsCl

-81.3 -56.4 -41.8 -34.5 -28.2

AG° ,kJ/mole r

-36.9 -14.1 -0.3 6.2 9.5

Logi K 0

6.5 2.5 0.1 -1.1 -1.7


p



240

On the other hand, literature references dating as early as 1869 have shown that alkali metal chlorides can be converted to alkali metal fluorides in the presence of HF. Gore (12) described the reaction between the chlorides of lithium, sodium, potassium, and ammonium as strong in action with effervescence and solution. Fredenhagen (73) has shown that potassium fluoride can be produced in quantitative yield from the reaction of potassium chloride with excess hydrogen fluoride. Hood and Woyski (14) gave the first indication that the regeneration reaction could be carried out more efficiently at room temperature rather than at elevated temperatures. For example, they reported that sodium chloride reacted with anhydrous hydrogen fluoride at room temperature to form solvated sodium fluoride, while no reaction was observed at 200 °C. Results - Organic Reaction We have found that alkali metal polyhydrogen fluoride melts possess a wide range of synthetic utility including substitution reactions, HF addition reactions, and H X elimination reactions (15-17). The utility of a specific melt is governed primarily by the cation and the value of "n" of the melt. Generally, Cs/K mixtures are more effective than pure Cs melts, which are better than pure Κ melts, which are in turn better than Cs/Na melts. Sodium bifluoride does not melt but decomposes on heating to sodium fluoride and hydrogen fluoride. The value of "n" of a melt determines its melting point and HF vapor pressure at a given temperature. This value is also useful in generalizing the kinds of reactivity towards organics that a specific melt possesses. If it is desired to affect a halogen exchange, i.e. fluoride for chloride, then values of V in the range of 0.5 to 1.5 are usually most effective. If it is desired to affect the addition of HF to an olefmic species, then values of "n" greater than 1.5 are usually most effective. If it is desired to affect the elimination of HX to form an olefmic species, then values of "n" less than 0.5 are usually most effective. It should be noted that optimization of the desired transformation depends on the cation, value of "n," temperature and most importantly on the properties of the organic substrate. Optimum reaction conditions to effect a substitution reaction on one substrate may give predominantly elimination products when a different substrate is used. Table IV Table IV. Substitution and Addition Reactions

Substrate

Cation

CF -CH C1

Cs/K Cs CH C1-CH C1 Cs CHF -CC1 H Cs CFH=CC1 Κ CC1 =CC1 Cs 3

2

CF3-CH2CI 2

2

2

2

2

2

2

V 0.8 1.0 1.0 1.0 21.0 1.4

Temp (°C) Product 300 300 204 204 160 300

Conversion Selectivity

CF -CH F CF -CH F CH F-CH F CHF -CF H CFH -CFC1 CF -CC1 H CF -CHFC1 CF -CF H 3

2

3

2

2

2

2

2

2

3

2

3

3

2

2

100 % 98% 90% 23% 32% 100 %

96% 99% 70% 15% 96% 51 % 12% 6%



241 Alkali Metal Polyhydrogen Fluorides

illustrates this point. Mechanistically it is reasonable to assume that the reaction of tetrachloroethylene to HCFC-123, HCFC-124 and HFC-125 is really a series of HF additions and HC1 eliminations, i.e. the well-known addition/elimination pathway. This is shown in Figure 1 below. CFC-1111, HCFC-123, HCFC-123a, HCFC-124 and HFC-125 are the observed products in this reaction when starting with tetrachloroethylene, and their presence lends credence to this hypothesis.


Results - Inorganic or Regeneration Reaction As mentioned earlier, the inorganic or regeneration reaction (eq. 3) has very little specific literature precedent. If one assumes that the regeneration reaction takes place through the intermediate heterobihalide anion FHC1" (18-26), then it is possible to carry out a Born Haber cycle calculation such as that shown in Figure 2 in order to ascertain in which direction the heterobihalide anion might decompose. Using standard thermochemical data (6,27-29), a remarkable dependence on the choice of the cation is observed. ΔΗοι - ΔΗο2 turns out to be negative only when M = Li and Na; the values for ΔΗοι - A H D 2 are given in parentheses following the respective alkali metal (see Figure 2). Although this analysis indicates that the transformation in equation 3 should be favored for sodium over potassium and cesium, there are several other factors to be considered. First, as was pointed out for the previous thermodynamic analyses, no allowances have been made for entropy changes (25) nor differing heats of solution for MF and MCI in a fused salt medium (30). In fact, these effects must be important as it is well-known that the solvolysis of each of the chlorides LiCl, NaCl, KC1, RbCl, and NH4CI by HF generates a solution of the respective fluoride (31,32). Bulkier cations would also be favored based on the fundamental stabilization of a complex anion as a salt by the use of a bulky cation (6,27). With this uncertainty, it was of interest to broadly determine the optimum conditions for equation 3 with respect to cation or mixtures of cations, cation concentration, and percent conversion. Other interests included studying the reaction equilibria, kinetics, mechanisms, and phase behavior. Statistical Design I. In order to optimize the conversion of MCI to MF we chose to use a statistically designed experimental approach. This approach is particularly useful when a large number of variables are involved over a rather large reaction space. Essentially a statistically designed experiment is produced by defining the reaction space with variables that may have some bearing on a desired result. This result must be quantifiable and is referred to as the response surface. The completed "experiment" then describes the response of interest as a function of the variables, i.e. a surface in "n" dimensional space. As physical variables temperature was varied between room temperature and 500 °C and pressure from 1 to 20 atmospheres. As compositional variables the cation was varied from 100% cesium to 100% potassium to mixtures including sodium, and the HF/cation ratio was varied between one and twelve. The response surface would then be % conversion.



242


^

|CCIF=CC1 | 2

CCIF2-CHCI2 HCFC-122

CHCIF-CC! F HCFC-122a 2

CF2=CCIF CFC-1113

Figure 1. Proposed mechanistic pathway for the reaction of tetrachloroethylene with a CsF-nHF melt. All of the products shown in the boxes were observed spectroscopically.

MF + HCI (S)

|AH

(g)

—• M

+ (g)

+ F'

(g)

+ HCI — • M* + F + HCI + e (g)

(g)

(g)

I

D 1

[FHCI]

iyr + Cl (g)

(S)

u

ΜΟφ + HF(g)

Overall:

A H

MCI

.._

Ε d

(g)

+ H

(g)

D

HCI

( G :

I

DHF

- H ^ M+fo) + Clfo) + HF(g) —^^(g) + Cl ' + HF(g) + e (g)

D

1

-

AH

D 2

= (U

MCI

-u ) MF

+ (^-^-(PHCI-PHIO

Li (-28kJ/mole) Na (-22kJ/mole) Κ (+54kJ/mole) Rb (+68kJ/mole) Cs (+64kJ/mole) Figure 2. Born-Haber Cycle for the formation/decomposition of M[FHC1].





243

A set of 50 runs or reactions made up the first statistical design. The reactions were run starting with 100% MCI as a solid in either a sealed, evacuated Monel or Inconel autoclave (static method) or a nickel boat in an inert gas (N2) purged nickel tube reactor (dynamic method). The autoclave or tube reactor was brought to the desired temperature and a preweighed amount of anhydrous HF was added in one portion to the autoclave or metered through the tube reactor with the aid of a mass flow controller . The autoclave was held at the desired temperature for 30 minutes with stirring, and the pressure was monitored. The volatile contents of the autoclave were then vented into an aqueous potassium hydroxide solution of known concentration. The effluent from each tube reaction was also scrubbed with an aqueous potassium hydroxide solution of known concentration. Following the addition of HF, each tube reaction was purged for 30 minutes with dry nitrogen with the effluent again going to the scrubber. The liquid and/or solid remaining in the autoclave or boat was then dissolved in another aqueous potassium hydroxide solution of known concentration. All solutions were quantitatively analyzed for chloride and fluoride by specific ion electrodes and/or ion chromatography. The percent conversion was then calculatedfromthese data for each run with the chloride concentration remaining in the solid being the primary basis for % conversion. Several experiments were run in the opposite direction, i.e. starting with 100% MF, in order to assure that equilibrium conditions had been established. Results. The results or response surfaces for several specific cases - starting with 100% CsCl, KC1, and NaCl, respectively - are shown as 3-dimensional plots in Figures 3-5. In each figure, % conversion is plotted as both a function of temperature and the HF to metal chloride mole ratio. These results indicate that high conversions of metal chloride to metal fluoride (bifluoride) are possible for all three alkali metal chlorides. When the three figures are considered together the trends on going from cesium to sodium are obvious with the highest conversions being favored by a) the larger cation, b) lower temperatures, and c) higher HF to metal chloride mole ratio. It is interesting to note that again we see trends contrary to thermodynamic predictions. Also noteworthy is the trough like feature on each plot in the 300 to 400 °C range. This is a real phenomenon and may represent the transition between a MCI to MHF pathway to a direct MCI to MF pathway as M H F 2 salts are unstable at elevated temperatures. 2

Statistical Design II. The second statistically designed experiment was focused on the more realistic case where a simulated "spent" melt (from the organic reaction) was regenerated with HF. This "spent" melt would be expected to only be partially depleted to the metal chloride, the bulk remaining as the metal polyhydrogen fluoride solvent. The first statistically designed experiment indicated that lower temperatures were most beneficial. The temperature in this experimental design was fixed at 150 °C as a compromise between the preferred higher temperature of the organic reaction and the preferred lower temperature of the inorganic regeneration reaction. Obviously, substantial cost savings could be realized by minimizing the difference between the two reaction temperatures, if this process were commercialized. The compositional variables in these experiments were mole fraction of M X as metal chloride (variedfrom0.1 to 0.4), mole fraction of M as potassium (variedfrom0.2 to



244

INORGANIC FLUORINE CHEMISTRY: TOWARD T H E 21ST CENTURY

Figure 3. Statistical design I results after starting with 100% cesium chloride.

Figure 4. Statistical design I results after starting with 100% potassium chloride.




245


1.0 with the remaining M being cesium), and the HF to metal cation ratio "n" (varied from 1.0 to 5.5). Again the response surface was % conversion of the dissolved metal chloride to the metal fluoride or polyhydrogen fluorides. All 27 trials in this statistically design experiment were run in the Monel autoclave as previously described above (static method). Pressure was allowed to vary, but was monitored. The effect of pressure from added nitrogen was also examined in some reactions. Results. Figure 6 displays a contour plot of % conversion as a function of the mole fraction of potassium and the HF to total cation mole ratio "n" at a constant temperature of 150 °C. Again, very high conversions of MCI to MF can be obtained at all molefractionsof potassium when between 4.5 and 5.5 mole equivalents of HF are used. Removal of Excess HF. Once the regeneration has been carried out, the resulting melt has a value of V « 5.0, i.e. much too high to recycle directly to the organic reaction. Therefore, it was necessary to study the thermal stability of these high V melts in order to determine the best way to return the melt to the desired "n" value of 0.8 - 1.0. For these studies a Cahn Instruments C-1100 pressure balance was utilized. This equipment allows thermogravimetric analyses (TGA) to be carried out in a stream of anhydrous hydrogen fluoride. A schematic diagram of the experimental setup is shown in Figure 7. In a typical experiment, a known amount of alkali metal fluoride is placed in the sample pan. The sample chamber of the balance is then heated to 500 °C while purging with dry nitrogen in order to dry the sample. This step is complete when a constant weight is obtained for the sample. After drying, the sample is heated at the desired reaction temperature while the nitrogen purge is replaced with an anhydrous hydrogen fluoride purge. The sample weight is continuously monitored by the computer. When a maximum weight is attained and noted, the hydrogen fluoride purge is replaced with a dry nitrogen purge. The weight loss is then followed and recorded versus time. Repeating this experiment at several different temperatures allows a plot of isotherms of stability to be determined as a function of the value of "n" with time. An example of this type of plot is shown for potassium fluoride in Figure 8. The effect of different cations or mixtures of cations on the thermal stability of various melts at a given temperature can be compared and contrasted. The plot shown in Figure 9 compares the thermal stability of a cesium-, a potassium-, and a cesium/potassium polyhydrogen fluoride based melt at 150 °C. As expected, the cesium melt is the most stable (top line), the potassium melt the least stable (bottom line), and the cesium/potassium melt of intermediate stability (middle line). Additional information on the freezing points (33), vapor pressures (34), densities/viscosities (35), solubilities (36), phase diagrams (33,37), and thermal stabilities (38) of the alkali metal fluoride-hydrogen fluoride systems is available in the literature. This information is useful when developing halogen exchange chemistry around these systems. We are in the process of measuring similar data for the ternary system cesium fluoride-potasssium fluoride-hydrogen fluoride. For example, one can



246

I N O R G A N I C F L U O R I N E C H E M I S T R Y : T O W A R D T H E 21ST C E N T U R Y

Figure 5. Statistical design I results after starting with 100% sodium chloride.

Figure 6. Statistical design Π results after starting with a molefractionof 0.2 chloride.


14.

FERNANDEZ AND THRASHER


To air


pressure balance controller

computer

strip chart recorder

laser printer

To computer

Figure 7. Experimental setup for TGA studies utilizing a Cahn Instruments C-l 100 pressure balance.


247

248


gain insight into the ternary phase diagram not only through experimentation, but also by careful examination of all available binary phase diagram information (33,37,39). The results of these studies will be the subject of a later publication.


1

1

1

-

CsFnHF

:

0.25KF0.75CsF nHFKFnHF

ι

c

-

:

0

1

1

1

31.25

62.5

93.75

125

Time (min)

Figure 9. A comparison of the thermal stability of CsF-nHF, 0.25KF0.75CsFnHF, and KFnHF at 150 C.

Conclusions Alkali metal polyhydrogen fluoride melts are tunable, participatory solvents for halogen exchange, HF addition, and HX elimination reactions. High conversions with excellent selectivities to the desired fluorocarbon products are possible with the appropriate melt. The resulting alkali metal chlorides can be easily regenerated with anhydrous hydrogen fluoride in a separate step.



Alkali Metal Polyhydrogen Fluorides 249

Acknowledgments We gratefully acknowledge the assistance of John F. Kook, John E. Miller and R. Bertram Diemer of the Du Pont Company and Liang Hu, H. P. Sampath Kumar, Jian Sun, Jeff Choron and Carl Williams all of the University of Alabama.


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

Wilkinson, J. A. Chem. Rev. 1992, 92, 505 and reference therein. Meshri, D. T. In Fluorine: The First Hundred Years (1886-1986); Banks, R. E.; Sharp, D. W. Α.; Tatlow, J. C., Eds.; Elsevier: New York, 1986; Chapter 10. Chambers, R. D. Fluorine in OrganicChemistry;John Wiley & Sons: New York, 1973 and references therein. Hudlicky, M. Chemistry of Organic Fluorine Compounds: A Laboratory Manual with Comprehensive Literature Coverage, 2nd (revised) ed.; Ellis Horwood: Chichester, 1976 and references therein. Young, J. A. Fluorine Chem. Rev. 1967, 1, 359. Steudel, R. Chemistry of the Non-Metals; Walter de Gruyter: Berlin, 1977; Chapter 5. Buckley, G. S.; Rodgers, A. S. J. Phys. Chem. 1983, 87, 126. J. Phys. Chem.Ref.Data 1985, 14, Suppl. 1. J. Phys. Chem. Ref. Data 1982, 11, Suppl. 2. Chen, S. S.; Rodgers, A. S.; Chao, J.; Wilhoit, R. C.; Zwolinski, B. J. J. Phys. Chem. Ref. Data 1975, 4, 441. Barin, I. Thermochemical Data of Pure Substances; VCH Publishers: New York, 1988. Gore, G. J. Chem. Soc. 1869, 22, 368. Fredenhagen, H. Z. Anorg. Allg. Chem. 1939, 242, 23. Hood, G. C.; Woyski, M. M. J. Am. Chem. Soc. 1951, 73, 2738. Cassel, W. R.; Fernandez, R. E.; Mader, F. W. U.S. Patent 4 990 701, 1991. Cassel, W. R.; Fernandez, R. E.; Mader, F. W. U.S. Patent 4 990 702, 1991. Fernandez, R. E.; Gumprecht, W. H.; Kaplan, R. B. U.S. Patent 5 045 634, 1991. Evans, J. C.; Lo, G. Y-S. J. Phys. Chem. 1966, 70, 543. Ault, B. S. J. Phys. Chem. 1979, 83, 837. Fujiwara, F. Y.; Martin, J. S. J. Am. Chem. Soc. 1974, 96, 7625. Evans, J.C.;Lo, G. Y-S. J. Chem. Phys. 1966 70, 11. Smart, R. St.C.;Sheppard, N. J. Chem.Soc.,Chem. Commun. 1969, 468. Smart, R. St.C.;Sheppard, N. Proc. Roy. Soc Lond. 1971, 320A, 417. Fujiwara, F. Y.; Martin, J. S. J. Phys. Chem. 1972, 56, 4091. Salthouse, J. Α.; Waddington, T. C. J. Chem.Soc.,A 1964, 4664. King, D. L.; Herschbach, D. R. Faraday Dicuss. Chem.Soc.1973, 55, 331. Douglas, Β. E.; McDaniel, D. H.; Alexander, J. J. Concepts and Models of Inorganic Chemistry, 2nd ed.; John Wiley & Sons: New York, 1983; p 229. CRC Handbook of Chemistry and Physics; Weast, R. C., Ed.; CRC Press: Boca Raton, Florida, 1986; pp D-100 - D-114.


250 29. 30. 31.


32. 33. 34. 35. 36. 37. 38. 39.


Greenwood, Ν. N.; Earnshaw, A. Chemistry of the Elements; Pergamon Press: Oxford, 1984; pp 94-97. Thompson, N . R . ; Tittle, B. In Halogen Chemistry; Gutmann, V., Ed.; Academic Press: New York, 1967; Vol. 2. Dove, M. F. Α.; Clifford, A. F. In Chemistry in Nonaqueous Ionizing Solvents; Jander, G.; Spandau, H.; Addison, C. C., Eds.; Pergamon Press: New York, 1971; Vol. 2, Pt. 2, pp 119-300 and references therein. Jache, A. W. In Fluorine-Containing Molecules: Structure, Reactivity, Synthesis, and Applications; Liebman, J. F.; Greenberg, Α.; Dolbier, Jr., W. R., Eds.; VCH Publishers: New York, 1988; Chapter 9. Cady, G. H. J. Am. Chem. Soc. 1934, 56, 1431. Yusova, Y. L.; Alabyshev, A. F. Russ. J. Phys. Chem. 1962, 36, 1506. Semerikova, I. Α.; Alabyshev, A. F. Russ. J. Phys. Chem. 1962, 36, 1507. Cady, G. H. J. Phys. Chem. 1952, 56, 1106. Winsor, R. V.; Cady, G. H. J. Am. Chem. Soc. 1948, 70, 1500. Opalovsky, Α. Α.; Fedorov, V. E.; Fedotova, T. D. J. Thermal Anal. 1970, 2, 373. Sangster, J.; Pelton, A. D. J. Phys. Chem. Ref. Data 1987, 16, 509.

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