PHYSICOCHEMICAL PROPERTIES OF ALKALINE EARTH METAL FLUORIDES

E. K. ZAVADOVSKAYA, N. M. TIMOSHENKO

Fluorine compounds have long been used in technology, for example, as working substances in refrigerating machines and as catalysts for many processes in chemical technology [1]. Progress in science and technology has increased interest in fluorine compounds. At present, fluorine compounds have found broad application in organic chemistry in the preparation of chemically active liquids and plastics [2, 3]. Interest has grown in fluoride compounds of the actinides, due chiefly to their use in the technology of processing nuclear-power raw materials, and also directly as nuclear fuel [4]. Fluoride compounds are used as important ingredients in the production of enamels, ceramics, and in the manufacture of special optical glasses (coated optics [5]).

Artificial single crystals of fluorides of alkali and alkaline earth metals, such as LiF and CaF₂, have become widely used as optical media in optical instruments designed for operation in the ultraviolet and infrared regions of the spectrum [6].

Recently there has been increased interest in CaF₂ crystals with additions of rare-earth elements, used as materials for quantum-mechanical systems (lasers and masers). This circumstance has also predetermined interest in the chemical analogs of calcium fluoride, SrF₂ and BaF₂.

The data available in the literature on fluoride compounds by no means provide a complete picture of the properties of these materials and the possibilities for their use. Investigations of the properties of these compounds are of definite interest both for science and for practice, and they expand our knowledge of this very interesting and distinctive group of compounds.

Features of Alkaline Earth Metal Fluorides in Comparison with Other Halides of the Same Metals

Interest in fluoride compounds of metals (both alkali and alkaline earth metals) was due, above all, to their chemical stability in comparison with other halide compounds of the same metals. Fluoride compounds owe their stability to an extraordinarily—

…chemical activity of fluorine, evidence of which is the fact that fluorine is not found in nature in the free state [1]. Compounds of fluorine with all elements are known, except for the inert gases.

The chemical activity of fluorine is due to its high electronegativity. Fluorine is the most electronegative of all elements, including oxygen [7] (Table 1). On the other hand, fluorine ions have a smaller radius than the ions of the other halogens (Table 1). This circumstance accounts for the stronger bond of fluorine ions with metal ions, and, along with this, the higher crystal-lattice energy for fluorine compounds than for other halogens.

Table 1

Values of electronegativity and ionic radius of halogens and oxygen analogs [1, 12]

The combination of the very high electronegativity of fluorine and the small size of its ion accounts for the sharp difference in the structure and properties of fluorides from other halides. Table 2 clearly demonstrates the change in properties in a series of compounds having one and the same cation but different anions.

Table 2

In the series of alkaline-earth-metal halides of Ca, on going from fluorine to iodine there is an increase in the interionic distance and a decrease in the lattice energy caused by the action of Coulomb forces. With a decrease in the lattice energy of the compound, i.e., on going from fluorides to iodides, the physicochemical properties of the compounds also change. Thus, CaF₂ has the highest melting and boiling temperatures, and CaI₂ the lowest. If CaF₂ is practically insoluble in water, the other halides possess good solubility, increasing as the lattice energy of the compound decreases. Analogous changes are also observed for the halides of strontium and barium.

Fluorides of alkaline-earth metals, in contrast to other halides, do not form crystalline hydrates. The impossibility of the stable existence of crystalline hydrates of calcium fluoride is also indicated by the theoretical calculations of Yatsimirsky [9]. The small radius of the fluorine anion and the associated high lattice energy hinder the coordination of water at the cation. For the same reason, CaF₂ does not form ammoniates [8]. If CaF₂ and CaCl₂ have a fluorite-type lattice, then CaBr₂ has the cesium chloride lattice, and CaI₂ has a complex lattice [10].

A characteristic feature of metal fluorides is their lower volatility and greater toxicity than the corresponding other halides. A notable difference between fluorides and other halides is characteristic of the elements of all groups of the periodic system [11].

Properties of alkaline-earth fluorides in relation to their chemical composition

When considering the difference in the properties of fluorides and other halides, a regular change in properties in the series of halides as a function of the lattice energy of the compound is observed (using calcium halides as an example, Table 2). A comprehensive study of various properties and their interrelation for a group of ionic structures made it necessary to choose a characteristic that would make it possible scientifically to systematize the properties of ionic crystals and their variation. In works [12] and [13], a common characteristic was chosen for all ionic compounds—the energy of the crystal lattice—which, on the one hand, is a thermodynamic quantity with the properties of generality inherent in such quantities when the crystal is considered as a continuous medium; on the other hand, the magnitude of the lattice energy, proportional to the product of the ionic charges and inversely proportional to the distance between them, is related to the details of the molecular structure and its chemical composition.

Let us consider the properties of fluoride compounds of the elements of the first four groups of Mendeleev’s table (Table 3).

For fluorides of alkali metals, which have the same lattice type, a regular change in the properties of the compounds is observed as the lattice energy changes. Some discrepancies with the lattice energy are observed for the melting and boiling temperatures of LiF.

For fluorides of metals of the 2nd group, a regular change in properties with the lattice energy is observed only for fluorides of alkaline-earth metals having the same lattice type. MgF₂ has lower thermal stability and greater solubility, although its lattice energy is greater than that of CaF₂. At the same time, MgF₂ has a lattice type different from the lattice of the fluorides of alkaline-earth metals. MgF₂ dissolves more strongly in acids than CaF₂, and is also less resistant to the action of melts of other fluorides, forming double salts with them. These circumstances confirm the assumption that the energy of the crystal lattice is a characteristic that determines changes in properties only in a series of compounds with the same structure and lattice type [12]. BeF₂ differs markedly in its properties from fluoride compounds of metals of the second group. When BeF₂ is heated, it gradually softens: it has no fixed melting temperature. BeF₂ begins to evaporate noticeably at 800°C. In melts with lithium fluoride, BeF₂ undergoes two polymorphic transformations, corresponding to different lattice types.

Fluoride compounds of elements of the 3rd and 4th groups are mainly gaseous substances. As the atomic weight of the element in the group increases, the state of the fluoride compound changes—

Table 3

...compounds. If \(CF_4\) and \(SiF_4\) are gaseous compounds, then \(TiF_4\) exists in the initial state as a powder.

Analysis of the data in Table 3 makes it possible to conclude that, from the standpoint of thermal and chemical stability (and this circumstance is important for radiation materials science), the fluorides of the metals of group 2, and especially the fluorides of the alkaline-earth metals, are of the greatest interest.

The study of the laws governing the change in properties in the series of alkaline-earth-metal fluorides is of interest, since these compounds have the same structure and lattice type. All three crystals—\(CaF_2\), \(SrF_2\), and \(BaF_2\)—have a fluorite-type lattice: the ions \(Ca\), \(Sr\), and \(Ba\) form a face-centered cubic lattice, while the fluorine ions are located at the center of each \(1/8\) of the cube. Each metal ion is surrounded by 8 fluorine ions in a cubic arrangement at a distance

\[ d_{Ca\ldots F}=\frac{a}{4}\sqrt{3} \]

[15].

As the radius of the metal ion increases, and along with it the ratio \(\frac{R_{\mathrm{cat}}}{R_{\mathrm{an}}}\) increases, the bond energy of the ions decreases, i.e., the energy of the crystal lattice in the series of alkaline-earth fluorides decreases (Table 4). The lattice parameter also increases in going from \(CaF_2\) to \(BaF_2\). In the same direction there is an increase in the molecular weight and density of the compounds.

The decrease in lattice energy in the series from \(CaF_2\) to \(BaF_2\) is accompanied by a decrease in the thermal and chemical stability of the compounds. The melting and boiling temperatures of \(CaF_2\) are higher than those of \(BaF_2\). For \(SrF_2\), values of the order of \(1190^\circ C\) have been cited [6]. Data from other authors [1, 16], as well as our experience in growing \(SrF_2\) crystals [17], cast doubt on this value.

The high lattice energy of the fluorides is the reason for the low solubility of the fluorides, which is also explained by the theoretical conclusions of Kapustinskii and Yatsimirskii [18].

As is seen from Table 4, calcium fluoride has the lowest solubility. The solubility of \(CaF_2\) and \(BaF_2\) in water at high temperatures, up to those exceeding the critical temperature, has been studied [19]. The solubility of \(CaF_2\) has a sharp maximum (0.06 g/1000 g of water) at \(360^\circ C\); the solubility of \(BaF_2\) decreases continuously.

Alkaline-earth fluorides do not dissolve in organic solvents. \(BaF_2\) has the greater solubility in acids. Nevertheless, the fluorides of the alkaline-earth metals dissolve only slightly in acids, which makes the use of these materials in chemically aggressive media promising.

The chemical inertness of alkaline-earth fluorides at relatively low temperatures under atmospheric conditions is accompanied by an increase in activity at elevated temperatures. When \(CaF_2\) crystals are heated in air, diffusion of oxygen into the crystal is observed. Jure [20] suggested that the reaction proceeds as follows:

\[ CaF_2+H_2O \to CaO+2HF \]

or

\[ CaF_2+\frac{1}{2}O_2 \to CaO+F_2. \]

Slight penetration of oxygen into the crystal up to temperatures of \(800^\circ C\) is accompanied by an increase in the depth of penetration at higher...

Table 4

Some properties of alkaline-earth metal fluorides [6, 12, 15]

temperatures [21]. This circumstance does not allow the process of growing crystals to be carried out in air, but requires a neutral medium or a vacuum of the order of \(10^{-8}\)—\(10\) mm Hg [22].

When crystals of CaF\(_2\) and SrF\(_2\) were heated in a stream of oxygen or air at a temperature of 1000°C for two hours, a light-scattering layer was formed with a depth of 0.93 mm for CaF\(_2\) and 1.34 mm for SrF\(_2\) [23]; the results presented make it possible to judge the decrease in resistance to the action of oxygen in the series of fluorides as the lattice energy decreases.

The specific heat capacity and thermal conductivity of CaF\(_2\) are greater than those of BaF\(_2\). The expansion coefficient of CaF\(_2\) crystals up to 650°C is described by the equation [24]:

\[ d_t = 1.851 \cdot 10^{-5} + 1.481 \cdot 10^{-8}t + 2.152 \cdot 10^{-11}t^2, \]

The Young’s modulus and microhardness (see the present collection) decrease in the fluoride series from CaF\(_2\) to BaF\(_2\).

Crystals of CaF\(_2\), SrF\(_2\), and BaF\(_2\) have found wide application in instrument engineering as optical media, owing to their high transmittance over a broad wavelength interval and their low light dispersion. The transmission limits of the crystals, on going from a compound with higher lattice energy to a compound with lower lattice energy, shift both in the short-wavelength and in the long-wavelength regions toward greater wavelengths. The crystals reach maximum transmission in the short-wavelength region in the order of the cation sequence in the Mendeleev table. Thus, CaF\(_2\) reaches 90% transmission at a wavelength of about 200 mµ, SrF\(_2\) at 280 mµ, and BaF\(_2\) at 350 mµ. As the cation radius increases in the series of alkaline-earth fluorides, the packing density of ions in the lattice decreases, which is reflected in the value of the relative free volume \(\Phi\). For alkaline-earth-metal fluorides we have derived a formula for calculating \(\Phi\):

where

\[ \Phi = 1 - \frac{2.52\left(R_{\text{cat}}^{3} + 2R_{\text{an}}^{3}\right)\cdot d}{M}, \]

\[ R_{\text{cat}} \text{ and } R_{\text{an}} \text{ are the radii of the cation and anion, } (\text{Å}); \]

\[ d \text{ is the density, } M \text{ is the molecular weight.} \]

The value of the relative free volume is an important characteristic, especially in evaluating the radiation resistance of a material, since it influences the diffusion processes occurring in the substance both under the action of radiation and after its cessation.

The electron density increases as the atomic number of the cation increases. The electron density \(N_e\) was calculated per 1 cm\(^3\) by the formula

\[ N_e = 6.02 \frac{d}{M}(z_{\text{cat}} + 2z_{\text{an}})\cdot 10^{23}\ \text{cm}^{-3}, \]

where \(z_{\text{cat}}\) and \(z_{\text{an}}\) are the atomic numbers of the cation and anion. The electron density is responsible for absorption of radiation energy. In BaF\(_2\)—a material with higher electron density—at the same exposure dose, a larger part of the radiation energy is absorbed than in CaF\(_2\) (calculated per 1 cm\(^3\); see the present collection).

REFERENCES

1. I. G. Ryss, Chemistry of Fluorine and Its Inorganic Compounds. Goskhimizdat, Moscow, 1956.

2. D. H. Simons, T. J. Brice. Fluorine and Its Compounds. Vol. 2, IL, 1956.

1. Fluoroorganic Electrical Insulating Materials, ed. B. M. Tareev. GEI, 1957.

2. I. V. Tanaev, N. S. Nikolaev, Yu. A. Luk’yanychev, I. F. Alenchikova. Chemistry of Fluoride Compounds of the Actinides. Publishing House of the Academy of Sciences of the USSR, Moscow, 1963.

3. V. A. Weil. Fluorine and Its Compounds. IL, vol. 1, Moscow, 1953.

4. E. M. Voronkova, B. N. Grechushnikov, G. I. Distler, I. P. Petrov. Optical Materials for Infrared Technology. Moscow, 1965.

5. L. Pauling. The Nature of the Chemical Bond. Goskhimizdat, 1947.

6. K. B. Yatsimirskii. ZhOKh, 17, 2019, 1947.

7. V. Bil’ts. Uspekhi Fizicheskikh Nauk, 5, 356, 1925.

8. C. Gottfried. Schlossberger F; Strukturbericht, III, 1937.

9. H. D. Emileus. Fluorine and Its Compounds. Vol. 1, IL, 1963.

10. A. A. Vorob’ev. Physical Properties of Ionic Crystalline Dielectrics, part I, Publishing House of Tomsk University, 1960.

11. A. A. Vorob’ev, E. K. Zavadovskaya. Electrical Strength of Solid Dielectrics. GTTI, 1956, Moscow.

12. A. V. Novoselova, Yu. P. Simonov, E. I. Yarembash. ZhFKh, 26, 1224, 1952.

13. Landolt-Börnstein, vol. 4, Kristalle, 1955.

14. Ruff, Boucher L. Z. Anorg. Chem., 219, 376, 1934.

15. N. M. Timoshenko. Growth of Crystals of Alkaline-Earth Metal Fluorides. Report, TPI, 1963.

16. A. F. Kapustinskii, K. V. Yatsimirskii. ZhFKh, 22, 1271, 1948.

17. Booth, Bidwell. J. Am. Chem. Soc., 72, 2567, 1950.

18. Ure R. W. J. Chem. Phys., 26, 1363, 1957.

19. Phillips W. L., Hanlon I. E. I. Amer. Ceram. Soc., 46, 9, 441, 1963.

20. I. V. Stepanov, M. A. Vasil’eva. In: Crystal Growth, Publishing House of the Academy of Sciences of the USSR, 1957.

21. Bruch H., Görlich P., Karras H., Lehmann R. Phys. Stat. Sol., 4, 3, 685, 1964.

22. Charma. Proc. Indian. Ac. Sci., 31A, 261, 1951.