Calcium chloride hydrate, melt

Calcium(II) chloride hexahydrate melts, supercooling phenomenon, 39 430, 432 Calcium(II) chloride hydrate, melt, 39 426-432... 

A comparison between structures of calcium(II) chloride hydrate crystals and melts, which can be used as a heat storage material, has been carried out by means of X-ray diffraction (52). The phase diagram of... 

X-ray diffraction analysis of the structure of the hydrate melts reveals that the calcium(II) ion in the CaCl2-8.6H20 (25°C) is surrounded by six water molecules with a Ca-OH2 distance of 245 pm (see Fig. 10). Chloride ions tend to enter the first coordination sphere of the calcium(II) with decreasing water concentration in the melt to form ion... 

Thus, the supercooling phenomenon in the calcium(II) chloride hexa-hydrate melts may be explained in terms of the dissimilarity in the structure around the calcium(II) ions in the melt and in the crystal, and the substitution of water by chloride in the coordination sphere of the calcium(II) ion in the melt should be completed during cooling. On the other hand, the CaCl24H20 hydrate melt has a very similar... 

In completion it may be remarked that we have arrived in two different ways at the conclusion that the lowering of the melting point of hydrated calcium chloride by addition of water or calcium chloride is less than normal. Both arguments proceed essentially from the same fundamental assumption the first from the partial dissociation of the hydrate for equilibrium in the fused state the second on the existence of a pressure of water vapour for the fused hydrate, which is intimately bound up with the partial dissociation of the hydrate. Hydrates, therefore, that when melted have no vapour pressure of water must give the normal lowering of melting point on addition of cither component that is very approximately the case for bodies like sulphuric acid, and completely so for the carbohydrates. 

Although in taking up the discussion of the equilibria between calcium chloride and water, it was desired especially to call attention to the form of the solubility curve in the case of salt hydrates possessing a congruent melting-point, nevertheless, for the sake of completeness, brief mention may be made of the other systems which these two components can form. 

F. T. Frerichs observed no precipitation and F. Kuhlmann found that calcium carbonate is only partially converted into the chromate by a soln. of potassium chromate hme-water gives no precipitate with a soln. of potassium chromate. F. Mylius and J. von Wrochem obtained the anhydrous chromate by heating the hydrated salt—G. N. Wyroubofi said at 300°—or by warming a supersaturated soln., containing 15 to 20 per cent. CaCr04, over 36°. The yellow powder consists of fine needles. L. Bourgeois obtained crystals of the anhydrous chromate by melting a mixture of 2 eq. of calcium chloride with an eq. each of potassium and sodium chromates. D. Vorlander and H. Hempel observed no transformation into... 

Anhydrous calcium chloride CaCli is a white, very hygroscopic powder. At room temperature, at a relative humidity of 60%, 1 kg of calcium chloride can absorb 1.6 kg of water. It is, therefore, widely used for desiccating and dehydrating gases. Solutions of the hydrate CaCl2-6H20 have a melting point of - 55 °C and are used as a coolant liquid. 

In 1885 C. A. von Welsbach isolated two elements as oxides, praseodymium (the word meaning green twin ) and neodymium (meaning new twin ), from a mixture of lanthanide oxides called didymia. The oxides can be transformed to fluorides by reaction with HF at 700°C (1,292°F), or with NH4HF2 at 300°C (572°F). The hydrated fluorides are then dehydrated in vacuo in a current of HF gas. The metals praseodymium and neodymium are obtained via metallothermic reduction with calcium at approximately 1,450°C (2,642°F), or via electrolytic reduction of the melts. The metals can also be obtained from anhydrous chlorides, obtained via reaction of the oxides with ammonium chloride at 350°C (662 °F), which are then reduced with lithium-magnesium at approximately 100°C (212°F). 

Sodium chloride, gypsum, and potassinm salts are all important minerals that are recovered as evaporites remaining from the evaporation of seawater and from brines pumped from below the ground. Sodium chloride in the form of mineral halite is used as a raw material for the production of industrially important sodium, chlorine, and their compounds. It is used directly to melt ice on roads, in foods, and in other applications. Potassium salts are, of course, essential ingredients of fertilizers and have some industrial applications as well. Gypsum, hydrated calcium sulfate, is used to make plaster and wallboard and is an ingredient in the manufacture of portland cement. 

Note that the molar enthalpies of melting of the salt hydrates are generally rather larger than those of similar anhydrous salts melting at a much higher temperature. Compare, for instance the Am///kJ moF of lithium salt trihydrates with the anhydrous salts nitrate 36.4 vs. 26.7 and perchlorate 40.6 vs. 17.0, the sodium salt decahydrates with the anhydrous salts carbonate 72.0 vs. 30.0, sulfate 78.7 vs. 23.0, and the calcium salt hexahydrates with the anhydrous salts chloride 37.2 vs. 28.5, bromide 35.6 vs. 28.9. 

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