THE PHYSICAL CHEMISTRY of METALS in THEIR MOLTEN HALIDES THESIS Presented for the Degree of DOCTOR of PHILOSOPHY in the UNIVERSI

THE PHYSICAL CHEMISTRY OF METALS IN THEIR MOLTEN HALIDES

THESIS presented for the degree of DOCTOR OF PHILOSOPHY in the UNIVERSITY OF LONDON By Lars-Ingvar Staffansson, Bergsingenjor, (Stockholm).

London, December 1959, ABSTRACT

The phase diagrams of the systems calcium- calcium chloride, calcium-calcium bromide and calcium- calcium iodide have been established by solubility measurements and differential thermal analysis. Each of the systems was found to have a eutectic a monotectic and a large miscibility gap. The eutectics are 3.3, 3.2 and 6.8 mole % calcium respectively, and the corresponding temperatures are 763, 728 and 760°C. The monotectics are at 99.5, 99.6 and 99.7 mole % calcium respectively and the corresponding temperatures are 826, 828 and 831°C. The consolute points are at 1338, 1335 and 1380°C and the corresponding compositions are 62, 64 and 74 mole % calcium. Discrepancies between this and previous work are mainly attributed to composition changes of the phases during quenching. A method has been developed that overcomes this difficulty. The depression of the freezing point of strontium by strontium chloride has also been determined. From the results of this work a model for the mutal solubility of metalS and their molten halides has been suggested. CONTENTS Page

CHAPTER 1: INTRODUCTION 1 1.1 General Introduction 1 1.2 Relevant previous work on metal-molten salts. 4 1.2.1 General 4 1.2.2 Alkaline earth metal-metal halide systems. 12

CHAPTER 2: EXPERIMENTAL 24 2.1 Programme 24 2.2 Preparation and purity of materials 25 2.2.1 Calcium 25 2.2.2 Strontium 26 2.2.3 Calcium Chloride 26 2.2.4 Calcium Bromide 28 2.2.5 Calcium Iodide 28 2.2.6 Strontium Chloride 29 2.2.7 Pure Iron 30 2.2.8 Gases 31

2.3 Dry-box. 32

2.4 Solubility measurements 46 2.4.1 Furnace assembly 46 2.4.2 Gas analysis apparatus 50 2.4.3 Initial experiments 53 Pau_

2.4.4 Segregation 59 2.4.5 Modified programme 64 2.4.6 Development of new method 65 2.4.7 Experimental procedure. 71

2.5 Differential thermal analysis. 76

2.5.1 The principle of the method 76 2.5.2 Initial experiments 78 2.5.3 Crucibles 80 2.5.4 Furnace and differential analysis assembly. 83 2.5.5 Experimental procedure 93

CHAPTER 3: RESULTS 99

3.1 Description 99 3.1.1 The Calcium-Calcium Chloride system. 99 3.1.1.1 Solubility measurements 99 3.1.1.2 Differential thermal analysis 102

3.1.2 The Calcium-Calcium Bromide system. 107 3.1.2.1 Solubility measurements. 107 3.1.2.2 Differential thermal analysis 107

3.1.3 The Calciuri-Calcium Iodide system 111 3.1.3.1 Solubility measurements 111 3.1.3.2 Differential thermal analysis 111 Page

3.1.4 The Strontium-Strontium Chloride system 114

3.1.4.1 Solubility measurements 114

3.1.4.2 Differential thermal analysis 115

3.2 Errors 119

3.2.1 Errors in the solubility measurements 119

3.2.1.1 Temperature measurements 119

3.2.1.2 Weighing 120

3.2.1.3 Gas analysis 120 3.2.1.4 Other errors. 120

3.2.2 Errors in the differential thermal analysis 121 3.2.2.1 Temperature measurements 121 3.2.2.2 Weighing 121 3.2.2.3 Other errors 121

3.3 Derived data 123 3.3.1 The heat of fusion of the investigated salts123 - 3.3.2 The monotectic compositions in the Ca Can2' 126 - Ca-CaBr2 and Ca CaI2 systems. nIAPTER 4: DISCUSSION 128 4.1 Comparison with previous work 12.8

4.2 Interpretation of the results 136

4.3 The nature of metal-molten halide solutions 142 ACKNOWLEDGMENTS 149 REFERENCES 150 LIST OF FIGURES Fig. no. Page 1. Phase diagram of the system Ca/CaF2. 15 2. Dry-box. 34 3. Glove port for dry-box. 35 4. Purification train for dry-box. 38 5. Equilibration furnace. 4.7 6. Gas-analysis apparatus. 51 7. Preliminary results in system Ca/CaBr2, 55 8. Photograph of section of quenched crucible. 60 9. Photograph of section of quenched crucible. 62 10. Crucibles for solubility measurements. 70 11. Glass envelope for premelting of salts. 73 12. Thermocouples for D.T.A. 77 13. Crucibles for D.T.A. 81 Y,-.1 Wilson seal. 86 15. Crucible arrangement for D.T.A. 87, Biasing circuit. 91 17. Temperature and differential curves. 97 18. Solubility versus tine in the system Ca/CaC12. 101 19. Phase diagram of the system Ca/CaC12. 106 20. Phase diagram of the system Ca/CaBr2 110 21. Phase diagram of the system Ca/CaI2. 112 22. Depression of freezing point of SrCl2 by Sr. 118 23. Depression of freezing point of Ca versus anion radius. 138 24. Electronegativity of anion versus consolute 141 composition.

CHAPTER 1 1. INTRODUCTION

1,.1 General Introduction. Molten salts have long been of importance for the electrochemical production of the alkali and alkaline earth metals and aluminium. In recent years the importance of the fused salts has increased with the interest in metals such as titanium, molybdenum, beryllium, zirconium, thorium and uranium. The rapid progress in the field of nuclear power generation is the main reason for the sudden great demand for these metals. The metals are produced either by fused salt electrolysis or by reduction of their halides or oxides with either 1.2.3. sodium, magnesium or calcium In the reduction processes the salts are either reaction products or added deliberately to act as a flux. The development of these methods has been largely empirical. Much funda- mental research is therefore needed for a fuller understanding of the principles underlying these processes. If e.g. uranium is produced by reduction of uranium tetra- chloride with calcium according to the reaction

UC1 + 2Ca -> U + 2CaC12 the activities of the calcium and the calcium chloride will be influenced by their mutual solubilities. A knowledge of the interaction between the calcium and its chloride in the 2.

temperature region where the reaction occurs is therefore important for the understanding of this reduction process. The solubility of metals in their fused salts is also of interest from the point of view of current efficiency in electrolysis. The current efficiency is often rather low in these fused salt electrolysEs and it was at one time thought that Faraday's law was not valid for these processes. However, work by Helfenstein4 b.nd Lorenz5'6 showed that this assumption was wrong and that 100% efficiency could be obtained if certain precautions were taken. The main reason for the low yield is that metal produced at the cathode dissolves in the salt and migrates to the anode or to the surface of the bath where it reacts. Even a very small solubility of the metal can thus cause great losses if stirring and convection currents allow the dissolved metal to react at the bath surface or the anode. To obtain further information for the theoretical background to processes such as those mentioned above, as well as to increase our knowledge of the nature of the solutions of metals in molten halides, research on metal-molten salts was started 3.

some years ago in this laboratory by Rogers? and Taylor8. This research is a continuation of their work and is complementary to work on electrical properties of metal-molten halides carried out in this laboratory9. The object of this work was to make an extensive study of the solubility of metal in the metal halide systems of those Group II:a metals which had not previously been studied in detail. Initial experiments on the calcium-calcium bromide and calcium-calcium chloride systems, previously determined in this laboratory, showed, however, the earlier results to be in error. These systems had therefore to be redetermined before further work was carried out. 4.

1.2. Relevant previous work on metal-molten salts.

1.2.1 General. 10 Since H. Davy's observation of the dispersion of potassium in molten potassium hydroxide in 1807 many investigations on the solubility of metals in molten salts have been carried out. idest of the early work was carried out by Lorenz and his co-workers at the beginning of this century and was largely qualitative and concerned with the question whether the dispersion of the metal in the salt was a conoidal suspension or a true solution. That part of their work carried out before 1926 is summarised in the book "Pyrosole" by Lorenz and Eitelll and in a chapter contributed by Lorenz to Alexander's book "Colloid Chemistry"12. According to Lorenz the metal dissolved in the melt in a colloidal form with a particle size > 10001. It is surprising that the colloidal theory could survive Aten's important observation in 1910 that the melting point of cadmium,chloride was depressed by the addition of cadmium, as this indicated a true solution13. Although Eitel and Lange14 placed doubt on the veracity of the colloidal theory, from their investigations of metal-salt melts with a high temperature 5.

ultra microscope, it was not until the work of Heymann and his associates that the colloidal theory was finally rejected. In their work on the distribution equilibrium of cadmium between molten cadmium chloride and a molten bismuth phase Heymann and Friedlander15 found a linear relation in the distribution of cadmium between the two phases. As bismuth did not react or dissolve in cadmium chloride, and because cadmium in bismuth was an atomic solution, this indicated that the solution of cadmium in cadmium chloride was also a true solution. The earlier investigations by Heymann and his co-workers also involved measurements of the solubility of cadmium in cadmium chloride at reduced vapour pressure of cadmium. Assuming the cadmium vapour above a molten cadmium-bismuth phase to obey Raoult's law they found the solubility of cadmium in a cadmium chloride melt in equilibrium with this vapour phase to obey Henry's law. Their work up to 1937 is summarised in a review by Heymann16 in which he also gives three possibilities for the solubility of cadmium in cadmium chloride, viz. Cd.nCdC1 1. Solvation: Cd + nOdCl2 2 6.

2. Bimolecular subchloride: Cd + CdC12 -> Cd2C12

3. Monomolecular subchloride: Cd + CdC12 -> 2CdC1

He points out, however, that m,asurements of the susceptibility of cadmium-cadmium chloride solutions by Farquharson and Heynann17 showed these solutions to be diamagnetic. As the monomolecular subchloride would be paramagnetic this mechanism for the solution has therefore to be ruled out. This possibility of a cadmium monomolecular subchloride has also been precluded as a result of 18 E.L.F. measurements by Karp achev and Stromberg . They constructed a concentration cell in Ihhich a graphite electrode was in contact with a solution of cadmium in a fused cadmium chloride-alkali chloride mixture. The other el,.ctrode Ivas in contact with a cadmium saturated cadmium chloride-alkali chloride mixture. The potential between these electrodes varied with the cadmium concentration according to the equation

E = 1E0 — RTnF In NCd where n is the number of faradays transferred in the half cell reaction. From a plot of the E.M.F. versus the concentration they obtained a value of n of 2 which indicated that cadmium was dissolved as a Cd atom or a 7

++ . ca.2 ion. The same value of n has also been obtained on -S.M.F. measurements in the cadmium-cad7ium chloride system by Crawford9 in this laboratory using a different cell arrangement. Grjotheim, Gronvold and Krogh-Moe19 repeated earlier work on the depression of the freezing point of cadmium chloride by cadmium and used a method that allowed very accurate results to be obtained. They could not detect any solid solution of cadmium in cadmium chloride and assumed that the solution was ideal and calculated the heat of fusion of cadmium chloride from the relation:

In N ++ AHf Cd n T Tf) where N Cd++ is the cation mole fraction of Cd++ in the mixture and T f and T are the liquidus temperatures of pure cadmium chloride and of the mixture respectively. They assumed three different models for the dissolution of the metal: 1. Atomic solution. 2. Formation of Cd2++ ions. 3. " Cd+ ions. 8.

As NCd++ varies depending on which of these models is chosen the calculated heat of fusion will also be different. Thu heats of fusion they calculated for these models were respectively 4.6 ± 0.3 Kcal/Mole, 5.4 ± 0.3 Kcal/mole and 10.4 Kcal/Mole. As the known heat tf fusion is 5.3 Kcal/Mole they ruled out the possibility of Cd+ ions being formed. Although the atonic solution cannot be excluded, they are in favour +-I- ions in the melt as this would of the formation of Cd2 20 also explain the phenomenon observed by Wirths 1 that electrolysis of saturated cadmium-cadmium chloride solutions can give a current efficiency exceeding 100%

if only Cd++ cations are considered in the transport process. As these metal-molten salt solutions are all intensely coloured, GrjotheiLl et al. also managed to show that the colour associated with the dissolved metal moved teiards the cathod in an electric field and thus is connected with positively charged particles. Corbett and von 'inbush21 measured the solubility of a number of metals in their salts. The methods they used were: A. Direct analysis of the quenched salt; B. Direct measurement of the loss in weight of 9. metal when equilibrated with a knm.n amount of salt at a temperature above the melting point of the metal; C. By weight of dissolved metal. Their results are given in Table 1. TABLE 1 Dissolved metal System Temperature °C moIT-7 Method - Sn SnC12 500 0.0032 - Sn SnBr2 500 0.068 Pb-PbC12 600 0.020 tl 700 0.052 If 800 0.123 Tl-T1C1 550 0.009 650 0.009 Ag-AgC1 490 0.03 700 0.06 - Ga GaC12 180 1.92 A Al- AlI3 380 0.0 A 423 0.3 A All the metal-salt solutions were intensely coloured except the gallium-gallium chloride and the aluminium- aluminium iodide which were colourless. For this reason they suggest that a subhalide in these two systems seems reasonable. They compare their results with results obtained in other systems and point out that there is in general a decreasing solubility with 1 0 .

increasing electronegativity of the metal. Using the same methods Corbett, von '!inbush and Albers22 later measured the solubility of the corresponding metals in lead iodide, antimony chloride, zinc chloride, zinc iodide, cadmium iodide and gallium bromide. From their results they favoured the formation of subhalides as accounting for the solubility in these systems. The sodium-sodium halide systems were investigated by Bredig, Johnson and Smith23. To prevent error in the results due to segregation during quenching they designed crucibles, made of stainless steel, in which they could separate parts of the phases at the equilibrium temperature. A part of the saltphase was sealed off by allowing a steel ball, initially retained in a sidearm of the crucible, to fall dovm and close off part of the bottom phase. The metal phase which floats on the salt was decanted into a sidearm by tilting the furnace. _hen the crucibles had been quenched the whole of these separated samples were analysed. Their results showed that each of these systems had a mono- tactic below 5 mole % sodium and a miscibility gap. They could not investigate the systems over the entire composition range but as the solubility increased rapidly with temperature they estimated from extrapolation a 11. consolute temperature between 1000 and 1100°0 in all these systems. 24 Bredig, Broniltein and Smith , using the same sampling technique,investigated the potassium- potassium fluoride and the cesium-cesium halide systems. The cesium-cesium halide systems did not exhibit any miscibility gaps but shoved instead a continuous decrease in the liquidus temperature on addition of metal to the salts. The potassium-potassium fluoride system had a miscibility gap with a consolute point at 910°C and about 30 to 40 mole % potassium. In a later investigation of the potassium' potassium halide systems Johnson and Bredig25 used a thermal analysis technique which permitted much more accurate results to be obtained. The potassiun- potassium fluoride system reinvestigated in this way had a consolute point at 904°C and 20 mole % potassium. The main features of the potassium systems are summarised in Table 2. 12,

TABLE 2 25 The K-K halide systems according to Johnson and Bredig. fa:atm K-KF K-KC1 K-KBr Consolute te:Iperature °C 904 790 728 717 rr compositio mole % K 20 39 44 50 Melting point of salt °C 858 769.5 734 681 Monotectic temperature °C 849 751.5 708 658.5 composition mole % K 5.5 11 18 13

Johnson and Bredig point out that the asy=etry in these miscibility gaps is reduced if the volume fraction is plotted instead of mole fraction. A component with a small molar volume should therefore be more soluble in the liquid phase of the other component than vice versa.

1.2.2 Alkaline earth metal-metal halide systems. Early work on alkaline earth metal-metal halide systems was mainly concerned with the formation of subhalides. Thus many investigators claimed to have prepared solid subhalides of these metals. However, no evidence has so far been found that these "subhalides" are anything but intimate mixtures of metal and the normal halide obtained on quenching from a temperature where the -]etal has an appreciable 13. solubility in its salt. As late as 1953 Ehrlich and Gentsch claimed to have proved the existence of a calcium monochloride26. X-ray investigations showed this 27 compound to have a metal-metal bond structure . However, they later published a paper in which they stated that the claimed CaC1 instead was CaHC1 produced by the mistake of passing hydrogen instead of argon through the furnace. One of the first quantitative investigations on molten alkaline earth metal halide systems was Zhurin's measurements of the solubility of magnesium in magnesium chloride and magnesium,chloride-alkali chloride mixtures29. He heated the mixtures in steel bombs for 21/2 hours and then quenched in water. The amount of metal in the saltphase was determined by measuring the volume of hydrogen evolved from the sample on addition of hydrochloric acid. He found the solubility of magnesium in pure magnesium chloride at 800°C, 1000°C and 1300°C to be 0.88, 1.23 and 1.68 mole % respectively. This work has recently been repeated by Rogers7 who found the solubility to be only 0.30 mole % at 800°C and 0.69 mole % at 1000°C. He explains that this discrepancy might be due to Zhurin's high magnesium oxide content in the salt, 14. which could lead to too high solubility values due to stabilisation of a magnesium emulsion. Rogers also investigated the influence of reduced activity of magnesium on its solubility in magnesium chloride. He found that the solubility was approximately proportional to the magnesium activity up to saturation. He also established the calcium- calcium fluoride phase diagram by differential thermal analysis and solubility measurements. The system was found to be completely miscible below the melting point of calcium fluoride, the melting point of which he found to be 1411°C. The consolute point of the shallow miscibility gap was at 1322°C and about 47 mole % calcium. The system has a monotectic at 1290°C and 25.5 mole % calcium and a eutectic at 821'(' and 98.6 mole % calcium. The melting point of his pure calcium was 837 ± 1°C. The system can be seen on Fig. 1. In his solubility measurements a small pure iron crucible containing the mixture was heated up inside a mild steel crucible which was sealed by arc- welding under a rough vacuum. After quenching in water the saltphase was exposed by sawing the crucible laterally. The sample was subsequently analysed for calcium by heating it up to 980°C under a vacuum of 15;

1500 x DIFFERENTIAL THERMAL ANALYSIS. ® SOLUBILITY.

1400\

Noe x`sc 1300 — \,1Pc --"`•

U• 1200— o

1100—

a. 2 1000-

900

x_x 800

700 0 50 100 COMPOSITION MOLE %Ca.

FIG.I. THE CALCIUM —CALCIUM FLUORIDE SYSTEM. (ROGERS) 16.

10 5 mm Hg and the weightloss measured. The vapour pressure of calcium at this temperature is about 10 mm Hg. He found that all the metal volatilised in two hours. For the differential thermal analysis he held the weighed mixture in a similarly welded mild steel crucible which was fitted with a thermocouple well. An identical empty crucible served as a thermal reference. As the melts always supercooled very much only heating curves were employed. To disperse the phases as much as possible before the experiment he first heated up the crucible above the estimated liquidus temperature and then rapidly cooled it down. Rogers? also made an attempt to measure the solubility of magnesium in calcium chloride at 1100°C. The results he obtained showed that the solubility is less than 0.05 mole %. Work by Heymann and Weber3° has also indicated that metals are only soluble in their own salts. They found e.g. that tin, bismuth, antimony and gold are not soluble in sodium bromide. However, if alloyed with sodium they pass into the saltphase together with the sodium, probably as intermetallic compounds. Cubicciotti and Thurmond31 investigated the calcium-calcium chloride, strontium-strontium bromide, strontium-strontium iodide, barium-barium chloride and 17. barium-barium bromide systems. The phase diagrams they have drawn from their measurements of these systems are all remarkably similar in nature. Each of them consists of one eutectic and one monotectic and a large miscibility gap which does not close within the temperature region investigated. In the calcium-calcium chloride system the highest temperature at which measurements were made was 1200°C, where the solubility of calcium in calcium chloride was found to be about 30 mole %. The solubility curve in this system passes through a minimum of about 15 mole % at 1000°C, and then increases again to about 20 mole % at 825°C, which is the monotectic temperature. The solubility of the salt in the metal is of the order of 1 to 2 mole % and does not change much with temperature. In their solubility measurements they contained the melt in a sealed pure iron crucible which was equilibrated in a furnace whose temperature was maintained constant within ± 10°C. After the crucible was quenched by dipping its lower part into water, it was cut into three transverse sections. The top section was analysed as the metal rich phase and the bottom as the salt rich phase, and the centre section was discarded. The 18. metal and the salt from the top and the bottom phases were dissolved out in water and analysed for halide ions. For the depression of the freezing point of the salt by addition of the metal, and vice versa, they employed a cooling curve technique with an open pure iron crucible containing the metal salt mixture. The thermocouple recording the temperature was fitted into a pure iron thermocouple well and placed in the centre of the melt. They found a eutectic at 767°C and 2 mole % Ca and a monotectic at 825°C and about 99 mole % Ca. 32 Schafer and Niklas re-investigated the barium-barium chloride system and found - contrary to Cubicciotti and Thurmond - that this system had a consolute point at about 55 mole % barium and 1010°C, above which there was complete miscibility. Their monotectic is at 15 mole % barium and 878°C with a solid solution of barium in barium chloride at this temperature of 7 mole %. In their investigation they used a thermal analysis technique which permitted accurate values for the phase boundaries to be obtained. According to these investigators the method used by Cubicciotti and Thurmond would cause the composition of the phases to change during quenching. For this reason not only the 19.

barium-barium chloride but also the other systems investigated by them would be expected to be in error. Taylor8 re-investigated the calcium-calcium chloride system as the phase diagram given by Cubicciotti and Thurmond seemed very doubtful from two points of view. The first was the change in sign of the temperature coefficient for the solubility of the metal in the salt at about 1000°C, and the second was the large discontinuous change in slope of the solubility curve at the monotectic temperature. This change should be rather small as it corresponds to the heat of fusion of the calcium which is only about 2 Kcal/Mole. In his investigations Taylor equilibrated the salt-metal mixtures in a pure iron crucible held inside a mild steel crucible which was sealed with a screw lid. After equilibration, the crucible assembly was rapidly quenched in water and the two phases in the inner crucible analysed. To make sure that no segregation occurred during quenching Taylor divided the saltphase into several sections, which he analysed. The differences he obtained were negligible. He also checked his rosults at 1000°C and 1200°C by a high temperature sa:pling technique similar to that used by 23 Bredig, Johnson and Smith . For the determination of 20. the depression of the freezing point of the salt he used a thermal analysis method. The metal-salt mixture was loaded into an iron crucible provided with a thermocouple well. The crucible was sealed by arc-welding under a vacuum of 0.1 mm Hg. As the melt supercooled very much all his measurements were made on heating curves. His phase diagram has a eutectic at 5.25 mole % Ca and 755°C with a solid solubility of calcilm in calcium chloride at this temperature of 0.6 mole % Ca. Table 3 gives his values for the solubility of the metal in the salt. TABLE 3 Solubility of Ca in CaC12 according to Taylor. Temperature °C Solubility mole % Ca 900 6.992 950 7.15 1000 7.475 1050 7.46 1100 8.10 1150 8.74 1200 9.22 1250 9.82 The monotectic in his diagram is at 0.6 mole % CaC12 in Ca and 838°C which is the same as the melting point for his pure calcium. The solubility of the salt in the metal increases with constant slope to 1.75 mole % at 1250°C. 21.

Using the same methods, Taylor also investigated the calcium-calcium bromide system. He found the eutectic to be at 4.5 mole % Ca and 729°C, with a solid solubility of 0175 mole % Ca in CaBr2 at this temperature. No solubility of calcium bromide in calcium was detected between 900°C and 1150°C. His solubility values for calcium in calcium bromide are given in Table 4. TABLE 4 Solubility of Ca in CaBr2 according to Taylor. Temperature °C Solubility mole % Ca 900 8.45 950 9.05 1000 9.7 1050 10.35 1100 11.0 1150 11.6 A check on the solubility of calcium in calcium chloride was made at 1100°C by Rogers7, who varied the activity of the calcium by alloying it with aluminium. He found that Henry's law was obeyed up to a calcium activity of one. He found the solubility at this activity to be 8.41 mole % Ca which is to be compared with Taylor's 8.10 mole %. Measurements of the solubility of calcium in 22. calcium chloride haVealso been carried out by 0stertag33 in her investigation of the equilibrium. Ca + MgCl2 < Mg -1 CaCl2

She only measured the solubility of the pure calcium in the pure calcium chloride at 1000°C, at which temperature her values vary from 11.2 to 17.8 mole % Ca. The metal salt mixture was kept in a welded iron crucible. To obtain equilibrium between the two phases she agitated the crucible for 1 hour in a horizontal position in a furnace maintained at constant temperature. Separation of the phases was achieved by holding the crucible in vertical position at the same temperature. The crucible was thereafter quenched in water. In a recent investigation of the equilibrium of barium with calcium chloride Peterson and Hinkebein5 also checked the binary systems b-,rium-barium chloride and calcium-calcium chloride. Their results from thermal analysis on the barium-barium chloride system agreed very well with the phase diagram given by Schafer and Niklas32, while solubility measurements on the calcium-calcium chloride system at 900 and 950°C gave 3.8 and 4.2 mole % Ca in the salt phase and 1.02 and 1.56 mole % CaCl2 in the metal phase respectively, 23.

They also reported a monotectic at 820°C and 99.5 mole % Ca and a eutectic at 768°C and 2 mole % Ca. Their procedure for the solubility measurements was similar to that used by Ostertag. 214,,

CHAPTER 2 EXPERIMENTAL

2.1. Programme. To obtain further information, concerning the mechanism of solution of metals in their molten salts, it was decided to make solubility and depression of freezing point measurements in the metal halide systems of those Group II:a metals which had not previously been studied in detail. It was decided to start with solubility measurements in the strontium-strontium chloride and calcium-calcium iodide systems, adopting a technique previously used in this laboratory by Taylor. To check the experimental procedure it was planned to carry out the initial experiments on the calcium-calcium bromide system previously investigated by Taylor.

25.

2.2. Preparation and purity of materials.

2.2.1 Calcium. The calcium metal was given by the U.K.A.E.A. Risley, and was in a granular form. It was supplied with the following impurity figures:

Cl Mg C Na Percent by weight 0.2 0.02-0.08 0.05 0.02

Parts Si Al N Ni Mn Fe Li Cu Cr Cd per million 20 <20 <20 19 11 8 5 3 2 0.05 <0.03

Eu Sm Gd Total rare earths <0.05 <0.04 <0.03 <0.4

The melting point of this calcium -was found to be 838.5 ± 1.0°C. This is about 12pC lower than the accepted value. From the impurity figures one would not expect this metal to have a melting point more than 5°C too low. The oxygen impurity figure is not given, however, and the metal was, for this reason, redistilled under vacuum to exclude possible oxide contamination. This treatment did not alter the melting point. The metal was kept in a closed glass tube, filled with argon. Only bright, clean grains of the metal were selected for the experiments. 26.

2.2.2. Strontium. The strontium metal was supplied by New Metals and Chemicals Ltd., who gave the following analysis in weight per cent: Sr metal Fe N Al Alkaline metals 99.85 0.09 0.02 trace nil However, a cooling curve of the metal gave a slight kink at 769°C and a definite break at 719°C as compared

to the accepted melting point of 770°C_ This indicated thp4- the metal probably was less pure than stated in the given analysis. A qualitative test for nitrogen showed this element to be present to an extent probably greater than that shown in the analysis. Reliable results • could thus only be obtained with this metal when it was used in small concentrations, where it was thought that this impurity would have little effect. The metal was kept in a closed glass jar filled with argon, and it was only handled in the argon atmosphere of the dry-box. It was supplied in the form of a cast bar 1" in diameter. Pieces of this bar were cut off with a hacksaw and cleaned with a file before use. 2.2.3. Calcium chloride.

Dried analar calcium chloride was supplied by Hopkin and Williams Ltd. It contained about 25 wt.% 27. of water, which was removed by heating in vacuum. The maximum limits of impurities in weight per cent were:

wt. per cent Sulphate 0.01 Nitrate 0.004 Phosphate 0.002 Silicate 0.004 Heavy metals (Pb) 0.002 Iron 0.001 Barium and strontium 0.02 Alkalis (Na) 0.06 Arsenic (As203) 0.0002

The dehydration was carried out in a glass envelope of the same shape as the one seen in Fig 11. At each time a batch o about 40 grams of dried calcium chloride was dehydrated. Taylor found that by heating the salt up to about 400°C over about half an hour under vacuum and leaving it at this temperature for 4 hours, the salt could be dehydrated without any hydrolysis occurring. This method was adopted although the temperature of the salt was raised more slowly. A temperature of 400°C was reached after about 5 hours and then the salt was left at 400°C overnight. The hydration water was always collected in a liquid nitrogen trap and its pH tested with B.D.H. lithmus paper to check that no 28. hydrolysis had occurred. The dehydrated salt was kept in the closed glass envelope in the dry-box.

2.2.4. Calcium Bromide. The calcium bromide could not be obtained in an analar grade. A pure grade supplied by May and Baker Ltd. contained, however, no chloride and no oxide. Taylor found on analysis that this salt contained 79.8% bromine and 19.92% calcilm compared to theoretical ?9.96% bromine and 20.04% calcium. This salt was dehydrated in the same way as the calcium chloride, without any oxybromide being formed.

2.2.5. Calcium iodide. The calcium iodide was supplied by Whiffen and .moons Ltd. It was manufactured from ferrous iodide and calcium hydroxide. The following analysis was given: Iodate nil Iron 5 P.P.M. Lead nil Silica <10 P.P.I. 85.5 wt. per cent. CaI2 the remainder being water.

The salt was supplied in the form of white co slightly yellowish white pieces. These were broken up 29. and dehydrated in a glass envelope under vacuum. The temperature of the salt was raised slowly over a period of 8 hours to 120°C and held there for 10 hours. The major part of the water was expelled at this temperature. It was then slowly raised to 400°C and held at this temperature for 10 hours. The dehydrated salt was tested for water with Karl Fischer reagent but none could be detected. To ensure that no oxyiodide was formed during dehydration and premelting, 10175 mgs of the premelted salt was dissolved in water to 10 c.c. , and the solution titrated with 0.014-N HCl. 2.8 c.c. was needed for neutralisation. A second solution of the hydrated salt was made to have the same concentration of iodide. In this case 2.7 c.c. of the same HCl solution was needed for neutralisation. This corresponds to 0.02 mole % oxide being formed.

2.2.6 Strontium Chloride. Analar hydrated strontium chloride SrC12.6 H2O was supplied by the British Drug Houses Ltd. with the following maximum limits of impurities: 30.

Wt._2er cent. Sulphate (SO4) 0.005 Nitrate (NO3) 0.002 Heavy metals (Pb) 0.001 Iron 0.0005 Barium 0.2 Calcium 0.1 Alkalis 0.03 The salt was easily dehydrated in the same way as the calcium chloride and was found to be the least hygroscopic of these salts.

2.2.7 lure Iron. Swedish Remko iron supplied by Ernest B. Vestman Ltd. was used for all the reaction crucibles. It was found not to be attacked under the conditions of the experiments, while e.g. nickel was heavily attacked by both calcium and strontium at 1000°C. The following analysis was given:

Wt. per cent. Carbon 0.03 Silicon 0.01 Manganese 0.12 Phosphorus 0.004 Sulphur 0.010 31.

2.2.8. Gases. The gases used were argon, hydrogen and forming gas. They were all supplied by British Oxygen Gases Ltd., who gave the following analysis. Argon 99.95cL Impurities in P.P.M. by volume

N2 02 CO2 H2 500 5 5 5 H2O < 0.1 g/M3 withdrawn from full cylinder (200 cu.ft.) H2O will rise to 0.5 g/m3 at the end of the cylinder. When this argon was used for experiments at high temperature it was always first pdrified. The water was removed by silica gel and magnesium perchlorate and the other impurities by passing the argon through a vertical stainless steel tube furnace containing calcium turnings. The hottest part of this furnace was at a temperature of 680-700°C. High_2urity hydrogen Impurities in P.P.M. by volume

02 N2 CO2 CO 500 500 10 5

Hydrocarbons estimated as CO2 5 P.P.M. H2O < 0.15 g/M3 rising to 1.0 g/m3 towards the end of the cylinder. 32.

Forming gas was only used for the protection of molybdenum from oxidation. It contained approximately 10% hydrogen and 90% nitrogen.

2.3. Dry-box. Experimental procedure in this laboratory involving the handling of calcium and its halides was earlier carried out in a standard commercial dry-box, made by Towers of. Widnes, and filled with dry nitrogen. Although this was sufficient foi. handling calcium and its chloride and bromide for short times it did not give the necessary protection for strontium which was rapidly attacked. It was therefore decided to build a dry- box with a circulating argon atmosphere that could be continuously purified. In this way, oxygen, nitrogen wad water vapour, entering the box by diffusionl couldsbe kept to such a low level that the reactive materials could be handled for a reasonable time, with only negligible attack. The size of the dry-box was chosen so that every corner of the bottom could be reached by the gloves, yet the pressure changes due to the movement of the gloves did not cause too much resistance. A suitable size was found to be 2' x 3' x 2' high. The box, which 33.

can be seen in Figure 2, was built of 3/8" Perspex sheets. They were glued together with Perspex cement and the joints reinforced with screws threaded into the Perspex. The Perspex walls ensured that the interior of the box was well lit and easily visible, although they provided large areas for diffusion of gases into its inert atmosphere. The bottom of the box was covered with an aluminium tray. The front wall of the box had two glove ports 6" in internal diameter, whose centres were 19W apart and 8" from the bottom of the box. The ambidextrous gloves used were 15 to 30 thou t'_ick and made of latex or neoprene. They were of shoulder length and beaded at the shoulder so that they could easily be clipped onto the double-grooved ring of bakelite which was attached to the Perspex around the ports. The arrangement can be seen on Figure 3. The main contamination of the argon atmosphere seemed to be caused by diffusion through the glove's. It was therefore essential that the gloves were not in contact with the argon longer than necessary. The glove ports were therefore sealed off with screwbungs when the dry-box was not in use. These were made of aluminium and provided with an 0-ring seal. See Figure 3. The gas trapped between the gloves and the screwbungs Fig. 2 Dry-box 35. PERSPEX WALL. Q-COMPOUND. STEEL CLAMP. RUBBER GASKET. RUBBER 0-RING.

4 RUBBER PROTECTION. GLOVE.

TUBE FOR EVACUATION OF THE GLOVE. NOT TO SCALE.

FIG.3. GLOVE PORT FOR DRY-BOX. 36. was always removed by evacuation before the ports were opened. The gloves, when not in use, were rolled up against the screwbungs and covered with pouches made of an old pair of gloves cut off and sealed. These pouches contained silica gel which kept water vapour away from the gloves. Materials 'v ere transferred into the box via an ante-chamber system, which could be evacuated and then filled with argon from the box. This system consisted of two cylindrical ante-chambers made of brass, each provided with two screwbungs of the same type as those used for the glove ports. The larger chamber was 10" long and 5 3/4" in internal diameter, the smaller was 5" long and 21/2" in diameter. The larger ante-chamber had a brass flange at one end, which was fitted around a hole in the side wall of the box. The smaller chamber was attached to the larger one and could therefore only be used when the larger one was filled with argon and open to the dry-box. Then the dry-box was not in use it was always maintained at a pressure slightly greater than atmospheric. This pressure was measured on a dibutyl phthalate manometer attached to the side of the box. On introducing the operator's hands into the box the pressure rises due to the decrease in volume. If this pressure rise were 37. not to cause too great a resistance to entry some of the argon in the box had to be removed. The size of the larger antechamber was so chosen, however, that its volume was slightly less than that of the inflated gloves. On most occasions when work was commenced it was necessary to transfer objects into the dry-box via the large ante-chamber thus no adjustment for pressure was needed as the evacuated ante-chember was filled with argon from the box. The argon gas entered the box through holes in a one inch brass tube slightly shorter than the box and placed along the bottom of the front wall. The holes - 31 in number and 2 mm in diameter - were evenly spaced along the tube, and directed into the corner between the bottom and the front wall. A similar tube to remove the gas was placed in the far upper corner of the box. The tubes were fitted to the Perspex wall with rubber 0- ring seals. The purification train for the argon can be seen on Figure 4. The separate parts of this train were connected with glass tubing., joined together by small pieces of thick rubber or P.V.C. tubes. At these points screw clips could be attached, as can be seen on Figure 4. The continuous circulation G

ARGON H ‘If FROM DRY- ' '------BOX

WATER #OUT • /( PURIFIED ARGON r TO DRY-BOX. WATER IN. • CA FIG.4. CIRCULATION AND PURIFICATION OF ARGON FOR DRY-BOX. CP 39.

KEY TO FIG. 4

Oil trap. B. Cotton wool filters. C. `unification furnace. D. Condenser. E. Bronze filter. F. Pressure switch. G. Pressure switch. H. Tap for inlet of argon. 40. of the gas through the dry-box and the purification train was first maintained with a neoprene diaphragm pump. This seemed, however, to be too weak for continuous running and was replaced by a refrigerator pump, which had roughly the same capacity - about 26.5 cu.ft./hr. The atmosphere of the box was thus changed slightly more than tTice an hour. This pump was oil sealed and the gas was therefore slightly contaminated with oil after it had passed the pump. To remove the oil mist from the argon an oil trap A and two filters B, which consisted of two large columns of hard-packed dried cotton wool, were found to be effective. The gas then passed through a vertical purification furnace C. The furnace consisted of an aluminous porcelain tube wound with a 20 gauge Kanthal wire. The resistance was 45 ohms. The tube was placed in a Steel casing and was surrounded with alumina powder for thermal insulation. The reaction tube of this furnace consisted of an 18 gauge 30 inch long stainless steel tube filled with about 500 gms of calcium turnings. The tube had an external diameter of 52 mm. On the top of the stainless steel tube was brazed a brass cone and a cooling coil of copper tubing. A B-50 standard glass socket was joined on to this cone with picein 41.

In the bottom of the tube was fitted a rubber bung through which a glass tube was passed. Tio.y bottom end of the reaction tube was kept cool by a lead cooling coil. The gas was let in through the bottom of the furnace and passed upwards through the calcium turnings which were held at a temperature of about 68000 in the hottest zone. It was found that a vertical furnace was superior to a horizontal one, where a free path always developed above the calcium turnings. After the impurities - oxygen, nitrogen and water vapour - had been removed by reaction with the calcium the gas was cooled down in the condenser D. Any entrained calcium oxide and calcium nitride were filtered off by the filter E before the gas entered the dry-box. The filter was a sintered bronze disc 2 mm thick and 76 mm in diameter. It filtered off particles bigger than 7 to 15 11. If the gas for some reason became heavily contaminated the heat evolved on reaction between the impurities and the calcium raised the temperature in the furnace so high that the calcium turnings agglomerated and caused obstruction in the gas floc.. The increase in pressure between the pump and the obstruction frequently damaged the whole dry-box system. For this reason the mercury switches F 42.

and G were introduced. They consisted of U-tube manometers made of glass and containing mercury. A tungsten electrode was sealed into the bottom of each U-tube. On a set distance from the surface of the mercury another electrode was placed so that if the pressure increased to this level, contact between the two electrodes was made. This caused a Londex relay to switch off the power supply to the pump and the furnace. Argon could be withdrawn from, or let in to the system through tap H whenever the pressure in the dry-box needed adjustment. This tap was also used when the gas train was evacuated prior to refilling. When the dry-box was to be filled, the air was first displaced by inflating a meteorological balloon, placed inside the box, with argon, until it virtually filled the box. Only a little free space was left in the corners and round the inlet and outlet tubes. The balloon and the dry--box were then closed and the free space flushed with argon via the inlet and outlet tubes. The balloon was then deflated by letting the argon it contained go into the box. The purification train of the system was easy to fill as it could be evacuated. The effectiveness of the dry-box for protection of the materials handled was indicated by the 4.3. time a piece of strontium stayed bright inside the box, i.e. before any interference colours appeared on the scratched surface of the metal. This time was of the order of 1 hour when the box was in good working condition. The expected amounts of oxygen, nitrogen and water vapour diffusing into the box were calculated from estimated values of the permeability constants for Perspex, given by the manufacturer, Imperial Chemical 35 36 37 Industries Ltd. and for rubber from Barrer and Buddery. As can be seen from Table 5 , the amounts of oxygen and nitrogen diffusing into the box were very small compared with the amount of water vapour. TABLE 5 Diffusion of oxygen and nitrogen into the dry-box through Perspex walla and rubber gloves at room temp. Surface areas: Perspex 19500 cm2, gloves 4900 cm2

11 ft Thickness: 9.5 mm 9 0.75 mm. Permeability constant Diffusion in cc x mm/cm2 x sec. x cm.Hg. cc/hour Perspex Gloves Perspex Gloves -10 -6 -2 Oxygen 10 0.020 x 10 10 7.5 -2 Nitrogen 10-10 0.0071 x 10-6 10 10 Water 3 x 10-6 1.6 x 10-6 20 at a 45 at a Vapour relative relative humidity humidity of 50% of 50% 44.

The relative humidity inside the gloves rose immediately to 100 per cent when the operator's hands were introduced. For this reason double pairs of gloves were always used. A quantitative test for water vapour was made by measuring the weight increase of a glass dish containing phosphorous pentoxide placed inside the box. The dish was covered with a watch glass during the weighing and transfer outside the box.

TABLE 6 Water vapour measured by pick up by phosphacous pentoxide. cc/hour Purification furnace on: glove ports closed 3

It off: If If

tf If on: glove ports open but 8 no hand in the gloves

As can be seen from the table, the measured amounts of water were much smaller than those calculated. The reason for this might be that the values for the permeability constants were uncertain and the relative humidities were not known. lithether the purification furnace was on or off did not very much alter the pick up of water vapour by the phosphorous pentoxide when the glove ports were closed. However, if the difference in 45. this pick up was assumed to be due to the complete removal of the water vadour from the argon when it passed the purification furnace, the concentration of water vapour in the argon when it left the box can be calculated. The volume of argon that passed the purification train was 7.5 x 105 cc/hour and consequently the concentration of water vapour in the argon was 1 x 106 1.3 P.P.M. by volume, when the 7.5 x 105 glove ports were closed. 46.

2.4. Solubility measurements.

2.4.1. Furnace assembly. The equilibration of the crucibles for the solubility experiments was performed in a vertical Kanthal wound furnace. The arrangement is shown in Fig. 5. The heating element was a Kanthal A 18 gauge resistance wire. It was wound on to a mullite tube which was 600 mm long and 60 mm in external diameter. Approximately 115 turnings of wire were used, giving a resistance of 26 ohms at room temperature. The wire was wound on the tube in such a way that the turns were closer together towards the ends of the tube tir,n in the middle. This had the object of producing a longer zone of constant temperature than would be obtained with a uniform The winding was kept in place by a coating of refractory cement and the tube was encased in a steel cylinder. Syndanyo plates formed the ends of the cylinder which was filled with alumina powder. The furnace tube was kept in position in the casing by means of a layer of asbestos string wound on to the tube, which was then forced into the hole cut in the Syndanyo plates. The inner mullite reaction tube was 900 mm long and 38 and 45 mm in internal and external diameter 47.

ARGON OUT 7

DIBUTYL PHTHALATE COOLING 0 0 COIL 0 0

COOLING o 0 0 COIL 0 .$. ••lrk. NOT TO SCALE.

ARGON VACUUM -61- THERMOCOUPLES FIGS. FURNACE FOR SOLUBILITY EXPERIMENTS. 48. respectively. It was fixed in position with asbestos string and supported by a brass ring at the bottom end. Both ends were cooled by water passing through lead cooling coils. The rubber bungs at either end of the reaction tube were protected from radiation by shields made of porous refractory bricks. The bricks were attached to the bungs with iron wires. A satisfactory constant temperature zone was obtained by putting a 15 cm long steel tube in the hot zone of the furnace. This steel tube was supported by means of a length of mullite tube and a refractory brick. A platinum/platinum 13;o rhodium thermocouple in a mullite sheath was passed through the bottom rubber bung into the centre of the furnace. The mullite sheath was so situated that the tip of it just touched the bottom of the crucible which was suspended by a Kanthal wire from the top bung. A nickel radiation shield consisting of three nickel discs separated by small alumina tubes was suspended on the wire above the crucible. The power to the furnace was supplied through a Variac auto-transformer, and the temperature controlled by a Kent Multilec potentiometric controller which operated a Eunvic vacuum switch. This switch was connected across the ends of a resistor in series with the furnace winding. -,.hen the furnace temperature fell below the temperature pre-set on the Kent the circuit through the controller closed and the heater in the vacuum switch became hot. This caused the switch to close and hence the furnace current to by-pass the resistor. The resistor was so chosen that it reduced the furnace current by about 20% when in circuit. The E.M.F. to the Kent was supplied by a platinum/ platinum 13% rhodium thermocouple, whose junction was placed between the reaction and furnace tubes in the hottest part of the furnace. The power input to the furnace was adjusted so that the on and off times of the Kent were equal (about 2 minutes at 1000°C). The variation of the temperature with time in the centre of the furnace was found to be 10.25°C at 1200°C. The inert atmosphere was maintained inside the furnace by passing purified argon in at the bottom and out at the top via a bubbler containing dibutyl phthalate. The way the argon was purified has already been described under gases. 50.

2.4.2. Gas analysis ap-oaratus. A gas volumetric techni4ue for the determination of alkaline earth metals dissolved or mixed in salts was 38 earlier developed in this laboratory. This method was used for the determination of calcium and strontium in their halides. The metal dissolved in the salt is determined by measuring the volume of hydrogen given off when a sample of known weight is dissolved in hydrogen saturated water. The method is simple and has the advantage of measuring directly the excess quantity of the metal. A diagram of the apparatus used is shown in Figure 6, and the analytical procedure is as follows. The two reservoirs A and B are filled with distilled water. The water in A is saturated with hydrogen by bubbling the gas through it for 20 minutes just before an analysis is carried out. The bores of tape C and D are thereafter filled with the water by applying pressure with a blow bulb on the reservoir A. These taps are then closed ante the burettes E and F evacuated through the tap H, the taps L, I and J being closed. The sample tube containing the sample and a small magnetic stirrer is then joined to the socket K and the tap J opened to evacuate the sample chamber.

TO AIR TO VACUUM 51. TO AIR a1:111 H r X _7 4

F 2m1 BURETTE E 50m1 BURETTE G

WATER BATH

FIG. 6. GAS-ANALYSIS APPARATUS. 52.

When sufficiently evacuated, the tap J is closed and tap I and H opened to the atmosphere. Water from A and B is then forced up to the zero marks on the top of burette E and F; taps I and H being open to air. After tap I is closed, water is allowed to flow on to the sample througla tap D, which is then immediately closed and tap J opened. The evolved hydrogen rises into the burettes and adhering bubbles are released by agitating the stirrer by means of a magnet. The water in burette E is then adjusted to an appropriate integral ml level by adjusting the pressure on bulb A, and the level in burette F and tube G are made equal by adjusting the pressure on bulb B. The gas volume giver off can now be obtained by adding the readings on E and F. The atmospheric pressure and the temperature of the water bath surrounding the burettes are read, and the gas pressure corrected for saturation vapour pressure of water. The apparatus was checked against weighed pieces of magnesium metal. These pieces were analysed in the apparatus using dilute hydrochloric acid. The apparatus has been shown to be accurate to ± 0.1%, with a standard deviation of a single measurement of

- 0.7%. 53.

2.'+.3. Initial experiments. In determining the solubility of a metal in its molten halide it has been shown that great care should be taken in order to obtain a true value of the solubility. For instance, cooling an equilibrated mixture of calcium and calcium chloride from a temperature where two immiscible liquid phases exist will cause the compositions of the two phases to change due 8 to segregation. Taylor found, however, in his investigation of the calcium-calcium chloride and calcium- calcium bromide systems that a rapid water quench of the equilibrated crucible was apparently sufficient to prevent segregation. The preliminary experiments were intended to be a check on the experimental procedure. As will be clear from the following these experiments failed, however, to reproducible results. Every effort was therefore made to improve on the procedure and the techniques until at last it was found to be radically wrong, and had to be abandoned.

The first results shown below in Table 7 were all obtained from experiments in which three thin pure iron crucibles were held inside a mild steel screw lid crucible. After equilibration and ouenching in 54. water the bottoms of these inner crucibles were sawn off and samples of the salt phase crushed out and analysed for metallic calcium by the volumetric method already described. TABLE 7 Solubility of Calcium in Calcium bromide initial aperiments.

Ca mole % Temp. 'C. in CaBr2 Equilibrium time

809 5.6 12 hr. 55 min. 839 6.5 47 hr. 6 min. 839 6.7, 6.4, 5.4 20 hr. 5 min. 856 5.9 5 hr. 2 min. 8?9 5.2, 5.5, 6.1 16 hr. 43 min. 885 3.5, 3.8, 3.9 18 hr. 10 min. 896 5.9 45 hr. 10 min. 896 5.8, 5.9, 6.1 72 hr. 23 min. 986 3.5, 2.9, 3.8 16 hr. 35 min. 986 4.q-, 4.0, 3.9 16 hr. 0 min. The results from Table 7 are plotted on Fig 7. As can be seen these results are much lower than those obtained by Taylor, and the scatter is very large. At this stage it was thought that the discrepancy was due to contamination of the melt during the equilibration caused by leakage through the outer 55.

0 TAYLOR. X THIS WORK,

1000-

X X XXX x

900-

X xx

0 TURE

ERA 800 TEMP

700 0 5 10 MOLE % Ca IN Ca Br2

FIG.7. SOLUBILITY OF CALCIUM IN CALCIUM-BROMIDE INITIAL EXPERIMENTS, COMPARISON WITH TAYLOR.

nen twora,droarnira•••••••.,..q•-“,, 41.s...m. env., ...••• • ...1•P 564

steel crucible. The outer crucible was therefore redesigned, to give a tighter fit, and only one, slightly larger, inner crucible was used. The following experiments were all carried out in purified argon. In spite of these alterations the solubility figures still remained lower than Taylor's, as can be seen from Table R„ TABLE 1 Solubility of calcium in calcium bromide initial experiments.

Ca mole % Tem2_°C in CaBr2 Equilibrium time

893.0 3.28, 3.28 18 hr. 6 min. 975.0 3.61, 3.55 16 hr. 40 min. 980.0 2.99 19 hr. 6 min. 985.0 3.46, 3.46 16 hr. 40 min. 1088.4 3.46, 3.0 18 hr. 11 min. 1184.7 3.83 5 hr. 12 min. To eliminate any possibility of contamination and at the same time increase the quenching rate it was decided to use only one crucible which was spotwelded under vacuum. The first run of this type gave a solubility at 1093°C of 7.08 mole % Ca which, although higher than the earlier values, is still lower than 57.

Taylor's. Three further runs were carried out at 1007°C using different amounts of Ca metal loaded into the crucibles. The amount of salt was the same and approximately 600 mgs. TABLE 9 Solubility of calcium in calcium bromide,initial experiments. Ca mole % Amount of Ca metal, mgs. Temp. °C in CaBr2 1007.0 8.5 13.8 1007.0 5.11 26.1 1007.0 7.47 43.0 As the scatter here is rather large it was thought that impurities in the Ca metal could have partly caused this discrepancy. The metal was therefore purified by distillation. With this purified calcium the following results were obtained. TABLE 10 Solubility of calcium in calcium bromide,initial experiments. Ca mole % aluilibrium time. Temp °C. in CaBr2 898.0 4.06 0 hr. 42 min. 899.0 4.45 1 hr. 12 min. 899.0 4.38 2 hr. 11 min. 900.0 4.74, 4.96 17 hr. 40 min. 902.0 4.3 19 hr. 50 min. 58.

As these results were still inconsistent with those of Taylor it vas decided to make some measurements on the calcium-calcium chloride system at 1100°C, as comparison of the results with those obtained by both Taylor and Rogers7 would serve as a check on the techniques employed. In this system Taylor found a solubility of 8.1 mole Ca at 1100°C, while Rogers at the same temperature obtained a value of 8.4 mole % Ca using the same technique. The followin[ results were obtained in the calcium-calcium chloride system: TABLE 11 Solubility of calcium in calcium chloride at 1100°C, Initial experiments. Temp °C Ca mole % Equilibrium time. in CaCl2

1100 4.8 6 hr. 50 min. 1100 3.8 23 hr. 30 min. 1101 6.11 3 hr. 11 mi'a. 1100 6.25 3 hr. 15 min. 1100 6.78 8 hr. 4 min. 1100 6.8 2 hr. 50 min. 59.

The first two of these experiments were carried out using a small pure iron crucible inside a steel screw lid crucible while the last four were performed with a single spotwelded pure iron crucible. The first two experiments show a significant lower solubility. The only possible explanation for this discrepancy appeared to be the different quenching rates employed and accordingly the effect of quenching rate on the segregation of calcium was studied.

2.4.?]-. Segregation. In the study of segregation a metallographic technique was employed in which the crucible was sawn in half along its axis and then carefully polished on emery paper under liquid paraffin. The first crucibles were quenched in water and showed small globules of calcium in the CaCl2 matrix very near the boundary with the calcium. To find out whether this was due to mechanical mixing, another crucible was quenched at a much faster rate, using brine as the quenching medium. Fig 8 shows the section. The crucible was short and wide to give a large thermal gradient across the salt phase during quenching. This represented optimum condition for segregation with a given quenching rate. As can be seen from the figure,

60.

calcium globule

Salt phase

FIG. 8 Section of crucible containing Ca/CaC12 mixture quenched from 1100°C in a NaCl water solution. (magnification 5x). 61.

a small globule of calcium is present on the top of the salt phase, some distance from the rest of the calcium, which is mainly present in the acute angle of the welded seal. As this globule may not necessarily have been formed by segregation, a narrower and longer crucible was prepared. This crucible was quenched in a 10% solution of NaOH which, according to Seybolt and 39 Burke, should have given a still higher quenching rate. The outside of the crucible was actually quenched from 1100°C to below red heat within approximately one second. Fig 9 shows the section of this crucible. As can be seen, the calcium metal has solidified very rapidly forming a bridge across the crucible and the salt has solidified more slowly, as one would expect from the difference in their thermal conductivities. A well defined shrinkage cavity has been produced in the top of the salt phase. At the bottom of this cavity, a large globule of calcium may be seen and it is apparent that this calcium must have come out of solution during the quenching, after the main body of the calcium had solidified. This would imply that samples taken from the lower part of the crucible would produce solubility figures which were much too low. This effect would be increased by slow quenching rates, such as those employed

62.

Calcium

Void.

Segregated Calcium

Salt phase

FIG. 9 Section of crucible containing Ca/Ca012 mixture quenched from 1100°C in a 10% NaOH water solution (magnification 5x) 63.

by Taylor and Rogers, because the small crucible containing the metal salt mixture, was contained in an outer crucible and thus cooled largely by radiation. This would decrease the temperature difference between the metal and the salt phases, allowing segregated calcium to re-enter the molten-metal phase. A possible explanation for the fact that Taylor and Rogers obtained consistently higher values than in this work seems to be that in their analysis they must have crushed out and analysed the whole salt phase, while in this investigation only the lower part of the salt phase was analysed. If one could be sure that during the quench no segregated calcium had re-entered the main calcium metal phase, analysing the whole salt phase should give the true value of the solubility. For this reason one experiment was carried out at 1250°C in which the calcium-calcium chloride mixture was held in a long and thin pure iron crucible, spotwelded under vacuum. After esiuilibration the crucible was ra-,idly quenched in a 10% LiC1 water solution. The whole salt phase was crushed out and analysed and gave a solubility of 17 mole % calcium. Taylor gives at this temperature a solubility of 9.82 mole %. From this it appears that in his case part of the segregated calcium had re-entered 64. the main metal phase. The microscopic investigation of the crucible on Fig 9 also showed that the size of the calcium particles precipitated from the salt phase increased towards the centre of the crucible, i.e. the particle size increased with decreasing cooling rate. It appears from this that the particles were not present as an emulsion, but were deposited from the solution during cooling. The salt phase from the other half of this crucible was divided into four sections and analysed for free calcium. The following figures were obtained, starting from the bottom: i+.99; 3.74; 4.42; and 31.7 mole % These four samples together give a weighted average solubility of calcium in calcium chloride of 16.8 mole % at 1100°C. Not much importance can be attached to this figure, however, as the crucible may not have been sectioned exactly in the middle.

2.4.5. Modified Programme. From the preliminary investigations described in the foregoing it appears that the solubilities of calcium in calcium chloride and calcium in calcium bromide measured earlier in this laboratory were erroneous due to precipitation of the metal from the 65.

solution during quenching. The phase diagrams for these two systems had therefore to be re-determined. A new technique which overcomes the effect of segregation had to be developed. Results from these new solubility measurements in the calcium-calcium chloride system indicated so high solubilities of the metal that one could expect a consolute point at a temperature below 1400°C. It was therefore decided to try to establish the entire solubility curve in this system with the aid of differential thermal analysis. If this proved possible, attempts would then also be made to establish the phase diagrams of the calcium-calcium bromide, the calcium-calcium iodide and the strontium - strontium chloride systems.

2.1-6. Development of new method. From the preliminary experiments in the calcium- calcium chloride system it became obvious that metal precipitated out fromlhe salt phase during the quench. As the precipitated metal was much lighter than the salt it tended to float up and join the main metal phase. Analysis of the quenched salt phase thus gave a solubility figure for the metal that varied depending on the quenching rate and from which part of the salt phase the sample was taken. If the whole of the salt 66. phase was analysed it was obvious that in order to obtain a true value of the solubility of the metal any material transfer between the two phases, after equilibrium was reached, had to be prevented. This was obviously impossible for the whole of the salt phase, but could be achieved for part of the salt phase if it were contained in a separate compartment. This compartment had to be connected to the main salt phase in such a way that the salt it contained could come into equilibrium with the metal but could be isolated completely after equilibrium was reached. To achieve this the following arrangement was tried. Two pure iron crucibles were prepared. One was ,12 mm long and 5.5 and mm in external and internal diameter respectively. The other one had corresponding dimensions of 60, 8 and 7 mm. The small crucible had two 3.5 mm holes drilled through the wall with their centres 3 mm from the top. This crucible was placed upside down inside the bigger crucible. The crucible assembly was filled with dehydrated calcium chloride and the salt slowly premelted under vacuum. It was thought that the salt would fill up the inner crucible during premelting. If calcium metal was then loaded into the outer crucible and the crucible closed 67.

and equilibrated, the metal would first saturate the outer salt phase and then start to diffuse in to the inner crucible through the holes in its wall. After sufficient time this part of the salt phase would also become saturated with the metal. If the crucible was then quenched it was thought that the metal precipitated from the salt in the inner crucible would be trapped. However, the inner crucible did not fill with salt on premelting, probably due to gas trapped within it and to the surface tension of the salt. Another attempt was therefore made in which the inner crucible was slightly longer. Salt was now first premelted into this crucible which was then cut off so that the salt completely filled the remaining part of it. Two slots, 3 by 1 mm each, were sawn in the crucible wall instead of the drilled holes in the previous design. Sawing was more convenient because the work had to be done in the dry- box. The slots were sawn 1 mm from where the crucible was cut and were parallel to the bottom. The outer crucible had the same dimensions as before. After the salt had been premelted intothe outer crucible the small crucible was placed upside down on top of the salt phase. The whole assembly was now heated so that the salt melted and the small crucible sank down to the bottom of the 68„

outer crucible. The metal was added and the crucible closed. After equilibration the crucible was sawn open. It was found that the inner crucible was no longer on the bottom but was stuck half way up the salt phase. It was obvious that the inner compartment had to be held firmly in its position. For this reason another arrangement was tried. The outer crucible was the same as before but the inner crucible was changed to a tube slightly longer than the outer crucible. In one end of this tube two vertical slots, three by one mm, were sawn. This end of the tube rested on the bottom of the outer crucible. The other end of the tube was cut into three portions which were bent round to press on the wall of the outer crucible. This served to hold the inner tube in place during premelting of the salt. The inner tubel as well as the space between it and the crucible wall were filled with salt and the salt premelted. The solidified salt kept the inner tube in position so that the bent portions on its top could now be cut off and the tube closed. The outer compartment around the tube was now loaled :letal and the crucible 7cchanic$111y closed and spotwolded. when the crucible was cut open after equilibration, and quenching, it was found that the inner tube was displaced from the centre and touched the crucible wall at the bottom. Calcium metal 69.

had crept down to the bottom, between the inner tube and the wall where they touched. It was thus impossible to say whether part of the metal had crept into the inner tube or not. To avoid these difficulties the outer crucible was made slightly wider and shorter and a shallow hole was drilled in its bottom in which the tube fitted tightly. Instead of the slots two 1.5 mm holes were drilled through the wall of the inner tube so thatt when it was in place i the holes were just above the floor of the outer crucible. With these alterations the method worked satisfactorily. It frequently happened, on premelting, that when the salt in the inner tube started to melt, gas trapped in the lower part of it caused the salt to jump out. It was found that this could be avoided by heating the loaded crucible under vacuum for some time prior to premelting. This caused the absorbed gas to be expelled. To avoid any transfer of material between the inner anc outer 'compartments, caused by possible pressure differences set up duririg'the quenching, the inner tube was pressed together with a pair of pliers just above the top of the salt phase. The sequence in which this double crucible was assembled and filled can be seen on Figure 10. After equilibration the crucible was quenched in a 10 per cent lithium 70,

O 0 A B C D

METAL 71 ,...., .1 I. ill SALT II !i; d 1.,r) I,

E F G

INNER TUBE 5mm. 0.D., 4mm. I.D. , HOLES 1.5 mm. OUTER CRUCIBLE 12mm. O.D., I I mm. ID., BOTTOM 2.5mm.

SCALE I/ I

FIG. 10 CRUCIBLE FOR SOLUBILITY EXPERIMENT. 71. chloride water solution as this (according to Seybolt and Burke 39) should give a very high quenching rate. As the two holes for the diffusion of the metal into the inner compartment were situated so near the bottom, they should be blocked immediately, 107 solidified salt, on quenching, thus preventing transfer of material between the two compartments. After quenching the inner tube was removed and its outer surface thoroughly cleaned in the dry-box. The whole content of the inner tube was then analysed for calcium.

2.'1-.7. Experimental procedure. In the experiments to determine the solubility of the metal in its molten halides the double crucible technique already described was used in all cases except in the preliminary investigations. The procedure was as follows. The inner tube and the outer crucible were heated overnight at 1000°C in an atmosphere of hydrogen in a small horizontal alumina tube furnace reserved for this purpose. This was done to remove surface oxide and lower the impurity content of the iron. The crucible was then assembled as described on Figure 10. Both the inner and outer compartment were loaded with the powdered salt in the dry-box. The crucible was then placed in a glass 72. envelope as shown in Figure 11. The envelope was transferred through the ante-chamber of the dry-box to a vacuum pump. Under a vacuum of approximately 10-3 mm of Hg the crucible was heated up in a small vertical mullite tube furnace to about 400°C and kept there for about 3 hours. This was done to drive off all gases that might have been absorbed on the salt and the crucible. Still under vacuum the envelope was placed in the coil of a high frequency induction furnace, where the salt was premelted. The pressure was measured during the premelting on a 2irani gauge. No pressure increase during the premelting indicated that all gases had already been expelled from the salt. After the envelope containing the crucible was cooled it was transferred into the dry-box. There the inner tube of the crucible was pressed together with a pair of pliers to decrease th volume above the salt phase in the inner tube. This procedure also increased the volume of the outer compartment of the crucible which was now filled with the metal. The top of the crucible was mechanically closed in a vice. The crucible was thereafter taken out of the dry-box and the closure spotwelded. The spotweld was polished on emery paper and examined under a binocular microscope for flaws. A small hook was then 73.

CRUCIBLE SCALE I/I

ALUMINA SUPPORT

FIG.II. GLASS ENVELOPE FOR PREMELTING UNDER VACUUM. 74. spotwelded on to the top of the crucible. Suspended on a Kanthal wire by this hook the crucible was lowered into the centre of the equilibrium furnace, after remaining in the top while the furnace was evacuated and filled with argon. 1.then the crucible had been in the centre for about ten minutes the temperature was read on a Tinsley potentiometer. After sufficient time had elapsed to allow the metal to come into equilibrium with the entire salt phase in the crucible the temperature was read again. This was done to make sure that the temperature had not changed during the equilibration. The crucible was thereafter removed from the furnace and rapidly quenched in a 10)6 lithium chloride water solution contained in a beaker placed on the top of the furnace. The beaker was placed so that the rim of it was in line with the top of the reaction tube. The crucible was then transferred to the dry-box where it was sawn open. Usually the inner tube could easily be broken loose from the bottom part of the crucible after the adhering salt had been scraped off. The outside of the inner tube was thoroughly cleaned with a file and then cut off just above the salt phase. The remaining part of it containin: the sample was placed in a sample tube containing a magnetic stirrer and fitted with a 75. bung. The sample tube had earlier been weighed with the bung and the magnetic stirrer. It was now weighed again after the addition of the sample. The procedure for the analysis of the metal in the salt has already been described under the gas analysis apparatus. After the gas analysis was carried out the empty inner tube was cleaned in water and dried with alcohol and acetone and then weighed, so that the actual weight of the sample could be calculated. 76.

2.5. Differential thermal analysis.

2.5.1. The principle of the method. A schematic drawing of a differential thermocouple can be seen on Figure 12. It consists of two platinum wires joined together with a piece of platinum 13% rhodium wire, thus forming two junctions A and B. On the middle of the platinum alloy wire is joined an extension wire of the same material. 01 D and E go to the cold junctions. The M.F. between C and E gives

the temperature difference between A and B, while the temperature of .. is liven between C and D. A is in thermal contact with the sample and B with an inert reference body of the same thermal mass. The two are placed in a uniform temperature enclosure, whose temperature is either raised or lowered at a uniform rate. Any heat effect occurring in the sample cause the value of th- differential E.M.F. to change. By plotting both the temperature of the sample and the differential E.E.F. against the same time- axis the temperature corresponding to this change on the differential r.M.7. can easily be found. If the differential U.M.F. is amplified very small heat effects can be detected which would not be possible with only

Pt C

TEMPERATURE. Pt 13% Rh. • D

DIFFERENTIAL. Pt

FIG.12. DIFFERENTIAL THERMOCOUPLE. 78. heating or cooling curves. The area under the differential curve due to its deflection corresponds to the amount of heat evolved or absorbed in the sample crucible during the reaction. If the conditions under which two samples are heated or cooled remain the same the ratio between the heat effects in the two samples can be calculated from the ratio between these areas.

2.5.2. Initial experiments. In order to determine the depressirm of the freezing point of strontium chloride by strontium and calcium iodide by calcium it was decided to use a differential thermal analysis technique, starting with the strontium-strontium chloride system. A Johnson Matthey platinum wound furnace was used for heating the 11/2" alumina reaction tube. It was placed in vertical position and the crucible assembly was hung in its centre suspended on an iron- wire from the top bung. Argon was let into the tube through the bottom bung and let out through the top bung. The crucible assembly consisted of two pure iron crucibles, one being placed 1 cm above the other. The upper one contained the sample and the lower one 79. was empty and acted as a reference. The thermocouples were placed in axial tubes in the crucibles so that the junction in the sample was located in the centre of the melt, and that in the reference was placed in a corresponding position. The differential 1.M.F. from the thermocouples was recorded on a Honeywell Brown two point electronic recorder which also recorded the temperature of the sample. Stray currents from the furnace caused heavy disturbances on the recorder trace even at temperatures below 80)°C. This made the interpretation of the differential curve impossible. Therefore results could only be obtained from cooling curves when the current was switched off, which meant that the cooling rate could not be varied. At this stage of the research the solubility measurements in the calcium-calcium chloride system indicated that a consolute point could be expected in this system at about 1300-1350°C. It was therefore decided to build up a more sensitive differential thermal analysis sytem with a molybdenum furnace which did not give any spurious results due to stray currents. this new system it was also thought to be possible to establish th, entire solubility curve in the calcium-calcium chloride system. 80.

However, even with this new system heavy disturbances on the recorder trace due to electrical leakage from the furnace winding were encountered although not until temperatures above 1200'C. This difficulty was overcome by shielding the sample and try:: thermocouples in the hot zone with a thin, earthed molybdenum shield.

2.5.3. Crucibles. For the initial experiments crucibles designed 40 by D.H. Bell of this laboratory were used. The body of the crucible was pressed out of 20 gauge pure iron sheet. The lid was machined out of pure iron rod and was welded on to the crucible body. Holes were then drilled in the lid and the bottom and a central thermocouple tube, slightly longer than the crucible, and a fil:ler spout were welded in. The crucible can be seen on Fig.13, A and B. The welds of this crucible were rather thick, mainly because ordinary oxyacetylene gas welding with filler rod was used. This welding was particularly difficult around the filler spout because of the interference of the welds from the lid snd the central tube. The welds were frequently contaminated with iron oxide which was reduced when the crucible was heated 81.

--FILLER SPOUT.

INNER TUBE.

A B

C D

FILLER SPOUT 6mmO.D. 5mm I.D. CRUCIBLE BODY 17mm O.D. WALL 20 GUAGE INNER TUBE 4mm OD. 2mm I.D. SCALE I/I FIG.I3. CRUCIBLES FOR THERMAL ANALYSIS. 82. up in a hydrogen atmosphere. This le, d to porosity and leaks in the crucible, which had to be rejected. As in the thermal analysis it is very important to have a low crucible to sample weight ratio, it was desirable to make the welds as thin as possible yet strong and sound. For this reason an argon arc welding technique was used instead of the ordinary gas welding and the crucible design was changed so that no double welds were involved. The size of the crucible body remained the same but instead of the central axial tube a diametrical one for the thermocouple junction was welded in slightly below the centre of the crucible body. In this way the filler spout could be placed in the centre of the lid. The crucible design is shown on Fig. 13, C and D. The numbers on Fig. C indicate the order and the arrows the approximate angles under which the welds were done. In this welding technique no filler rod was used which made the welds considerably thinner. It was, however, difficult to get an even and deep weld. Satisfactory results were achieved by making a rotating device to hold the crucibles during welding. A deeper weld was achieved by using a thin thoriated tungsten pen-electrode. 83 .

Most of the crucibles welded in this way leaked in the welds after they had been reduced in a hydrogen atmosphere. Microscopic examination of sections of the welds revealed porosity due to reduced inclusions of iron oxide. However, with an addition of 6 to 7% of hydrogen to the welding argonlsound welds were obtained. After the crucibles had been welded they were cleaned in hydrochloric acid and washed with distilled water. They were then heated to 1000°C in a hydrogen atmosphere overnight to reduce any inclusions of iron oxide. Before use the crucibles were tested for leaks by subjecting them to a pressure of 40 p.s.i. of hydrogen under water. After the crucibles had been loaded they were evacuated and the filler spout closed and spotwelded. The spotweld was polished and microscopically inspected for flaws. The sealed crucible can be seen on Fig. 13, D.

2.5.4. Furnace and differential analysis assembly; The final arrangement for the differential thermal analysis consisted of a vertical tube furnace, and a Honeywell Brown two point electronic recorder, connected to the thermocouples via a biasing circuit. 84.

The heating element of the furnace was a molybdenum tape 20 thou thick and 1/8" wide. It was wound on to a recrystallised alumina tube which was about 2' long and 2" in diameter. 'The winding tube was sealed into a steel casing by means of a gland at either end, packed with graphite-oil impregnated asbestos string. The furnace casing was filled with alumina powder as thermal insulator. The heating element was protected from oxidation by a reducing atmosphere of forming gas. This gas was passed through the casing which %as therefore made gas-tight by sealing all the joints with a calcium fluoride-sodium silicate cement. The design of the furnace has previously been described . The furnace was placed on a welded steel frame so that its bottom was 55" from the floor. The alumina reaction tube had an internal and external diameter of 38 and 46 mm respectively. It was 900 mm long and was fixed in position by a clamp fastened on to the steel frame. A 10" long earthed shield was fitted tightly on to the reaction tube covering its hottest part. The shield was made of 2 thou molybdenum sheet, spotwelded together with an insert of a thin platinum wire to facilitate welding. It protected the thermocouples from electrical disturbanc=-tes caused by the furnace winding. The space between the furnace and reaction tubes was flushed continuously with forming gas to avoid oxidation of the shield. For this reason either end of the furnace tube was closed against the reaction tube with silastomer. The ends of the reaction tube were water cooled by lead coils. The upper end of the reaction tube was joined with araldite to a glass socket into which fitted a closed glass cone from which platinum radiation shields were suspended by a platinum wire. The bottom end of the reaction tube was connected to a brass tube by an 0-ring seal. This had a Wilson seal fitted at its lower end. A 9 mm alumina tube 900 mm long passed through this seal. The arrangement can be seen on Fig. 14. The brass tube had the object of providing a long cool zone so that the crucible assembly, resting on the top of the alumina tube, could be lowered down while still hot without burning the Wilson seal. The crucible arrangement can be seen on Figure 15. A was the reaction tube and B was the alumina tube to which the alumina container C was joined by alumina cement. D was another alumina container in which the sample crucible F and the reference crucible G were held. This container had four 86, ALUMINA --, RE ACT ION TUBE / / 0-RING SEAL

r, I-. / / / / / / r / r

/ /

% ....MIRO 1 / / / / r SCALE I / I / / / / / / / / /

WILSON SEAL

I ALUMINA TUBE FOR INLET OF ARGON I THERMOCOUPLES AND SUPPORT OF CRUCIBLE- ASSEMBLY. FIG. 14. H 87.

/.. 1. di, 2' •• 1 , z J. 1 41"

_1-

SCALE I/I 1

i I

1 1 i 1

H 'A

FIG.I5. CRUCIBLE ARRANGEMENT FOR DIFFERENTIAL THERMAL ANALYSIS. 88.

KEY TO FIG. 15

A. Alumina reaction tube B. Allnlina supporting tube C. Alumina container D. AluEina cont-liner E. Alumina lid F. Sample crucible G. Reference crucible H. Radiation shields I. Platinum wire J. Platinum 13% rhodium wire K. Platinum 13% rhodium wire 89.

holes drilled in its wall. Two alumina insulators passing through the tubes in the crucibles, were fitted into these holes. The thermocouple junctions were placed in the centres of these alumina insulators. K was the common platinum-13% rhodium wire and I and J the pure platinum wires. The wires were separated from each other by alumina insulators and passed out of the furnace throu-:h the tube B. The space between the crucibles and the container D was packed with alumina powder and the container was centered with the aid of a lid f E , made of alumina cement. The convection and radiation shields ,H,were made of 5 thou platinum sheet. The four upper were suspended from a platinum wire fixed to the top, while the four lower ones were spotwelded on to two platinum wires which were fitted into slots cut in the alumina tube B. This alumina tube passed through the Wilson seal and its lower end was fitted to a 2" long glass tube. Argon was let into the furnace through the bottom of this glass tube and the thermocouple leads entered through three side-arms attached to it. Small rubber bungs, each containing a small glass capillary tube, were placed in the ends of the side-arms and seals for the thermocouple leads were effected by passing the leads through the glass 90.

capillaries and sealing with Apiezon ,;c--compound. The furnace reaction tube could be evacuated through the stopcock joined to the brass tube. See Fig. 14. The power to the furnace was supplied through a 50-B variac auto-transformer. To permit a fine adjustment of the input to the furnace, part of the winding of the variac v,as tapped off over a 100-R variac so that its maximum output could be varied continuously between 230 and 270 volts. Before the E.M.F.:s from the thermocouples were fed into the amplifying unit they had to be biased as the full scale deflection on the recorder only corresponded to 1 millivolt. The biasing unit was almost identical to one designed and used by 42 G.P. Jones of this laboratory. The circuit can be seen on Fig.16. The points 9, 10 and 11 were the cold junctions of the thermocouples and T.C.1 and T.C.2 were the two inputs of the biased E.M.F.:s to the amplifying unit of the recorder. When the three-pole switch ABC was in position I, T.C.1 was short-circuited, while the bias E.M.F. for the differential was fed to T.C.2 and could be recorded. In position II the biased E.M.F. of the sample thermocouple wac fed to T.C.1, while the biased M. F. of the differential

91.

I' 0 T.C.I. © A 0-0 0 III

@ T.C.2. @ I B 41 o L.,..4 "-en si f 1:.4. — • in/

1 Ls 80SL S 150K SL •

+ POT ® 0 — 10.SL

200 SI- 50 St

MAW 7 5 0 0 ..rt 1.5 V

FIG.I6. BIASING CIRCUIT. 92.

KEY TO FIG. 16

1 - 2: 1.5 volt battery for biasing of differential E.M.F. 3 - 4: Potentiometer. 5 - 6: 1.5 volt battery for biasing of temperature E.M.F. 7 - 8: Recording on T.C.1 9 - 10. - 11: Thermocouple cold junctions. 12 - 13: Recording on T.C.2.

A B C: Three pole switch Position I: Short circuit on T.C.1 Bias E.M.F. for differential on T.C.2- Position II: Biased temperature E.M.F. on T.C.1 Biased differential E.M.F. on T.C.2. Position III: Biased temperature E.M.F. on T.C.1. Potentiometer on T.C.2. 93'.

thermocouple was fed to T.C.2. In position III the biased E.M.P. of the sample thermocouple was fed to T.C.1 while to T.C.2 was fed the E.M.F. provided by the Cambridge potentiometer (POT on Fig. 16) biased by the same E.M.F. Thus with the potentiometer switched on the recorded biased E.M.F. of the sample thermocouple could be compared to a known E.M.F. biased to the same extent.

2.5.5. Experimental procedure. The cleaned and hydrogen reduced crucible was weighed and transferred into the dry-box where it was loaded with salt. It was then placed in a glass envelope and taken out of the dry-box. The salt was premelted under vacuum in the same way as in the solubility experiments already described. Then the crucible was cool the glass envelope was filled with argon and the crucible talcen out and rapidly weighed, before it was once again transferred into the dry-box to be filled witil a weighed amount of metal. In the case of strontium the metal was weighed in a sealed glass tube of known weight, so that the metal was never in contact with the laboratory atmosphere, while the

calcium could be weighed without this precaution, 94.

provided the time it was in contact with the air was not too long. The amount of metal-salt mixture in a crucible varied with composition but was about 60- 70 m.moles in the calcium-calcium chloride and calcium-calcium bromide systems, and about 30-50 m.moles in the calcium-calcium iodide and strontium- strontium chloride systems. After the crucible had been filled with the metal its filler spout was joined to a short piece of rubber pressure tubing which was connected to a closed stopcock. The crucible could now be handled in air without its content being contaminated. The other end of the stopcock was connected to a glass flask which could be evacuated through another stopcock. This system was now evacuated with the two stopcocks open. The glass flask increased the volume of the system so that when it was evacuated, closed and detached from the vacuum line any minor leaks in the rubber connections only caused negligible increase in pressure. The filler spout was mechanically closed near the crucible lid. The closure was spotwelded and the superfluous part of the spout sawn off. After the spotweld had been microscopically examined for flaws the crucible was weighed and placed in the top of the alumina container 95.

D in Fig. 15 which already had the empty reference crucible in its lower part. Alumina powder was packed around the crucibles and the thermocouple junctions were fitted into the alumina tubes going through the centres of the crucibles. The thermocouple junctions could be made very thin by butt- welding the wires together in a small oxy-hydrogen flame. Vqien the thermocouple wires had been covered with alumina insulators they were passed through the alumina tube B in Fig. 15 and the crucible container D was placed in its position in container C. The whole crucible assembly was placed in the lower part of the furnace and the 0-ring seal fitted on to the reaction tube, which was evacuated and filled with argon. The crucibles were slowly raised into the even hot zone of the furnace by pushing up the supporting alumina tube through the Wilson seal. The crucibles were equilibrated at a temperature considerably above the expected reaction temperature and their cooling and differential curves then recorded. Cooling curves were always employed in the miscibility gaps of the metal salt systems as it was only possible to detect the phase boundaries when a second liquid phase began to separate out from a homogenous phase. As this separation 96. was a continuous process with the decrease in temperature the evolution of heat was also continuous. On the recording trace this appeared as a small but distinct change on the differential curve, indicating a sudden temperature increase in the sample crucible compared to the reference crucible. Fig. 17 A and B show typical cooling curves from the calcium-calcium chloride system. The first one indicates the separation into two liquid phases at about 128000 and a composition of 29.44 mole percent calcium, while the second shows the precipitation of the solid calcium metal from the monotectic liquid at about 826°C. The cooling rate employed varied between 1 to 3°C per minute. Because of the large supercooling of calcium chloride, its melting point could only be determined on heating. This was established at a very slow heating rate (about 0.1°C per minute). After the experiments the crucibles were always weighed to ensure that no leak had occurred. However, even minor leaks could easily be detected after a run due to the reaction of the escaped metal with the alumina powder surrounding the crucible. When leaks had occurred all the alumina insulators were

97.

DIFFERENTIAL. TEMPERATURE.

v. x 14.000 M.V. 14.400 M.V. t 1278.6-C C.J. 25.5°C.

i x 8200 MY.

I- B C.J. 24.5°C.

E.M.F. IN MILLIVOLTS.

FIG.17. DIFFERENTIAL AND TEMPERATURE CURVES FROM THE SYSTEM CALCIUM -CALCIUM CHLORIDE. 98. renewed to ensure that no contamination of the new thermocouples took place. The thermocouple wires were always either completely changed or its top parts renewed after each run.

99.

CHAPTER 3

RESULTS

3.1. Description.

3.1.1. The Calcium-Calcium Chloride system.

3.1.1.1. Solubiliti measurements. The time to reach equilibrium, using the double crucible technijue, was established at two temperatures. The results are given in the Tables 12 and 13 and in Figure 18. TABLE 12 Temperature °C Mole % Ca Time in CaCl2 minutes.

1005.0 5.82 315 1005.0 6.18 319 1005.0 6.58 454 1005.0 6.76 608 1005.0 7.18 750 1005.0 7.33 942 1000.5 7,19 x 1642

€ Correction to 1005.0°C gives a value of 7.31 mole %. The time to reach eJuilibrium at 1005.0°C was thus about 15 hours and the solubility of calcium in calcium chloride at this temperature 7.32 1 0.10 mole %. 100.

TABLE 13 Temperature °C Mole % Ca Time in in CaCl2 minutes 1101.0 8.98 143 1101.0 8.90 173 1101.0 9.85 272 1101.0 9.80 395 1104.0 10.13 K 447 1101.0 9.72 692 Correction to 1101°C gives a value of 10.05 mole %. The time to reach equilibrium at 1101.0°C was thus about 5 hours and the solubility at this temperature ± 9.86 0.15 mole %. Further measurements were made at higher temperatures. The crucibles in these experiments were always left in the furnace overnight and so remained at equilibrium temperature for a time that far exceeded the e4uilibrium time. The obtained values are given in Table 14. TABLE 14 Temperature °C Mole % Ca in CaCl2 1149.0 11.4 1201.0 16.4 1201.0 17.4 1247.6 23.9 1250.4 22.8 6 CC M 0 I 1 I i I I I I I I I I I I I N 12 if; cc 1 n II 0 I tn 3( )C

w 8 -i 1005°C '0 x M 7

›-1— 6 J_ Cl 5 / D -J 0 4 111 / 3

2 I

I I I I I 1 1 I I I 1 i I I I 1 i I 0 500 1000 1500 TIME IN MINUTES. 8 FIG 18 SOLUBILITY OF CALCIUM IN CALCIUM CHLORIDE AT 1005°C & 1101°C 102.

The scatter in the results was greater at these higher temperatures. This might partly be due to the difficulty in completely preventing transfer of liquid between the inner and outer compartments of the crucible during quenching. For this reason no solubility measurements were made above 1250°C. As the temperature coefficient of the solubility at this temperature was quite large, the solubility curve above this temperature was established with differential thermal analysis instead.

3.1.1.2. Differential thermal analysis. As the pure calcium chloride supercooled very much, its melting point could only be determined on heating. This was established at a very slow heating rate (about 0.1 °C/Minute). Even the eutectic supercooled to a great extent (up to 55'C). To overcome this effect a mixture of approximately the eutectic composition was held at a temperature between the melting point of the salt and the expected eutectic, until most of the salt had formed a molten solution with the metal. On cooling the remaining solid salt prevented supercooling. The most suitable cooling rate for obtaining the phase boundary in the miscibility gap was 103.

found to be about 1°C per minute. The melting point of the pure calcium metal was taken as the average of the values obtained on a heating curve and a cooling curve. The results are given in Table 15 below. The monotectic temperature was only recorded on a few experiments. TABLE 15 Differential thermal analysis of the calcium-calcium chloride system. Composition Temperature of inflexion °C Remarks of mixture mole % Ca Liquidus Monotectic Eutectic 0 774.4 0 774.6 0 Average 774.5 4.35 - - 763.6 4.35 - - 762.8 4.35 Average - - 763.2

29.44 1278.6 826.6 29.44 1280.9 827.2 29.44 Average 1279.7 826.9 48.55 1327.5 48.55 1325.8 48.55 Average 1326.7 104.

TABLE 15 (continued) Differential thermal analysis Of the calcium-calcium chloride system.

Composition Temperature of inflexion °C Remarks of mixture mole % Ca Liquidus Monotectic Eutectic 64.32 1336.8 64.32 1339.0 (819.9) 64.32 Average 1337.9

Amm• MEM 84.32 1305.1

84.32 1304.5 825.8 84.32 Average 1304.8

95.34 (1248.0) 824.2

100 839.8 - Heating curve 100 839.1 It It 100 837,7 - Cooling 11 100 837.4 11 100 Average 838.5 As received

100 839.6 - Heating curve 100 836.8 - Cooling 11 100 Average ' 838.2 - Redistilled

As the melting point of the calcium metal was lower than expected from the impurity figures given it was redistilled under vacuum. This was done to exclude any possible oxide contamination, as the oxygen impurity 105. was not given in the analysis. As can be seen from the table this treatment did not alter the melting point. The average monotectic temperature was 826.0°C if the extreme loy. value 818.9 is excluded. The inflexion on the differential curve for the composition 95.34 mole % calcium was very weak and uncertain due to the small temperature coefficient on this part of the solubility curve. This temperature figure is therefore placed in brackets. The eutectic composition, 3.3 ± 0.4 mole % calcium at 763°C, was obtained by extrapolation on a plot of the logarithm of the solubility versus the reciprocal of the absolute temperature, taking into account the change in slope due to the heat of fusion of the calcium metal at the monotectic temperature. The consolute point is at about 62 mole % Ca and 1338 t 5°C. The monotectic is at 826°C and 99.5 mole % Ca. The phase diagram of the calcium- calcium chloride system is given in Fig. 19. 106,

1500 X DIFFERENTIAL THERMAL ANALYSIS.

® SOLUBILITY. 1400

1300

1200

u o 1100 URE T I000 RA E MP

TE 900

800- Nt

700 I 0 50 100 COMPOSITION MOLE °/° Ca.

FIG.19. PHASE DIAGRAM OF THE SYSTEM CALCIUM /CALCIUM CHLORIDE.

107.

3.1.2 The Calcium-Calcium Bromide system.

3.1.2.1 Solubility measurements. Only two solubility measurements were made in the calcium-calcium bromide system, as the major part of the diagram was established by differential thermal analysis. The results are given in Table 16 below. The crucibles were held at equilibrium temperature for about 20 hours to ensure that equilibrium was reached. TABLE 16 Solubility of calcium in calcium bromide.

Temperature 'C Mole % Ca in CaBr2

983.3 9.63

1103.5 15.45

3.1.2.2. Differential thermal analysis. Even in the calcium-calcium bromide system the eutectic supercooled very much (about 50°C). To establish the eutectic temperature on a cooling curve a mixture of a composition less than the expected eutectic was therefore held at a temperature below the liquidus in the same manner as in the calcium-calcium chloride system. The cooling rate used for obtaining the phase boundary in the miscibility gap was about 3°C per minute. The results are given in Table 17. The bee, monotectic temperature was only recorded on two of the mixtures. TABLE 17 Differential thermal analysis of the calcium-calcium bromide system. 0omposition Temperature of inflexion °C of mixture mole % Ca Liquidus Monotectic Eutectic 1.98 727.6 29.99 1252.1 29.99 1257.6 29.99 1253.1 29.99 1255.5 29.99 1251.8 29.99 1255.6 827.7 29.99 Average 1254.3 ± 2.1 40.44 1304.0 40.44 1308.3 40.44 1307.7 40.44 Average 1306.7 ± 1.9

54.08 1321.0 54.08 1328.3 54.08 1329.4 54.08 Average 1326.2 ± 3.7 69.61 1332.6 69.61 1333.0 69.61 Average 1332.8 ± 1.0 87.23 1287.5 828.3 87.23 1288.9 87.23 1288.3 ••• 87.23 Average 1288.2 0.6 109.

The phase diagram of the calcium-calcium bromide system can be seen on Fig.. 20. The consolute point is at about 64 mole% Ca and 1335 ± 5°C. The monotectic is at 828°C and 99.6 mole % Ca. The eutectic is at 728°C and 3.2 ± 0.4 mole % Ca. The eutectic composition was extrapolated in the same way as in the calcium-calcium chloride system. The melting point for pure calcium bromide,742.5°C,was taken from Taylor's work.

110.

1500 X DIFFERENTIAL THERMAL ANALYSIS. 0 SOLUBILITY 1400

1300

1200

1100 0()

URE 1000 AT

MPER 900 TE

800

700 0 50 100 COMPOSITION MOLE % Ca

FIG 20 THE PHASE DIAGRAM OF THE SYSTEM CALCIUM / CALCIUM BROMIDE.

3.1.3 The Calcium-Calcium Iodide system.

3.1.3.1 Solubility measurements. Only three solubility measurements were made in the calcium-calcium iodide system. The crucibles were equilibrated for about 20 hours. The results are given in Table 18. TABLE 18 Solubility of calcium in calcium iodide.

Temperature °C Mole % Ca in CaI2 1095.3 15.2 1103.5 14.5 1204.2 22.8

3.1.3.2 Differential thermal analysis. The pure calcium iodide as well as the eutectic did not supercool more than a few degrees at the most. The samples thus recovered the equilibration temperature when the solidification started. The results are given in Table 19 and Fig. 21. The consolute point is at about 74 mole Ca and 1380 5°C. The monotectic at about 99.7 mole % Ca and 831°C. The eutectic at 6.8 mole % Ca and 760°C. FIG.21. THEPHASE DIAGRAMOFTHESYSTEM 0 U TEMPERAT URE 700 0 0 SOLUBILITY. X DIFFERENTIALTHERMALANALYSIS. CALCIUM /CALCIUM IODIDE. COMPOSITION MOLE 50 1 I

% Ca. 100

112, 113.

TABLE 19 Differential thermal analysis of the calcium-calcium iodide system. Composition of mixture mole % Ca Liauidus Monotectic Eutectic o 779.6 o 778.7 o 778.6 0 Average 779.0 t 0.5 2.39 772.0 759.2 2.39 771.4 758.5 2.39 Average 771.7 758.8 4.37 767.0 4.37 767.6 761.1 4.37 Average 767.3 30.10 832.9 760.0 49.99 1330.5 49.99 1330.7 49.99 1330.7 830.7 758.6 49.99 Average 1330.6 64.91 1370.6 831.7 64.91 1369.3 830.3 753.2 64.91 Average 1370.0 84.05 1367.4 84.05 1370.0 84.05 Average 1368.7 114.

3.1.4 The Strontium-Strontium Chloride System.

3.1.4.1 Solubility measurements. Six solubility experiments were carried out at about 1000°Cin the strontium-strontium chloride system. They failedl howeveri to give reproducible results due to leakage in the spotwelds in the top of the crucible. The solubilities obtained varied between 9.? to 21.8 mole 'X7 strontium. As it was thought that the miscibility gap in this system might be closed in the vicinity of this temperature, three more experiments were carried out at about 920°C. In all these cases the crucibles leaked in the top weld of the outer compartment, so that the inner compartment of the double crucibles contained very little salt. It was thought that the boundary of the miscibility gap could be more easily established with the aid of the differential thermal analysis. The solubility experiments were therefore postponed. However, the differential thermal analysis showed the strontium metal to be so impure that only experiments were reliable where the strontium metal content was low. For this reason no further solubility measurements were carried out. 115.

3.1.4.2 Differential thermal analyais. The first three experiments with 0, 0.74 and 1.63 mole percent strontium were carried out in the platinum wound furnace described in the initial thermal analysis experiments on page 78 . All other differential thermal analysis experiments were carried out in the apparatus described under Furnace and Differential Assembly. The results obtained are given in Table 20. TABLE 20 Differential thermal analysis of the strontium- strontium chloride system. Composition Temperature of inflexion of mixture 00 Mole % Sr First Second 0 872.4 0 872.5 0 872.5 0 872.6 0 872.5 0 Average 872.5 t 0.1

0.74 868.9 0.74 867.7 0.74 868.7 4•01. 0.74 868.1 0.74 Average 868.4 ± 0.5

1,63 863.2 116.

TABLE 20 (continued) Differential thermal analysis of the strontium- strontium chloride system.

Composition Temperature of inflexion of mixture 'D C Mole % Sr First Second 4.64 845.5 838.2 4.64 844.3 839.0 4.64 Average 844.9 838.6

7.50 836.2

19.05 825.0

32.45 801.0 754.o 100 769.3 719.o

As can be seen from the table the experiments failed to give any indication of a miscibility gap above the monotectic temperature. Instead the monotectic temperature decreased with increase in strontium content. The determination of the melting point, of the metal gave a slight kink in the cooling curve at 769.3°C and a definite break at 719°C compared to the value of 770°C given by Kubaschewski and Evans43. This indicated that the metal was probably less pure than stated in the analysis given by the suppliers. A qualitative test for nitrogen with Nessler's reagent 117•

showed this element to be present to a great extent. It was thought, however, that the impurity of the metal would have little effect on the depressing of the freezing point of the strontium chloride at compositions less than the monotectic, which is at about 5.5 mole /'strontium. This part of the strontium-strontium chloride system can be seen on Fig. 22. 900 ...

Li 0 URE

T 850 A R x MPE TE

800 1 I I I I 1 1 1 0 I 2 3 4 5 6 7 8 COMPOSITION MOLE 0 Sr

FIG. 22. DEPRESSION OF FREEZING POINT OF STRONTIUM CHLORIDE. 119.

3.2 Errors. The limits of error attached to the figures in the tables of results are the standard deviations from the arithmetic mean. The standard deviation 5 is given by the equation

= 7:(X )2 n-1 where is the deviation from the arithmetic mean and n is the number of observations. The possible contributions to the errors are analysed below.

3.2.1 Errors in the solubility measurements.

3.2.1.1 Temperature measurements. It is estimated that the fluctuations in the crucible temperature during equilibration was not greater than ± 1°C. The temperature coefficient between 1000 and 1100°C is about 0.03 mole %/°C in the Ca/CaC12 system and about 0.05 mole %/°C in the Ca/CaBr2 and Ca/CaI2 systems. The contributions to the errors by fluctuations in this temperature range would thus be ± 0.03 mole % in the Ca/CaC12 system and ± 0.05 mole % in the Ca/CaBr2 and C;a/CaI2 systems. 120.

3.2.1.2 Weighing. The accuracy of the weighing was about 0.5 mgs and the weight of the samples about 400 mgs. At a solubility of Ca in CaCl2 of about 8 mole % this would give a contribution to the error of t 0.01 mole %.

3.2.1.3 Gas analysis. The standard deviation of the gas analysis apparatus was ± 0.7 /,which gives a contribution to the error at 8 mole of ± 0.06 mole %.

3.2.1.4 Other errors. If the above mentioned accidental errors are added up they give a maximum error of ± 0.10 mole % calcium in the calcium-calcium chloride system between 1000 and 1100°C. This is equal to the calculated standard deviations in this system at 1005°C. However at 1100°C the standard deviation was ± 0.15 mole /0. It is possible that this discrepancy was due to another error, caused by transfer of material between the two compartments in the crucible during the quench. This effect would probably increase with temperature and would thus explain the larger deviation from the smooth solubility curve of the few measurements at temperatures 121. above 1100°C. The error at this higher temperature was estimated to be about ± 0.5 mole %.

3.2.2. Errors in the differential thermal analysis.

3.2.2.1 Temperature measurements. On the cooling curves the temperature was determined by extrapolation between the points of known &L P. printed on to the chart from a Cambridge portable potentiometer. An accuracy of about ± 0.5 °C is estimated on these temperature determinations, taking into account the inaccuracy of the thermocouples.

3.2.2.2 Weighing. The influence on the composition of the mixtures due to weighing errors was probably very small and would not have any influence on the phase boundary temperatures as the differential thermal analysis was only used over the regions where the solubility curves had large temperature coefficients.

3.2.2.3 Other errors. The standard deviations from the average temperatures in the tables of results are much greater than expected from the above mentioned accidental 122. errors. The discrepancy may partly be due to the difficulty in determining the exact point on the cooling curve when the corresponding deflection on the differential curve was small. An accuracy better than ± 5.0°C canIfor this reasoninot be claimed on this part of the phase diagrams. The agreement in the solubility figures obtained by the two different methods is strong evidence for the assumption that only negligible systematic errors were involved in the techniques. The impurity of the materials was, however, a source of systematic errors that was independent of the technique used. The impurity of the calcium metal lowers its melting point as well as the monotectic and the miscibility curve in the calcium-calcium halide systems. If it is assumed that the metal has a melting point that is 5°C too low, the monotectic temperature will also be about 5*C too low, while the influence on the miscibility curve will decrease towards the salt rich side of the diagram. 123.

3.3 Derived data.

3.3.1 The heat of fusion of the investigated salts. In a phase diagram with coexisting solid and liquid phases the equilibrium condition for the solvent can be expressed by the equation

AHf AT In a(1) a(s) R x Tfx T where am and a(s) are the activities of the solvent in the liquid and solid phases respectively. The standard state in the liquid is pure liquid solvent and in the solid)pure solid solvent. Tf and T are the liquidus temperatures in °K of the pure solvent and the solution respectivelyl and AT is their difference.

AHf is the heat of fusion of the pure solvent and is assumed to be constant between Tf and T. With the aid of this equation the heats of fusion for calcium chloride, calcium bromide, calcium iodide and strontium chloride have been calculated from the obtained values of the depression of the freezing points of these salts by their metals. The calculations have been carried out assuming the activities of the dilute solutions correspond

124. to ideal Temkin ionic mixtures and that the activities of the solid phases are one. According to Temkin44 the activities of any particular ion species in an ideal ionic solution is equal to the ratio of the number of such ions to the total number of ions of like charge. This ratio is called the ionic fraction and the activity of a salt is thus equal to the product of its cation and anion fractions. The following three different possible models for the solution of the metals in the salts have been considered: 1. M + M++ -=v 1/1° + M++ or 2M++ + 2e- ++ 2. M + M++ M2 3. M M++ 2M+ These three models give the following three different activities for the salt Mx2 a Nu 2 1. Mx = N )( Nx- 2 M++ Nmx2 m

a N A Nx- Ifix2 NM 2. M 2 = NM++ '1111x2 a 2 NMx2 - NM 3. M N _ 2 NM++% x Nmx2 Nm 125.

The calculated heats of fusion for these three different possibilities are given in Table 21, together with the corresponding known values taken from the literature. TABLE 21 Calculated heats of fusion for CaCl2, CaBr2, CaI2 and SrC12 in Kcal/mole.

Modal :1 2 3 .6Hf from literature Ref. CaCl2 6.3 6.5 12.8 6.8 ± 0.1 43 CaBr2 4.5 4.7 9.2 4.18 45 CaI2 8.0 8.7 16.6 $.2 x SrC12 4.2 4.5 8.7 4.1 ± 0.6 43 x(Calculated from McCreary's phase diagram of the system 46 CaI2 - CaF2 . An estimation of LHf for CaI2 from the area under the differential curve in the thermal analysis gives, however, a value of about 7 Kcal/mole.) It is difficult to assess any limits of error to these calculated heat values as the amount of solid solution of the metal in its salt is noi known. However, the other influencing factors, temperature and composition, would probably not give rise to a larger error than ± 1 Kcal/mole for the CaI2 and theSrCl2. The significance of these different heats of fusion will be dealt with in the Discussion. 126.

3.3.2 The monotectic compositions in the Ca-CaCl2 - - Ca CaBr2 and Ca CaI2 systems.

From the measured depressions of the freezing point of calcium the activities of the calcium at the monotectic temperatures have been calculated in the calcium-calcium chloride, calcium-calcium bromide and calcium-calcium iodide systems. The same equation for the relation between the activity and the depression of the freezing point as in the calculation of the heat of fusion of the salts have been used. The heat of fusion for calcium used in these calculations was the value 2,070 cal/mole given by Kubaschewski and Evans43. From these calculated activities the monotectic compositions were calculated assuming a salt molecule gives rise to two foreign particles on solution in the calcium metal. It has been assumed that there is no solid solubility of the solute in the solvent. For comparison the compositions corresponding to one and three foreign particles entering the metal on solution vvere also calculated. The result is given in Table 22.

127.

TABLE 22 Calculated values of the monotectic compositions - in the Ca-CaCl2, Ca-CaBr and Ca CaI2 systems. 2

No. of foreign Monotectic composition particles entering in mole % the metal Activity - - - Ca CaCl2 Ca CaBr2 Ca CaI2

N Ca 1 99.00 99.16 99.40 NCa + NCaX2 4- NCa NOaX2 2 99.50 99.6o 99.72 N + aa CaX2

N NOa CaX2 3 99.70 99.76 9982 N 2N Ca CaX2 128.

CHAPTER 4 DISCUSSION

4.1 Comparison with previous work. The phase diagram for the calcium-calcium chloride system obtained in the present work shows two remarkable differences from that by Taylor8. The first is the complete miscibility between the metal and the salt above 1338°C and the second is the depression of the freezing point of the metal by the salt. The highest solubility of calcium in calcium chloride given by Taylor was 9.82 mole % at 1250°C. As can be seen from Fig. 19 this is considerably lower than the value obtained in the present work. The reason for Taylor's lower results was found to be the segregation of the metal from the saltphase during the quench, and this problem has already been discussed in detail in connection with the initial experiments in Chapter 2. The depression of the freezing point of the metal by the salt was 12°C. The monotectic composition corresponding to this depression was calculated to be 99.5 mole % Ca under the assumption that there was no solid solution of calcium chloride in calcium and that the solution of one molecule of 129.

the salt gives rise to two foreign particles entering the metal phase. This calculated composition agrees well with Taylor's analysed value of 99.4 mole % Ca although he gave no depression of the freezing point. The eutectic composition of 5.25 mole % Ca given by Taylor is obviously too high, as one would not expect a quenched sample to give a solubility figure much less than the eutectic. The lowest solubility figures obtained from the quenched samples in the initial experiments were 3.74 and 3.8 mole % Ca. Taylor measured the eutectic and the liquidus temperatures on heating curves. In order to obtain a correct value of the liquidus on a heating curve, the liquid would have to be in equilibrium with the solid during the passage through the two phase region, so that the last remaining part of the solid disappears immediately when the liquidus temperature, corresponding to the overall composition, is reached. With a heating rate of 1 to 1.5 'C/Minute it is doubtful whether equilibrium could be obtained. Non-equilibrium condition would lead to too high results. Taylor's eutectic temperature is 8°C lower than in this work. This can be explained by the difficulty in interpreting 130. heating curves in two phase regions. The eutectic temperature was taken as that point on the heating curve where it first showed a reduced slope. This change of slope occurs so gradually that it is difficult to locate it accurately. In their work on the calcium-calcium chloride system Cubicciotti and Thurmond31 found a solubility of calcium in the saltphase of about 20 mole % at 825°C. The solubility then decreased to a minimum of about 15 mole % at 1000°C and then increased to about 30 mole % at 1200°C. The disparity between their work and the present appears to be due to differences in sampling techniques and to segregation. After equilibration and quenching they cut the crucible into three transverse sections. The upper section containing the metal was analysed as the metal phase. The middle section containing the metal-salt interface was discarded and the lower section was analysed as the salt rich phase. That this method of sampling can lead to high values can be understood from Fig. 9, which shows that the metal has crept down between the salt and the crucible wall. This was appreciated by Taylor as a result of his preliminary experiments. Using the same sampling method as 131.

Cubicciotti and Thurmond he obtained irreproducible values on dissolving the whole section in water, as his samples contained varying amounts of the metal phase. Taylor later avoided this source of error by crushing out the salt from the crucible. Other workers have only investigated part of the calcium-calcium chloride system. Thus Ostertag33 in her investigation of the solubility of calcium in calcium chloride at 1000°C obtained values ranging from 11.2 to 17.8 mole % Ca. To reach equilibrium between the phases she agitated the welded iron crucible containing the mixture for 1 hour in a horizontal position in a furnace maintained at constant temperature. To allow the phases to separate she kept the crucible in vertical position at the same temperature before quenching in water. It seems strange that the values she obtained are so high in spite of her quenching method. However, her sampling technique is not described in detail and her results may therefore suffer from the same errors as 31 Cubicciotti's and Thurmond's particularly as the scatter is so large. Peterson and Hinkebein 32 in their recent investigation of the calcium-calcium chloride system 132. found a solubility of 3.8 mole % Ca at 900°C and 4.2 mole % Ca at 950°C. They used more or less the same method as Ostertag but had their crucibles made of stainless steel tubes. The crucibles were agitated for 2 hours at the equilibrium temperature and then held 1 hour to allow separation of the phases before quenching in water. As only these two values are given for the solubility it is difficult to estimate the scatter in their results. It is, however, most probable that their results are too low due to segregation of the metal from the saltphase during quenching, particularly as their crucibles were rather large (internal diameter 7/16"). The saltphase they analysed would therefore not represent the composition at the equilibrium temperature, but was instead a salt-metal mixture with a composition close to the eutectic. Their thermal analysis gave a eutectic at 2 mole % Ca and 768°C. They also recorded a monotectic at 820°C and 99.5 mole % Ca, i,e. a depression of the melting point of their pure metal by 16°C. The composition of their monotectic is thus very close to the figure given by Taylor and corresponds exactly to that calculated from the measurements in this work. The depression of the 133.

freezing point of calcium chloride according to Peterson and Hinkebein would, however, lead to a heat of fusion of the salt of from 22 to 45 Kcal/mole depending on the choice of the model for the solution. As this is obviously too high, doubt must be thrown on their results. The melting point of pure calcium chloride obtained in this work was 774.5°C. This is to be compared with 770°C obtained by Taylor, and by. Peterson and Hinkebein. Seltveit and Flood47 in a recent investigation obtained a value of 771.6°C. The only previous work on the calcium-calcium bromide system is that by Taylor. Even in this system there is a discrepancy between his work and the present investigation in the monotectic. He gives no solubility of calcium bromide in calcium and no depression of the freezing point of calcium. The eutectic temperature, 728°C, obtained in this work on cooling is in close agreement with his 729°C obtained on heating. The discrepancy in the solubilities at higher temperatures has been discussed in detail in connection with the initial solubility measurements in Chapter 2. No previous work on the strontium-strontium 134. chloride and calcium-calcium iodide systems has been found in the literature. The melting point of pure strontium chloride obtained in this work, 872.5°C, is in good agreement with the value 873°C recommended by Kubaschewski and Evans43. The melting point of the pure calcium iodide, 779°C, is somewhat lower than the value given by McCreary461 783.7°C. The melting point of calcium obtained in this work, 838 ± 1°C, is in agreement with the value 837 ± 1°C obtained by Rogers7 and by Bevan48 using calcium from the same supplier. Peterson and Hinkebein34 give a melting point of 836°C and Cubicciotti and Thurmond31 828°C. The value recommended by Antropoff and Falk49 from their investigation in 1930 is 851°C. This value was obtained by extrapolation to zero content of nitrogen. Hoffmann and Schulze5° in 1935 repeated the experiments by Antropoff and Falk with calcium of a purity of 99.9%. They arrived at a value of 849°C and claimed an accuracy of ± 1°C. However, in view of their 0.1% of impurity they recommended the value given by Antropoff and Falk. In their thermal analysis the cooling curve did not give a definite break at the freezing point 135. but showed instead a gradual change in slope. They ascribe this small change as being due to the small heat of fusion. The freezing point was therefore taken as that point on the curve where it first showed a deviation from a straight line. This point would not correspond to the freezing point if the simple had a large thermal gradient. This possibility cannot be excluded, particularly when one considers the small change on their cooling curve. From the impurity figure of the calcium in the present work one would not expect this metal to have a melting point more than about 5°C too low. An analysis for nitrogen before and after experiments showed no indication of the presence of this element. In view of this the value 851°C recommended seems too high. 136.

4.2. Interpretation of the results. The different heats of fusion for the investigated salts given in Table 21 on page 125 were calculated considering three possible mechanisms for the solution of the metal in its molten halide, viz: 1. M + M++ -› M° m++ +-I- 2. M + M -> M2 3. m + M++ -› 2M+ From the calculated values it is not possible to differentiate between Model 1. and model 2. for the solution. However, the large difference between the actual recommended heat of fusion for strontium chloride and the calculated value assuming model 3. is an indication that this model for the solution is very unlikely. In the calcium-calcium iodide system the value 8.2 Kcal/mole for the heat of fusion of calcium iodide was calculated from the phase diagram CaI2 - CaF2 given by McCreary-46. This value is somewhat smaller than the value: calculated from model: 2. and much smaller than that calculated from model 3. It would also appear here, that model 3. is unlikely, especially as the value for the heat of fusion of calcium iodide estimated from the differential thermal analysis is about 7 Kcal/mole compared with 8.0 and 8.7 Koalimole 137. calculated from models 1 and 2, and this discrepancy is further reduced if one assumes that calcium is dissolved in solid calcium iodide. If the dissolution of strontium in strontium chloride and calcium in calciun iodide can thus be explained by either model 1 or model 2, it is reasonable to assume that the same is the case for calcium in calcium chloride and calcium in calcium bromide. In that case, the comparison of the heats of fusion for calcium chloride and calcium bromide, calculated from the extrapolated eutectic compositions, using either of these models, and the values given in the literature, serve as a check on these eutectic compositions. The monotectic compositions in the calcium-calcium chloride, calcium-calcium bromide and calcium-calcium iodide systems were calculated under the assumption that one molecule of the salt on dissolution in the metal gives rise to two foreign particles. The value thus obtained in the calcium-calcium chloride system agrees 8 very well with the analysed values by Taylor Rnd by Peterson and Hinkebein34. As the halide ion entering the metal phase is a comparatively large particle one would expect a decreasing solubility of the salt in the metal with increasing size of the anion. That this is the case can be seen from Fig. 23, where the temperature

138,

I I I I I I I 1 1 I 20 I I •

. ... U 0

OPP

URE. _ Ca /Ca Fa - T ERA MP 15 TE - _

- 0 Ca /Ca Cla _

- o Ca /Cobra -

- I I I 1 - I I I 1 iCa/C a Ia I I I I I I I I Li-) i - - 43 ;5. fa .13 II II Ai - 11 i I i I I 1... I

i i I ▪ I I I I I i I I ▪ I 1.0 1.5 2.0 ANION RADIUS IN ANGSTRoM UNITS.

FIG.23. THE DEPRESSION OF THE FREEZING POINT OF CALCIUM VERSUS THE ANION RADIUS OF THE ADDED SALT. 139. differences between the freezing point of the pure metal and the monotectic temperatures have been plotted versus the ionic radius of the different halide ions, taken from Pauling51. The displacement of the consolute composition from the salt rich towards the metal rich side of the phase diagrams with increasing size of the anion of the salt is a feature of the calcium-calcium halide systems that was also found in the potassium-potassium 25 halide systems. Johnson and Bredig interpreted this as being due to changes in the difference between the molar volume of the components and showed that the asymmetry in the potassium-potassium halide systems was reduced if volume fraction was plotted instead of mole fraction. In the calcium-calcium halide systeirs the molar volumes are not known at the consolute temperatures, but qualitatively one can say that the asymmetry in these systems would also be reduced if the volume fraction was introduced, using the values of the molar volumes at lower temperatures. The significance of the displacement of the consolute composition with the difference in the molar volume is, however, not so easily understood, particularly if one considers that there are several 140. other properties besides the molar volume that are related to the radii of the metal and the halide ions, e.g. polaris ability of the ions and the ionic 51 character of the bonds. According to Pauling the percentage ionic character of the halide bonds is given by the difference in electronegativity between the corresponding halogen and metal. If there is a relation between the molar volume and the consolute composition one wouldl thereforel also expect a relation between the electronegativity and this composition. That this is in fact the casei can be seen from Fig. 24, where the electronegativity of the anion has been plotted versus the consolute composition of the potassium-potassium halide and the calcium-calcium halide systems. FIG. 24.THERELATIONBETWEEN THEANIONELECTRO— I— ELECTRONEGATIVITY 0 =LLI NEGATIVITY ANDTHECONSOLUTE COMPOSITIONIN THE CALCIUM— CALCIUM HALIDEANDPOTTASIUM— POTTASIUM— HALIDE SYSTEMS. MOLE c t OFMETAL. 50 141. 142.

4.3 The nature of metal-molten halide solutions. The mutual solubility of metals and molten halides gives rise to an interesting series of solutions which are strongly ionic at one side and completely metallic at the other side of their phase diagrams. As the complete miscibility between metal and salt has not been appreciated until recently, attempts to explain the solubility has mainly been restricted to the salt rich solutions. Thus Cubicciotti52 ) on the basis of his investigations of the solubility of metals in their molten chlorides has proposed a theory for these solutions that implies that the metal goes into solution in octahedral holes between the anions of the molten salt. The valence electrons of the dissolved metal are thought to be transferred from the metal to bands belonging to the salt system as a whole. As the anions usually are larger than the cations they determine the main structure of the melt, which according to Cubicciotti can be regarded as a close packed arrangement of anions. According to this theory a salt with a larger number of available holes, e.g. a halide of a trivalentmetal should have a larger solubility of its metal than a halide of a divalent metal. Furthermore, the ratio of 143. the radii of the cation to anion in the salt would be important. ITVhen this theory was proposed there was some support for its validity. However, most of the phase diagrams on which the theory was based have been shown to be in error. Bredig et a23.24, from their investigations on alkali metal halide systems, suggest that the alkali metal enters the salt as an ion and that the valence electron, either takes the place of a halide ion5 or is + confined to a molecule M2 If the electron takes the place of a halide ion it might be mobile enough to contribute electronic conductance to the solution as a whole and thus give rise to development of conduction bands. On the metal rich side they suggest that the salt dissolves in an ionized form and that the anions arc substituted for electrons in the metal. If the salt goes into the metal as anions one would expect the solubility of the salt to be dependent on the size of the anions, as it would be more difficult for a large anion to enter the metal phase than a small one. A comparison of the solubilities of the calcium halides in calcium at the melting point of the metal supports this view. The solution of a metal in its halide as an 144. ionized double ion has also been suggested by Grjotheim et al.19 from their investigation on the cadmium-cadmium chloride system. The agreement between the heat of fusion of the chloride, calculated from the depression of its freezing point by cadmium using this model, and the known value supported this view. They also showed that the coloured metal-salt solution was associated with positively charged particles. Further- more, the solution was found to be affected by the addition of foreign cations. As the stability of the Cd2+4- ion would be dependent on its anionic environment, it was thought that the solubility was dependent on the contrapolarising effect of these added ions. Thus the addition of trivalent cerium ions which are strongly contrapolarjs'ng would stabilise the Cd21-+ ions and thus increase the solubility. This has also been found to be the case, while the addition of e.g. e ions lowers the solubility. Grjotheim et al. are thus of the opinion that the formation of

Cd2++ ions is responsible for the solution, and that thes,-) ions have a metallic type bond, similar to the one existing between the silver ions in the solid Ag2F. This view is also shared by Crawford9. He 145.

+4- has calculated the radius of the Cd2 ion from his density measurements in the cadmium-cadmium chloride system. According to his calculations the radius for 0 this ion is about 1.41. This is to be compared with 0 the single bond radius of 1.413A for Cd given by Pauling51. The solubility of a divalent metal in its ++ halides as M2 ions is also supported by the cryoscopic measurements of the present work. Calculations show that the solution of strontium in its chloride and calcium in its iodide by the formation of le ions is very unlikely. Of the two remaining possibilities the one with the metal going into the salt as atoms is also unlikely, as one would not expect an atom to remain unionised in a strongly ionic invironment. The formation of M24-4- ions in the melt should only be regarded as a first step towards metallic character of the solutions. Although at 50 mole% metal there would statistically only be M24-1- ions and halide ions present in the melt one would expect an interaction between these cations to occur before this concentration is reached. The interaction between the M2++ ions would, with increasing metal composition, successively lead to a completely metallic solution in which the anions are each 146. co-ordinated by several metal ions. This picture of the solution mechanism is supported by Crawford's conductivity measurements9. He found that the specific as well as the equivalent conductivity of cadmium in cadmium chloride solutions decreased to a minimum at about 10 mole % Ccl. On further increase in cadmium composition the conductivity rapidly increased. Taylor7 has compared the solubilities of different metals in their halides with the ratio of the radii of the cation to the anion in the salts. He points out that the solubility generally increases with this ratio, probably because the cations are closer together and thus make it easier for bonds to be formed with the added metal atoms. However, it is not easy to compare solubilities of metals in different metal-salt systems because of the difficulty in choosing a suitable temperature at which comparison can be mr..de. Even within metal-metal halide systems having a common cation, as e.g. in the calcium-calcium halide systems, it is not easy to make a comparison, when the salts have different melting points. Thus if the solubilities of calcium in its halides are compared at the melting points of the salts the calcium will be completely miscible in the fluoride, while the solubility in its other halides is very limited. If, however, the solubilities are compared just below the consolute points of the systems the calcium will be least soluble in its fluoride and the solubility will increase with the anion radius, i.e. increase with decrease in the electronegativity of the anion, as can be seen on Fig. 211. If the radius ratio of the cation to anion was the only determining factor for the solubility it would appear that a metal would dissolve in salts of other metals, e.g. magnesium would be expected to dissolve in calcium chloride. This is generally not the case. However, Heymann and Weber30 found that some metals, which are not soluble in a foreign metal salt, go into solution in the salt, if they are previously alloyed with the metal of the salt, probably as intermetallic compounds. Thus, although pure bismuth does not dissolve in sodium bromide at 770°C it goes into solution when alloyed with sodium. However, cadmium alloyed with sodium does not go into solution in the sodium bromide at the same temperature. The compound Na Bi is stable up to 775°C and has 3 a heat of formation of 11.4 Kcal/Mole, while the compound Na0d2 is only stable up to 385°C and 148. has a heat of formation of only 2.7 Kcal/mole. In conclusion, it may be said that the results from the present work support the view that calcium goes into solution in its molten halides by interaction ++ with the cations, probably forming Ca2 cations at dilute solution. It is plausible that the valence electrons at dilute solution are confined to these ions. However, as the concentration of the metal increases ++ the Ca2 ions successively start to interact. As the valence electrons are probably not strongly bonded, this interaction might lead to development of conduction bands so that in the region of complete miscibility there is a continuous change from a strongly ionic to a completely metallic solution, in which the anions are not confined to any particular metal ion. However, much more work is required before a model, such as the one suggested above)can be regarded as established. 149.

ACKNOWLEDGMENTS

The author is indebted to the United Kingdom Atomic Energy Authority, Harwell, for their financial support of this research and to Professor F.D. Richardson, Director of the Nuffield Research Group, for the provision of research facilities. He particularly wishes to thank his supervisor, Dr. J.W. Tomlinson, Nuffield Research Fellow, whose help has always been available when needed. Thanks are also due to his colleagues, in particular to Messrs. G.P. Jones, A.W.D. Hills, and D.H. Bell for many valuable discussions, and to the permanent staff of the Nuffield Research Group.

Nuffield Research Group in Extraction Metallurgy, Royal School of Mines, Imperial College, London, S.W.7. 150.

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