THEPLATINUM ELECTRODE IN HALIDE MELTS
Oct., 1956
0.16RT. Discrepancies between observed and calculated excess free energies of mixing of the order of O.1ORT may be due to any number of causes, which are ignored in the derivation of the regular solution-solubility parameter equations. In this light, therefore, a value of 0.16RT for the C2Fs-CHF3 system is seen to be in some disagreement with theory. I n our first paper1 we suggested that the difference in physical properties of CF4and CHF3as compared to the great similarity in properties of CC14 and CHC13 might lie in the possibility of hydrogen bonding in CHFa. The results for the C2F6-CHF3 system would seem to level support to this point of view. The possibility of hydrogen bonding in fluoroform is also evidenced by its behavior with ethane and xenon. With these two substances it forms two phases below a consolute temperature of about 186°K. From equation 3 this gives an experimental (61 - a2) difference of 3.9 and 4.1 for CzHs-CHF3 and Xe-CHF3 mixtures, respectively. The “thermodynamic” 6 difference for these systems are 1.6 and 1.5; from 6 differences of this magnitude one would expect rather small deviations from ideality. The discrepancy between the observed and calculated excess free energies of mixing for the CH2F2-CHF3 is within the limits of experimental error. This result together with the high 6 value for CH2F2would seem to suggest hydrogen bonding in this compound also. If the hydrogen bonding between like and unlike pairs is about equivalent, then there would be no net contribution to the
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energy of mixing, and the solution of CHF3 with CH2F2 would be nearly ideal, as it is. If the thermodynamic 6 values for the CH2F2C2H6, CH2F2-C2F6, CHZFz-CF4 and CHF3-CH4 systems are substituted in equation 3 consolute temperatures well above the boiling point of the lowest boiling component are expected. ThiR is borne out by the experimental results, all these systems having two phases above the boiling point of one of the components in the mixture. On the basis of the foregoing discussion it would be expected that the consolute temperatures for CH2F2C2H6 and CHF3-CH4 would be well above the calculated consolute temperatures. The systems CH2F2-CF4 and CHzFz-CzFswould probably form one phase somewhat above the calculated consolute temperatures. It is unlikely that any of these consolute temperatures could be attained without the use of high pressure apparatus. We have no information on the volume changes occurring during mixing for these systems and little information on the ionization potential for most of these compounds studied so that we are unable to apply the corrections of Simons and Dunlaps or Reed.6 Acknowledgments.-This work was supported by the Atomic Energy Commission under project 13 of contract AT(l1-1)34 with the University of California. We wish to thank the Jackson Laboratories of E. I. du Pont de Nemours Co. for their generous gift of the fluorochemicals used in this work. ( 5 ) J. H. Simons and R. 0. Dunlap, J . Chem. Phus., 18, 335 (1950). (6) T. M. Reed 111, THIS JOURNAL, 69, 425 (1955).
ELECTRODE POTENTIALS IN FUSED SYSTEMS. 111. THE PLATINUM ELECTRODE I N SOME HALIDE MELTS’ BY KURTH. STERN Department of Chemistry, University of Arkansas, Fayetteville, Arkansas Received June 4, 1066
The galvanic cells Ag/A C1, KCl; KCl/Pt, Ag/AgBr, KBr; KBr/Pt, and Cu/CuCI, KCl; KCl/Pt have been studied in the temperature range 880-o-OOOo. It is shown that the K+-K(g) equilibrium operates on the platinum electrode a t pressures near 10-Ia atmosphere.
Introduction In the previous paper of this series12a study of the cell Ag/AgCl, KC1/Cl2 showed that for concentrations of AgCl less than 0.05 mole fraction the reaction Ag KC1 = AgCl K proceeds spontaneously. It was suggested that the change of potential with time could be used for a study of reaction kinetics. Indeed, a preliminary investigation showed that over a considerable time span for a cell which was Ag/KC1/CI2 initially dE/dt was nearly constant, suggesting first-order kinetics. A kinetic study of this sort requires a reference electrode. One possible objection to the chlorine
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(1) This research was supported by the United States Air Force, through the Ofice of Scientific Research of the Air Research and Development Command. ( 2 ) K. H. Stern, THIS JOURNAL, 60, 679 (1956).
electrode used in the previous study for this purpose is that the dissolved C12 in the melt may react with the metal under consideration and thus give spuriously high rates. Hence an inert. metal electrode offers an attractive possibility as reference electrode, particularly if the nature of the reaction taking place on the electrode can be established. For halide melts such an electrode is platinum. Fusion of alkali halides in platinum dishes, for example, has long been standard procedure for the removal of traces of moisture in the preparation of ultra-pure salts. This can be considered evidence that chemists generally recognize the inertness of this metal in fused salts. The present work is concerned with elucidating the electrode reaction on platinum in halide melts. Results are presented for the cells Ag/AgCI,KCl;
KURTH. STERN
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KCl/Pt, Ag/AgBr,KBr;KBr/Pt, and Cu/CuCl, KCI;KCl/Pt, where the semicolons represent liquid-liquid junctions between solutions of nearly the same composition. Hence liquid junction potentials are virtually negligible. Experimental Part Materials .-Copper and silver electrodes were made from spectroscopically pure No. 10 B. and S. gage wire. Platinum electrodes were made by braiding several strands of No. 18 wire into a small cable. All salts were Mallinckrodt Reagent Grade. Except for CuCl they are better than 99.9% pure. The assay on CuCl is 90%. KC1 and KBr were dried at 110” for several days before use. AgCl and AgBr were dissolved in NHaOH and reprecipitated with HC1 and HBr, respectively. They were dried at 110’ and stored in a desiccator. CuCl was used without further purification. Experimental Methods.-Vycor cells were made by bending lengths of 25 mm. diameter tubing into U-shapes in such a way as to constrict the bend of the “U” to capillary dimensions. Cells were about 10 cm. high. While the opening between the legs of the “U,” the electrode compartments, was large enough to permit electrical contact through it, it was small enough to prevent diffusion between the compartments for several hours. The furnace was similar to the one described previous1y.z Measurements were made both under argon and in melts open to the atmosphere. No significant differences were observed. In the case of copper, oxide formed on the electrode above the melt, but the electrode remained bright in the melt, even when cells were open to the atmosphere. The temperature was measured with a chromel-alumel thermocouple whose hot junction was kept near the bottom of the furnace touching the cell. The couple was calibrated at the melting point of KC1 and kept in a nitkel protection tube. The temperature was constant to f 2 Runs were carried out by heating the Vycor cells containing 15 g. of the alkali halide in each leg to the desired temperature, dissolving a quantity of AgCl, AgBr or CuCl in the proper compartment, immersing the electrodes, and measuring the potential when tem erature equilibrium had again been reached. Since the celE were open to the atmosphere salt concentration in them could be varied by adding the salt through a funnel. Measurements were made as rapidly as possible to minimize diffusion between compartments. Five or six concentrations were prepared in approximately one hour. A t the conclusion of a run the electrodes were withdrawn and the cooled cells examined for evidence of diffusion between the compartments, as shown by a violet color in the case of silver (after exposure to bright light) or a greenish-brown color for copper. If diffusion had occurred measurements from these cells were discarded. In many cases analyses of the melts were carried out by se arating the electrode compartments, i.e., breaking the c e i at the bend of the “U,” and dissolving their contents in water. Analytical Procedure. Silver.-Silver was determined spectrophotometrically as the colloidal sulfide when present in concentrations of a few parts per million. Large quantities were precipitated as chloride. Copper.-It was initially determined that essentially all of the copper present in the melt was in the cuprous state. This was done by dissolving the quenched melt in a solution of potassium thiocyanate. Copper( I ) immediately formed the insoluble thiocyanate and copper( 11) was not found to any appreciable extent in the filtrate. I n subsequent analyses of the melt copper(1) was oxidized to copper(11) and determined by one of two procedures. Small quantities, of the order of a few parts per million, were determined spectrophotometrically by the diethyl dithiocarbamate method.3 Larger quantities were determined ae the sulfide.
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to be the metal, for AgCl and KC1 the pure salt, and for the (provisionally) unknown reaction on the platinum the reduced species at a to be determined concentration. TABLE I
E.M.F.’s
FOR THE
CELLAg/AgCl, KC1;KClJPt t = 875”
Ni
E (v.)
EQS
0.000814 .00146 .00377 .0121 .0207 .0353
0.784 .747 .680 .565 .517 .465
0.080
.IO1 .128 .128 .I33 .I34
EO2 values tend to constancy near 0.01 mole fraction AgC1. This is in faic agreement with the results of the previous study2 on the cell Ag/AgCl, KCl/C12 in which the reaction Ag KCI = AgCl K was shown t o go spontaneously when the concentration of AgCl was below 0.05 mole fraction. EO2 values for the very dilute solutions are thus thermodynamically meaningless. I n Table I1 are given the ,302 values for the three cells of this study a t several temperatures. Each of the EO2 values is the result of a run such as is shown in Table I and represents the thermodynamically reversible value.
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EO2
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TABLE I1 VALUESFOR SEVERAL GALVANIC CELLS EO’ (volts)
705 827 862 875
Ag/AgCl, KCl; KCl/Pt 0.147 f 0 . 0 0 2 .140
.130 .131
760 832 865 008
Ag/AgBr, KBr; KBr/Pt 0.303f0.002 ,276 .269 .241
820 830 840 895
Cu/CuCl, KC1; KCl/Pt 0.414f0.01 .433 .440 .418
Results and Discussion Table I shows the results of a typical run on the cell Ag/AgC1(N1),KC1;KC1/Pt. N 1 = mole fraction AgCI, E = measured potential, EO2 = the standard potential, taking the standard state of Ag
Taking the results of the previous study2it seems reasonable to adopt the hypothesis that the electrode reaction on the platinum electrode is the reduction of potassium ion, ie., K++ e- = E(, and proceed t o calculate the pressure of potassium vapor at which the equilibrium operates. The method will be illustrated for the cell containing AgCl but is analogous for the other cells. The standard potential EO1for the reaction Ag KC1 = AgCl K can be calculated from recently published data on the Eo of formation of AgCl and KC14 by the method outlined there. The standard state of EO1 differs from that of EO2 only in that for EO1 the pressure of potassium vapor is taken as one atmosphere. Hence E o 2 = EO1 RT/S In P K and
(3) E. B. Sandell, “Colorimetric Determination of Traces of Metals,” Interscience Publishers, New York, N. Y., 1944, p. 221.
(4) W. J. Homer, M. S. Molmberg and B. Rubin, J . Electrochem. soc., 103, 8 (Is5a).
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Oct., 1956
THEPLATINUM ELECTRODE IN HALIDE MELTS
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hence the pressure of potassium vapor operating on the platinum electrode can be calculated. For the cell involving AgBr EO1 values were calculated from the table of thermodynamic functions prepared by the Bureau of standard^.^ It is instructive t o compare the pressures of potassium vapor with that available from the dissociation of the potassium halide. The latter is calculated from the Eoformation = - Eodisaociation by the equation
within a little over one order of magnitude. This is not too surprising, considering the approximations that are made in the calculation of thermodynamic functions of substances at high temperatures. These approximations are discussed in the references cited. Thus, for example, an error of 3 kcal. in one of the AFO's used would change PK by an order of magnitude. At any rate, the results imply that between 92 and 99% of the potassium vapor available from the thermal dissociation of KCl distils into the atmosphere, the remaining 1 to 8% is adsorbed on the platinum electrode. The increase of the dissociation pressure of the I n all the cells the fraction of potassium vapor on potassium halide with rising temperature parallels the electrode is between 0.001 and 0.1 of the disso- the rise of equilibrium vapor pressure of potassium ciation pressure of the potassium halide. The re- in the reaction Ag KCl = AgCl K. No exsults of the calculations of potassium vapor pres- act calculation of the latter is possible because the sures are shown in Table 111. equilibrium ratio NAgCl/NKCl is only known within an order of magnitude. Taking this ratio as 0.01 TABLE I11 and the activity of solid silver as one, PK can be POTASSIUM VAPOR PRESSURES (ATM.)IN VARIOUS CELLS estimated4 from the equilibrium constant &. For example, a t 800" Q = 7 X 10-l2, from which PK = 7 X 10-lo. At 875" Q = 7.6 X lo-" and PK = Ag/AgCl, KCl; KCl/Pt 7.6 X lovg, assuming no change in the equilibrium ratio of AgCl to KCl. These are equilibrium pres795 1.15 2.20 5.2 sures attained only in closed systems. In open 827 3.15 5.00 6.3 systems the actual pressures are much lower and 862 13.7 15.4 8.9 the reaction is driven by the distillation of potas875 20.2 21.2 9.5 sium, as shown previously.2 Ag/AgBr, KBr; KBr/Pt It seems reasonable t o conclude that platinum 760 0.279 1.0 2.8 can be used as reference electrode in fused alkali 832 7.96 6.4 1.2 halide systems, even though the potassium pres865 19.0 29.0 15.5 sure at which it operates is known only within an 908 373.0 50.0 75.0 order of magnitude. Within this range its potential, whatever it is, remains constant and this is the Cu/CuCl, KCI; KCl/Pt chief requirement of a reference electrode. 820 0.24 4.5 0.53 At the present time no potential can be assigned 830 0.81 5.9 1.4 to any half-cell reaction in molten system since 840 1.1 7.8 1.4 there is as yet no common agreement as to a zero of 895 9.5 38.0 2.5 potential analogous to the H+-Hz couple in aqueI n general, the fraction of the total potassium ous systems. vapor participating in the electrode reaction inAcknowledgment.-I would like to thank Dr. creases with rising temperature, but the results Jack K. Carlton, now with the Chemistry Departwithin each of the different cells show variations ment, Georgia Institute of Technology, for the (5) Circular 500, National Bureau of Standards. analytical determinations.
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