2488
KOTES
parisoii of the relative aciditlies of Br02+and NOs+ in fused I(xO3-T\TaSO3 could be made provided no oxides of nitrogen are formed during the course of the Cr201-2 BrOs- reaction. This turned out to be the case. The reactions involved are
+
Vol. 67 TABLE I RECIPROCAL P S E U D O H A T E C O S S T A N T WITH BARICM IONCOXCENTRATION [CrzO?-Z] = 0.1 m
\7ARIATION O F
T, OC.
10.5 50.8 76. 5.13 6.65 12.7 17. 2.2 4.1 8.5 14.
K
+ BrOs- , I Br02+ + 2Cr04-2 IC BrOz+ -+ products (Brz + 02)
Crz07-2
250
Experimental Materials and Apparatus.-A.C.S. reagent grade chemicals were used. All determinations were performed using a system similar to that described by Duke and L a n ~ e n c e . ~ Procedure.-KzCr2O7 was added to a solution of Ba(lU03)Z and NaBrOs in fused KN03-NaN03 solvent. The rate of reaction was studied by collecting evolved bromine in a sulfit solution and titrating with standard Agh'03 using eosin as an indicator. The reaction was allowed t o go to near completion. The rate of appearance of bromine was determined, and from this the rate of disappearance of total bromate (i.e., TB?= &OsBrOz+) may be calculated. The concentrations of dichromate and barium ions Tere always in excess of bromate; however, barium ion concentration was varied from run t o run to enable separation of the equilibrium constant from the rate constant. In all cases concentrations are expressed in molality units.
+
Results and Discussion Preliminary runs allowed selection of proper temperature and concentration ranges for convenient study. The Ba+2,because it precipitates the chromate ion but not the dichromate ion, controls the concentration of chromate through the solubility product. The experimental data gave a first-order plot in disappearance of total bromate, indicating decomposition of bromyl ion by some first-order mechanism to bromine and oxygen. T'arying the rate of the sweep gas, nitrogen, had no effect on the rate of the reaction; therefore, there appears to be no gaseous intermediate as found in the nitrate reactiona2 Upon substitution [BrOa-] = Tsr - [Br02+]into the equilibrium expression and replacing B r 0 ~ +in - dTn,/dt. = k [BrOz+]one obtains an expression
260
TABLE I1 SOLCBILITY OF BaCrOl IN EQUIMOLAR MIXTURE OF KNOI A N D NaNOs T,OK.
Solubility, m
532 522 512
2 . 0 x 10-8 1.8 x 10-3 1 . 6 x 10-3
TABLE I11 EQUILIBRITJJI ASD RATECONSTAKTS AS
+
Since [Cr207-2]and [Ba+2]are in excess k' =
IcK [Cr207-2] [Ba+212 KBP2
+
K[Crz07-2] [Ba+I2
the pseudo rate constant, remains constant for a given run. The reciprocal of k' is l/k'
=
l/k
1 + kK [ C r ~ 0 7 - ~[Ba+2]2 ] K,p2
A plot of l / k ' us. l/[Ba+2]2gives a straight line from which k and R are evaluated. The reaction was studied a t 230, 260, and 260'. The pseudo-first-order rate constants obtained from In T B vs. ~ time plots are listed in Table I. Solubility
A
FUYCTIOK OF
TEMPERaTCRE
T,O I i .
k, mk-1
503 523 533
0.13 .33 .58 TABLE
kK[Crz07-] [Ba+2]2 dTBr TB~ dt K[Crz07-2][Ba+2]2 K S p 2
0.10 .025 ,020 .070 .050 ,030 ,025 .10 .05 .03 .02
data of BaCr04 in NaNO3-KNO8 solvent are listed in Table 11. The solubility appeared to be independent of the relative amounts of potassium ion and sodium ion present but varied with temperature to give a straight line when the log of the solubility was plotted against reciprocal temperature. Table I11 lists the temperature dependence of the equilibrium constant and the rate constant to give the AH of reaction as 65 kcal. and activation energy as 55 kcal., respectively. Also both const,aiits varied with cation ratio of the solvent. The values obtained for variation in solvent are listed in Table IV.
\rARIATIOX IK
Using the solubility product for BaCr04, Ksp,one obtains
[Ba +z]
l/k'
230
K , mole-'
1.05 x 10-8 3.50 x 10-0 4.80 X
Iv
RATEAND EQUILIBRIUM C O N S T A N T S W I T H SOLVENT COMPOSITION AT 250"
Cation ratio (K/Na)
k , rnin-1
K , mole-'
60/40 50/50 38/62
0.30 .33 .38
3.2 x 3 . 5 x 10-8 4.8 X
THERMOPOTENTIAL MEASUREMENTS FOR NOLTES CADMIUM CHLORIDE, CADMIUM BROMIDE, A.ND LEAD CHLORIDE BY JACOB GREENBERG. DONALD E. WEBER,AND LAWREXCE H. TH4LLER Lewzs Research Center, Xalzonal Aeronaufzca and Space ddministralzon, Cleveland, Ohio Recezued -%fag 1 , 1965
In order to obtain some information concerning the possible use of molten salts as thermoelectric materials, the thermopoteiitials developed when molten CdCle, CdBr2, and PbClz are subjected to a thermal gradient were measured. These fused salts were pipetted into a Vycor 1T-tube, the two arms of which were heated to two different temperatures. The open-circuit voltage,
NOTES
Nov., 1963
which was recorded as a function of the temperature difference at the electrodes, was directly proportional to this temperature difference and appeared to be independent of the average temperature for the range and the temperature differences observed. It can be shown that the thermopotential is a function of the reduced heats of transfer and the transport numbers of the species involved. By using this relation, it is possible to estimate the entropy of transport for the salts. Experimental A piece of 8-mm. i.d. Vycor tubing was bent t o form a U-tube dpproximately 13 cm. in length with side arms that were 7 cm. long. The tube then was uniformly wrapped with nichrome ribbon and covered with Alundum cement. One side arm was then wound again with t,he nichrome and another layer of Alundum cement applied. The cell was heated by the use of two Variacs connected to the nichrome windings. After the cell had been heated to Borne temperature above the melting point, the salt was pipeited into it. Two platinum-platinum-13% rhodium thermocouples with 5-mm. platinum disks welded to the tips then were inserted into predetermined positions in the melt. The Variac settings were adjusted for an approximate temperature gradient, and the cell was allowed to reach a steadystate condition for several hours. Reagent grade salts were melted, washed with the appropriate halide gas, and filtered through Pyrex glass wool before being pipetted into the cell. Each datum point represents a t least five readings taken a t 20min. intervals. The data were considered significant if the temperatures did not vary by more than 0.1% and if the opencircuit voltage did not vary by more than 1%. The voltage readings were made from the platinum leads of the thermocouples by means of a K-3 Leeds and Xorthrup potentiometer with an electronic galvanometer.
Results The open-circuit voltage observed was directly proportional to the temperature difference a t the electrodes and appeared to be independent of the average
F
2489
= 1 faraday
A@ = open-circuit voltage
transference number of species k algebraic valence &k** = reduced heat of transport, T(&* - S k ) Sk * = entropy of transport of species k S k = partial entropy of species k A S = change in entropy due t o heterogeneous electrode renction when 1 faraday of electricity has passed through cell = reduced heat of transport of electrons in electrodes &el** tk =
Zk =
r-
loo
-
E,,/AT SALT
rndk.
CdCI,
0.600 CdBr2 .846 PbCI2 .349
---I
i E
YBr2
601 L
1
i, D
W
0 Fig. 1.-Plot
20
40 60 AT, "C.
80
IO0
of therrnopotentials of fused CdC12, CdBrz, and PbC12.
TABLE I
THERMOPOTEXTIAL DATAFOR MOLTENCdC4, CdBr?, AND PbClz Seebeck coefficient,
Salt
Range of av. temp., OC.
s
=
2= AT
mv./'C.
CdClz CdBrz PbCh
6 17-75 1 594701 529-594
0.600 ,846 .349
e, AT % deviation in s
10.96 f .25 1 .43
temperature for the temperature differences and range observed. The higher temperature electrode had a negative polarity. The plot of the thermopotentials observed against the temperature gradient, shown in Fig. 1, indicates a linear relation, which extrapolates to zero voltage a t zero temperature change for the cadmium systems. A t higher values of AT, the potentials observed for the lead chloride system showed positive deviations from linearity. The slope of the PbClz system was estimated from the first three points closest to the origin. These results are summarized in Table I. The primary sources of error were the failure of the systems t o reach steady-state temperatures under the experimental conditions and the possibility of reactions a t the electrodes. The general equation for the potential of a thermocell has been given by deGrootl and has been interpreted for molten salt systems by Sundheim.2 This equation is (1) H. Holtrtn, P. RZasur, and S. R. deGroot, Phyeic,, 19, 1109 (1953). (2) B, R. Elundhsiin and J . Roaenntreiah, J. Phum Chem., 68, 419 (19bQ)b
Since Qel**/T will be a t least an order of magnitude lower than the other values, we can consider its effect as being negligible. The value of AX is defined as the change in entropy due to the heterogeneous electrode reaction when 1 faraday of electricity has passed through the cell. In our case, this value would represent the change in entropy due to the decomposition of the salt into its elemental components. I n essence, AX refers to the difference in entropy of the ions as they exist in the liquid phase and at the electrode. This contribution mill not be zero a t vanishing AT if the decomposition products are present at the electrodes. Thke presence of two platinum electrodes immersed in a molten salt constitutes a decomposition cell. Electrolysis of the salt will occur when, in the process of determining the potential of the cell, the e.m.f. applied by the potentiometer is either above or below that of the salt itself. When the e.m.f. of the potentiometer is above that of the thermopotential of molten PbC12, for example, lead will be deposited a t one electrode and chlorine released a t the other. In the case where the voltage applied is below that of the salt system, the release of decomposition products will be reversed anld therefore tend to cancel the previous deposits. In the event that there remains an excess of electrolysis products, these are soluble in the melt and will have a tendency to diffuse from the electrodes. The presence of these products is easily detected since their accumulation a t the electrodes will create a back e.m.f. This effect will lower the readings observed and eventually
2490
NOTES
36 -
Vol. 67
F-
AT
=-
F -A@ =
2
_-
[SBalt*
- S s a ~ t l- 7
tM+a [&alt*
- Sealtl
+ Qx-**
NOTES
Nov., 1963
2491
lead ions leads to a change of approximately 15% in the value of the Seebeck coefficient for lead chloride. Acknowledgment.-We wish to express our appreciation to Dr. A. E. Potter for his valuable assistance.
THE THERMODYNAMICS OF ASSOCIATION OF THE TRISETHYLENEDIAMINEPLATINUM(IV) ION WITH VARIOUS AKIONS BY DONALD C. GIEDT AND C. J. NYMAN Department of Chemzstry, Washington State University, Pullman, Washington Received M a y .90, 1868
--t m + 2
Fig. 3.-Plot
*
of F( A@/AT) against - t M + 2 / 2 . Ssalt] -67 e.u.
-
Slope = [8*sa~t
potential in ionic crystals.* This latter equation is based on the models of Frenkel and Schottky for lattice disorder aind permits a more detailed analysis of the system. Since the structure of a fused salt corresponds to that of tht: crystal, the conduc1,ion process in the melt can be considered to be a function of ions moving through vacancies. Only those ions with sufficient kinetic energy to overcome the coulombic forces at their lattice position will participate. The negative polarity at the hot junction of the thermocell can be considered to be established by the immigration of cations to the cooler junction. Once in the lower temperature zone, they lose the energy necessary for their return. The thermopotential, as well as the conductivity of these systems, is a function of the number of mobile ions present and their ability to move through the liquid. If we assume that, a t about 25' above the melting point, the number of vacancies is much greater than the number of mobile ions, then the thermopotential is solely a function of the number of mobile species present. In this case the slope of the plot of A@ against AT (Seebeck coefficient) will vary in the same manner as the rate of formation of mobile cations. If we take the slope of the last three points in Fig. 2 for the PbCL system, the value of F ( A @ / A T ) is 9.31. This is a 15y0 change over the value of 8.09 given in Table 11. When the value of 9.31 is used to calculate QM +2**/Tin eq. 2, the value obtained is 68.6 e.u. (equal to S p b i 2 * - SPbt21. If Spb+Z* is relatively the same (-4 e.v.), this represents about a 3% change in S p b t2. Thus, a 3% change in the rate of formation of mobile (8) R. E. Howard and A. B. Lidiard, Phil. Mag., 2,1462 (1957).
The previously reported1 stability constants for the 1:1 outer-sphere ion pairs formed by trisethylenediamineplatinum(1V) with various anions were 11 for C1-, 8 for Br-, 3.3 X lo3for S04-2, ca. 0.8for r\;O3-, and 0 for C104-. These values supported King, Espensen, and Visco's2 claim that the stability constants for outer-sphere complexes reported by earlier workers were too high. Evans and IL'ancollas3plotted values of A S of formation for the ion pairs vs. A S of hydration values for the anion and found that these plots for the [CO(NHJ~]+~, [ C ~ ( e n ) ~ ]and + ~ ,Fe+a(aq.) cations were parallel lines. By means of an entropy cycle they intt:rpreted this close linear relationship to mean that t'he magnitudes of the entropies of association for ion pairs of the same cation but different anions were dependent upon the differences between the entropy of hydration of the anion and the entropy of hydration of the ion pair. The present work was undertaken to study the effect of temperature on the stability of the outer-sphere ion pairs and to establish values for the thermodynamic functions AH', AF', and AXo for the formation rea,ctions of [Pt (en)3]+4 with various anions. Experimental Solutions.-All stock solutions were prepared as previoudy reported.' It was necessary to deoxygenate the water used in the preparation of the iodide and thiocyanate solutions by sweeping it for 30 min. with 2 stream of nitrogen prior to use. The solutions prepared for the [Pt(en)3f4,C1-] ion pair study contained 1.017 X lO+M [Pt(en)l]+4and1.003 X lO-aM [H*], with a chloride concentration varying from 0.005 to 0.1 Id. The solutions prepared for the other anions studied containled 1.016 X IObs M [ P t ( e r ~ ) ~ ]1.208 + ~ , X lO-lM [H+], and 1.208 X M [C104-]. The concentrations of the complexing anions in these solutions variied from 0.001 to 0.05 M . The acidities of all final solutions were maintained a t p H 3 or less to suppress acid dissociation of the [Pt(en)r]+4. The ionic strength varied depending on the composition of the solution. Procedure.-A Cary Model 14 recording spectrophotometer was used to measure the absorbance of the solutions in matched 1cm. quartz cells. The apparent molar extinction coefficients D were 9easured a t 2600 A. for the C1- and Br- anions and a t 2700 A. for the I- and SCN- anions and at temperatures of lo', 25', and 40". A constant temperature in the cell compartment of the spectrophotometer was maintained by circulating water of (10.0 f O.Z"), (25.0 f 0.2'), or (40.0 f 0.2') as required.
Results and Discussion The spectrophotometric data were evaluated by the same method as employed by Bale, Davies, and Monk4 and later by Kyman and P1ane.l Activity coefficients (1) C. J. Nyman and R. A. Plane, J . Am. Chem. Soc., 82, 5787 (1960). (2) E. L. King, J. H. Eaipensen, and R. E. Visco, J . Phys. Chem., 63, 765 (1959). ( 3 ) M. G. Evans and G. H. Nancollas, Trans. Faradaw SOC.,49,363 (1853). (4) W.D.Bale, E. W. Davies, and C . B. Monk, ibi.d, 52,816 (1956).
