RALPHP. SEWARD AND JEROME P. MILLER
3156
Galvanic Cells with Molten Bisulfate Solvents
by Ralph P. Seward and Jerome P. Miller' Department of Chemistry, The Pennsylvania State University, University Park, Pennsylvania (Received April $1, 1066)
Fused bisulfates, mainly ammonium bisulfate, have been used as solvents in galvanic cells which appear to operate reversibly with silver and mercury(I), mercury(I1) platinum electrodes. The latter electrode with excess solid mercury(1) sulfate and mercury(I1) sulfate present served as the reference electrode. The dependence of cell e.m.f. on mercury(I1) concentration has been investigated. The silver electrode is shown to behave ideally with respect to silver ion concentration but to be dependent on the acidity or basicity of the solvent in an unexpected way.
While galvanic cells have been most fruitfully employed in evaluating the thermodynamic properties of fused salt solutions, their use has largely been confined to cells having fused nitrates or halides as solvents. The investigation reported here was undertaken with the object of finding other fused salts which would serve as solvents in reversibly operating cells. Measurements on cells with bisulfate solvents, mainly ammonium bisulfate, are described below and the results are discussed. The strongly acidic nature of fused bisulfates precludes the use of more active metals as electrodes, but silver has been found to be satisfactory and platinum inert. Solute salts have been limited to sulfates since most common anions would react with the solvent forming volatile or unstable acids.
Experimental Materials. Mercury(1) sulfate was prepared by the reaction of aqueous mercury(1) nitrate and sulfuric acid solutions. Other salts were commercial reagent quality chemicals. Ammonium bisulfate after 2-3 hr. at 110" under about 1 mm. pressure melted at 145'. Potassium bisulfate as received melted at 212". Sodium bisulfate, obtained as the hydrate, was dehydrated under vacuum at about 110". It then melted at 183". When analyzed for bisulfate hydrogen and ammonium by the method of Kolthoff and Stenger2 the ammonium bisulfate was found to contain excess acid, in three different lots amounting to 0.42, 0.48, and 0.66 wt. % when calculated as HzS04. The excess sulfuric acid content of the sodium bisulfate was 2.1%. Some maximum freezing point (100%) sulfuric acid The Journal of Phy8ka-l Chemisfry
was prepared in order that the effect of excess sulfuric acid and of excess base (ammonium sulfate) on the cell e.m.f. values could be explored. Electrodes. A platinum wire immersed in fused bisulfate in contact with an excess of both solid mercury(1) sulfate and solid mercury(I1) sulfate served as the reference electrode. The reference electrode, a glass tube about 10 X 1.0 em., was suspended in a larger tube (3.5 em. diameter) where it was surrounded by a solution of variable content with which it was in electrolytic contact through an asbestos fiber junction. Dipping into the solution in the outer tube were also a thermistor probe, a thermocouple junction, a mechanical stirrer, and a second electrode. Two cells designated as A and B have been investigated. In cell A, the outer electrode was the same as the reference electrode except that the solution in contact with the platinum wire was unsaturated, and of variable concentration in mercury(I1) sulfate. In cell B, the electrode in the outer chamber was a piece of silver foil. Operation. Weighed amounts of the desired materials were placed in the larger tube, brought to 160", and stirred continuously during measurements. Excess mercury(1) and mercury(I1) sulfate was stirred in fused ammonium bisulfate and this mixture then added to the reference electrode. After 0.5 hr., the silver foil electrode was dipped into the melt and within 1 (1) From the Ph.D. Thesis of J. P. Miller, The Pennsylvania State University, June 1964; supported by the U. S. Atomic Energy Commission under Contract AT(30-1)-1881. (2) I. M. Kolthoff and V. A. Stenger, "Volumetric Analysis," Interscience Publishers, Inc., New York, N. Y., 1947.
GALVANIC CELLSWITH MOLTENBISULFATE SOLVENTS
3157
Table I: E.m.f. of Cell A as a Function of HgSOd Concentration at 160" lOs( HgSO4)
E, v.
1.134 0.06900
1.295 0.06620
1.590 0.06083
1.788 0.0583
min. stable e.m.f. measurements could be made.
The silver foil was not kept continuously in the melt but pulled out and allowed to cool except when meaaurements were being made. This was done to minimize the attack of the solvent on the metal. An experimental determination showed that silver immersed in ammonium bisulfate at 150-160' lost 0.17 mg./cm.2 hr. Under the conditions of cell operation this is not sufficient to alter the silver content in the surrounding solution by a significant amount. Moist nitrogen was bubbled through the cell to prevent decomposition of the solvent into pyrosulfate and water. The thermistor probe actuated a relay to control current in a furnace which surrounded the cell, maintaining the temperature at 160 1". Temperatures were obtained from a chromel-alumel thermocouple and all of the e.m.f. values were measured With a beds and Northrup Type K3 potentiometer, a Leeds and Northrup Type E galvanometer serving as null point detector. Solubility Measurements. Excess solid was stirred in molten ammonium bisulfate for about 2 hr. at 160' and, after standing 0.5 hr. to permit excess solid to settle, samples of clear solution were removed by means of B preheated pipet. After cooling and weighing the samples, they were dissolved in water and mercury(1) precipitated and weighed as Hg2Cl2, mercury(I1) as HgS, and silver as AgC1. Solubilities in moles of solute per mole of ammonium bisulfate at 160' were found as Hg8O4 0.00065, HgS04 0.0085, and fig2804 0.0215. V6cosities. The viscosities of certain solutions were measured, using Cannon-Fenske viscometers. A Westphal balance was employed to obtain the densities needed for calculation of the viscosity.
ptl
f
0.06
0.05 2.6
l
2.7
2.8
-Log (HgSOr).
2.8
Figure 1. Dependence of e.m.f. of cell A on HgSO4 concentration.
various concentrations are given in Table I and graphically in Figure 1. Concentrations throughout this paper are the ratios of the number of moles of the species whose formula appears in parentheses to the number of moles of bisulfate ion in the solution, The effect of variation in sulfate ion concentration was studied by adding ammonium sulfate to the cell. The results are presented in Table 11.
Table 11: Variation of E.m.f. of Cell A with Changes in Sod*- Concentration (HgS04) = 2.23 X 10-8 lO'(SO~*-) 3.41" 7.08 9.61 12.5 26.6 E, v. 0.0487 0.0637 0.0668 0.0694 0.0868
Excess acid; the figure iS that of (H2S04).
It is assumed that junction potentials are a negligible
HgzSOi(c), HgSOi
NH4HS04
2.250 0.0484
Q
a
reference electrode
2.211 0.04950
0.07
*
Results and Discussion Cell A may be formulated as
1.839 0.05581
(A)
where (c) indicates the crystalline state and the reference electrode is the one described above. Hence the left-hand electrode difFers from the reference electrode only in having a lower and variable concentration of HgSO,. Electromotive force values for
factor in all of the measured e.m.f. values. Transport through the junction should be essentially all by NH4+ and HSOI- ions and the solutions sufficiently dilute that the activities of these ions on each side of the junction should not W e r by a si&cant amount. If the electrode reaction is Hgz2+ = 2Hgz+ 2e-, the e.m.f., for ideal solute behavior and with 2.3RTIF = 0.0861 v. at 433°K.) should be
+
RALPHP. SEWARD AND JEROME P. MILLER
3158
E
=
E"
(Hg2+), + Eref- 0.0861 -log 2 (Hgz2+) ~
(1)
In the presence of excess solid HgzS04, (Hg22+) = Ksp/(S04,-). If complete dissociation of HgS04 and no reaction of SO4,- with the slightly acidic solvent to form HSO4- are assumed, (Hg2+) = (HgSO,). Hence
This equation is not consistent with the rise in e.m.f. with increasing (SO4,-) observed and shown in Table 11. If incomplete dissociation of HgS04 is assumed, (Hg2+) = (HgS04)Kep/(S042-)which gives
E
=
E"
(HgSO,) 2K2eq + Eref- 0.0861 2 log (S042-)Ksp
(3)
The slope of a plot of E 21s. log (HgSO4), if (S042-) is assumed constant, should be 0.0861 v. The line shown on Figure 1, drawn with this slope, is consistent with the experimental values at higher concentrations of HgS04. Equation 3 predicts qualitatively the dependence of E on (S042-) which was observed. The slope of E us. log (S042-) is 0.031 rather than the predicted value of 0.043. This disagreement could be the result of nonideal behavior, related to that shown below for silver sulfate solutions. If the concentration of HgS04 on the left side of cell A is increased to saturation, the two electrodes become identical and the e.m.f. must become zero. If the straight line in Figure 1 is extrapolated to E = 0, the concentration corresponding to this point is found to be 0.0083, which agrees with the measured solubility within the experimental uncertainty of the latter. Cell B had a silver electrode and silver sulfate in solution in the outer electrode chamber. With the same reference electrode as cell A, cell B may be formulated as Ag~YH4Hso4~ lreference electrode
(B)
The e.m.f. of cell B was measured as a function of silver sulfate concentration including the special case of saturation, and as a function of the acidity or basicity of the solution as altered by adding varying amounts of H2S04and (NH&SOI. The measurement with excess solid Ag2S04was made with the idea that the e.m.f. might be calculated and agreement with the observed value would show that the electrodes were truly reversible. The cell reaction in this case is The Journal of Physical Chemistry
+
+
2Ag(c) 2HgS04(c) S042-(anode) + Hg,SO4(c) Ag,SO4(c) S042-(cathode)
+
+
(4) Using values of standard enthalpies of formation and entropies as given by Latimer3 for the four solids involved, AG"za8 = -15.6 kcal., which corresponds to an e.m.f. of 0.338 v. The observed value was 0.3335 v., in reasonable agreement. Since the heat capacity of HgS04 has not been measured and the entropy of HgS04 was only an estimated value, a more accurate comparison, taking into account the difference in temperature and the S042- activities, could not be made. When ammonium sulfate was added to the solution around the silver electrode the e.m.f. did increase, as it should according to the cell reaction equation. The increase, however, was much smaller than that predicted on the assumption that sod2activity is proportional to S042- concentration. The effect of added sulfate is described further below. Since the solvent originally contained a small excess of sulfuric acid, ammonium sulfate was added to neutralize it. As the cell e.m.f. was changed when this was done, the effect of adding sulfuric acid and ammonium sulfate was investigated extensively with solutions having several different silver sulfate concentrations. For any solution containing a definite quantity of silver sulfate, addition of sulfuric acid increased the cell e.m.f. and addition of ammonium sulfate decreased it. These changes were reversible in that the e.m.f. may be altered by adding either the acid or base and then restored to its original value by addition of an equimolar amount of the other. The silver electrode thus acts as an indicator for this titration although why it does so is not clear. This behavior is shown graphically in Figure 2, where the parallel curves become linear with a slope of -0.30 at higher sulfate concentrations and become slightly steeeper on the acid side. All sulfate concentrations have been calculated on the assumption that the reaction HzS04 S042- = 2HSO4- is complete. Hence the number of moles of Sod2-is assumed to be the sum of moles of sulfate added, as ammonium sulfate and silver sulfate, less the number of moles of sulfuric acid present or added. Where the acid is in excess, the concentration is recorded as (H2S04). The decrease in E with increasing sulfate concentration is opposite to what had been expected. A decrease in E corresponds to an increase in silver ion activity. Any complex ion formation with silver ion or simply the effect of the presence of doubly charged anions around the silver ions would be expected to decrease the silver ion activity.
+
'
(3) W. M. Latimer, "Oxidation Potentials," Prentice-Hd, Inc., New York, N. Y., 1952.
GALVANIC CELLSWITH MOLTEN BISULFATE SOLVENTS
The dependence of the cell e.m.f. on silver ion concentration was investigated by taking values of E at three arbitrarily selected sulfate concentrations from each curve of Figure 2. When these E values were
0.47
0.45
0
10 lO'(HBO4)
10
20
30
lOa(SOr'-)
Figure 2. Dependence of e.m.f. of cell B on acid and base concentration: 10*(Ag+)for I, 0.80; 11, 1.18; 111, 1.68; IV, 2.58; V, 2.88.
plotted against the logarithms of the silver ion concentrations, three parallel straight lines were obtained. The slope of these lines was -0.086 f 0.002 v., in agreement with the value predicted for ideal behavior of the silver ion. To include nonideal behavior as well, the cell e.m.f. may be described by the equation
E = E"
+ ER - 0.086 log (Ag+)yAg+
(5)
where E" is a standard potential for the silver electrode, EX the reference electrode potential, and y an activity coefficient. If y is assumed to be unity in those solutions for which the calculated sulfate concentrations are zero, on substitution of the measured e.m.f. and silver ion concentration, a numerical value of E" ER may be obtained. When this was done for six solutions with silver ion concentrations varying from 0.8 X to 2.88 X and e.m.f. values from 0.487 to 0.441 v., values of E" EE from 0.2200 to 0.2221 v. were obtained, the mean value being 0.2206 v. Activity coefficientsfor the silver ion in solutions containing excess H$04 or SO4+ may be obtained from the curves of Figure 2. For the region of excess sulfate ion concentration where the curves are very close to straight lines,the activity coefficient is empirically
+
+
log y = 4.5(SOd2-)
(6)
The utility of this relation will be illustrated by a calculation of the solubility of AgzS04 in pure ammonium bisulfate. When excess solid AgZSOr was
3169
present and enough ammonium sulfate was added to neutralize the original excess of sulfuric acid, the cell e.m.f. was 0.3339 v. In this solution (Agf) = 2(S02-) and log y = 2.25(Ag+). On substituting this for log y in eq. 5 and solving for the silver ion concentration, (A@;+)= 0.0392, and thus (Ag9SO4) = 0.0196. This value is thought to be more nearly correct than the measured value of 0.0215. The slight excess of acid in the solvent which had been used for the solubility measurement had not been neutralized. This acid, by reaction with the sulfate ions from silver sulfate, would cause the silver ion concentration to be higher than it would have been in a neutral solvent. As the changes in cell e.m.f. which accompanied increase in sulfate ion Concentration could not be accounted for in terms of a simple silver-silver ion electrode reaction, various changes in the solution were investigated. A cell was set up identical with cell B except that the ammonium bisulfate solvent was completely replaced by an equimolar mixture of sodium bisulfate and potassium bisulfate. The titration curve here was of the same nature as that for the ammonium salt solvent, decreasing e.m.f. with increasing sulfate content. A cell having (Ag+) = 12.6 X had an e.m.f. of 0.3860 v. at the point corresponding to zero sulfate ion content. The e.m.f. for this concentration of Ag+ when calculated by eq. 5 is 0.3841 v. It is concluded that the cell characteristics are not significantly dependent on the presence of the ammonium cation. Other substances which might form in the cell owing to solvent decomposition are water and pyrosulfate. The effect of addition of each was investigated. Passing water vapor into the cell did not produce any detectable change, nor did addition of potassium pyrosulfate. The viscosity of ammonium bisulfate with various additions of ammonium sulfate was measured at five different temperatures in the 160-180O range. As a bisulfate melt is a favorable situation for hydrogen bonding, the structure of the melt may well be quite different from simple random mixing. It was thought that the anomalous effect of sulfate ion on the silversilver ion electrode might be caused by a change in the structure of the medium accompanying addition of sulfate ion. Should this be so, evidence of such a change might show up in viscosity. Experimentally, it was found that the viscosity of the melt increased on the addition of ammonium sulfate, and also on the addition of silver sulfate. The temperature coefficient of the viscosity, however, was found t o be independent of sulfate concentration. To f2%, the observed viscosities in centipoises are given by 71 = 0.00197. [l 10(5042-)]ea'O/T. Although the viscosity in-
+
V d u m 69,Numbsr 9 Sortamber 1966
3160
L. BARTON, S. K. WASON,AND R. F. PORTER
creases caused by rather small increases in sulfate ion content do suggest a significant change in the struc-
ture of the melt, why this should increase the silver ion activity is not evident.
Thermochemistry of Interconversion of H2B203 (g) and H,B,03(g) '
by Lawrence Barton, Satish K. Wason, and Richard F. Porter Departntent of Chemistry, Conell University, Ithaca, New York (Received April 86, 1966)
The thermochemical stability of gaseous H2B2Oa has been investigated. Alternate routes in the preparation of H2B203were employed to establish upper and lower limits in AHo for '/3B203(S). The methods the reaction HsB303(g) 1/z02(g) = H2B203(g) l/~B&(g) of chemical preparation involve the reaction of boroxine with oxygen and the reaction of diborane with oxygen under electrical discharge conditions. The molecules HzBz03and H3B303 have been shown to be precursors to each other. Mechanisms for the interconversion reactions have been suggested from results of isotopic substitution experiments in the H3B303-02and H2B203-B2Hsreactions using infrared and mass spectrometric techniques. From these observations, a mechanism for the formation of boroxine in the explosive reaction of pentaborane-9 with oxygen is proposed.
+
Introduction Reactions of gaseous boranes with oxygen are frequently explosive and lead to the formation of stable boron oxide. For this reason, it has been difficult to isolate partially oxidized compounds having both B-H and B-0 bonds. One interesting intermediate was observed in the reaction of BaHs with oxygen by Bauer and Wiberley.2 Later, Ditter and Shapiro3" identified the compound as H2B203 from its m a s spectrum. This compound is also observed as a product in the reaction of gaseous boroxine (HJ33Oa) and 0 2 . 3 b When mixtures of B2H6, H2B203, and 0 2 are heated, an explosive reaction occurs and boroxine is observed as a product.* The high temperature stability of boroxine was demonstrated from its preparation in a reaction of HzO(g) and elemental boron at temperature of about 1400°K.6 Taken collectively these observations suggest that Ha203 and HsBaOaare precursors to each other under different sets of experimental conditions, In the work described in this paper, certain aspects of tthethermochemistry and reaction mechanism The Journal of PhySical Chemistry
+
+
of the conversion of HzB20ato H&Oa and HJ3303 to HzB~O are ~ examined. Two procedures for the prepa ration of HzB20rB2Ha mixtures used in these studies are described below. Experimental It has been reported that H2B20ais unpredictably explosive in the condensed state. Thus, although very low pressures were used in this work, caution was observed at all times. The first method for the preparation of HzB203was (1) Work supported by the U. S. Army Research O5ce (Durham) and the Advanced Resesrch Projects Agency; preaented at the 149th National Meeting of the American Chemical Society, Detroit, Mich., April 1965. (2) W. H. Bauer and 8. E. Wiberley, Abstracts of Papers, 133rd National Meeting of the American Chemical Society, San Francisco, Calif., April 1958, p. 13L. (3) (a) J. F.Ditter and I. Shapiro, J. Am. Chem. SOC.,81,1022 (1959) ; (b) 9. I(.Wason and R. F. Porter, J. Phys. C h m . , 68, 1443 (1964). (4) G. H. Lee, W. H. Bauer, and 8. E. Wiberley, ibid., 67, 1742 (1963). (5) W. P. Sholette and R. F.Porter, ibid., 67, 177 (1963).
