Present Status of Cerium(IV)-Cerium(III) Potentials EARL WADSWORTH, FREDERICK R. DUKE, and
C. A. GOETZ
lnstitufe for Atomic Research and Depurfmenf of Chemisfry, Iowa Stute College, Ames, Iowa
b The variation of potential of the cerium(1V)-cerium(ll1) couple with the type of anion present is reviewed. The spectrophotometric data on sulfate complex formation and on hydrolysis correlate nicely with the potentiometric data, yielding a value of 1.74 volts for €0 at an ionic strength of 2. More work needs to b e done to elucidate the chemistry of cerium(lV) in nitrate and chloride solutions.
-
T
of the cerium(1V)cerium(II1) couple is very dependent upon the nature of the anion present. This variation in potential and the cerium(1V) and cerium(II1) complex formation constants determined independently have not been brought together to show how thoroughly the chemistry of the two cerium oxidation states is understood. The present work is an attempt to correlate the pertinent thermodynamic data on this system and to point out what work needs to be done to understand the chemistry of cerium(1V) and cerium(111) in the presence of the common anions. HE POTENTIAL
REVERSIBILITY
OF
CERIUM(IV)-CERIUM(III)-
PLATINUM ELECTRODE
Once it is established that the electrode responds to the aqueous system present, the condition for reversibility is that the current drawn in measuring the potential be negligible compared with the exchange current a t the electrode. A test for this condition is to measure the potential as a function of current; if the potential is invariant a t very small currents, reversibility is indicated. Vetter (20) has studied the cerium(1V)-cerium(II1) system in sulfate and nitrate solutions from this point of view; these solutions appear to act reversibly a t platinum electrodes. Another useful method for determining the exchange current is the use of radioactive tracers to follow the electron exchange on the electrode. Duke and Parchen ( 6 ) ,using tracers, failed to find any exchange catalysis by platinum in 6M perchloric acid solutions of cerium(1V)-cerium(II1). Fronaeus and Ostman (8) have shown, however, that platinum does catalyze the exchange a t lower acidities; they found the surprising result that a hydrolyzed dimer of cerium(1V) is the species involved in this catalytic exchange. Thus, it is 1824
ANALYTICAL CHEMISTRY
possible that, a t high acidities, the couple is not reversible a t platinum surfaces; but a t lower acidities it is almost certain that reversibility is attained. In chloride solution, it is highly improbable that the electrode is reversible to the cerium(1V)-cerium(II1) couple. The potential (18), -1.28 volts, obtained in 1M hydrochloric acid solution would be a consequence of a great stability of chloride complexes of cerium(IV) if the reversible cerium(1V)cerium(II1) potential were being measured. Duke and Borchers (4) have shown in a kinetic study of the oxidation of chloride by cerium(1V) that chloride complexes of cerium(1V) cannot possibly be this stable. The likelihood is that the ClrC1- couple preferentially reacts a t the platinum electrode and equilibrium in the oxidation of chloride by cerium(1V) is not attained. POTENTIAL MEASUREMENTS
The potentials using the platinum electrode have been measured by a number of investigators (1, 12). Kunz (13) found the standard potential in sulfate solutions to be -1.444 volts; Noyes and Garner (16) determined the potential in nitric acid solution, - 1.61 volts, very nearly independent of the acid concentration between 0.5 and 2M nitric acid; Smith and Getz (18) repeated part of this work and extended the measurements to perchloric and hydrochloric acid solutions, finding the values listed in Table I. They qualitatively attributed the wide spread in potentials to complex ion formation. Sherrill, King, and Spooner ( 1 7 ) investigated perchlorate solutions and found that the potentials were only slightly dependent on the cerium(1V)cerium(II1) ratio; they confirmed the results of Smith and Getz on the effect of acid concentration, and showed, by substitution of sodium perchlorate for perchloric acid, that the potentials are independent of perchlorate concentration. Thus, these authors attributed the variation in potential with acid concentration to the presence of hydrolysis products of cerium(1V). COMPLEX ION FORMATION
The formation of cerium(1V) complexes has been suggested by several authors (14, 15). The best data available appear to be those of Hardm%4i and Robertson (10, 1 1 ) ; their work was
done a t an ionic strength of 2. Thus, the potentials of cerium(1V)-cerium(111) in various acid media are best compared a t this ionic strength. Hardwick and Robertson (11) determined the following constants spectrophotometrically:
If the approximations are made that the ionic strength, p, equals 2 and [HS04-] = [H+] in Z M sulfuric acid, then
The complex ions of cerium(II1) and sulfate have been studied by Newton and Arcand (15), Fronaeus (7), and others (3, 19). The work of Newton and Arcand includes a value for the first dissociation constant a t p = 2 and 25" C.:
and a value for
Then, with the approximations that [H+] = [HS04-] in 2.11 sulfuric acid and p = 2, [Ce+++] = 0.5[Ce(III)t,bJ The higher complexes of cerium(II1) and sulfate are neglected, in view of the fact that a t [SO,--]= 0.02 (as in 2M sulfuric acid), Newton and Arcand found it unnecessary to include the higher complexes to explain their data. At equal concentrations of cerium(II1) and cerium(1V) and 25" C., p = 2,
Taking E = - 1.43 in 231 sulfuric acid, Eo = - 1.74 volts. In perchloric acid, Hardwick and Robertson ( I O ) found hydrolysis of ctriuni(1S') and dimerization of the hydrolysis product. They found K = [ Ce(0H) ++I [H-1 = 5 . 2 +
[Ce +dl
and
Table I.
both a t 25" C. and p = 2. Since [Ce(IV)totall= [Ce(OH)+31 2[CeOCe+6],in 0.025M[Ce(IV)t,t,~], Ce+4 = 0.0054, and
+
+
0.0054 -1.71 = Eo - 0.59 log 0.025 Eo = -1.75 volts
This agrees well with the value calculated from the sulfate data. Hardwick and Robertson showed that changes in ionic strength have little or no effect on the hydrolysis constants in perchloric acid. Under these conditions it is possible to calculate Eo as a function of ionic strength (Table 11). Duke and Bremer (6) have correlated these potentials empirically with the water activity in perchlorate solutions. Data on complexing of nitrate ions by cerium(TV) are not available. Noyes and Garner (16) showed that the potential in the range from 0.5 to 2M nitric acid is very nearly independent of acidity and ionic strength. Their value agreed with the results of Smith and Getz (18). To explain the acid independence and a t the same time reconcile these results with the Eo's calculated above, one can write: Ce+4
+
+
H 2 0 7 2 [Ce(N03) H+ (OH)] [Ce(KOs) (OH)++] [H+] K = [CeA4][NO3-] SO3-
++
+
This constant must have a value of about IO2, to agree with the results in perchlorate and sulfate media. The results of Blaustein and Gq-der (2) indicate strongly that the situation in nitrate solutions is much more complicated than this, n-ith dimerization of this species and other reactions occur-
Acid Normal-
Cerium(ll1)-Cerium(lV) HalfCell Potentials
Measured E HC1Od
ity 1
-1.70 -1.71 -1.75 -1.82 -1.87
2 4 6 8
HC1
"03
-1.61 -1.62 -1.61 -1.56
...
-1.44 -1.44 -1.43
-1.28
..,
-1.42
anions mould clarify the complexing and hydrolysis picture; and when used in conjunction with the work of Blaustein and Gryder ( 2 ) )such spectrophotometric data should allow a detailed description of the status of cerium(1V) in nitrate solution. It is probable that a t low temperature cerium(1V) in the presence of chloride is sufficiently stable to lend itself to a similar spectrophotometric study.
Table 11. Standard Potential as a Function of Ionic Strength P 1 2 4 6
8
LITERATURE CITED
[H+I
Eo
Bauer, E., Glaessner, A,, Z . Elektro-
1 2 4 6 8
1.75 1.76 1.7R 1.82 1.87
Blaustein, B. D., Gryder, J. W., J . Am. Chem. SOC.79, 540 (1957). Connick, R. E., Mayer, S. W., Zbid.,
chem. 9, 534 (1903).
73, 1176 (1951).
Duke, F. R., Borchers, C. E., Zbid., 75, 5186 (1953).
ring (9). There is little doubt that the lower potential a t higher nitric acid concentrations is a consequence of substitution of nitrate for hydroxyl in the complex, but no quantitative data apart from the potentials are available. If the thermodynamic (zero ionic strength) potential is to be obtained, potential as well as complex and hydrolysis constants must be measured a t much lower ionic strengths. There are indications that such work might be very difficult, as perchlorate solutions of cerium(1V) hydrolyze and polymerize to the extent of being insoluble in low concentrations of acid; and in sulfate solutions a t low ionic strengths, both hydrolysis and polymerization would probably occur. However, the utility of the true potential is very limited, as cerium(1V) is not used practically at very low ionic strengths. Further studies of cerium(1V) in the presence of nitrate and chloride ions need to be done; particularly, work similar to that of Hardwic$ and Robertson (10, 11) in the presence of these
Duke, F. R., Bremer, R. F., ANAL. CHEW23, 1516 (1951). Duke. F. R.. Parchen. F. R.. J . Am. C h e k Soc: 78,1940 11956): (7) Fronaeus, S., Svensk Kem. Tidskr. 64, 317 (1952).
(8) Fronaeus, S., Ostman, O., Acta Chem. Scand. 10, 769 (1956). (9) Gryder, J. W., Dodson, R. W., J. Am. Chem. SOC.71, 1894 (1949);
73. 289n ~ ---- ( I R . I (IO) Haidwick, T. J., Robertson, E., Can. J . Chem. 29,818 (1951). (11) Ibid., p. 828. (12) Jones, E. G., Soper, F. G., J. Am. Chem. SOC.57,802 (1935). (13) Kuna, A. H., Zbid., 53, 98 (1931). (14) Moore, R. L., Anderson, R. L., Zbid., 67, 167 (1945). (15) Newton, T. W., Arcand, G. M., Zbid., 75, 2449 (1953). (16) Noyes, A. A., Garner, C. S., Zbid., I
\ - - - - I .
58. 1264 (19361. - - - ~ \----,(17) ShlFrill, M. S., King, C. G., Spooner, R. C., Ibid., 65,170 (1943). I
(18) Smith, G. F., Getz, C. A,, IND. ENG.CHEM..AKAL. ED. 10. 191 119381. (19) Spedding, F. H., Jaffe, S., J . Am. Chem. SOC.76,882 (1954). (20) Vetter, K. J., 2. physilz. Chem. 196, 360 (1951).
RECEIVED for review December 31, 1956. Accepted July 24, 1957.
Anodic Stripping Voltammetry Using the Hanging Mercury Drop Electrode RICHARD D. DeMARS and IRVING SHAIN Chemisfry Department, University o f Wisconsin, Madison 6, Wis.
b Anodic stripping voltammetry using the hanging mercury drop electrode offers a rapid and convenient method for analyzing extremely dilute solutions of metals which form amalgams. The anodic stripping is performed using the techniques of voltammetry with continuously varying potential, and the anodic peak current is a function of the concentration of the ion in the
solution and the cathodic plating time. As examples the method was applied to cadmium and thallous ions in the 10-6 to 10-9M concentration range and to mixtures of cadmium and zinc in the to 10-8M concentration range.
T
of certain metal ions into a mercury cathode a t conHE PLATING
trolled cathode potential has been used many times to concentrate and separate these metals from rather dilute solutions ( 5 ) . Subsequent anodic stripping of these amalgams may be used to concentrate the sample for some conventional method of analysis (10). The anodic stripping process also has been the basis of direct coulometric determinations ( I , 3, 8). Nkelly and VOL. 29, NO. 12, DECEMBER 1957
1825
