1314
Y. MARCUS
pronounced formation of mixed complexes will affect the self-association equilibria in such a way that the average degree of self-association is lowered. On this basis one can understand why addition of either alcohol to the other results in a lowering of the viscosity. When TFE is added to EtOH, the viscosity decreases only slightly because the greater molecular weight of TFE largely compensates for the decrease in average association number, and the mass of the average kinetic unit appears to remain almost constant. When EtOH is added to TFE, these effects work in the same direction and the viscosity decreases markedly.
Vol. 62
The very complex behavior of the partial molal volumes (Fig. 1) must remain uninterpreted at this time. It is worth noting, however, that the partial molal volume of very dilute TFE in EtOH is markedly smaller than the molar volume of the liquid TFE, while the partial molal volume of very dilute EtOH in TFE is larger than the molar volume of liquid EtOH. It is also worth noting that the minimum in the viscosity occurs a t about the same mole-fraction of TFE, 0.24, as the extrema in the partial molal volumes. On the other hand, the maximum boiling point occurs a t 0.386 mole fraction of TFE.
THE OXIDATION-REDUCTION COUPLES U(1V)-U(V1) AND Fe(I1)-Fe(II1) I N PHOSPHORIC ACID SOLUTIONS BY Y. MARC US^ Contributionfrom the Israel Abmic Energy Commission Laboratories, Rechouot, Israel Received February 68,1868
The otentials of the U(1V)-U(V1) and the Fe(I1)-Fe(II1) couples in 1.2 to 7.5 M phosphoric acid were measured a t 25' anfformal potentials E' were calculated. Plots of E'u and E ' F vs. ~ log U H ~ P O ~the , phosphoric acid activity, were linear, with slopes of (+3.2 =k 0.2) X 0.60/2 mv. and (-1.3 f 0.1) X 60 mv., res ectively. For 1 M phosphoric acid the values are E'u = -475 & 10 mv. and E ' F ~= -555 f 10 mv. us. the normal h Xrogen electrode. Solutions of U(IV), in sharp contrast to Fe(II), in phosphoric acid are very stable to air oxidation. $he behavior of iron and uranium ions, singly or mixed, in phosphoric acid solutions was found to be as might be predicted from the single formal potentials. At hosphoric acid concentrations above 2.4 M Fe(I1) reduces U(VI), and at concentrations below 1.2 M Fe(II1) oxidizes U ( 1 3 .
Recently, Baes2 has published -data on the oxidation-reduction potential of the Fe(I1)-Fe(111) couple in phosphoric acid, and on the oxidation-reduction equilibrium between this and the U(1V)-U(V1) couple. His solutions, however, contained 0.36 M sulfuric acid in addition to phosphoric acid. Previously, Bock and Herrmann3 had measured the potential of the iron couple in phosphoric acid both with and without sulfuric acid, and obtained results which were not in agreement with those of Baes, nor with the earlier results of Carter and Clews.4 Clarification of the issue is desirable; in particular it would be desirable to obtain values for the uranium couple, measured in the absence of sulfuric acid.5 Measurements were made on the uranium and iron couples separately and in mixed solutions, the only anion present being that of phosphoric acid. A few equilibrium measurements in 6 n/r phosphoric acid were also made, for comparison with values calculated from the potentials. Consider the oxidation-reduction process of the iron couple in phosphoric acid where the Fe(I1) and the Fe(II1) species are complexed by phosphoric acid. The position of the equilibrium will depend on the activity of phosphoric acid and hydrogen ions. (1) Taken from the thesis submitted by the author to the Hebrew University, Jerusalem, 1955,for the degree of Ph.D. (2) C. F. Baes, Jr., J . Phys. Chem., 60,805 (1956). (3) R. Bock and M. Herrmann, 2. anorg. allgem. Chem.. 273, 1 (1953). (4) S. M.Carter and F. H. Clews, J . Chem. Soc., 126, 1880 (1924). (5) The work described here waa carried o u t in Israel in 19541955.before the results of Brtes, U. S. Atomic Energy Comm. AECD3594 (19531,cf., Nucl. Sei. Abstr., 8, 190 (1954),became known.
Fe(I1)
+ zH3POrI _ Fe(II1) + e- + gH+
(1)
For a constant phosphoric acid concentration, provided the iron concentration is small compared to the phosphoric acid concentration also, the hydrogen ion activity will be constant, and the potential (in mv. at 25") will be given by
where [ ] are stoichiometric concentrations and y are stoichiometric activity coefficients. Evaluating the expression E - 60 log [Fe(III)I/ [Fe(II) ] from measurements a t constant phosphoric acid Concentration will yield values of E', the formal potential of the couple, defined as its potential vs. the standard hydrogen electrode when the ratio of the stoichiometric concentrations of the oxidized and reduced forms of the substance equals unity E' = E
- 60 log [Fe(III)]/[Fe(II)]
(3)
For the analogous uranium oxidation-reduction reaction, the formal potential at constant phosphoric acid concentrations will be given by E'
=
E
- (60/2) log [U(VI)l/[U(IV)I
(4)
The same situation exists for the oxidationreduction reaction between the uranium and the iron couples U(V1) 2Fe(II) + xH3P04I_
+
U(1V)
+ 2Fe(III) + zH+
(5)
for which an apparent concentration quotient may be calculated for every phosphoric acid concentration from the formal potentials
Oct., 1958 log K = (2/60)(E’u
POTENTIALS OF THE
URL4NIUM(IV)-(V1)COUPLE
- E ‘ F ~=)
If constant values are obtained for the formal potentials E’ at a given phosphoric acid concentration, that would mean that the approximation of a constant ratio of activity coefficients of the metallic species holds. Experimental Materials .-Uranium( VI) stock solution was prepared by precipitating uranyl ammonium phosphate from uranyl sulfate solution by ammonium dihydrogen phosphate, washing the precipitate until no sulfate could be detected in the washings, and then dissolving- it in the minimum amount of phosph&ic acid. Uranium(1V) stock solution was prepared by boiling the UfVI’, solution with zinc amalgam. The UfVI) remaining after ’this treatment was neglizble. It wa9‘ observed thax passage through a Jones reductor did not result in quantitative reduction, nor could reduction be achieved by electrolysis using either platinum or mercury cathodes, because in solutions of sufficient acidity for non-formation of a U(IV) phosphate gel, hydrogen was evolved. Iron(111) stock solution was prepared by precipitating hydrous iron( 111) oxide from iron( 111) ammonium sulfate solution by ammonium hydroxide and dissolving the washed precipitate in concentrated phosphoric acid. On standing, a pink solid crystallized (Anal. Calcd. for Fez03. 2P&5HzO (dried at 110”): Fez03, 29.9; P z O ~ ,53.2. Found: FezOl, 30.2; P2OS, 54.0). Jameson and Salmon6 described a similar compound, but richer in water, viz., Fez03,2P206, 8H20, as crystallizing from similar solutions. The supernatant liquid was clear and pink and was approximately 0.5 M in iron(II1). Iron( 11) stock solution was prepared by dissolving iron powder in warm concentrated phosphoric acid in an atmosphere of carbon dioxide. The color of this solution was brown and persisted even with dilution to 3 M acid and 0.05 M iron( 11). All materials were of analytical reagent grade. The uranium solutions contained as “impurities” ammonium and zinc ions a t concentrations similar to that of uranium. The only anion present in the solution is the phosphate anion. Method.-The potentials E (in mv.) were determined in a constant temperature bath regulated a t 25.0 f 0.3’ in a cell Pt/Fe(II), Fe(II1) and/or U(IV), U(V1); C M H3POd/sat. KCl/S.C.E. -
+
the total concentration of iron or uranium being less than C/100. The cells were 20-ml. beakers closed with stoppers supporting the electrodes. In cases where iron(I1) was present, an atmosphere of carbon dioxide was maintained above the solution, after removal of dissolved oxygen by bubbling carbon dioxide through the cell, which was then sealed with collodion. The S.C.E. (saturated calomel electrode) was a Beckman fiber type electrode. Potentials of the half cell://sat. KCl/ S.C.E. have been given as -242,T -244* and -245O mv. to the nearest mv. The last named figure was arbitrarily selected, and all measurements are referred to the normal hydrogen electrode, N.H.E., obtained by subtracting 245 mv. For comparison, published potentials also are referred to the N.H.E., and adjusted for the difference in the assumed potential of the S.C.E. where necessary. The platinum electrode was bright metal of area 5 sq. mm. or more. Because of the high resistance of the fiber type S.C.E., and the low sensitivity of the available galvanometer, the Leeds and Northrup Model 7552 potentiometer at first used for measurement did not give good results. It was re(G) R. F. Jameson and J. E. Salmon, J . Chem. Soc., 28 (1954). (7) W. M. Latimer, “Oxidation Potentials,” 2nd Ed., PrenticeHall, New York, N. Y.. 1952,p. 177. (8) R. G. Bates, Chem. Reus.. 42, 1 (1948). (9) D. A. McInnes, D. Belcher and T. Shedlovsky, J . A m . Chem. Soc.. 60. 1094 (1938).
I N PHOSPHORIC
ACID
1315
placed by a Beckman Model H-2 pH-meter, which gave better performance, but limited the accuracy obtainable. Equilibrium was reached rather slowly: for the iron couple about one day was necessary before a steady potential could be observed, and for the uranium couple as many as four days. The slowness in reaching equilibrium has been observed previously in other solutions of uranium by Ahrland and Larsson,lo who recommend sodium indigo tetrasulfonate as a catalyst for charge transfer for the uranium couple. Unfortunately, neither this, nor any similar reagent was available. The slowness of other reactions of uranium in phosphoric acid also was observed.11 Curves of potential us. time were drawn, and the equilibrium potential could be estimated to within f 3 mv. Analytical.-The metal concentrations of iron and uranium solutions were obtained from volumetric dilution values. Stock solutions of iron(II), uranium(1V) and (VI), the latter after boiling with zinc amalgam, were titrated with potassium permanganate. Uranium( VI) concentrations also were checked by the thiocyanate-acetone method1lt12 and were found t o agree well with volumetric values. The concentration of iron( 111) was determined likewise by a colorimetric thiocyanate-acetone method. Phosphoric acid concentrations were determined by titration between p H 4.50 and 8.90. I n solutions containing both uranium and iron, the concentrations of uranium( IV) and (VI) were obtained spectrophotometrically by the method of Andrews, Schaap and Gates,ls following which the test solution was diluted with concentrated phosphoric acid in the volume ratio 1 :9, the extinction of the solution then obeying Beer’s law. The extinction a t 630 mp is a measure of the U(1V) content and a t 410 mp, corrected for the presence of U(IV), a measure of the U(V1) content. Total reducing substances, U(1V) Fe(11), were determined by potassium permanganate titration, and the iron(I1) and (111) were then obtained by difference, the total iron content being known.
+
Results Air Oxidation.-Uranium(1V) solutions in phosphoric acid are very stable to air oxidation. Identical solutions, one of them exposed to the atmosphere for one week, the other kept under carbon dioxide, showed no significant difference between measured potentials. This may be contrasted with the behavior of iron(I1) in phosphoric acid: solutions exposed to the air became oxidized at an increasing rate, although to a small extent only on the first day. The Uranium Couple.-The potential of the U(1V)-U(V1) couple was measured at four ratios of the concentrations of oxidized to reduced uranium, and five phosphoric acid concentrations C. The formal potential E’, in mv., was calculated from eq. 4,and the data are given in Table I. Fairly constant E’u values were obtained for every C value indicating that equilibrium has been reached. Values for the activity of phosphoric acid U H ~ P O were , calculated from the data of Mason and Blum14 and the values of E‘u were plotted os. log a H s p o , (Fig. 1). The points lie approximately on a straight line. A least squares treatment shows that E’u = 475 f 10 mv. at log UH3P04 = 0.00 (where approximately C = l.OO), and that the slope is ($3.2 f 0.2) X 60/2 mv. The Iron Couple.-The potential of the Fe(I1)(10) S. Ahrland and R. La’rsson, Acto Chem. Scond., 8 , 137 (1954). (11) Y.Marcus, Ph.D. thesis in chemistry, Jerusalem, 1955. (12) C. E. Crouthamel and C. E. Johnson, A n d . Chem., 24, 1780 (1952). (13) L. J. Andrew, W. B. Schaap and J. N. Gates, Jr., U. S. Atomic Energy Comm. CD-4014 (1945); cf. C. J. Rodden, Ed., “National Nuclear Energy Series,” McGraw-Hill Book Co., New York, N. Y., 1950, Vol. VIII-I, p. 78. (14) C. M. Mason and W M. Blum, J . Am. Chem. Soc., 69, 1246 (1947).
Y. MARCUS
1316 I
720
I
t
3601
I
0
C , total HSPOl concn., M 1.2
~HBPOI
3.0
0.71
4.5
1.10
6.0
1.50
7.5
1.90
log
0.11
-E -E‘ -E --E’
-E -E‘ -E --E’
-E --E’
2 Log 0ypo4
log [U(VI) I/IU(IV) -1.70 -0.20 +O.GO +1.90 Potential, mv. 440 485 515 555 499 497 491 491 480 515 590 560 532 542 531 521 525 575 605 535 578 587 576 581 580 620 650 690 632 632 631 626 610 650 720 670 661 656 662 652
1 o , oo Av. E’ 495
*5
530 & 5 580 f 5 630 = t 5 660
5
Uranium-Iron Oxidation-Reduction Equilibria. -Mixtures of phosphoric acid solutions of uranium and iron were prepared, containing both uranium(IV) and uranium(V1) to ensure that both reduced and oxidized species were present. The nominal compositions are given in Table 111. I n preliminary experiments it was found that solutions containing U(V1) and U(1V) in a ratio of ten were bright yellow, that when the ratio was around
Vol. 62
unity the solutions were yellowish green, and when it was one tenth the solutions were bright green. Gel formation was observed only in solutions in which the uranium was mostly in the U(IV) state. Solutions in 6 M phosphoric acid were subjected to direct analysis for uranium and iron. The results of these observations on the mixed solutions and the measured potentials of cells containing them are reported in Table 111. The measured poten$ials were applied in eq. 6 and using E’ values taken from Fig. 1the values of [U(VI) J/[U(IV)] and [Fe(III)]/[Fe(II)] given in Table I11 were calculated. Although the “observed” values are for the main part only orders of magnitude, they show agreement with the “calculated” values, so that the values of E’ should be essentially correct. General conclusions from the data are that (a) Fe(I1) reduces U(VI), while Fe(II1) is unable to oxidize U(V1) a t C 2 2.4 M ,and (b) Fe(II1) oxidizes U(IV), while Fe(I1) is unable to reduce U(V1) a t C 5 1.2 M . This is borne out in Fig. 1, where it is seen that the E‘ curves cross around C = 2.2 f 0.5 M . Discussion The Oxidation-Reduction Behavior of Uranium. -Predictions from the measured potentials may be compared with the behavior of uranium in phosphoric acid solutions as far as it is known. In sulfuric aLid solutions, as is well knownl15iron(II1) oxidizes uranium(1V) rapidly. Schreyer and BaesI6 found that in 1.6 M sulfuric acid and 1.0 M phoaphoric acid iron(II1) (18 d) still oxidizes uranium(1V) (2 mM) quantitatively although slowly. The potentials (Fig. 1, disregarding the sulfuric acid) show an 80 mv. difference for the free energy change in the reaction. In hydrochloric acid, tin(I1) chloride reduces uranium(1V) only on prolonged heating”; even in 6 M hydrochloric acid at 97”, the reaction is not complete (E’u about -330 rnv.).l8 Making the solution 0.6 M in phosphoric acid brings the reaction to completion (E’u = -450 mv.), in the presence of iron(III), which is also reduced. After dilution to about 4 M hydrochloric acid and 0.4 M phosphoric acid, mercury(I1) chloride is able to oxidize only excess tin, but neither iron nor uranium (E’u = -430mv., E ’ F ~= -590mv.). After dilution to 3 M hydrochloric and 0.3 M phosphoric acids, iron ( E ’ F ~= -600 mv.) is able to oxidize the uranium (E‘u = -415 mv.) rapidly and completely. Iron(II1) (EO = -770 mv.) is reduced by hydroxylamine in acid solution (EO = 50 mv. for oxidation to dinitrogen monoxide) but not U(V1) , according to Strubl,lg while Canning and Dixon found that it oxidizes U(IV) in dilute acidz0(EO = -330 mv.). These reactions still occur in the presence of 0.8 M phosphoric acidz0 (E’u = -465 (15) I. M. Kolthoff and J. J. Lingane, J. A m . Chem. SOC.,66, 1781 (1933). (16) J. M. Schreyer and C. F. Baes, Jr., Anal. Chem., 26, 644 (1953). (17) E. F. Kern, J . A m . Chem. Soc., 2 3 , 605 (1901). (18) A. R. Main, Anal. Chem., 26, 1507 (1954). (19) R. Strubl, Coll. Czech. Chem. Comm., 10, 466 (1938). (20) R. G. Canning and P. Dixon, Anal. Chem., 2 T , 877 (1055).
POTENTIALS OF THE URANIUM(IV)-(VI) COUPLEIN PHOSPHORIC ACID
(ht., 1958
1317
TABLE I1 FORMAL POTENTIAL OF THE Fe(I1)-Fe(II1) COUPLE IN PHOSPHORIC ACID C,total
1fsPo4 concn., M
UHSPOI
1.2
0.11
log
3.0
0.71
4.5
1.10
-E -E’ -E -E‘ -E -E’ -E -E’
1.50
6.0
P O T E N T I A L S AND E Q U I L I B R I U M
C,
HsP04 (M)
6.0
2.4
log aH8P04
1.50
0.53
U(1V)
- 1.50
-1.00
465 555 415 505 380 470 350 440
495 555 440 500 415 475 375 435
0.4 10.0 10.4 0.4 10.0 0.16 4.0 4.16 0.16 4.0 0.08 2.00 2.08 0.08 2.00 0.016 .40 .416
545 551 510 516 465 47 1 435 441
TABLE TI1 RATIOSOF U R A N I U M AND
Nominal composition, mM U(V1) Fe(I1) Fe(TI1)
10.0 0.8 10.8 10.0 0.8 4.0 0.32 4.32 4.0 0.32 2.00 0.16 2.16 2.00 0.16 .40 .032 .432
log IFe(II1) I/[Fe(II) 1 -0.10 +1.00 Potential, mv.
40 40 40
16 16 16
40 40 40
IPON
590 530 570 510 525 465 495 435
MIXTURES IN
625 535 585 495
505 i 10
560 470
470 i5
530 440
440 i5
545 i 10
PHOSPHORIC
ACID
Obsd: potential, -mv.
375 440 450 559 650 a
470b 16 16 16
0.00 Av. E‘
$1.50
(I
< -3”J < -3”f < -3°F
-8.2
-6.0 -5.7 -2.0”’ -2.3 1.08‘’d 1.0
-1.15“ 0.11’ 0.30‘ , 1.60” >2.30”
-1.1 0.0 0.2 1.8 3.5
0
< - 1*J
-1.8
-0.6‘
-0.8
LI
0.3O 1.1 >Id 1.8 580 630 >Id 3.5 >IP 1.9 Id 9.0h .032 .40 1.6 675 a Gelation of the solution through formation of U(1V) phosphate gel. Slight gelation observed. Values obtained from analysis. d Observed color, bright yellow. e Observed color, yellowish green. f Observed color, bright green. 0 CalCalculated asculated from nominal content, the equilibrium oxidation state (color) and the stoichiometric reaction 5. suming linearly extrapolated values E‘u = -405 mv. and E ’ F ~= -610 mv.
mv., E ’ F ~= -560 mv.), but in 4.8 M phosphoric acid, in the presence of copper as catalyst, hydroxylamine reduces U(V1) (E’u = -590 mv.). Hydrogen peroxide (EO = -1770 mv.) was found to oxidize uranium(1V) a t this acid concentration.20 Although standard potentials should not be used for concentrated solutions7 as detailed above without due caution, the observed behavior is nevertheless seen to be predictable from them using the experimental formal potentials. The Oxidation-Reduction Potentials.-Previously reported formal potentials of the uranium couple in phosphoric acid are those of B a e ~ ,made ~.~ in the presence of 0.36 M sulfuric acid. (Woody and co-workers21 reported a value of -0.5 v. found for about 4.7 M phosphoric ‘acid by the titration of 10 m M uranium(1V) with potassium permanganate, but this is not an equilibrium value.) The interference of sulfuric may be twofold: it may complex with any of the uranium and iron species
and the hydrogen ions it furnishes may suppress the ionization of the phosphoric acid or of the complexes, thereby affecting the quantitative relat,ionships among the species. These effects should diminish with increasing phosphoric acid concentrations. Figure 2 gives iron couple potentials taken from the literature. The data of Carter and Clews4 were determined a t 17” with an “N.C.E.” (quoting the authors) as reference electrode, for which the authors allowed -560 mv. to convert the data to the N.H.E. basis. It is, however, not easy t o understand why this value was selected. If their data are recalculated, using the value for the 0.1 M calomel electrode, -338 mv., for conversion, then the values would be in line with those of Baes2for 4 and 7 M phosphoric acid, and also with the present measurements for 4.5 and 6 M phosphoric acid (see Fig. 2). A single value g.iven by Lingane22 at 1.0 M phosphoric acid also agrees with the present values. At phosphoric acid con-
(21) J. H. Pannell, R. J. Woody and J. M. Grandfield, U. 8. Atomio Energy Comm. MITG-248 (1951).
(22) J. J. Lingane, “Electroanalytical Chemistry,” Interscience Publ., New York,N. Y., 1953, p. 124.
Y. MARCUS
1318
Vol. 62
E' rn V.
600
,500
8 A 400
300 I
I
-1.0
0.0
I
0.0
LDQCH,p04
0.5
1.0
at$ po4
1.5 I
1.0
Fig. Z.-Formal potential E' of the Fe(I1)-Fe(II1) couple as function of phosphoric acid concentration C published this work, 25'; a,B a e ~ , ~ 0 . 3M6 k,S04, 25"; values: 0, 0, Carter and Clews,4 17'; X, Lingane,2225"; +, Bonner and RomeyqZ30.15 M Na2HPOd solution. All other symbols refer to work of Bock and Herrmanng a t 20°, viz.: A, no HzSO4; - - - -, about 0.05 M H2S04; _ _ , about 0.25 M Hz'SO4, ---------, about 2 M HzSOd.
centrations below 3 M the values of Baes are somewhat lower than the present values, which should be due to the presence of sulfuric acid. Baed values for 0.36 M sulfuric acid are not very far from those of Bock and Herrmann3for 0.34 M sulfuric acid. Bock and Herrmann state that they could not observe any dependence of the formal potential on phosphoric acid concentration in solutions without sulfuric acid, and give the potential a t 0.33 M as -434 mv. On the other hand, in the presence of only 0.09 M total sulfate, their potential curve does show a decrease from about -590 mv. at 0.033 M phosphoric acid to -495 mv. a t 0.33 M and to about -370 mv. a t 0.83 M. Further, the difference in temperature does not explain the differences between the data of Baes a t 25' and of Bock at 20". A single determination in 0.5 mM sulfuric acid and 0.15 M disodium hydrogen phosphate made by Bonner and RomeynZ3is not directly comparable. (23) W. D. Bonner and H. Romeyn, Ind. Eng. Chem.. Anal. Ed., 8, 87 (1931).
The Oxidation-Reduction Equilibrium Constant. -There are two sets of values for the parameter K , eq. 6, for reaction 5 , to compare with values calculated from the measured E' values. These are shown in Fig. 3, and it is apparent that there is no quantitative agreement, probably because of the different conditions in the experiments (no sulfuric acid was employed in the present work, while Baes2 used 0.36 M sulfuric acid and Canning and Dixon19about 0.1 M sulfuric acid in their solutions. I n the case of the latter, it seems that the equilibrations were performed in hot solutions, and the values given are calculated from extinction data and the nominal content of iron). However, it is evident that a t some concentration C between two and three the true value of log K is zero, i.e., the potential curves E' u and E'pe intersect, as was found above. Acknowledgment.-Thanks are due to the Director, Israel Atomic Energy Commission Laboratories, for his help and encouragement.
