nec. ,?, 1955 CHLORIDE COMPLEXING AND DISPROPORTIONATION O F PU(IV) [CONTRIBUTION FROM
THE
IN
HYDROCHLORIC ACID 6145
UNIVERSITY O F CALIFORNIA, LOS ALAMOS SCIENTIFIC LABORATORY]
Chloride Complexing and Disproportionation of Pu(1V) in Hydrochloric Acid' BY S. W. RABIDEAU AND H. D. COWAN RECEIVED MAY18, 1955 From measurements of the formal potentials of the Pu+"Pu+' couple in hydrochloricperchloric acid solutions a t constant total acidity and ionic strength it has been shown that the complex ion P u C ~ is + formed. ~ The dissociation quotient of this complex a t 25' has been found t o be 1.77 =k 0.10 a t unit ionic strength and 1.70 0.03 a t an ionic strength of two. P u + ~ C1- are A F = -0.34 At 25" in molar hydrochloric acid the thermodynamic quantities for the reaction PuClia kcal./mole, AH = - 1.9 kcal./mole and A S = - 5 e.u. Equilibrium and rate measurements were made in hydrochloric acid solutions at 6, 25, 35 and 45' in a study of the disproportionation reaction ~ P u + 2Hz0 ~ F? 2Pu+s PuOz+2 4H+. I n accord with this reaction, a fourth power hydrogen ion dependence of the equilibrium quotient was observed. The disproportionation reaction rate constant exhibited an inverse third power dependence upon the hydrogen ion concentration. \-slues of 39 kcal./mole and 53 e.u. were obtained for the heat of activation and for the entropy of activation, respectively, for the disproportionation reaction in molar hydrochloric acid. The change in the heat content a t 25" for the reaction Pu+a H + e Pu+4 l / Z H Ewas calculated t o be 14.32 kcal./mole in molar hydrochloric acid and 13.63 kcal./mole in molar perchloric acid from measurements of the formal potential of the Pu+"Puf4 couple as a function of temperature. The P u + L Pu02+2potential in molar hydrochloric acid was calculated from the measured Puf"Puf4 formal potential and from the disH + 3 / 2 H 1 in proportionation equilibrium q;otient a t each temperature. For the reaction Puf3 2Ht0 Pu0zf2 molar hydrochloric acid a t 25 , A F = 70.8 kcal./mole, AH = 78.4 kcal./mole and A S = 25.4 e.u. The ionic entropies of puO2+2and Pu+4 have been calculated to be -28.9 and -87 e.u., respectively.
+
+
+
+
+
+
+
Introduction From the previously reported2 values of the PU+S-PU+~ couple in molar perchloric and in molar hydrochloric acids i t is possible to calculate the extent to which P u + ~is complexed by chloride ion if the assumptions are made that P u + ~is not complexed to any significant degree and that the for~ is known. Measurements mula of the P u + complex have been made in this work which show that the formula of the complex is P u C ~ +up~ to an ionic strength of two. Connick and McVey3 calculated an apparent fifth power hydrogen ion dependence for the disproportionation equilibrium quotient and a - 3.5 power hydrogen ion dependence for the disproportionation rate constant in hydrochloric acid solutions of varying ionic strength and chloride ion concentration. I n the present study the expected fourth power hydrogen ion dependence of the equilibrium quotient was obtained from measurements a t constant chloride ion concentration and a t unit ionic strength. The disproportionation rate constant has been found to exhibit an inverse third power hydrogen dependence in hydrochloric acid solutions as has been previously shown to be the case in perchloric acid solutions. 3,4 The disproportionation of Pu(1V) has been studied as a function of temperature to obtain both rate and equilibrium data from which thermodynamic functions have been computed. Experimental Reagents.-The plutonium solutions were prepared by dissolving a known weight of oxide-free metal in the requisite weight of standardized Baker and Adamson reagent grade hydrochloric acid, followed by dilution to the desired acidity with water redistilled from alkaline permanganate. When necessary, the ionic strength was maintained a t unity with the addition of purified sodium chloride. This salt was precipitated from a saturated solution with concentrated hydrochloric acid, washed with water, dried and ignited in an electric furnace a t 600' overnight. The perchloric acid was Mallinckrodt 71% reagent grade. The potassium dichromate used in preparing a plutonium solution of known (1) This work was done under the auspices of the A.E.C. (2) S. W. Rabideau and J. F. Lemons, THISJOURNAL, 73, 2895 (1951). (3) R. E.Connick and W. H. McVey, i b i d . , 76, 474 (1953). (4) S. W.Rabideau, i b i d . , 76, 798 (1953).
+
+
Pu(IV)/Pu(III) ratio was Mallinckrodt "primary standard" quality. Apparatus.-The cell assembly, the potentiometric equipment, the folding-scale recording potentiometer and the thermostat have been described p r e v i o u ~ l y . ~The ~ ~ temperature of the plutonium solution in the cell was obtained to zk0.001' from measurements of the resistance of a type 20D Western Electric thermistor. Over a period of several The therhours the temperature variation was =tO.Ol'. mistor was calibrated against a Beckmann thermometer which had been calibrated against a platinum resistance thermometer certified by the National Bureau of Standards. The hydrogen used for the hydrogen electrode was electrolytic grade purified by passing it through a uranium furnace heated to 600'. An atmosphere of helium was maintained over the plutonium solutions throughout the experiments. Both the hydrogen and the helium were saturated with water vapor a t the cell temperature to minimize concentration changes because of evaporation. Analyses.-The concentrations of Pu + 3 , P u + and ~ PUOZ+~ were obtained as a function of time in the disproportionation studies from measurements of the potential of a cell , NaCl(c2); HCl(cl), of the type: P t ; P u + ~P, u + ~HCl(cl), c p = 1. As has been deNaCl(c2); Hz, P t , where c1 scribed p r e v i ~ u s l y ,from ~ measurements of the cell potential, the mean oxidation number and the total plutonium concentration i t is possible to calculate the concentrations of Pu+3, Pu+4 and PuOz+2. In general, the experiments were devised such that the PuOS+concentration was negligibly small. This was brought about bv keeDinP the Pu +4- concentration relatively -large in comparison- with P u + and ~ Pu0~+~.
+
Results Chloride Complexing of Pu(1V).-If a plutonium(IV) chloride complex of the formula PuClf3 is formed, the dissociation of the complex may be represented by the reaction PUC1+3
PU+4
+ c1-
and the equilibrium quotient is given by the expression K = [ P u + ~ [Cl-]/ ] [PuClf3]. The uncomplexed species, P u + ~is, related to the total Pu(1V) concentration, ZPu(IV), by the equation PU+'
=
kZPu(IV)/(K
+ [Cl-1)
It is necessary to make an assumption concerning , for this work the chloride complexing of P U + ~and is uncomplexed by chloride it is assumed that P u + ~ , ~ asion. (As has been discussed p r e v i ~ u s l y this sumption may not be strictly valid. Further studies ~ P U O * +are ~ of the chloride complexing of P U + and in progress.) Also, both Puf3 and P u + are ~ con-
sidered to be uncomplexed in molar perchloric acid. and the fast reaction The difference in the formal potentials of the P u + ~ k3 P u + couple ~ in molar perchloric acid and in the perPU(V) PU(IV) J_ PU(III) PU(VI) kc chloric-hydrochloric acid solutions ([H+] = 1) at 25' is related to the dissociation quotient of the In perchloric acid solutions it was necessary to concomplex by the equation sider in addition to these two steps the diminution of the mean oxidation number as a result of the aAEO' = 0.0591 log K / ( K + [Cl-1) particle reduction. The rate law for the disproThe formal potential of the PLI+~--Pu+~ couple was portionation reaction both in perchloric and in hydetermined over a tenfold change in the ZPu(1V) drochloric acid solutions as derived3from the above P u + ratio ~ for each of the perchloric-hydrochloric reactions is acid solutions. The results are given in Table I.
+
- d(Pu(IV))/dt
TABLE I THE DETERMINATION OF THE DISSOCIATIO?: QUOTIENT FOR PUc1+3AT 25' ICl-1, rnoles/La
0.0000 .1990 .2988 .3987 ,4988 .5990 .6994 .7999 .9006 1.002
EO'* v.
AEO'. v . X I O 3
K
-0.9819 ,9796 .9784 .9767 .9750 .9739 ,9730 ,9725 .9712 ,9701
0.0 2.3 3.5 5.2 6.9
.. 1.90 2.04 1.78 1.62 1.64 1.69 1.81 1.74 1.72
8.0
8.9 9.4 10.7 11.8 Mean
1.77rt0.10
[H+] = 1.000 in each case.
+
+ d(Pu(V))/dt = 3kiPu(IV)' -
3ki [Pu(III)*Pu(VI)/PU(IV)K']
+
CY
where K' is the equilibrium quotient, P u ( I I I ) ~ P u (VI)/Pu(IV) *,without acidity or hydrolysis corrections and a is the rate of a-reduction in the solution. Kasha and Shelines reported a decrease in the mean oxidation number of 0.0038 day-' in 0.95 M hydrochloric acid a t 25'. This rate is approximately 40Y0 of the a-reduction rate observed in perchloric acid. However, in the present work it has been found that in hydrochloric acid solutions of plutonium a t 2 5 O , with the ionic strength held a t unity with added sodium chloride, the a-reduction rate is essentially zero. -4dditional details on a-reduction studies in perchloric and in hydrochloric acids as well as in solutions of these two acids will be given in a later publication. Thus the "a" term in the rate law may be omitted in the disproportionation experiments in hydrochloric acid solutions. I n general, the Pu(V) concentration was less than 1% of the total plutonium concentration, and thus the term d(Pu(V))/dt may also be omitted from the rate law. Values of the specific rate constant for the forward reaction, kl, can be obtained from the expression
From the mean value of the dissociation quotient of the complex ion, PuCI+~,it can be calculated that in molar hydrochloric acid approximately 36y0of the total Pu(1V) is in the form of the complex ion. The measurements of Ahrland and Larsson5 indicate that U+a is complexed t o about the same degree (ca. 33% a t p = 1) by chloride ion, whereas Waggener and Stoughton6 found that the - d( Pu( IV))/dt/( 1 - K*/K')3Pu(IV)* kl first chloride complex of thorium is of somewhat where K* is the apparent equilibrium quotient, greater stability, about 60y0 existing in the comP U ( I I I ) ~ P U ( V I ) / P U ( I Va) t~ , a given time. In plexed form a t unit ionic strength. From the difTable I1 are listed the disproportionation equilibference in the heat of reaction for the Pur3-Pui4 couple in molar perchloric and in molar hydrochlo- rium quotients together with the fourth power hyric acids, -0.70 kcal./mole (vide znfra), together drogen ion dependence of this quotient. In the with the dissociation quotient of P u C I + ~it, can be fourth column of Table I1 the results have been corcalculated that for the reaction P u C I + ~iS P u + ~ rected for the effect of the hydrolysis of P u + ~ . ' C1- in molar hydrochloric acid a t 2 5 O , AF = TABLE I1 -0.34 kcal./mole, AH = -1.9 kcal./mole and DISPROPORTIOSATION EQCILIBRIUM QUOTIENTS I N HYDROA S = -5 e.u. The P u + ~ - P u + formal ~ potential CHLORIC ACID AT 25" measurements were also made in perchloric-hydrochloric acid mixtures a t a constant acidity of 2 M and a t an ionic strength of two. The value of the 1.000 1.7~3f 0 . 2 x 10-3 1.92 x 10-3 0.000 dissociation quotient under these conditions was 0,500 ..ino 2.40 i 0.04 x 10-2 1.80 x 10-3 1.70 f 0.03 a t 25'. Thus it appears that the single .zoo ,800 7 . 7 i 0 . 1 X lo-' 1.90 X complex, P u C ~ + is ~ ,sufficient to explain the data The results given in Table I1 for the quantity even in 2 M hydrochloric acid. V)~ the mean values Disproportionation of Pu(1V) in Hydrochloric P U ( I I I ) ~ P U ( V I ) / P U ( Irepresent Acid Solutions.-Connick' has shown that the of about five separate equilibrium experiments at over-all disproportionation reaction can be con- each acidity. Equilibrium was approached from both disproportionation and reproportionation disidered to consist of a rate-determining step rections with good concordance in the equilibrium ki quotient values. The reproportionation experiPu(1V) Pu(1V) Pu(II1) Pu(V) ments were carried out by heating the plutonium kr solution a t disproportionation equilibrium to 70" (5) S. Ahrland and R. Larsson, A c t a Chim. S c a d . , 8, 137 (1954).
+
+
+
(6) W. C. Waggener and R. W. Stoughton J . P h y s . Chem., 66, 1 (1952).
(7) R . E. Connick, THISJ O U R N A L . '71, 1628 (1949).
(8) M. Kasha and G. E . Sheline, "The Transuranium Elements." National Nuclear Energy Series, Div. IV, 14B, >IcGran-Hill Book C o . , Inc., New York, N. Y . , 1949, p. 180.
Dec. 5, 1055 CHLORIDE COMPLEXING AND DISPROPORTIONATION OF Pu(1V) IN HYDROCHLORIC ACID6147 for about ten minutes, then cooling the solutions to 25'. This short heating period shifted the equilibrium to produce concentrations of p ~ and + ~ P U O ~ which + ~ were greater than the equilibrium values a t 25'. Inasmuch as a fourth power hydrogen ion concentration dependence of the disproportionation equilibrium quotient in hydrochloric acid solutions is found, i t appears that the stoichiometry of the disproportionation reaction in hydrochloric acid is the same as has been found to be the case in perchloric acid solution^.^^^ Temperature Dependence of the PU+~-PU+* Couple.-Because the ratio [Pu+']/ [ P U + ~con] stitutes an essential part of the analytical method used in the disproportionation studies, the formal + ~ in molar hydropotential of the P U + ~ - P Ucouple chloric acid was measured as a function of temperature. These results permit the calculations of the thermodynamic functions to be made for the reaction
+
+
P u + ~ H + J' P u + ~
HZ
TABLE I11 THE TEMPERATURE DEPENDENCEOF THE Pu +"Pu FORMAL POTENTIAL IN 1 M HCl T,
E Q J volts ,
1 / T X 10*
EO'/T X 108
-0.94815 .97019 ,98208 ,99356
3.5782 3.3539 3,2425 3.1421
3.3927 3.2539 3.1844 3.1218
OK.
3.26
3.42 3.58 103. Fig. 1.-Evaluation of heat content change in 1 M hydrochloric acid for the reaction: P u + ~ H + F! P u + ~ l/zHz. AH288 = 14.32 kcal./mole.
I/T
I n Table I11 are given the formal potentials vs. temperature for one atmosphere of hydrogen, and for equal and small concentrations of P u + ~ and Pu'4.
279.47 298.16 308.40 318.26
3.10
+4
x
+
+
The value of AH = 78.4 kcal./mole was obtained from the slope of the line in Fig. 2. This change in heat content is for the reaction Pufa
+& +a / ~ H ~
+ 2H20
PUOZ+~
in molar hydrochloric acid a t 25'. The other thermodynamic functions a t this temperature are A F = 70.8 kcal./mole and A S = 25.4 e.u.
From the slope of the plot of E Q f / Tvs. 1/T, a value of A H = 14.32 kcal./mole was calculated and a t 25' the values of A F = 22.37 kcal./mole and A S = -27.0 e.u. for the reaction P u + ~ H+ Puf4 l / z Hz (see Fig. 1). The temperature dependence of the P u + ~ - P u +couple ~ in 1 M perchloric acid has been measured p r e v i o ~ s l y ,and ~ the results which have been obtained in the present work are in good agreement with the earlier findings. The values found in this work are AH = 13.63 kcal./ mole, A F = 22.6 kcal./mole, and A S = -30.2 e.u. Temperature Dependence of Pu(1V) Disproportionation.-The disproportionation equilibrium quotient was measured in molar hydrochloric acid as a function of temperature from 6 to 45". From I I I I I I I the disproportionation equilibrium quotient and 3.0 3.2 3.4 3.6 + ~ potential a t a given temperthe P U + ~ - P Uformal I / T x 103. ature, the formal potential of the P U + ~ - P U O ~ + ~ Fig. 2.-Temperature dependence of the formal potential couple can be calculated (see Table IV). for the reaction P u + ~ 2H20 e PUOZ+* H+ 3 / ~ H ~
+
+
TABLE IV VARIATIONOF DISPROPORTIONATION EQUILIBRIUM QuoTIENT WITH TEMPERATURE IN MOLAR HYDROCHLORIC ACID EQ"/T, 1/T ( p u + s ) i ( p u o * + nIH+l/E0,pU+9PU02+1 ) 1o
T,OK. 279.59 298.16 308.40 318 32
(Pu +4)
3.76 X 1.42 X 1.35 X 9.67 X
x
3
lo-'
-1.0301 1.0238 1.0202 1.0150
3.6843 3.4337 3.3080 3.1886
103
3.5767 3.3539 3.2425 3,1415
(9) R. E . Connick and W. H. McVey, THISJ O U R N A L , 1 3 , 1798 (1951).
+
+
+
in 1 M hydrochloric acid, AH288 = 78.4 kcal./mole, data derived from variation of disproportionation equilibrium ~ with temperature. quotient and P u + ~ - P u +couple
With Latimer's'O values of the entropies of H+, Hz and HzO it can be calculated from the entropy change for the P U + ~ - P U Ocouple ~ + ~ that the entropy difference, S p u a + 4 p u +=~ 12.1 e.u. Also, ~ in from measurements of the P u + ~ - P u +couple ( l o ) W. M. Latimer, "Oxidation Potentials," Prentice-Hall, Inc., New York, N. Y.,1952.
Second Edition,
molar perchloric acid as a function of temperature TABLE V it is found that Spu-i- SpuTa = -45.8 e.u. (The DISPROPORTIONATION RATE CONSTANTS IN HYDROCHLORIC ~ P u + obtained ~ entropy difference between P u + and ACIDSOLUTIONS AT 25' in perchloric acid was used to eliminate from this [H+I, jPiaC11, k l , moles-' moles/l. ZPu,moles/l. 1. hr,-1 kl[H+]a calculation the effect of the chloro complex of P u + ~ . moles/l. 1.000 0.000 5.623 X 0.47 0.47 However, any entropy contribution as a result of 6.410 .53 .53 + ~ been negthe chloride complexing of P U O ~ has 7.385 . 54 .54 lected.) A value of -41 e.u. has been estimated? 7.385 ,50" .50 for S p U + sin molar perchloric acid a t 25', and it is 7.798 .48 .48 assumed that this estimate is valid in molar hydro9.201 .58" .58 chloric acid as well. Thus, the entropies of Pu+' lo. 04 .63 .63 and P U O ~are + ~calculated to be - 57 and -28.9 e.u., respectively. This result for the entropy of the Mean 0.53 plutonyl ion is considerably more negative than the 0.31 10 ,300 7.138 X 4.76 0.595 entropy of the uranyl ion ( - 17 e.u.) which is some/ . 131 4.88 ,610 what surprising since from charge and size consid8.063 4.72 .590 erations alone, these oxy-cations would be expected Mean 0.60 to have comparable ionic entropies. Influence of Acidity and Temperature on Rate of 0.200 ,800 5,096 X 73.1 0.585 Pu(1V) Disproportionation.-The rate of dispropor5.573 61.1 ,489 6.954 74.9 ,607 tionation of Pu(1V) has been measured as a func7.192 77.5 ,620 tion of both acidity and temperature. Within a fivefold change in the hydrogen ion concentration >feat1 0.575 the specific rate constant, k l , has been found to Obtained from reproportionation studies. he inversely dependent upon the third power of chloric acid because the PuOz+concentration was not the hydrogen ion concentration, as shown in Table IT. This parallels the observations made in per- sufficiently small to be neglected during the early chloric acid solution^.^^^ However, it is of interest stages of the reproportionation reaction which were to observe that the rate of disproportionation of most suited to the determination of k l . The Pu(IV) is approximately five times greater in hy- method of analysis used in this work presumes that drochloric acid than in perchloric acid of the same the Pu02f concentration is negligibly small. Sufxidity. Connick and McVey3noted rather marked ficiently accurate values of the Pu02+-PuO2*? discrepancies in the value of kl in molar hydro- couple in hydrochloric acid solutions are not availchloric acid depending upon the direction from which able. These potentials together with the other the equilibrium was approached. This anomaly available analytical data would permit the concenwas not observed in the present study. The re- trations of all four ionic species to be coniputetl. The rate constant, k l , was obtained as a function proportimation results were within the experimental error of the determination of the rate constants. I t of temperature over the interval 6 to 4.3'. From a was not possible t o obtain accurate ralues of kl from plot of 2.303 log kl us. 1' T , the heat of activation for re!)rol,ortionation studies in 0 5 and in 0.2 J/ hydro- the disproportionation reaction was found to be 39 _ _ kcal. 'mole. The results are listed in Table V I and r--
4\
l
3.2
3.3 34 35 1 / T x 108. Fig. 3.-Determination of heat of activation for the 2H20 e ~ P u + ~ disproportionation reaction 3PuL4 P L I O ? + ~ 4H+ in molar hydrochloric acid, A H z ~ ~ *= 39 k c d . /mole.
3.1
+
+
+
TABLE \'I TEMPERATURE YARIATIOSOF DISPROPORTIONATION RATE IN MOLARHYDROCHLORIC ACID CONSTANT T,O R . I? I 2.303 log k i 1 / T X 10' 279.56 3.40 f 0 . 1 5 X -3.730 3.577 279,s: 2.79 =t0.10 x i n - * -3.580 3.577 -0.635 3.354 298.16 0 . 5 3 i 0.04 3.243 6.10 i 0.10 1.SO9 308.39 1.823 3.242 308.44 6.19 i 0.07 3.880 3.143 318.14 48.4 i 3 . 3 45.8 f 3 . 8 3.825 3.141 318.32
plotted in Fig. 3. Ah entropy of activation of 58 e.u. was calculated from this study of the effect of temperature upon the rate of disproportionation. Los ALAMOS, NEW MEXICO
