1812
J.
c. HINDMAN, J. c. SULLIVAN AND D. COIIEN
water and excess 8-quinolinol, was obtained analysis showed the composition NbO(CpH6ON)S. The same technique by which niobium 8-quinolinolate was prepared failed to produce a compound from a pure tantalum solution. When pure tantalum and niobium were mixed however, tantalum appeared to have formed a similar compound. When tantalum pentachloride and 8-quinolinol were combined in benzene solutions, a stable addition compound containing about 5 moles of 8-auinoliuol was formed. Kiobium Dentachloride appeared to form a compound similar t o the tantalum
Vol. 80
addition compound with 8-quinolinol but which was quite unstable. Niobium 8-quinolinolate shows a strong characteristic band at about 10.9 p in the infrared spectra t h a t apparently is due to the “NbO” gr0up.~J0 (9) R. G. Charles, H. Frieser, R . Friedel, I,. E . Hilliard and W. D. Johnston, S$ectrochim. Acta, 8, 1 (1956). (10) J. P. Phillips and J. F. I k y e A n d . C h i m . A r l a , 17, 233 (1957).
RCFFAT,O,N. Y .
[CONTRIBUTION FROM ARGONNE NATIONAL LABORATORY]
Kinetics of the Neptunium(II1)-Neptunium(V) Reaction in Perchlorate Solution1 BY J. C. HINDMAN, J. C. SULLIVAN AND DONALD COIIEN RECEIVED OCTOBER10, 1957 The kinetics of the Sp(II1)-Np(V) reaction in perchlorate solution have been investigated. The rate law for the forward reaction can be written as - d[Np02+]/3t = ~ I [ N ~ + ~ ] [ N ~ O L + ] At [ H 25” + ] . and p = 2.0, kl = 2590 1.2mole-* min.-l. The rate is slightly decreased in a deuterated medium. The experimental activation energy is 6.5 f. 0.5 kcal. in HCIOI solution and 5.6 2c 0.5 kcal. in DClOa solution. Possible mechanisms for the reaction are considered.
Reactions which involve the oxidation of Np(1V) to Np(V) appear to proceed through hydrolyzed species of the lower ~ t a t e . ~ Similar -~ behavior is noted for reactions of the analogous species of U(IV)5 and plutonium (IV).6 It has been suggested? that the mechanism of such reactions involves the formation of species which can readily be converted to the XO2+ or XOz++ ions. It is of considerable interest to examine the mechanism for the reverse situation, ;.e., the reduction of an oxygenated cation to an ion of a lower unhydrolyzed oxidation state. The reaction studied in the present investigation was the reduction of Np(V) by Np(III), yielding Np(1V) quantitatively. Two investigations have been made of the reduction of Np(V) by Fe(II).2r7 The rate law for the reaction appeared to be given by the expression -4 simid(NpOz+)/dt = k [NpOz+][Fe++][H+]. lar rate law has been indicated for the reduction of PuOz+ by P u + ~ .The ~ role of the hydrogen ion in these reactions is of particular interest. For this reason the present investigation has been carried out both in HzO and I 3 2 0 solutions. Experimental The general aspects of the experimental techniques and preparation of reagents have been described in previous paper^.^^^ Two methods were employed in the preparation of the Np(II1) solutions. T h e first method used was 3. reduction of the Np(V) stock solutions with hydrogen gas at a platinized platinum catalyst. Considerable difficulty was experienced in obtaining consistent results with these solutions. More reproducible (1) Work performed under the auspices of the U. S. Atomic Energy Commission. (2) J. R. Huizenga and I,. B. Magnusson, Tms JOURNAL, 73, 3202 (19.5 1). ( 3 ) J. C . IIindmnn, J. C. Sullivan and D. Cohen, ibid., 76, 3278 (1954). (4) J. C. Sullivan, D. Cohen and J. C. Hindman, i b i d . , 79, 4029 (1957). 1 3 ) R. H. Betts, Can. J. Clzetn., 33, 1780 (1955). (6) See R. E. Connick, Chapter 8, Vol. 14A, National Nuclear Energy Series, Div. I V , “The Actinide Elements,” SIcGraw-Hill Book Co., Yew York. S . Y . , 1049. (7) I.. R. Xlajinusson, J. C. Hindman and T. J. I,aChapelle, Paper It5.11, Vol. 14B. National Nuclear Energy Series, “The Transuranium F:lrnirnts,” M c G r n w - R i l l Rook Co., New York, N. Y., 1919.
results were obtained if the Kp(111) solutions were prepared by electrolytic reduction. The instability of Kp(111) with respect to air oxidatioii necessitated working under an inert atmosphere. Deaerated Np(V) solutions of the desired composition were transferred to five-cm. quartz absorption cells, modified by addition of an extra opening for the insertion of a gas bubbler assembly. Gas was kept flowing through the cell until temperature equilibrium was reached (ca. two hours) after which the Np( 111) was transferred under nitrogen to the absorption cell and mixed. The course of the reaction w p followed by spectrophotometric monitoring of the 7230 A. absorption band of Np(1V). The temperature in the absorption cell compartment of the Cary spectrophotometer wa5 maintained to zt0.1’.
Results and Discussion Molecularity of the Reaction.-The molecularity of the reaction with respect to the metal and hydrogen ions was determined by experiments in which the concentrations of the species were varied in the reaction mixture. The reaction was found to be bimolecular with respect to the metal ions. Data for a typical experiment are summarized in Table T. The apparent rate constant, k’, was calculated TABLE I Xp(II1) WITH x p ( v ) AT 16.1 0 1” = 2.0, [€I+] = 0.31 M , Np(II1) Kp(V) = 6.90 X 10-4 .if
THERATEO F RRACTION p
OF
Apparent rate constant
Apparent rate constant
k’,
Time, min.
7230
A.
I , mole-’ min - 1
Time, min.
0.25 .50 .75 1.00 1.25 1.50 1.75
0.092 ,170 ,233 ,282 ,331 ,371 ,405
614.5 623.8 619 7 603.0 610 0 GO9 0 004 7
2.00 2.25 2.50 2.75 3.00
Doh..
07
k’,
nab..
7230
A.
0.437 ,466
,492 ,518 ,540 ,960 Av.
1. mole-1 min. - 1
605.8 607.9 609.8 618.0 621.5
612.3 5 G . O
by the integrated equation for a bimolecular process. No correction for the back reaction is necessary since a t equilibrium the concentrations of Np(111) and Np(V) are negligible. Data on the effect
KINETICSOF Np(II1)-Np(V) REACTION IN PERCHLORATE SOLUTION
April 20, 1958
of changing the metal ion concentrations are given in Table 11. The data are consistent with the biTABLE I1 EFFECT OF METAL IONCONCENTRATION O N THE R A T E OF THE Np(II1)-Np(V) REACTION AT 16.1’ p = 2.0, [ H + ] = 0.31 )! lT Initial h’P(V), moles/l.
x
2.95 3.41 6.82 6.82 10.40 6.90 6.90
Apparent r a t e constant
k’,
1. mole-’ min. - 1
104
10.05 10.05 ;i. 03 5.03 2.51 6.90 6.90
Av.
588 554 536 517 555 612 606 567 f 31
molecular dependence of the rate on the Np(II1) and the Np(V) concentrations. Table I11 summarized the experiments a t varying hydrogen ion concentration.
1813
Mechanism of the Reaction.-The direct dependence of the reaction rate on hydrogen ion concentration indicates that hydrolysis proclucts of Np(II1) are not involved. The general correspondence of the stoichiometry suggests that the fundamental mechanism is probably the sxme for the Np(II1)-Np(V), Fe(I1)-Np(V) and Pu(V)Pu(II1) reactions. TABLE IV THE EFFECTOF DEUrERIUM AXD TEMPERATURE O U IXE RATEOF THE iVp(III)-Np(V) REACTIOX p = 2.0, [H+]i = 0.25 X , Kp(II1) = Np(V) G X lOP4AI ki(H) (mean), 1 . 2 mol-2 min. - I
Temp., ‘(2.
5.53 1190 16.10 1870 25.00 2590 a A 95% deuterium content.
Ki(D)a
-
(mean), 1.2 mole-2 min. - 1
k H/l: D
1060 1560 2060
1.13 1.20 1.25
The role of the hydrogen ion is not clear. A reasonable assumption would appear to be that the hydrogen ion is involved in a pre-equilibrium reaction with the NpOz+ ion to give NpOOH++. TABLE I11 EFFECTOF [H+] ON THE RATE OF TIIE ~ p ( I I I ) - ~ p ( v ) I n view of the positive charge on the Np(II1) ion i t would not appear likely that the activated complex is REACTION bridged through hydrogen. The relatively slight t = 15.90, = 2.0, [ ~ ~ ( I I I = ) ] [ K ~ ( V ) I 6 x 10-4 M effect of deuterium on the reaction is supporting Apparent rate cons,tant evidence that the hydrogen is not directly involved k , [H’l, in the reaction. The low activation energy for the 1. mole-1 m i n - 1 k‘/[H+] moles 1. -1 reaction suggests that extensive rearrangements 0.109 226 2070 are unlikely in the formation of the activated com,108 209 1940 plex. Two configurations that would be consistent ,260 509 1960 with the data are ,275 572 2080
-
,523 ,523 1.019 1.018 1.91 1.91
980 1020 1645 1780 3720 3290
1880 1950 1610 1750 1930 1720 1870f115
Np( V)
Np( 111)
A
Np(V)
Sp(II1) B
For configuration B, one of the two possibilities, ;.e., an oxygen atom transfer mechanism, would appear unlikely since a single electron reaction step Analyses of the data by least squares yields a should be favored. There are three remaining hydrogen order of [H+]’.05 * O . O z . The rate law possibilities : electron transfer across an oxygen or for the forward reaction can therefore be written as hydroxyl bridge and a hydroxyl atom transfer. If the mechanism is by electron transfer via an oxygen bridge there would not appear to be any reason for the involvement of hydrogen ions in the reacThe value of kl is computed from the observed bi- tion. Indeed, the increase in net charge for the molecular rate constants by the relation activated complex due to the hydrogen would be ki = k’/[H+] (2) expected to decrease the probability for reaction. The k1 value computed from the hydrogen ion data On the other hand, a mechanism which involves the formation of a hydroxyl bridged intermediate 115 1.2 mole-2 min.-‘. in Table I11 is 1870 This is in satisfactory agreement with the mean would provide an explanation for the appearance value of 1830 f 180 L2 mole-2 min.-’ obtained of hydrogen ion in the stoichiometry of the reaction. The data cannot distinguish between a procfrom the metal ion order data in Table 11. ess in which an electron transfer is considered to The Effect of Deuterium on the Rate -As shown in Table IV, the rate is decreased slightly in occur by a radiationless transition between the metal ions which are bridged by a hydroxyl ion the presence of deuterium. The Effect of Temperature.-Least squares and a process in which the electron transfer occurs analysis of the data on the temperature coefficient by way of an intermediate state in which the hyof this reaction rate on HzO and DzO yields activa- droxyl bridge approaches the configuration of a 0.53 and 5.62 f 0.54 kcal., hydroxyl atom. The small isotope effect could be tion energies of 6.52 respectively (95y0confidence level). The apparent attributed to the stretching of the 0-H bond in the activation energies are the lowest thus far found formation of the activated complex. “Solvent effects” have also been used to explain isotope for any neptunium oxidation-reduction system. Mean
*
LEONARD NEWMAN,
1814
JOAO DE
0. CABRAL AND DAVIDN. HVME
VOI.
80
effects of this order of magnitude.8 The N P + ~ tion of the dismutation products of the activated (aq.) ion presumably is formed by further reac- complex with hydrogen ions. (8) Cj.. V. Gold and D. P. N. Satchell, @rrart. Reo.. 9, 51 (1955).
(COSTRIBUTION FROM
THE
LEMONT, ILLINOIS
DEPARTMENT OF CHEMISTRY AND LABORATORY FOR NUCLEAR SCIENCEO P
THE
MASSACIIUSETTS
INSTITUTE O F TECHNOLOGY]
A Polarographic Study of Mercuric Cyanide and the Stability of Cyanomercuriate Ions' BY LEOSARD N E W M A N , JOAO DE 0 . CABRAL
AND
DAVIDN. HUME
RECEIVED OCTOBER 28, 1957 The reduction of mercury cyanide and the oxidation of mercury in hydrogen cyanide were shown to be polarographically irreversible in acid solutions, becoming reversible in basic solutions. In basic solutions the cyanide released by the reduction can complex mercuric cyanide a t or around the surface of the drop. The species being reduced are Hg(CN)*, Hg(CN)?The logarithms of th: over-all formation constants of these species are 33.9, 38.1 and 40.6, a t an ionic and Hg(CN)a--. strength of 2.0 and a temperature of 30
.
Introduction The reduction of mercuric cyanide a t the dropping mercury electrode was first extensively studied by TomeY who found in the range pH 6.0 to 9.5 a reversible wave corresponding to the over-all electrode reaction Hg(CN)*
+ 2H+ + 2e-
--f
Hg
+ 2HCN
I n basic media the reaction became Hg(CN)2
+ 2e- -+ Hg 4-2CX-
The half-wave potentials in the pH range above 9.5 were observed to deviate from predicted values, and this was attributed to complexation of mercuric cyanide by cyanide ions liberated in the electrode reaction. Tome3 offered no experimental evidence for the formation of higher complexes, and the deviations were in the opposite direction from that expected to be caused by such an effect. It has been the purpose of the present investigation to study the polarographic behavior of mercuric cyanide over a much wider range of PH values and determine the nature and effect of possible complex formation. Since the completion of this study,' Tanaka and hlurayama3have reported an extensive study of the anodic wave of mercury in cyanide medium which corroborates our findings regarding the formation of higher cyanide complexes of mercury a t the drop and leads to values for the formation constants which are in substantial agreement with those determined froni our cathodic waves of mercuric cyanide. About the same time, Anderegg4 published the results of a potentiometric determination of mercuric cyanide stability constants. Theory Tome3 derived for the equation of the wave, the expression
and pointed out that E l / , was not independent of concentration. I n place of El/, it is more conven(1) Taken in part from the Doctoral Dissertation of Leonard Newm a n , Massachusetts Institute of Technology, 1956. ( 2 ) J. Tome:, Collection C5cch. Chem. Commtrn., 9, 81 (1937). (3) N. T a n a k a and T. Murayama, 2 . physik. Chem. N . F., 11, 366 (1957). (4) G . Anderegg, Helv. Chim. Acta, 40, 1022 (I'J.77).
ient to use the concentration-independent potential on the wave, which we shall call Ei and which leads to Ede = Et
+ 0.030 log ( '9)
for the equation of the wave, It is shown readily that KA COH . f ~f f m - ( 3 ) Et = E' 0.060 log
+
+
+
KA
in which E' = Eo
d2 + 0.030 log 1 dnKzf2cN-
(4)
where K A is the ionization constant of hydrogen cyanide, COH+is the concentration of hydrogen ion a t the electrode surface, f H + and f C N - are the activity coefficients of hydrogen and cyanide ions, Eo is the standard potential of the mercury-mercuric ion couple, dl and d, are the diffusion coefficients of cyanide and mercuric cyanide, respectively, and Kz the formation constant of mercuric cyanide. It is seen that Ede is equal to Ef when (id i)/i2 is equal to unity and that Er is independent of mercuric cyanide concentration. It should be noted, however, that the relative position of Ef on the wave shifts with concentration. If all measurements are done a t a fixed, high ionic strength, activity coefficients may be assumed to be constant. The equation of the wave predicts a plot of Ef vs. pH to be linear with a slope of 60 mv. per #H as long as CoH+rH+fCN- >> KA. For values of C'H'fH'fCN
