URANYLIONHYDROLYSIS EQUILIBRIUM
277 1
into the hydration sheath of the cation is possible. (Woodward has since observed a Raman effect in the case of the CHZ-Hg-NO3 complex. 15) The ferrocyanide ion pairs with potassium and sodium may admit to a similar explanation in that these small cations are able to penetrate (or form covalent bonds with) the ferrocyanide ion, thus losing their identity as cations in the resultant complex. The kinetic result for the potassium and sodium may thus be in-
ions, whose behavior is "normal" in producing ion pairs in accord with the existing theories, would be expected to affect the rate. We intend to investigate this possibility.
Acknowledgment. The authors acknowledge partial support of this work through st grant from the donors of the Petroleum Research Fund administered by the American Chemical Society.
Rapid Reaction Rates in a Uranyl Ion Hydrolysis Equilibrium1
by David L. Cole, Edward M. Aris Silzars, and Ronald P. Jensen
Dennis T. Rampton,
Department of Chemistry, University of Utah, Salt Lake City, Utah 84112
(Received Sovember 29, 1966)
+
kl
Rate constants for the uranyl nitrate equilibrium UQZ2+ H2Q
U02QH+ 6-
+ H + in
1
aqueous solution a t 2.5' have been determined by the dissociation field effect relaxation method. The specific rate IC-1 = 1.65 =k 0.33 X 1O1O J1-l sec-l a t an ionic strength 5 4 x 10-4ni.
Introduction Stability constants for the equilibria present in an aqueous solution of uranyl nitrate have been proposed by Baes and 3 l e ~ e r . At ~ 25" and p = 0.5 nz ionic strength, adjusted with KNOa, their results are
U0z2+
+ HzO
ki
,
U020H+
i-
+ H+ K1 = 2 X
2U02'+
+ 2H20
kl
k-
+ 5HzO
(1)
112
(2)
+ 2H+
(U02)2(OH)22+ 1
K2 3UOz2+
112.
=
1.2 X lo-'
+
(U0~)3(0H),+ 5H+ K3 = 6 X lo-''
(3) Thcsc constants were used by Spinnler and Patterson4 111~
in interpreting absolute Wien effects5s6 in aqueous uranyl nitrate and have also been employed in temperature-jump' and pressure-jump8 relaxation rate studies of reaction 2 . At concentrations between (1) Acknowledgment is made to the donors of the Petroleum Resenrch Fund, administered by the American Chemical Society, and t o the University of Utah Research Fund for financial support of this work. (2) To whom all correspondence should be addressed. (3) C. F. Baes, Jr., and K. J. Meyer, Inorg. Chem., 1, 780 (1962). (4) J. F. Spinnler and A. Patterson, Jr., J . Phys. Chem., 6 9 , 500 (1965). (5) 31. Wien, Ann. P h g s i k , 83, 327 (1927). (6) >I. Wien and J. Schiele, 2.P h y s i k , 32, 545 (1931). (7) A i . P. Whittaker, E. M. Eyring, and E. Dibble, J . Phys. Chem., 69,2319 (1965). (8) P. -4.Hurwitz and G. Atkinson, Abstracts, 152nd National Meeting of the American Chemical Society, S e w I-orli, N. T., Sept 12-16, 1966.
Volume 71,Number 9
August I967
2772
D. COLE,E. EYRISG,D . RAMPTOX, A.
and JI uranyl nitrate, Spinnler and Patterson4 observed no relaxations in the electrical conductance in the 0.5-4-psec time range. Whittaker, et al.,' measured a relaxation time 7 = 7 msec in a 14 uranyl nitrate solution at 25' and an ionic strength p = 0.5 d l which with similar data a t other concentrations led to k, = 116 AI-' sec-I. Their apparatus would not permit a reliable determination of 7's shorter than 20 psw. Hurwitz and Atkinsons determined a kz = 66 sec-I at 25" and p = 0.06 M and, though limited by their apparatus to the measurement of 7's greater than -250 psec, they estimated a lower limit for kL1 of 3.3 x lo6 J 1 - l sec-'. The discrepancy in the above values of k , is acceptable considering the great difference in ionic strength for the two kinds of experiments. I n the present work we report 7's of the order of 2 psec for 1-20 x 10-5 JI uranyl nitrate solutions a t 25' determined by the dissociation field effect (DFE) relaxation m e t h ~ d . ~These measured 7's permit a calculation of kl, L1, and K1 at essentially zero ionic strength.
Experimental Section Apparatus and Techniyue. In the DFE method, a chemical equilibrium involving an incompletely dissociated electrolyte is perturbed by a suddenly applied dc field of -lo5 v/cm. The relaxation of electrolyte concentrations from their equilibrium values at zero field t o their equilibrium values in a high field gives rise t o a rapid conductance change measurable with a Wheatstone bridge. The largest dissociation field effect, or second Wien effectj6would be anticipated for equilibria involving charge neutralization, but if the ionic species on the two sides of a chemical equilibrium differ markedly in equivalent conductance as in reaction 1, a relaxation should still be observable. In our apparatus (see Figure l ) , a rectangular highvoltage pulse is applied to a Wheatstone bridge. Two arms of the bridge consist of identical platinum sealed in glass cells with electrode separations of 0.3 cm and a cell constant a = 0.03 cm-l. Of these two cells, the J I ) aqueous sample cell S contains a dilute solution of uranyl nitrate. The reference cell R contains a dilute solution of aqueous HCl such that the two cells have the same electrical conductance measured on an ES1 >Iode1291B impedRricebridge Operated strong electrolyte solutiorl in R is reat 103 cps. quired to cancel a first Wien effect5 occurring in S. The other t ~ arms ) of the bridge are c o n v e n t i o d carbon resistors equal to one another but lower in resistance th:in s and R , i.e.,-1030hms as opposed to -7 X lo3 ohms. Rectangular high-voltage pulses The Journal of Physical Chemistry
SILZSRS, AND
R . JENSEN
RG U DELAY C A B L E
m GP-15 SPARK GAP
POWER SUPPLY
B
TO SCOPE EXTERNAL TRIGGER
f
Figure 1. Schematic of the square-wave dissociation field effect relaxation method apparatus constructed a t the University of Utah.
are generated with 730 m of RGSU coaxial cable and a Kilovolt Corp. IW 60-5 power supply. Triggering of the pulse is accomplished with an E. G. & G. Xodel GP-15 spark gap. Having balanced the probes with pulses from a Tektronix 114 pulse generator previous to the application of the rectangular 30-bv pulse, the exponential process of going out of balance a t high fields can be followed on a Tektronix 535A oscilloscope. The difference in voltage a t points -4and B is detected by Tektronix P6015 probes connected to a Tektroriix Type W high-gain differential preamplifier. A typical oscilloscope trace is shown in Figure 2. Sample solution pH was measured with a Beclimari 1019 research p H meter incorporating a 41263 glass electrode arid a 39071 calomel electrode. 2lIateriaZs. Conductivity water having a specific conductance of 1.4 X lo-' 0hm-I cm-' was prepared by an electrophoretic ion-exclusion technique.'O The uranyl nitrate hexahydrate was 1Iallinclirodt reagent grade with a nominal molecular weight of 502.15 g. We determined the actual molecular weight to be 492.43 g by gravimetric analyses involving an ammonia precipitation." -411 weighings were performed on a Mettler X5SA balance. The hydrochloric acid used in the reference arm of the bridge was Baker and Adamson reagent grade. All sample solutions were prepared under an atmosphere of Linde high-purity dry nitrogen. (9) 11. Eigen and L. DelIaeyer, "Technique of Orgnnic Chemistry," Vol. V I I I , Part 11, S. L. Friess. E. S. Lewis, and A . Weissberger, Ed.. Interscience Publishers. h e . , New T o r k , N. T.,1963, pp 988-
lool,
\v.
(10) Hnller and H. C. Duecker, J . R ~ S.yat/. . Bur. Std., A64, 527 (1960). (11) W. R. Schoeller and A . It. Powell, "The h d y s i s of l\Iiirer:1ls :)!Id Ores of the Rarer Elements," 3rd ed, Hnfner Publishing Co., Xew T o r k , N. T., 1955, p 305.
2u.3 li.3 l3.!1 11.1 8.21, 6.33 4. I4 1.46
4 .!IO 4.3i
16.6 I:I.!I
1.37 1.22
3. X!I
11.1 !I. "!I
0.622
4.tIX
2.1,i 2.2!1 l.4.i Li4
G.tN 4.81, 2 . .i3 l.l!l
0.979 11.671 0.404
0.365 0.0.56
several urauyl species i i i an aqucnus uitrnte solution of iouic,strerigth p = 0.5 M a t 2.5' :irc uot kuowu :tnd their calculation from rxtcuded Dehye-Huckcl theory would he specious.'J Howcver, from cxperimcutal meiiu activity coeflicictits for aqueous solutions of calcium nitrate, lithium iiitr:rte, :ind uitric acid,".'5 we cau estimxte that yuol.+ = ~(ao,~.(ati),:-= 0.36 nnd yll + = yan,olr+ Y ~ U O . ) , ( O B ~ , + = 0.72 at 25" ntid p = 0.5 ,lf, Thus, for the uranyl inti hydrolysis (eq 1) \YC \voultl have
Results and Discussion Our measured DVK relaxatiou times T iiloug with the correspouding couceirtratious of the various species calculated from the equilihriurn coustauts of eq 1-3 are shown i n Table I. The molar hydrogen ion cow ceutration [H+] \vas cnlculated from the experimental pH usiug the rclatioul2 aHf= 10-0" [ H f ] ; at these very low iotiic streiigths the Debye-Hiickel limiting h v activity coefficient Y H + lies in the range 1.0-0.97; hence the omissiou of y e + from the equatiou UIE+ = y s [ H + ] is a quite satisfactory approximatinu. I'ersuasive argumcuts cau be rniacd i n support of the use of the0.5 I I I innic strength cnucei~trationequilibrium constants iu cnlculntiug conceutrntious of the various urariyl species at iouic strengths p I 4 X lo-' M BS was doire iu Table I. The activity coefficients of the +
or Kal = 2.9 X The cnrrespondiug equilibrium quotients corrected to zero ionic strcugth for reactions 2 and 3 are K@ = 1.7 X 10-6 and Ka3 = 1.S X lo-". As an illustratiou wc can recalculatc the various roncentratious of Table I for Co = 13.9 X M and [H+] = 3.89 X 10-5 .If. Wc find [U02*+] = 10.18 X d l and [UO?OH+] = 7.59 x M . Sow we note that the relaxation time is concentratiou dependent, thus precludiug the possibility that a first-order process is responsihle for the observed effect. Also, only one relaxatiou time is ohserved in this time range. (A particularly importaut advantage of the square(12) 11. G . Riites. "Determination of pH-Theory and I'metire." John \ W e y rind Sons. h e . . K e w York, N. Y..1964. pi' 75-77, 91. (13) See. for example. H. S. Fmnk. Diacuasiom Fardo# h.. 21, 66 (1957).
(14) 11. H. Stokes and 11. A. Itohinson. 3. Am. C h m . Sw.. 70. 1870 (1948). (15) E. A. Moelwyn-Hughex. "l'hysienl Chemistry." 2nd e d , I'ergnmon I'ress. New York. N. Y.. 1961. 1) 899.
D. COLE,E. EPRING, D. RAMPTON, A. SILZARS,AND R. JENSEN
2771
errors in 7 are too large to justify such a procedure anyway. The value IC-1 = 1.65 X 1Oln M - I sec-' is somewhat less than a limiting value of 3.7 X 10'" J 1 - I sec-' that we estimate for a diffusion-controlled reaction between H + and UOzOHf after the manner of Debye.16 This discrepancy suggests that not all angles of approach of the proton to the U 0 2 0 H + are equally favorable for reaction. Were the dimerization equilibrium 2 responsible for our DFE relaxation data, the appropriate expression for r would be
0
I
I
1
I
I
2
3
4
([H']
t [UO,OH+])
,
5
M
Figure 3. Plot of the reciprocal of the experimental dissociation field effect relaxation time T us. the sum of the hydrogen ion and UOZOH+ ion concentrations a t 25'. The points are experimental data and the straight line is a least-squares fit of the data. The slope of the line corresponds to the rate constant k-1 and its intercept to kl.
wave DFE technique over the dispersion of the amplitude methodg is the greater likelihood of recognizing multiple relaxation times with the former.) If equilibrium 1 is responsible for the observed relaxation time, the appropriate expression for 7 is 7-l
=
kl
+ k-,([H+] + [U020H+])
(4)
and since k1 = k-&1 we find using [U020H+] = 7.59 X M , etc., that k-1 = 1.50 x 1Oln AI-' sec-I. As we see from Table I, the value of Ll a t Co = 1.39 X M calculated from eq 4 and the concentrations obtained with 0.5 i n ionic strength equilibrium quotients is 1.59 X 1Olo I l l - I sec-'. Considering the speculative nature of the activity coefficients used, the dominance of [H+] in eq 4, and the significantly greater uncertainty in k-1 introduced by the experimental errors in 7 than by the slight discrepancy in [U020H+], the correction of the equilibrium quotients to the ionic strength used is clearly unwarranted. Our kinetic results have been plotted in Figure 3. The slope of the least-squares straight line is k-1 = 1.65 X 1 O I n J 1 - I sec-'. the intercept is kl = 1.73 X lo4 sec-l, and the quotient kl/kLl = K 1 = 1.05 X 10-6 dl (p 5 -1 X lop4X)in reasonable agreement with ill (p = 0.5 Ill) of Baes and the value K1 = 2 X I\I~yer.~ An iterative method of data treatment in which our kinetic k-1 is used to recalculate equilibrium concentrations, etc., is ruled out by the unavailability of similar "improved" kinetic values of Kz and KB for our low ionic strength conditions. The experimental T h e Journal of Physical Chemistry
This equation yields a poor fit of the experimental data. Furthermore, the least-squares straight line sec-I has a negative slope k-z = -4.43 X 1014 that decisively rules out equilibrium 2 as a major factor in the observed dissociation field effect. Thus, our results agree with the concluding statemerit by Spinnler and P a t t e r ~ o n ,based ~ 011 rionkinetic evidence, that "the effect of field is on the involvement between hydrogen ion and the uranyl complex, presumably UOzO H + . . . . " We have observed relaxation times over the entire 10-60" temperature range and confirm qualitatively both the larger amplitude of the dissociation field effect a t elevated temperatures and also its negative character; ie., the conductance decreases under the application of an intense field at GO", reported by these same ~ v o r k e r s . ~ 111 principle, we should be able to determine kl, k-1, and K1 at elevated temperatures by the same DFE technique and thus determine both forward and backward activation energies, enthalpies. and entropies for equilibrium 1. Presently, both the thermostating of the bridge cells and more important the measurement of sample solution pH at elevated temperatures are insufficiently accurate to warrant
Table I1 : Experimental Ilissociation Field Effect Relaxation Times in Aqueous Cranyl Nitrate at 60'
co.
f\.
x
10-5
4,06
1.48 1 83
2.46 I .43
1.96 2.32
8.10
(16)
T,
rsec
P.Debye, Trans. Electrochem.
SOC.,8 2 , 266 (1942).
REACTIONS IN
THE
IRRADIATION OF WATERVAPOR
such an exercise. We have found, however, that the experimental relaxation time does decrease significantly with rising temperature (see Table 11). This
2775
fact is consistent with our assignment of the measured D F E to equilibrium 1 since the concentration of U02O H + increases with t e m p e r a t ~ r e . ~ , ~
Reactions of Electrons and Hydrogen Atoms with Oxygen and Methyl Bromide in +radiated
Water Vapor
by G. R. A. Johnson and M. Simic Laboratory of Radiation Chemistry, School of Chemistry, The University, .l’ewcastle u p o n T y n e , England Accepted and Transmitted by The Faraday Society
~
(December 7 , 1966)
~~~
Oxygen and methyl bromide react efficiently with the electrons and hydrogen atoms produced by irradiation of water vapor-propane mixtures. I n the D 2 0 C3H8 0 2 system, the dependence of the H D and H2 yields on the ratio Oz/C3H8is consistent with a simple competition of 0, and C8H8for the D and H atoms produced radiolytically from D20 and CaHs, respectively. The rate-constant ratio [~I]IC~D+~*+J~)/~(D+C~H~) = 105 f 10 (11 = D 2 0 , 140”). The same value, within experimental error, was obtained for the corresponding ratio for H atoms. The D atom yield, extrapolated to zero concentrations of O2 and C3H8 (G(D),O = 5.5 0.5)) is consistent with efficient electron capture by 0 2 , which prevents the formation of D atoms by hydronium ion neutralization. I n the system DzO C3H8 -I- CH3Br, the dependence of the H D and CH3D yields on the ratio CH3Br/ CH3Br + CH3D Br, C3H8has been studied. It is concluded that the reaction, D competes with the dehydrogenation of C3HRby D atoms and that G(D),..! = 4.6 f 0.4 and JC(D+CH~B~)/~C(D+C~H~) = 2.9. It is suggested that electron capture by CH3Br gives CH3Br- and that neutralization involves reaction of this negative ion with D 3 0 + to give CH3D as a product. The formation of CH3D by this process is eliminated by SFs and by 02,which are more efficient electron scavengers than CH3Br.
+
+
Introduction I t has previously been shown that in the radiolysis of mater vapor, the formation of hydrogen atoms can be attributed to a t least two different One of these involves hydronium ion-electron neutralization; the other does not involve hydronium ions. X20 and SF6,which are known to be capable of capturing electrons to form negative ions, completely supress the formation of hydrogen atoms from hydronium ions. I t ha5 been concluded that neutralization of hydronium ions by the negative ions from these compounds does
+
+
+
not involve simple electron transfer, but that atomic rearrangement also O C C U ~ S We . ~ have now investigated the neutralization of hydronium ions by the negative ions formed when oxygen and methyl bromide are used as electron scavengers. These compounds have previously been shown to be effective as electron scavengers in the radiolysis of methanol vapor.4 We (1) J. H. Bxxendale and G. P. Gilbert, Sciencr, 147, 1571 (1965). (2) G. R. A. Johnson and M. Simic, Nature, 210, 1356 (1966). (3) G. R. A . Johnson and hi. Simic, J . Phys. Chem., 71, 1118 (1967).
Volume 7 1 , Xumber 9
August 1967
