HYDROLYSIS OF URANIUM(VI)
April, 1963
more negative than -15.5.64760 Kote that the lack of agreement among these studies is independent of the heat of formation of Si02. Consider also the two studies which appear to involve silica in t,he determination of the heat, of formation of silicon carbide. From the combustion mrasurrincnts of FIumphrey28for the reaction SiC(cub.)
+ 202
=
SiO,
+ CO,
(9A) me obtain AHfnZg8 (Sic, cub.) - AHf0298 (SiOz,a-quartz) = +196.86 f 0.88 kcal. From t,he equilibrium pressure measurements of TaylorG1for the reaction
Si02
+ 3C(graphite) = SiC(cub.) + 2CO(g)
(10A)
we obtain by the third law (after a minor correction for the Presence of in the gas) aHf0z9e (sic, Cub.) -AHf0298 (SiOz, a-quartz) = +201.9 1.1kcal. It is obvious that these values of the quantity, AHf(SiC) -
*
(00) .J. Smiltens, J. Phvs. Chem., 64, 368 (1960). ( 0 1 ) J. D. Baird and .I. Taylor, Trans. Faraday Soe., 64, 526 (1958); D. A. R. Kay and J. Taylor, ibid., 66, 1372 (1960).
821
AHf(SiOz), which is independent of the data for silica, do not agree with each other within their estimated prccisions. (For comparison with the results given in the previous paragraph the studies of Humphrey2*and of TaylorR1give, with OUT value of AIlf0298(SiO,, a-quartz) AWfoZg8(Sic, cub.) = -20.9 f 0.9 and -15.%j f 1.1kcal. mole-I, respectively). From the foregoing discussion it is clear that there exist unknown sources of error in some of the measurements leading to the above values of the heat of formation of silicon carbide. A critical redetermination of the heat of formation of silicon carbide by calorimetry should shed light 011 these discrepancies. IV. In conclusion. we find that the value for the heat of formation of a-quartz presented in this work is in good agreement with all the high temperature equilibrium measurements on substances which appear to be reasonably well characterized. Thus, the discrepancies discussed by Chipman (except those involving Sic) appear to be substantially resolved.
HYDROLYSIS OF URANIUM(V1) : ABSORPTION SPECTRA OF CHLORIDE AND PERCHLORATE SOLUTIOKS1 B Y RICHaRD
bf. RUSHAND JAMES 8. JOHNSON
Chemistry Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee Received September 21, 1962 Optical absorptions of hydrolyzed U(V1) solutions in 1 M chloride have been measured for U(V1) concentrations in the range 0.001-0.1 M and for values of n (average moles of hydroxide bound per mole of U(V1)) up to 1.3. A few measurements have been carried out for 1 N perchlorate solutions with n up to 1.0. The molar absorptivity E increases sharply with n; chloride complexing of the unhydrolyzed species and a t least one hydrolyzed species is indicated by higher values of E in chloride solutions than in perchlorate solutions with the same n. The chloride measurements are correlated with a hydrolysis scheme previously derived from ultracentrifugation and acidity measurements. These results are combined with acidity measurements and spectra in perchlorate media to give values of the formation quotients for hydrolytic species in perchlorate media. The results are consistent with the earlier indication that the species ( U O ~ ) Z ( O H ) which ~ + ~ , appears to be present in substantial amounts in 1 M chloride, is of little, if any, importance in 1 A+' perchlorate.
In a recent publicati~n,~ we presented a hydrolysis scheme for U(V1) in one molar chloride, which was consistent with acidity measurements and with ultracentrifugation estimates of molecular weights (25'). The major species proposed were (UOz)z(OH)2+2, ( U O Z ) ~ ( O H ) ~and + ~ , (UOZ)~(OH),+.Formation quotients were given for these species. A scheme, similar except that it did not include (TJ02)3(OH)4+2,based on literature acidity measurements3 carried out at 20', was also presented for one molar perchlorate solutions. Although the species as written do not indicate complexing of other ligands present in the solution, from the fact that (UO&(O€I)4+z appeared to be of importance only in chloride solution, presumably this species includes complexed chlorides and it seemed possible that other species also complexed anions. We have recently carried out optical absorption measurements of hydrolyzed uranyl solutions in chloride and perchlorate media. I n this paper we attempt to correlate the spectra of the chloride solutions with (1) This document is ba3ed upon work performed for the United States Atomic Energy Commission a t the Oak Ridge Xational Laboratory, operated by Union Carbide Corporation. (2) R. M. Rush, J. S. Johnson, and X. A. Kraus, I n o i v . Chem., 1, 378 (I 962). (3) S. Ahrland, Acta Chem. Scand., 8, 374 (1949).
the hydrolysis scheme earlier reported. From the species absorptivities thus obtained, plus the spectra and acidity measurements in the perchlorate media, we have estimated formation quotients for the hydrolyzed species in 1 M NaC104. Experimental and Computational Procedure Spectra of solutions having the following hydroxyl numbers n (average number of hydroxyls bound per uranyl) were measured: 1 M in total chloride (with Ea+): 0.1 M U(V1)-n = 0.38, 0.67, 0.90; 0.01 &f U(VI)---n = 0, 0.12, 0.21, 0.30, 0.42, 0.51, 0.61, 0.71, 0.79, 0.89, 0.98, 1.07, 1.21, 1.30; 0.001 M U(V1)-n = 0.45, 0.73, 0.99, 1.21. In 1 M total perchlorate (with Xa+): = 0.42; 0.01 M U(V1)-n = 0, 0.43, 0.62, 0.1 M U(V1)-n 0.81; 0.001 M U(VI)-n = 0.45, 0.64, 0.82, 1.00. Interpretations are based on measurements in the range 36505000 A., carried out on a Gary Model 14 PM spectrophotometer; cells were of 0.5, 1, 2, 5, or 10 em. path length. This instrument gives a recording of the absorbance of the solution, A = log ( I o / I ) ,I being intensity of light transmitted by the solution and 10, of that transmitted by a reference solution (1 M KaCl or 1 M NaC104). From these values, the molar absorptivities E = A / c b are obtained, where c is the total stoichiometric U(V1) concentration (moles/l.), and b is the path length in cm. The molar absorptivity is related to the species absorptivities, e,,,, by the equation
RICHARD AI. RUSHAND JAMES S.JOHNSOS
822
Vol. 67
where ci.j is the concentration of the species (UO,),(OH)j+(Zl- j) (complexing with other ligands being ignored), Fi,j = ici,j/c is the fraction of total uranium found in the (i,j) species, and d i , j = ~i,i/i. I n the computational procedure, the species fractions Pi,j are computed from the formation quotients of the scheme being tested, for the hydroxyl number and total U(V1) concentration of the solution being considered. The formation quotient kl,j is defined as
(kOi,j being the formation constant, G i , j the appropriate activity coefficient ratio, and brackets indicating concentration in moles/ 1.). For a given solution the hydrogen ion concentration CH and free uranyl concentration c1 ”, are obtained from the hydroxyl number n and the formation quotients ki,j by solving the following equations by the h-ewton-Raphson method
WAVELEVSTH, anqstroms.
spectra of 0.010 M U(V1) solutions in 1 1%’ ) and 1 M total perchlorate (- - - -) a t 25”.
Fig. 1.-Absorption total chloride (--The species concentrations c ~ and . ~ fractions F,,j are then calculated from CLO, CH, and the formation quotients. With these fractions for each solution in a given medium (chloride or perchlorate), values of a t a given wave length for each hydrolyzed species are calculated to give the minimum sum of squares deviation4 of the values of E for all of the solutions. The values of €1.0 are obtained from the spectra of an unhydrolyzed solution. The program also computes the standard error in each e l l t Jand the deviations between observed E and values computed by equation l . Computations were carried out on an IBM 7090 computer. The criterion for correlation is the degree of success with which the observed spectra of individual solutions can be reproduced from the values of e l L , ,obtained in the manner described. Equilibrium quotients for the same conditions as the measurements on perchlorate solutions (25”, 1M KaC104) were not available. Estimates of these values were made in a manner t o be described in detail below with the help of separate measurements of the solution acidities. The correlation of the absorption spectra was then tested as with the chloride solutions. Acidity measurements were carried out by the technique described previously2 with the cell
I
r
y
-
1
1
-I
80-
1
glass electrodel (UOz(OH),+(2-n), H+, Xa+)C104(1 Jl)jHCl (0.01 M ) , HClOd (0.01 M ) , NaC104 (0.98M), AgCl(s); Xg(s) Results and Discussion 1. Results-The general features of the absorption spectra can be seen in representative curvoes given in Fig. 1 5 (measurements were made to 5500 A. at which point the absorption was negligible indicating the absence of serious turbidity). The values of E increase sharply with hydroxyl number, and the position of maximum absorption shifts to longer wave length in both chloride and perchlorate media. Similar results have been reported in perchlorate by Sutton.6 Chloride complexing of the unhydrolyzed species is in(4) We are indebted to &I. H. Lietzke for use of his Fortran generalized least squares subroutine; see U. 8. Atomic EnerRy Commission, ORNL3259 (1962). (5) Those interested in a more detailed presentation of the primary data are referred to R. M. Rueh, J. S. Johnson, and K. A. Kraus, U. S. Atomic Energy Commission, ORNL-3278 (1963). (6) .J. Ruttcn, J . Chem. Soc., S275 (1949); see also National Reuparch Council of Canada, Atomic Energy Project, Division of Research, CRC 325 (K.R. C. No. 1612) (1947).
4000
4500
5000
WAVE 1. ENGTH , o nqs t rom s.
Fig. 2.-Species molar absorptivities for U(V1) species in 1 M total chloride: equilibrium quotients used are k z , ~= 6.7 X IO-’, ka,a = 4.7 X lO-l3, and kj,s = 1.0 X lo-’’; vertical lines indicate one standard error on each side of the value.
dicated by the difference in spectra a t n = 0, and of at least one hydrolyzed species by the greater absorption in 1 M NaCl than 1 M KaC104 a t a given n. Since free chloride is nearly constant in all experiments, the ratios of species of a given (i,j) which are complexed to different extents by chloride will be essentially constant and the values of E’i,j will be a composite for species (UOn)i(0H)jCl~+(2i-’-1). 2. Uranyl Spectra in 1 M Chloride Solutions.-In the interpretation2 of our acidity measurements of U(V1) hydrolysis, the simplest scheme giving an adequate fit involved the hydrolyzed species (UO&(OH)z+2,
HYDROLYSIS OF URARIIUM(VI)
April, 1963
( U ~ Z ) ~ ( O H )and ~ + ~(U02)3(OH)sC; , these species (as well as the unhydrolyzed uranyl) may further complex chloride ions. An equally satisfactory fit was obtained if the species U02OH+ was included. Although there is some evidence from other studies’ for this species, especially a t higher temperatures,$ it does not seem to us conclusive at the temperature and concentration range of our studies. I n any case, the formation quotients we obtained indicated that, if present, it did not constitute a major fraction of total U(V1) under the conditions of the present study, and we have therefore neglected it for the most part in this discussion. The formation concentration quotients which we obtained2 for the scheme without the (1,l) species are as follows: kz,z = 6.7 X lo-’, ka,r = 4.7 X and 1c3,5 = 1.0 X If the (1,l) species is included, the other quotients are modified slightly. The values of the species absorptivities, obtained by computing the concentration of the individual species with the above scheme, and finding the values of e ’ i , j which give the best fit to the observed spectra for all solutions (see Experimental section) are given in Fig. 2. Vertical lines given at representative wave lengths indicate one standard error in on each side of the “best” value. Correlation is illustrated in Fig. 3 as deviations of obsewed values of E from those computed for individual solutions. Most 0.1‘ the deviations represent less than one chart division (0.01 at A ) on the recorded spectrum. In the most notable exception, 0.1 M U(VI), n = 0.899, A was very high (up to 2 ) , and the computed values are within 2% in A where the deviation is greatest. Inclusion of the (1,l) species modified the values of e ’ i , j for the other hydrolyzed species shown in Fig. 2 slightly; the over-all fit corresponding to Fig. 3 was about the same. 3. Uranyl Spectra in 1 M Perchlorate Solutions.The equilibrium quotients2 derived on the basis of Ahrland’s acidity mearurements in perchlorate media3 are not strictly applicable to the interpretation of the spectra reported here, since Ahrland’s measurements were at 20”. Our interpretation of his results had indicated that the (3,4) species is not present in important amounts in perchlorate solutions. From this, it is reasonable to suppose that this species includes complexed chloride ions. If it is further assumed that the (2,2) and (3,5) species do not complex chloride, and ~ e’3,5 are the same in chloride that the values of E ’ ~ ,and and perchlorate media, the following expressions are obtained which ma:y be solved for the fractions F i , j of is known from measurements the various species (qO on an unhydrolyzed solution)
+ + + +
n Fw (5/3)F3,5 8 = Fi,oci,o F2,ze’z,z -I1 = Fi9o Fz,z F3,6
F3,56’3.5
(5)
Further, since
(7) See, e.& J. Rydberg, .Arkiv K e m i , 8 , 118 (1956). (8) (a) C. F. Baes, Jr., and N. J. Meyers, Inorg. Chem., 1 , 780 (1862); see also (b) K. A. Kraus, “Hydrolytir Behavior of the Heavy Elements,” Proceedings of the International Conference on the Peaceful Uses of Atomic Energp, Vol. 7, p. 245, Session 10B.1, P/731, United Kations (1956).
O P r
O
823
I=**””- *w*-
-06
0.980
n***-.e*
..”*
1.205
” .
3500
/
/
,
I
,
I
/
,
/
4000 4500 W A V E L E N G T H , angstroms.
,
!
I
1
5000
Fig. 3.-Deviations of observed molar absorptivities for U(VI) solutions in 1 M total chloride from values based on kz,z = 6.7 X k3,a = 4.7 X 10+s, and ka,s = 1.0 X lo-’’. Numbers to the right of each plot are the hydroxyl numbers; vertical lines at the left of each plot represent 0.005 in A , symbols represent M U(V1)as: A, 0.1; 0 , 0.01; m, 0.001.
from the optical absorption of a single solution at one wave length, values of ki,j may in principle be obtained, if the acidity of the solution is also measured. I n practice, of course, the solutions must be selected to have a n appreciable amount of all species present. We have carried out such measurements with eight solutions, and the quotients evaluated at 4300 A. are listed in Table I. Agreement between the values obtained from the individual solutions is about as good as could be expected. The acidity measurements on the
RICIUHD31. RUSHBND JAMES 8. JOHNSON
824
Vol. 67
TABLE I EQGTLIRRIUM QCOTIENTS OBTAINED FROM SPECTRA A X D ACIDITY MEASUREMENTS IN 1 M TOTAL PERCHLORATE
80
E
- ' W
GO
>
t 2
k lx 0
a [r a
40
r
w
gcn 2 0 n
I~
A
A
~
~
A
~
A
~
A
A
~
* *
7
*
-
*
...............
-~
V.L
~
1
A ~ ~A A ~~ A~ A ~~ , ~* ~ ~~ A ~. ~~ ~0,420 A~ A ~ ~
e
.
**'-
e
........................ ................... ..... .... . 3 .................
-02-
0.1000 ,0099s ,01007 .Ol005 .000996 .001004 .001002 .001011
0.420 .429 ,623 .806 ,454 .638 ,816 1.003
3.322 3.833 3.996 4.137 4.299 4.434 4.548 4.652
(4300
ke,n
A.)
13.26 14.58 20.95 28.46 17.17 24.00 31.64 39.76
x
10'
1.33 1.12 1.13 1.13 1.06 1.05 1.05 1.34
-
ka,a
x 10'7 4.07 4.06 4.05 3.87 4.28 3.54 3.35 4.16
3.92
'8.m.
,
0.816
m '. .
L
,
I
I
I
I
,
U
~
The assumptions made in interpreting the absorption data were tested by correlating the spectra for all wave lengths in the same manner as with the chloride solu~ X tions, with the constants k ~=, 1.15 and I C ~ = ,~ 3.9 X lo-''. Comparison of the values of d Z , 2 and d 8 , 5 obtained from chloride and perchlorate for all wave lengths is shown in Fig. 4, and the deviations between experimental and computed values of E for perchlorate solutions are shown in Fig. 5. The agreement is coiisistent with the assumption that the (2,2) and (3,s) species do not complex chloride. The values of kz,z and of k3,G are somewhat smaller in chloride than in perchlorate. This direction is exA pected A A A ~if ~UOz+z ~ ~ ! is ~ complexed by chloride, in agreement More quantitawith the conclusions of many p e ~ p l e . ~ tively, if the only important complex of the unhydrolyzed species is UOzC1+, if the values of Gi,j are assumed to be the same in the two media, and if no chloride complexing of the (2,2) or the (3,5) occurs, the ratio of the apparent quotients for the tu70 media should be
$0.638
mm .~ .';.,~a'.m~.
O
_I 0 429
........... S
-0.4
--log [R"I
of acidity measurements alone ( u = 0.0107 in n) 1.22 3.69 XOTE: At 4300 A. e1.0 = 5.233 from unhydrolyzed perchlorate solution and ~ ' 2 . 2 = 23.34, e'3,6 = 77.59 from measurements in chloride solution (Fig. 2).
cn
-02
n
Av. 1.15 Values obtained from least squares fit
-I 0
0 020T
F
I
1
1.003
where kl,o,l = [UOzCl+]/[U02f2] [Cl-1. An estimate of kl,o.l can therefore be made from comparison of the apparent values of ki,j in chloride and perchlorate media. From such a comparison of the quotients for the (2,2) species, kl,o,l = 0.35 is obtained; from the (3,5) species, kl,o,l = 0.52. These may be compared with the value of Nelson and Kraus,lo corrected to M = 1, kl,o,l = 0.50. With the assumptions made, the agreement between these values is good. optical absorption of hydro4. Discussion.-The lyzed U(V1) chloride solutions seems to be consistent with the species postulated, and the formation quotients evaluated, to explain acidity and ultracentrifugation results.2 Analysis of the chloride and perchlorate spectra support the conclusion based on acidity data that the (3,4) species, which is important in chloride, is not very important, if present a t all, in perchlorate. The values of the equilibrium quotients derived for perchloratmesolutions are reasonable. They (9) .J. Bjerrum, G. Schwarzenbach, and L. G. SillBn, "Stability Constants." Part 11, The Chemical Sooiety. London, 19.58. (10) F. Nelson and K, A. Hraus, J. Am. Chem. Xoc,, 73, 2157 (1951).
ApriI, 1963
IONIZATION O F a - P A R T I C L E S I S
are somewhat greater than those estimated from 20” acidity data; however, this is the direction expected, and of the approximate magnitude predicted, from studies of the temperature coefficient of U(V1) hydrolysis in nitrate medium.*a The question of the importance of other species which may be present in minor amounts is no more resolved by the present results than by the earlier study. For example, it was apparent that about as satisfactory an interpretation could be obtained with a scheme including UOzOH+, although there would be considerable uncertainty in the values of el,l. Similarly, our conclusions for the perchlorate solutions agree in general with those of SuttonJ6though he includes a small contribution of the (3,4) species. The most important point of disagreement with the interpretation of acidity measurements on a core-link
BINARY GAS3 h X T U R E S
825
modelll seems to be resolved, since the St>ockholm group now postulate an important contribution by the (3,5) species, which is not of the core-link type.l2 They still postulate the presence of several higher corelink species, for which we do not find any evidence in our results, but since the constants quoted for these species are rather low, the remaining disagreement is perhaps more in a conceptual model of hydrolysis rather than in a practical description of U(V1) solution chemistry. From a recent publication,13 it appears that even this disagreement may no longer exist. Acknowledgment.-We wish to express indebtedness to Kurt A. Kraus for many helpful and stimulating discussions and to Neva Harrison for technical assistance. (11) S. Ahrland, S. Hietenan, and L. G. SillBn, Acta Chem. Scand., 8 , 1907 (1954). (12) L. G. Sillh, private communication. (13) L. G. SillBn, Acta Chem. Scand., 16, 1051 (1962).
IONIZATIOK BY ALPHA PARTICLES IS BINARY GAS MIXTURES BY T. D. STRICKLER~ Health Physics Division, Oak Ridge National Laboratory,, Oak Ridge, Tennessee, and Department of Physics, Berea College, Berea, Kentucky Receiiied September 31, 196W The W value (average energy loss per ion pair) for a-particles has been measured in a number of binary gas mixtures of molecular gases as a function of the fract’ionalpressures. The W of the mixture ( Wij) can be represented in terms of the 14”s of the pure constituent gases Wi, Wj) and the fractional pressures ( P i , Pj) by the relation: Wij = (Wi -- Wj)Zij” Wj, where Zij” = Pi/(Pi f i j Pj),in whichfij is a constant determined empirically for each pair of gases. These constants verynearly satisfy the relationshipfij = fj/fi = (fj/fk)/(fi/jk) = jkj/jki,where i, j , and k refer to any three gases. Thus, if the followingj-values are assigned to the gases in this study (Nz, 1; COz,1.8; H,, 0.5; 02, 1.3; CH,, 1.8; CeH4, 3.4; GH6, 3.5; CJ&, 4.5; C,Hs, 4.5; C9H8,6.3) then the constant f i j determined from the ratio of any two of these will serve t o predict the W of any mixture of these two gases with an accuracy better than 1%. Slight departures from the W predicted by the above equation have been noted in the case of nitrogen mixtures, indicative of an effect similar to t’hat observed in the noble gases when small amounts of impurities are added.
+
+
Introduction
If an a-particle of kinetic energy Eo is completely stopped in a gas, a,nd in becoming stopped produces Ni ion pairs, then the mean energy lost per ion pair E,/Ni is commonly called the W of the gas for aparticles. The W of most gases lies in the range 20 to 46 e.v. per ion pair and in many cases is found to be practically independent of the energy of the initial ionizing particle. The practical importance of W (in the measurement of radiation dose, in the calculation of energies of particles in nuclear reactions, and in the interpretation of radiation-induced chemical reactions) and its theoretical significance have been pointed out in many recent publications on the ~ u b j e c t . ~ I n the case of binary mixtures of gases, two distinct phenomena have been observed. One is the marked increase in ionization, and consequent decrease in W, when small amounts of some gases are mixed with the noble gases. This has been studied by Jesse and Sadauski^^-^ using helium and neon, and by Melton, Hurst, (1) Dent. of Physics, Berea College, Berea, Kentucky. (2) Operated by Union Carbide Corporation for the U. S. Atomic Energy Commission. (3) (a) S. C. Curran and J. M. Valentine, Rept. Progr. Phys., 21, 1 (1958); (b) W. Binks, Acta Radiol., Splppl., 117, 85 (1954); (c) R. L. Platzman, NAB-National Researoh Council Publication 752, 109 (1960). (4) W. P. Jesse and J. Sadauskis, Phys. Rev., 88, 417 (1952). (5) W. P. Jesse and J. Sadauskis, kbid., 90, 1120 (1953). (6) W. P. Jesse and J. Saudauskis, ibid., 100, 1755 (1955).
and Bortner’ using argon. The effect is attributed, in part, to the excitation of the metastable level in the noble gases and the subsequent ionization of the impurity (by interaction with the excited atom), provided the ionization potential of the impurity is lower than that of the metastable state. This has been referred to as the “Jesse effect.”8 The fact that increased ionization occurs in argon, even when the impurity has an ionization potential greater than that of the metastable level, has been demonstrated by Melton, Hurst, and Bortner,’ but the explanation of the effect is not entirely clear. On the other hand, in the molecular gases, the W of the mixture is found to lie between the extreme values for the pure gases and to change smoothly from one limit to the other as the composition of the mixture is changed. This has been studied by Huber, et U Z . , ~ ~ ~ ~ and by Hurst, et uZ.11--13 It has been shown that the W of any mixture of two of these gases can be expressed (7) C. E. Melton, G. S. Hurst, and T. E. Bortner, ibid.,96, 643 (1954). ( 8 ) R. L. Platzrnan, “The Physical and Chemical Basis of Mechanisnis in Radiation Biology,” “Radiation Biology and Medicine,” W. D. Claus, Ed., Addison-Wesley Publishing Co., Inc., Reading, Mass., 1958, pp. 15-72. (9) P. Huber, E. Baldinger, and W. Haeberli, Helu. Phyg. Acta, 23, Suppl. 111 (1949). (10) W. Haeberli, P. Huber, and E. Baldinger, ibid., 26, 145 (1963). (11) T. E. Bortner and G. S. Hurst, Phys. Rev.,93, 1236 (1954). (12) H. J. Moe, T. E. Bortner, and G. S. Hurst, J. Phya. Chem., 61, 422 (1957). (13) G. S. Hurst and T. D. Strickler, NAS-National Researoh Council Publication 752, 134 (1960).
