August, 1958
KINETICB OF
THE
REACTION BETWEEN Pu(V1)
AND
U(1V)
943
THE KINETICS OF THE REACTION BETWEEN PLUTONIUM(V1) AND URANIUM(IV)l BYT. W. NEWTON University of California, Los Alamos Scientific Laboratory, Los Alamos, New Mexico Received February 1% 1968
The kinetics of the reaction between Pu(V1) and U(1V) have been studied in perchlorate media. The hydrogen ion dependence indicates that the reacting system passes consecutively through two activated complexes with nearly the same A F * values. The formulas and heats, free energies and entropies of formation of these complexes from the reactants were found to be: (HzO~U~OH.OPuO) +s, A H * = 17.6 kcal./mole, A F * = 16.6 kcal./mole and A S * = 3.4 e.u., and for (HO.U.OH.0PuO).+~,A H * = 21.4 kcal./mole, A F * = 16.0 kcal./mole and A S * = 18.1 e.u. The relation between the kinetics of this reaction and the similar reactions between Np(1V) and Np(V1) and between Pu(1V) and Pu(V1) are discussed. distillation was from alkaline KMnO4 in a Pyrex still. Introduction Uranium solutions for the individual runs were made by Although plutonium and uranium are associated diluting the stock solution and were analyzed for U(1V) a t in nuclear reactors and U(1V) is a good reducing the time of the rate runs by making use of its absorption at agent, only recently have data on the reduction of 6478 A. At this wave length the molar absorptivity of in 1M HC104was found to be 58.9 M-l cm.-l. plutonium by U(1V) been published. RydbergZ U(1V) Procedure.-Appropriate amounts of the various reagents has shown that U(1V) can be a useful reducing were pipetted into the compartments of a two chambered agent in connection with extraction processes. The absorption cell. This cell had a 10-cm. light path and was present study was undertaken to provide quantita- arranged so that in a vertical position the two portions of tive data and to allow coniparison with the formally liquid could be kept separate during temperature equilibrain a thermostat. After mixing the two portions, the analogous reactions between Np(1V) and N P ( V I ) , ~ tion cell was placed in the cell holder in the spectrophotometer. between Pu(1V) and Pu(VI), and between U(1V) This holder was filled with water and was thermostated by and U(V1). The kinetics of these last two reac- water pumped from the thermostat. With this arrangemenot tions must be inferred from the equilibrium quo- it was possible to maintain temperatures to within 0.2 . A Cary Recording Spectrophotometer, Model 14, No. 5 was tients and the kinetics of the reverse reactions, the used. The concentration of Pu(V1) was followed as a funcdisproportionation of Pu(V)*and of U(V).5 tion of time by measuring the absorption at 8302 A. Since
Experimental Part Reagents.-Pu( VI) stock solutions were prepared by dissolving weighed amounts of pure plutonium metal in standardized concentrated HC104. The resulting solutions were diluted and then oxidized by passing ozone for several hours. The oxidized solutions were diluted in volumetric flasks and weighed. Solutions for the individual runs were made up from weighed amounts of stock solution and appropriate amounts of HC10, and were reoxidized before use. Excess ozone was swept out with a stream of nitrogen. The U( IV) stock solution was prepared by the electrolytic reduction of a u O ~ ( c 1 0 4 solution )~ using a mercury cathode. The UOz( ClO4)zsolution was made by dissolving pure UaOs in hot concentrated HC104. The U(1V) solution was assayed by reduction with zinc amalgam, air oxidation and titration with standard ceric sulfate solution. The hydrogen ion concentration .was determined by titration with NaOH after removing the uranium with an ion-exchange column.6 The HClO4 solutions were prepared by diluting the concentrated acid and were standardized by titration. The concentrated acid was analytical reagent grade and was further purified by boiling a t atmospheric pressure and then again under reduced pressure. Solutions of LiClO4 and NaC104 were prepared by neutralizing the appropriate carbonate with HClO4, boiling out the COz and recrystallizing from water. The NaC104 solutions were standardized by drying and heating to constant weight a t 150”, and the LiC104 solutions were standardized by titrating the effluent from a cation-exchange column. The water used in the preparation of all solutions was doubly distilled; the second (1) This work was done under the auspices of the U. 9. Atomic Energy Commission. (2) J. Rydberg. J . Inorg. Nucl. Chem., 6 , 79 (1957). I n 1943 and 1944 some fragmentary results were given in various Manhattan Project Reports by H. W. Crandall, b y J. Halperin and W. J. Knox and by G. E. Moore and L. B. Werner. (3) (a) J. C. Hindman, J. C. Sullivan and D. Cohen, J . Am. Chem. Soc., 76, 3278 (1954);(b) J. C. Sullivan, D. Cohen and J. C. Hindman, ibid., 79,4029 (1957); ( c ) J. C. Hindman, personal communication. (4) S. W. Rabideau, ibid., 79,6350 (1957). (5) D. M. H. Kern and E. F. Orlemann, ibid., 71, 2102 (1949); H. G. Heal and J. G. N. Thomas, Trans. Faraday Soc., 46, 1 1 (1949). (6) The author is grateful to F. B . Baker for preparing and standardizing the U(1V) stock solution.
this absorption peak is very narrow it was found convenient to extend the suggestion of Moore and Kraus’ and determine the concentrations from the quantity: Assor - 0 . 5 (AS400 Aszoo), where A is the absorbance. When this is done, explicit correction for change in absorbance due to the other species need not be made. Catalytic Impurities.-Evidence was found for a catalytic impurity in the LiC104. When the salt was prepared by neutralizing analytical reagent grade LizCOa with HClO4 and recrystallizing from water, the rate constants obtained in solutions made from the crystals differed significantly from the constants obtained in solutions using the mother liquor. It was found that the same rate constant was obtained using LiC104 which had been recrystallized four times and that which had been recrystallized six times A special batch of LiC104wasprepared using a different source of lithium and a radically different purification procedure. Lithium metal was dissolved in water to give a LiOH solution. This was electrolyzed into a mercury cathode to give an amalgam which in turn was treated with HC104 to give LiC104.* This LiC104 was recrystallized twice from water to remove excess acid. This material was used in rate runs and the rate constants were identical with the ones obtained using the material which had been recrystallized six times. This agreement makes it very improbable that catalytic impurities are left in the purified LiC104. Recrystallization of the NaC104 had a negligible effect on the rate constants. Experiments have shown that chloride concentrations as high as low3M have no effect on the rate and that U(1V) prepared by dissolving ucl4 in HClO4 gives the same rate as the U(1V) which was prepared by reduction of U(V1).
Results and Discussion Stoichiometry.-The oxidation potentials involved indicate that U(1V) is capable of reducing Pu(V1) all the way to Pu(II1) and Rydberg2 has shown that with 0.05 M U(1V) the rates are such (7) G. E. Moore and K. A. Kraus, paper 4.22,“The Transuranium Elements,” edited by G. T. Seaborg, J. J. Katz and W. M. Manning, “National Nuclear Energy Series, Division IV,” Vol. 14B, McGrawHill Book Co., Inc.. New York. N. Y.,1949, p. 608. (8) The author acknowledge8 helpful discussions with Professor H. L. Friedman concerning this purification procedure.
T. W. NEWTON
944
that the reduction to Pu(II1) actually occurs. However, in an experiment preliminary to the present work, it was shown that when the Pu(V1) and the U(1V) were both initially 0.002 M in a 1 M HClO4 solution, Pu(V) was the principal plutonium oxidation state produced, along with small amounts of Pu(II1) and no detectable Pu(1V). It was further shown that if the Pu(V1) was initially 1.8 X M and the U(1V) was initially 0.62 X M, the amount of Pu(V1) which had disappeared when the reaction was over was 96Oj, of what would be required by the reaction 2Pu(VI)
+ U(1V) = 2Pu(V) + U(V1)
Pu(V)
+ U(1V) = Pu(II1) + U(V1)
TABLE I TYPICAL RATERUNS IN LiC104-HC10, SOLUTIONS AT 25' ANDu = 2 Pu(VI), 1.857 X lO-4M 1.803 X 10-4 M U(IV), 1.564 XlO-'M 0.902 X M H+, 0.60 A4 0.15 M Time, min.
1 2 3 4 5 G 7 8 9 10 12
(1)
This result shows that even when z/3 of the Pu(V1) is reduced, an insignificant amount of U(1V) reacts in some other way. I n treating the kinetic data, most of which were taken at less than z/3 reduction, these other reactions have been ignored. Since U(V) is probably an intermediate in reaction 1 the possibility that it reacts with traces of O2 in solution was examined. It was found that the same amounts of Pu(V1) were consumed in solutions which had been swept with O2 as in solutions which had been swept with Nz. This result indicates that the reaction of U(V) with O2 is too slow t o affect the stoichiometry and that the discrepancy mentioned above is probably due only to a reaction such as (2)
The Pu(V1) and U(1V) Dependences.-Reaction 1 was found to be first order in U(1V) and in Pu(VI) by making use of the usual plots of the integrated second-order rate equation. Thus for all the rate runs in which the concentration of U (IV) was significantly larger than half the concentration of the Pu(VI), plots of log[U(IV)] - log[Pu(VI)] versus time gave good straight lines. I n those runs in which the concentration of U(1V) was very nearly equal to half the concentration of Pu(V1) plots of the reciprocal of [U(IV)] l/z[Pu(VI)] (the effective average concentrationg) versus time gave good straight lines also. The slopes of these lines are equal to {2[U(IV)]o - [Pu(VI)]o}k'/2 loge 10 and k'/2, respectively, where [U(IV)]o and [Pu(VI)]o are the initial concentrations and k' is the second-order rate constant defined by the equation : k' = -d[Pu(VI)]/dt[Pu(VI)]-' [U(IV)]-'. Values for IC' were computed using these relations; the slopes were determined by least squares. Table I shows absorbance data from typical runs compared with values calculated from the best straight lines. The good agreement between observed and calculated values indicates that the reaction is first order in each of the reactants, Pu(V1) and U(1V). Further confirmation of the order with respect to U(1V) was obtained when it was found that the apparent second-order rate constant was the same within the experimental error when the initial concentration of U(1V) was varied by a factor of four. Hydrogen Ion and Temperature Dependences.A number of runs were made at different hydrogen ion concentrations a t 25.0' and an ionic strength of
+
(9) A. A. Frost and R. 0.Pearson, "Kinetics and Mechanism," John Wiley and S o w , Inc., New York. N. Y . , 1953, pp. 17-19.
Vol. 62
A (obsd.)
A (calcd. by 1)"
0.894 ,811 .741 .680 .621 .573 ,531 .494
0.895 ,811 ,740 ,677 .622 .574 .531 ,492
...
...
A (obsd.)
A (talc:. by 11)
0.757 ,625 .533 ,464 .412 ,370 ,339 .308 ,286 .265
0.756 .62G .534 .466 .413 ,371 ,337 .308 ,284 .264
.425 .426 .371 ,371 ... ,.. 14 .326 .326 ... ... I , calculated from the best straight line through log 2[U(IV)] - log [Pu(VI)] uersus t . 11, calculated from the best straight line through [Pu(VI)]-1 versus t .
two, made up with LiC104. Runs a t various hydrogen ion concentrations also were made at three temperatures other than 25". The results of these runs are given in Table 11. Mechanism.-The fact that the reaction is first order in Pu(V1) and in U(1V) at any particular hydrogen ion concentration implies that the rate determining step, without regard t o H + or HzO,is
+ U(1V) = PU(V) + U(V) Pu(V1) + U(V) = PU(V) + U(V1) Pu(V1)
(3)
This is followed by one of the fast reactions or U(V)
+ U(V) = U(1V) + U(V1)
(4) (5)
Reactions 4 and 5 cannot both be important, for if they were, the over-all rate would not be first order in each of the reactants. This is because Pu(V1) is reduced in reaction 4 but not in reaction 5.1° It seems quite likely that (5) is unimportant with respect t o (4)since it involves the breaking of a U-0 bond and (4)does not. The effect of H + on the rate is much more complicated. Since the rate increases with decreasing concentration, an inverse power is suggested. I n addition, significant hydrolysis of the U(1V) occurs in the hydrogen ion range investigated since the equilibrium hydrolysis quotient K is reported l1 to be 0.0236 at 25" and p = 2. If the activated complex is formed from the principal reactant species with the loss of one H+ the rate would be given by lc[PuOz+] [U+4] [€I+]-' and the quantity k ' . ([Hf] )'1 should be a constant a t any one tem-
+
(IO) If reactions 4 and 5 were of comparable importance, the rate would be given b y the expression
-d[Pu(VI)]/dt
=
ks[Pu(VI)][U(IV)]
+
which is not first power in each of the reactants. This expression was derived making the steady state assumption d[U(V)]/dt = 0 and letting ks, kr and ks stand for the second-order rate constants for reactions 3, 4 and 5. (11) K. A. Kraus and F. Nelson, J . Am. Chem. Soc., 71, 3901 (1950).
KINETICS OF THE REACTION BETWEEN Pu(V1)
August, 1958
AND
U(1V)
945
TABLE I1 SUMMARY OF APPARENT SECOND-ORDER RATECONSTANTS IN LiC10rHC104 SOLUTIONS AT 25'
2
AND p =
The units are M-1 mix-1. Hydrogen ion concn., M
0.10 .15 .20 .25 .30 .40 .50 . 00 .70 .80 1.00 1.25 1.50 Av. difference (obsd. Max. difference(obsd
- calcd.), %'
- calcd.), %
25.0'
Obsd.
Calcd.'
4161 2915 2214 1770 1446 1063 838 G77
4117 2876 2191 1758 1460 1076 842 684
..
..
481 368
486 369
..
..
219
217 0.9 1.4
perature. At 25" a plot of this quantity versus [H+]-' was not a horizontal line but showed distinct curvature between a value of 514 a t 0.1 M and 334 a t 1.5 M (see Fig. 1). It has been pointed out that a plot of log IC'. ([H+] K ) versus [H+] gives a satisfactory straight line suggesting that the lack of constancy is due t o a medium effect.12 This interpretatiori requires that the rate constant be about 520 min.-' in 2 M LiC104 but only about 280 min.-l in 2 dd HC104. The Bronsted principle of specific interaction of ions13would predict a medium effect of zero since the reaction involves only positive ions, and in addition an effect of this magnitude is quite unlikely in view of the results of the studies of the kinetics of the reactions between Np(1V) and Fe(III)14 and between Pu(1II) and Pu(VI).15 For these reactions only single reaction paths were indicated and a t a total ionic strength of unity the rate constants were found t o be independent of the relative amounts of HClO, and NaC104to within the experimental error. Thus it has been assumed that for the present reaction the lack of constancy in the quantity k' ([H+] K ) is due t o the operation of a more complicated mechanism rather than t o a medium effect. A mechanism involving additional parallel paths with different inverse hydrogen ion dependences would require that the plot have a curvature opposite to that actually observed. It was therefore concluded that a different sort of mechanism is required. It was suggestedle that a mechanism involving an inhibiting back reaction favored by H + might be satisfactory. A somewhat general mechanism for this has been formulated as
34.50
Obsd.
Calod.
..
..
..
..
3540
3499
.. ..
.. ..
..
..
2055 1458
15.5'
1676
Calod.
Obsd.
1639
..
.. .. .. ..
..
... ... ... ...
392
81.4
..
..
388
..
..
..
... ...
1451
201
202
..
..
961 736 581
962 733 582
,
.
..
126 94.8 74.6
0.7 1.8
2.6'
382
..
2092
..
Obsd.
Calcd.
376
... ... ... ...
83.0
*..
...
41.4
...
41.4
*. 128 94.9 73.6
26.0 18.6 14.5
25.8 18.8 14.4
1.1 2.3
...
0.9 1.9
+
+
(12) The author wishes to thank the referee for pointing this out. (13) J. N. Bronsted, J . A m . Chem. Soc., 44, 877 (1922). See also
R. A. Robinson and R. H. Stokes, "Electrolyte Solutions." Butterworths Scientific Publications, London, 1955, p. 425. (14) J. R. Huisenga and L. B. Magnuason. J. A m . Chem. Sac.. 7 3 , 3202 (1951). (15) S. W. Rabideau and R. J. Kline, T H i s JOURNAL,62, G17 (1958). (16) The author is very grateful to Professor Henry Taube for this
suggestion.
300I
+ I ' u O ~ ++~ 2Hz0
I
I
I
I
I
+ +
U+4
i=Int. ? (p q)H+ forward rate = k ~ [ U + 4 ] [ P ~ O ~ + e l [ H +(6) l-p reverse rate = k-~[Int.][H+]g (7) Int. = PuOZ+ ( T s)H+ rate = kz[Int.][H+]-' (8) PuO2'a uoz+ = PuOz+ UOZ+Z fast, rate limited by second step (9)
+ +
+
+
I n this scheme the symbol [Int.] stands for a metastable intermediate and p , q, r and s stand for integers, the sum of which equals four. Applying the usual steady-state approximation t o [Int. 1, defining IC' as above, and correcting for the hydrolysis of U+4: [U+4] = [U(IV)l[H+]([H+] K)-l, gives [k'([H+] IC)]-' = (2kl)-'[H+]P-' +
+
+
k-i(2kikz)-' [H +]P+Q
fr-1
(10)
As might have been expected, the form of the rat,e law gives the sum r+q rather than their individual values. This rate law was tested by plotting [k'([H+] K)]-'versus [ H f ] ; a t each temperature good straight lines resulted indicating that p = 1 and q 3- r = 1. These values of p and 9+r require that (K20.U.0H.PuOz)+6and (HO.U.OH.PUO~)+~ be the formulas (but not necessarily the structures) of the two activated complexes through which the reacting system passes, Appropriate values for K , the hydrolysis quotient of U(IV), were obtained
+
T. W. NEWTON
946
Vol. 62
by extrapolating the data of Kraus and Nel~on.~J’ the system cross the two barriers consecutively Best values for 2kl and kllcz/k-l were obtained from instead of by parallel paths. the straight lines by the method of least squares. Rates were calculated from the three rate laws The values found for these quantities a t the various and are compared with the experimental ones in temperatures are shown in Table 111and were used Table IV. It is seen that either inhibition mecht o obtain the calculated rate constants shown in anism fits the data somewhat better than the one Table 11. The good agreement between the cal- involving parallel paths. More precise data are culated and observed rate constants confirms that required before the two inhibition mechanisms can be distinguished. the rate law is satisfactory. The disproportionation reactions of Pu(V) and TABLE I11 U(V)( are both reported t o be first power in H+, RATECONSTANTS AT VARIOUS TEMPERATURES IN SOLUTIONSthis requires that the formulas of the activated OF LiC104 WITH p = 2 Complexes be Temp., OC.
UW) hydrolysis quotient, M
2.6 15.5 25.0 34.4
0.00508 ,0127 .0236 ,0420
Zkr, min.-l Calcd.
Obsd.
kika/k-i, M min.-1 Obsd. Calcd.
42.0 33.6 33.1 42.0 193.8 188.6 196.1 198.3 529.2 520.7 659.1 675.6 2155 2107 1305 1336
The activated complex in the reaction between P U O ~ and + ~ U 0 2 + (step 3, above) need not have definite linkage between the two reactants, all that is required is that a configuration be reached which allows electron transfer. However, the activated complexes in the first and second steps of the mechanism are more definite than this in that they involve the formation or breaking of a definite chemical bond which is present in the intermediate. Since the reaction Np(IV) Np(V1) = 2Np(V) is formally the same as reaction 3, it is of interest to compare the two. Hindman, Sullivan and Cohen have studied the reaction and suggest3b that it involves two parallel paths and has the rate law k’ = -d [Np(IV)]/dt [Np(IV) 1-1 [Np(VI)I-1 =
+
(HO.M.OaM02)+s
where M stands for either Pu or U. This activated complex is formally the same as the one in the r+q = 1 mechanism discussed above for the N p reaction. TABLE IV APPARENTSECOND-ORDER RATE CONSTANTS k’ FOR THE Np(1V)-Np(V1) REACTION AS A FUNCTION OF H + AT 25” ANDp = 2 Observed rates are from reference 3a. Hydrogen ion concn., M
1.994 1.046 0.532 0.227
Calcd. according to:= k’ (obsd.), I 111 M-1 min. 1 M-1 min.-1 M - l E i n . - l M-1 min.-l
0.581 0.661 0.580 0.590 2.60 2.43 2.61 2.50 10.47 9.64 10.51 10.40 54.5 56.6 54.1 56.1 I, k’ = 2.59[H+]-* 0.074[H+]-a; 11, [k’([H+] K ) ] - l = 0.3045[Hf] 0.0303[H+J3;111, [k‘([H+] K ) ] 4 = 0.2793[H+] O.0696[H+I2.
+ + +
++
It appears likely that in all four reactions: U( 1V)-U (VI), Np (IV)-Np (VI), PU(IV)-PU (VI) It has now been found that the inhibiting back re- and U(1V)-Pu(VI), the reacting system consecuaction mechanism given above can account for the tively crosses barriers corresponding to activated observed rates very satisfactorily. For the pur- complexes with the formulas: (H20.M.0H.M02)+6, pose of this calculation the hydrolysis quotient of (H0.M-OH.M02)f 4 and (H0.M.0.M02)+3. It is N P ( I V ) ~has * been taken as 0.039, the average of not clear, however, why a different complex corthe corresponding values for U(IV)ll and Pu(IV).lg responds t o the highest barrier in the different reThe best fit of the data is provided by the rate law actions. Thermodynamic Quantities of Activation.-In [k’( [H+] + K)]-l = 0.3045[H+] + 0.0303[H+]* (12) the reaction scheme presented above for the U(1V)where k’ is defined in rate law (11). This rate law Pu(V1) reaction, the second activated complex is corresponds to p = 2 and q r = 2 in the mech- formed from the metastable intermediate. Now anism with kl = 3.28 M min.-’ and kzlk-1 = 10.05 consider formally its formation from the original M 2 . I n addition another rate law was found which reactants, and write the net process as fits the data satisfactorily. This is (HO.U.OH.PUO~)+~2 H + 2.59[H+]-2
+ 0.074[H+]-a
(11)
+
[k’( [H+]
+ K)]-1 = 0.2793[H+] + 0.06961[H+]a
Uf4
(13)
The coefficients in this rate law give kl = 3.58 M rnin.-’ and k2/k-l = 4.01 M . The powers of [ H + ] correspond to p = 2 and q +r = 1 and thus the formulas of the activated complexes are (HOaNp. OH.Np02)+4 and (HO.Np.O.NpOz)+3. This mechanism involves activated complexes with the same formulas as those in the mechanism suggested by Hindman, et al. The essential difference is that the inhibiting back reaction mechanism requires that (17) K. A. Kraus and F. Nelson, J . Am. Chem. Sac., 77, 3721 (1955). (18) R. Sjoblom and J. C. Hindman, ibid., 73, 1744 (1951). show that Np(1V) is significantly hydrolyzed, hut no value of the hydrolysis quotient waB given. (19) S. W. Rabideau, ibid., 79, 3675 (1957).
+ PuOztZ + 2Hz0
+
The AF*for thisprocessisAF1* - AF-l* 4-AFz”. Writing IC = ( k T / h ) exp (-AF*/RT)*O we see that klk2/k-l = (kT/h) exp[-(AF*l) - AF-l* AF2*)/RT] and that the hypothetical rate constant for the process is klkz!k-l.. Thus the thermodynamic quantities of activation can be obtained from this quotient in the usual way. Log kl and log (klkz/k-l) were plotted us. 1/T and good straight lines resulted. The method of least squares was used to determine best values for the intercepts and best values and uncertainties for the slopes. These values were used to obtain the calculated rate constants given in Table 111.
+
(20) 8. Glasstone, K. Laidler and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., New York, 1941, pp. 195-199.
’
August, 1958
BONDDISSOCIATION ENERGIES OF POLYATOMIC MOLECULES
947
TABLE V THERMODYNAMIC QUANTITIES OF ACTIVATION AT 25" N e t process
U+4 U+4 Np+4
+ PuOz+* + 2H20 F? (HzO.U.OH.PuOz)+5 + H+ + PuOz+* + 2H20 e (HO.U.OH.PUOZ)+~ + 2 H + + NpOz+Z + 2H10 (HO.NP.OH.NPO~)+~+ 2H+
+
+ 2Hz0 i=? (RO.Pu.O.PuOz)+3 Activation energies of 18.2 i 0.4 and 21.9 * 0.3 P u + ~ PuO~+'
kcal. were calculated from the slopes. The uncertainties given are twice the standard deviations. The free energies, heats and entropies of activation were calculated according to the transition state theory.20 These values are shown in Table V. For comparison, corresponding values for the N p (1V)-Np (VI) reaction and the Pu (1V)-Pu (VI) reaction are given also. The Np values were recomputed from the original authors' data3&and a more recent value for the activation energy.3c The Pu values were computed from Rabideau's data on the disproportionation reaction4 and values for the over-all reaction, determined from published thermodynamic data.21 It isseen that the A S * values increase in a uniform way as the charge on the activated complexes decreases. The agreement between the A S * values for the two processes leading to complexes with similar formulas, (HO-U.0H-Pu02)+4 and (HO-Np-OH. Np0.J +4, indicates that their structures are probably similar also. This agreement also supports the suggested mechanism for the U(1V)Pu(V1) reaction. Ionic Strength and Medium Effects.-The main (21) S. W. Rabideau and H. D. Cowan, J. A n . Chsm. SOC.,77, 6145 (1955); 8. W. Rabideau, i b i d . , 78, 2705 (1956).
+3H+
AF*, kcal.
A H * , kcal.
AS*, e.u.
16.6
17.6 f 0 . 4 21.4f .3 24.7 f . 3 37.8 f 2 . 7
3 . 4 zk 1 . 5 18.1 f 1 . 0 18.4 f 1 . 0 37.6 f 9
16.0
19.2 26.6
series of experiments was done a t ionic strength two in LiC104 solutions. In order to estimate the ionic strength effect, a short series of experiments was done with p = 1 in LiC104 solutions. Toestimate the possible magnitude of the medium effect another series was done in NaC104 solutions with p = 2. Values for the rate constants in these solutions are given in Table VI, together with corresponding values from the main series for comparison. TABLE VI EFFECT OF IONIC STRENGTH AND MEDIUMAT 25" Solution LiClOd. p = 2 LiClO,, p = 1 NaC104, p = 2 2kl, min. -1 529 364 473 klhlk-1, M mim-1 659 562 785
Acknowledgments.-The author wishes to acknowledge helpful discussions with Dr. C. E. Holley, Jr., and especially with Dr. J. F. Lemons, under whose general direction this work was done. ADDEDIN PROOF.-After this article was submitted for publication Dr. J. C. Hindman sent the author some recent data obtained in his laboratory on the Np(1V)-Np(V1) reaction. Unlike the data in Table IV, these new results are more satisfactorily explained on the basis of a mechanism involving parallel paths, although it is not inconsistent with the assumption of consecutive reactions.
CORRELATION OF BOND DISSOCIATION ENERGIES OF POLYATOMIC MOLECULES USING PAULING'S ELECTRONEGATIVITY CONCEPT1 BY MORTON A. FINEMAN Department of Chemistry, Providence College, Providence 8, Rhode Island Received March 8, 1968
Bond dissociation energies of a set of compounds R-Y and a set of compounds R-R are correlated by the extension of Pauling's equation relating the extra ionic resonance energy to the electronegativities of atoms. Requirements necessary for this correlation are that (I) all the radicals R employ the same bonding atom and (2) the functional group Y be the same '/2D(R-R). Further extension for the entire set, R-Y. The linear relationship proposed takes the form D(R-Y) = C of this relationship makes it possible to correlate the bond dissociation energies of two sets of compounds R-Y and R-Z. The reasonably good fit of the experimental data to the proposed equations implies that a group of radicals using the same bonding atom has the same electronegativity. Furthermore, this electronegativity was found to be equal to that of the bonding atom as given by Pauling. These conclusions, as well as conflicting conclusions by other authors, are discussed.
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Introduction A knowledge of bond dissociation energies of chemical bonds in molecules is essential for understanding chemical processes. At the outset it should be emphasized that there is a distinction between a bond dissociation energy and an average bond energy. The distinction as defined by Szwarcz will be used in this article. The bond dissociation energy of a bond in a molecule R-Y is the energy (1) This work was supported in part by the U. S. Atomic Energy Commission under Contract AT(30-11-1670. (2) M. Szwarc, Chem. Reus., 47, 75 (1950).
needed to separate the radicals R and Y to infinity, each species being in its ground state. The average bond energy of a bond A-B is defined as l l n t h the energy needed to separate each of the atoms in a symmetrical molecule AB, to infinity, all species being in their ground states, that is, l/nth of the heat of atomization of the molecule. I n general, the bond dissociation energy plays a more important role in chemistry than the average bond energy. Both Steacie3 and Szwarc2 have (3) E. W. R. Steacie, "Atomic and Free Radical Reactions," Vol. I, Reinhold Publ. Corp., New York, N. Y . , 1954.
