Oct., 1961
1897
hTOTFS
Acknowledgment.-This work is being supported by the American Chemicd Socicty Yetroleum Ikwarch Fund.
TIIE REACTION BETWEEN URANIUM (IV) AND HYDROGEN PEROXIDE'
decomposition over the whole range of reactant concentrations studied. Similar determinations in 0.5 M HCI04 or in solutions containing up to 0.006 M Na2S04,0.006 hl TICl, 0.008 M Co(I1) or 0.009 U(VI), or a t a temperature of 2.4' gave consumption ratios in the range of those listed ilk Table I.
BY F. B. BAKERA N D T. W. NEWTON
TABLEI
Uniterdtv of Califorriia Loa Alamos Scientifi Laboratory. Loa Alamor, New Mezico Received Ala7ch 20 1961
The purpose of this note is to describe some experiments which were done to help understand rhe reaction betneen U(ITr) arid Hz02in aqueous HClOd solutions. The pertinent oxidation potentials are such that the net reaction to be expected is tJ+4
+ HtOs = UOL" + 2II+
(1)
The results to be given here show that the reaction does not proceed directly by this simple reaction but that it is in part, a t least, a chain reaction and that it is accompanied by a small amount of HzOs decomposition. Experimental Solutions of U(VI), IJ(IV), HClO,, NaCIO, and LiCI04 were prepared as before.? Other U(1V) solutions were made by the reduction of U(V1) with metallic Zn or Pb. U(1V) made in all three wayb showed essentially the samc rate behavior. The H?02 was from two murces, 3OW0 solution from the Mallinckrodt Chemical Co. and 90Yo sollition from the Becco Chemical Division of the Food Machinerv and Chemical Corp. Both of these solutions were furthrr purified by distillation under reducrd pressure and gave iitially the samc kinetic results. The second fraction from the 30% solution had a disagr~eahieodor and gave slightly higher rates; so subsequent runs were made with diluted 90% material. Cu(C10,)2 was made by fuming Cu(YOq)awith HClO, and then crystallizing from concentrated HClO,. The other perchlorate salts were from the G . F. Smith Chemical Go and were used without further purification, except for Co( C104)2which was recrystallized. The U(1V) solutions wrre analyzed by titration using standard Ce(1V) sulfate. H20t was analyzed as an oxidising agent against standard Fe(I1) and as a reducing agent against standard C'e(1V) with the same results. Mixtures of U(1V) and HZ02 were itnalyzed for the difference in their concentrations by adding aliquots to eycess standard Fe(I1) and back titrating with standard Ce(1V). Thc rate rnns were madr spertrophotometrically by following the absorhsncr at 6475 A. where C(1V) absorbs relatively strongly The apparatus and general procedure have been dwcribed prPviorislv.2
MOLESOF HtO, CONSUMEDFOR EACRMOLE OF U(IV) REACTED (CONSUMPTION RATIO) Conditions: 2M HC104,23-25', solutions deaeratrd with A r -Initial
P$'l'd; 3.04 3.08 3.03 1.54 2.76 1.35 2.69 1.35 1.48 1.48 1.48 1.48
values---
[HtOzl/[U(IV)]
0.54
.55 .55 1.1 1.1 2.5 2.5 5.0 8.3
10,s 18.8 18.0
Consumption ratio
Approximate rate constant, M - 1 min. - 1
1.024 0.997 1.017
48
1.014 1.010 1. O X $ 1 1. ( J U
57 68 95 111 114 ..
.o:n
1.035 1.029 1.019 1.029
49 47
.. ..
..
Reaction Rates.-Preliminary runs made a t 25" in air-saturated 2 M HClO, showed good reproducibility and were in accord with a rate law which was first order in U(IV) and in H202. &st values for the rate constants for these runs were determined by the use of a non-linear least squares program* wbich minimized the sum of the differences between the observed and calculated absorbancr values. The average serond-order rate constant was found to be 56.8 with a standard deviation of 1.85 and a maximum deviation of 2.9 M-l inin.-'. The second-order rate law fit the experimenQa1data quite well; the average difference between t,he calculated and observed absorbance values was only 0.002. The HC10, used in some of these runs was prepared by vacuum distillation, but no significant differences were observed. Similarly, the water used in some of the runs had been distilled but once, giving no detect.able effect. Alt,hough oxygen in the air reacts with V(IV) a t an insignificant rateSat room temperature in 2 AI Results HCIO,, it has been found that it has a significsnt Stoichiometry .-The stoichiometry of the re- effect on the kinetics of the U(IV)-H202 reaction. action was examined closely since it is known3 that Four additional runs were made in which the soluthe reaction between Fe(I1) and H20?is quantita- tions were deaerated with argon; these gave avertive when Fe(1I) ic. in excess but that a large age second-order rate constants ranging between amount of extra E120?is consumed when HzOzis in 38 and 44 M - I min.-'. Two additional runs using excess. The rt,oichionietry of reaction 1 was stud- air-saturated solutions gave 53 and 54 M-' min. ied by analyzing known mixtures after reaction in agreement with the previous series. Two runs $ahrated solutions gave apparent either for [H?O?]or for the [l?(IV)]-[H,O,] differ- made wing 0% ence. The reliults of some typical determinations of second-order rate constants of 56 and 61 M-' the number of moles of H2O2 consumed for each min.-l. The data from the solutions which had mole of U ( W ) which reacted (consumption ratio) (4) We are indebted to R . H. Moore and Ivan Cherry of the Theoare given in Table T. The data show that, unlike the retical Division of this Laboratory for computations involving the Computer: the former for modifying existing least square case of Fc(II), thcre is a small amount of H202 IBM-704 programs for our purpose (see R. 13. Moore and R. K. Zeigler. LA(1) This work was done under the auspices of the U. S. Atomic Energy Cornmiasion. (2) T. W. Newton, J . Phys. Cham.. 63, 1491 (1959). (3) J. H. Baxendale, in "Advances in Catalysis." Vol. IV. Academic Press Inc., New York, pr'. Y . , 1952, p. 46 ff.
8367) and the latter for slope calculations. (5) J. Halpern and 3. G . Smith, Con. J . Chem., S4, 1419 (IRRC,), report -d[U(IV)]/df ~ [ U ( I V ) I [ O I ] / [ H ~and ] show that in 0 . i .If HClO' a%30° and 0.96 atm. of 0 , the half-tiiiie for the oxidatirrii of U(IV) is about 130 min.
-
I50
0.01 M U(VI), 0.004 M Hg(I1) and 0.003 A I Pb(1I) were without significant effect. The rate u a s roughly tripled by 0.005 J I hg(1) and by 0.01 d1 Mn(I1) while 0.0002 AI Fe(I1) increased the rate about sixfold. The ions of Cu(1I) aiid Co(I1) nere found to inhibit the reaction. The decreases in rate observed with these ions are summarized in Table 11. It is to be noted that small amounts of Co(I1) cause less inhibition when Ro = 2 than when RO= 0.5, but that this effect is reversed at higher Co(I1) r omen tra tions.
,
'
f 100 P
I
=.
1
50
TABLE
been deaerated with argon fit a second-order rate law much less well than those from the air saturated solutions mentioned above. Apparent second-order rate constants were computed from the slopes of the absorbance versus time functionsJ; these rate couqtants were found to decrease uriiformly during the course of the run, having dropped about 20yGby the time the reaction was 7.5%) complete. It thus appears that oxygen in thf. solution.: increases the rate and allows better adherence to a second-order rate law. Consumption ratio determinations showed that there is a small amount of induced oxygen reaction during the U(IV) -H?O? reaction. The data obtained in all the runs made at 23' in deaerated 2 M IIClO, are summarized in Fig. 1. In order to make the data more readily comparable, the observed rates have been divided by tht, colicentration product [U(IV)] [HZ021 to give ic, thc apparent second-order rate coiistant, and plotted against R, the [H202]/[U(IT.')]ratio. The individual runs are shown as separate lines. Although there is considerable scatter among the various runs, it is clear that, the apparent second-order ratc, constants are larger in the high H202 region. At constant, R values there is no correlation betw-cen the apparent rate constants and the reactant concentration:t. l o r example, when R = 5, the H202concentrations for the runs hhown from top to bottom 111 Vig. 1 \?-ere 3.9, 3.9, 5 0. 2.3, I.!), 5 0, 2.5, 2.5. 3.6, 1.4, 1.9 and 4.5 mniolei per litel. Experimc~nt~ done a t 25' nith R,. the initial value of R, equal to 0.5 showed that decwasirig the HCIOl concmiration from 2 to 0.5 (ionic strength held a t 2 by the use of LiClO, or NnC104) mcieascd the rate by a factor between 2 arid 3. Some rates were measured a t 2.4, 9 8 and 34.2' ti weil its a t 25'. The temperature cocfficieiits of ond-order rate coii+tiits m:ide it ate the over-all activatioii energy fcir the reaction under various conditions. 111 2 11 HCIOa w t t , Bo = 0.57 the activatioii cncrg.rr n-as found to he 17 kc:tl 'mole and with Xo = 5 the corrcspondiiig valuc was 1 6 kea1 mole. 'In 0.5 ;II HC104 with I? = 0 5 , the iictivatiori criergy ma< found t o ti^ 19 kral 'mole. Duc to the knon II watter of the data these activatioii energies arc uneertaiii by a t least 1 kcal.jmole and it is not PO:sible to att:icli m y significance to the differences among the values The catalytic effect of several cations W : L ~invcstigatcd: in 1 Jf HClOd solutions with R,, := 0.5, 1
3
11
EFFECT OF Cu( 11) A V D Co( 11) O N THE RATEOF THE U( 1V)HzO2 REACTION Conditions: 2 j 0 , 2M HClOa. The tabulated values are the fractions of the corresponding uninhibited rates. Inhibitor
R~
CU(I1) Co(I1) CO(I1)
0 5 0 5 2 0
-----Inhibitor 10-4
085
concn -f% 10-3
io-?
10-1
073 83 05
060 56 50
0 33 18
The effect of Ro on the rates becomes smaller as the Co(I1) concentration j s increased; however, in 0.1 X solutions the rate is about 35% greater in solutions with the higher Ro. Small amount? of sulfate did not affect the reaction rate but 1.4 x 10-3 11.1 HC1 in 2 M HC104 iiicreased the rate by a factor of about 1.7. That the reaction is not catalyzed by glass surface w a q shown in an experiment in which the surfare area was tripled by the use of a coiled Pyrex rod placed in the reaction vessel. For these runs the reaction was followed by removing aliquots a t defiute times and quenching them 111 excess Ce-
(IT-). Photochemical effects are unimportant since the rate was not increased significantly when the light from a 100 W tungsten lamp was focussed onto the reaction vessel. The possibility of the formation of peroxide cbomplexes has been considered. Although U(V1) does not form peroxide complexes in acid solution6 the existence of Pu(IV)-peroxide complexes' makes U(T\')-peroxide complexes plausible A qtudy of such complexes is difficult because of :he rapidity of reaction 1. Preliminary experiment5 were made in 0.5 ;I1 IIC'104 at 2.4' 111which absortxuire. ohwved at 647S -1 xere extrapolated t o the t m e ot mixir~g At a con\fant U(IV) concentration of 2.6 >( 10 - I -If the extrapolated absorbance.s were found t o be L: nearly linear function of the H,02 roncent ration. A HzO? concentration of 4 G X 10 B I gax(' an extrapolated absorbance n hi(,h IT as 95' of its value in the absence of H a 0 2 This iiidiwtes that some complexing occurs but the effect i.s tim mlnll to make even LL qualitatix e eitiniatc I t I t s t'xteiit Discussion I n their study of the reaction hctwec-ri L(1Y) and dissolved oxygen, Halpern aiid Sniithj found that the reaction proceeds i n two stages. the firht a chain process leading to the net reaction
+ + 2H10 = UO?-' + H201 + 2H+ ( 2 ) 0 2
-___I
(6) J Corpel, BzlZZ. BOC chzm. Francs, 752 (1953). (7) R. E Conniek and W. H &IcVey, J Ani Chem %e., 71, 1334 (1949).
NOTES
Oct., 1961
1899
TETRACHLOROPHTHALIC ANHYDRIDEand the second reaction I , which they found to be instantaneous. The chain carriers which were AZAHYDROCARBON COMPLEXES proposed for reaction 2 are U(V) and HOa. BYMIHIRCHOWDHURX~ Our experiments indicate that reaction 1 also is a & Technology, 98, Upper C i r c d a r Rand chain process and that its rate is conveniently Uniaersitv College of ScienceCalcutta-9, India measurable.* The discrepancy in the reported Received March 81, 1961 rates of (1) is probably due to the fact that Halpern It was reported in a previous pap(+ that there and Smith used higher reactant and lower hydrogen was close correspondence between Anla, of T.C.P.A. ion concentration. The principd evidence that reaction 1 proceeds complexes of hydrocarbons and those of cwrrespondby a chain is the fact that it is strongly inhibited by ing azahydrocarbons. This suggest,ad that the Co(I1) and Cu(I1). This conclusion is supported elect.ron involved in the spectral jump came from by the lack of reproducibility which was encoun- the ?r-orbital of the azahydrocarbon. This does not tered in the rate measurements. It is interesting however necessarily mean tha,t t’he stability of the to note that the ions iMn(I1) and Co(I1) which had complex is also due to 7r-r interaction alone. In no effect on reaction 25 were found in the present order to see whether the stabilities also are close to work to catalyze and inhibit reaction 1, respec- each other, we have determined the equilibrium tively. Also the ions of Fe(II), Ag(I), Cu(1I) and constants of T.C.P.A. complexes of three azahydroC1- show the opposite effect in the two reactions. carbons by following a modified Benesi-Hildebrand These results imply that a t least one of the radicals relation. The data are summarized in Table I involved in reaction 1 is different from those in- and the plots are given in Fig. 1. It has not been possible to extend the study to other am-compounds volved in reaction 2 . It is plausible that the chain cariiers in reaction because of incipient precipitation. 1 are U(V) and HO and that in the absence of TABLE I catalysts or inhibitors the most important reaction< EQUILIBRIUM CONSTANTS O F CHARGE-TRANSFER COMPLEXES Inverse of a re Concn. of concn.
+ Hz02 = U(V) + HO TI(’$) + HZ02 U(V1) + HO U( [V) + HO = U(V) + H20 TJ( V) + HO = U(V1) + H20 U(1rV)
(3) (4)
Donor
Quinoline
+ = HOt + H20 C(1V) + HO? = U(V) + HzOz U ( \ ) + HOz C(V1) + H2Ot IIO + HO? Hz0 + 02 H202
+ 0,
U(V1) $- HOz
3.670
9.711 11.653 14.567 16.648 19.423
355
10.502 12.602 15,753 18.004
355
(6)
(7)
(8) (9) (10)
Since Co(I1) does not affect HOz radicals, their postulated existence in solutions with large R values is consistent with the observation that small concentrations of Co(I1) have very small inhibiting effectsin such solutions. The effect of oxygen may derive from the reaction I;( V )
of donor
(5)
The mechanism given by reactions 3-6 does not show all that happens since it fails to account for a-Methyldeviations from a second-order rate law, the slight quinoline lack of stoichiometry, and the increase in the apparent second-order rate constant a t high H20? concentrations. These complirations might tw explained by the additional reactions i,S-BenzoH’O
T.C.P.A.. moles/l. X 10’
(11)
Acknowledgment.-The authors wish to thank Helen D. Cowan for technical assistance in the rate measurements. They also acknowledge many helpful discusmms with Prof. Henry Taube, Dr. C. E. Holley, tJr., and especially with Dr. J. F. Lemons under whose general direction this work was done. (8) A recent Russian report describes some preliminary experim e n t s made in IliSO* solutions: E. A. Kanevsky a n d L: At Federova, Radiokhimiya, 2, 559 (1960).
quinoline
3.17i
(~/CD), I./mole
Wavr length, mF
21.400 29,270 36.590 48.780 58.550
K,
- C1 mole-’ X 102 1.
1.01,5 1 080
26
1.152 1.211 1 .‘ai5
If.79U
15
,887 978 I .ocit?
1.13:
21.004
2.819
-~
J?
875
0.759 .865 975
18
1.182 1.3iG
The equilibrium constants of azahydrocarbon complexes are somewhat higher than those of corresponding hydrocarbon This indicates that in addition to X--R interaction there occurs some specific interaction nith nitrogen non-bonding electrons. This also explains why thr stabilities of complexes of methylquinolirie and 7,8-benzoquinoline are less than that of quinoline in spite of thc higher ?r-ionizat.ionpotential of quinoline. Substituents in quinoline possibly offer hindrance t’othe localized interaction, thereby decreasing t,he stability. I t should be noted that the equilibrium constants of T.C.P.A.-hydrocarhoii and T.C.P.X.-azahydroc a r h i complexes are doser to one another than the rase of 12-hydrocarbon3 and T2-azahydrocarl)on coniplexes.* In the case of 12-complexesof aza-compounds, where n-electrons are supposed t o be iir(1) Whitmore Chemical Laboratory, Pennsylvania S t a t e Univers i t y , University P a r k , Pennsylvania. ( 2 ) M. Chowdhury a n d S.Basu, Traw. Faraday SOC.,6 6 , 335 (19tin). (3) R. B h a t t a c h a r y a a n d S. Basu. ibid.. 64, 1286 (1938). ( 4 ) .J. Nag Chaudhuri a n d S. Basu, ibid., 6 6 , 898 (1D60).
