May 5, 1961
OXYGEN
EXCHANGE IN NITRIC ACIDSOLUTIONS
a n induced exchange does not exclude the existence of such intermediates. In order to detect induced exchange, two requirements have to be fulfilled:
2035
(a) the XOOH should exchange with water a t a reasonable rate; (b) i t should exist in the system long enough for such an exchange to take place.
[CONTRIBUTION FROM THE ISOTOPE DEPARTMENT, WEIZMANN INSTITUTE OF SCIENCE, REHOVOTH, ISRAEL]
The Exchange of Oxygen between Hydrogen Peroxide and Water in Nitric Acid Solutions BY M. ANBARAND S. GUTTMANN RECEIVED OCTOBER 20, 1960 The kinetics of isotopic exchange of HnOI with water in nitric acid solutions have been studied a t 60 and at 100’. The dependence of the rate of exchange on the concentrations of nitric acid, nitrate ions and perchloric acid suggests nitrogen pentoxide as nitrating agent. A simultaneous isotopic exchange between hydrogen peroxide and nitric acid has been observed; on the other hand, no exchange between water and nitric acid is induced by the hydrogen peroxid: exchange. The Y I mechanism of exchange postulates pernitric acid as intermediate: N2O6 HiOz** + HNOi OtXOOH; OzNOOH
~ , 6
+
+
O ~ N ~ ~~ H6
The isotopic exchange of oxygen between hydrogen peroxide and water was subject to a number of investigations‘ but was not detected until recently. A careful study of this exchange in aqueous solutions highly enriched in 0lsfacilitated the determination of the upper limit of the specific rate of this exchange2R < 7 X a t 2.5’. A hydroxide ion catalyzed exchange was also observed2 in 1 N NaOH with a rate constant of 1.1 X lo-’ min.-’ at 25’. This nucleophilic substitution on the 0-0 bond implies that this -6
a,
bond eventually may be polarized to 0-0; this polarization, however, is a process of a very low probability. It seemed reasonable that if the 0-0 bond is pdarized by the inductive effect of a. substit-
-6
+
+
6 ~ .
6+
uent X t o form XOOH, the 0-0bond may be more vulnerable to a nucleophilic attack. Substituents having a strong negative inductive effect are for instants NO, C1, NO2 and SO,; the corresponding hydroperoxides ON OOH and ClOOH are only intermediates in hydroperoxide oxidation-reduction reactions,3-4 02NOOH-pernitric acid has been postulated in certain reactions,6 persulfuric acid 03SOOH,on the other hand, is a well established species. Isotopic oxygen exchange between hydrogen peroxide and water was observed on interaction of nitrous and hypochlorous acids with HzOz, but no quantitative conclusions could be derived owing to the transient nature of these hydroperoxides2 It was the purpose of this study to investigate the oxygen exchange between HzOzand water in a system where a stable hydroperoxide may be formed even if a t low concentrations. Nitric or sulfuric acid solutions of HzOz may contain such stable hydroperoxides. (1) E. R . S. Winter and A. V. A. Briscoe, J . A m . Chem. SOC.,73, 496 (1951); C. A. Bunton and D. R . Llewellyn, Rerearch, 6, 142 (1952). (2) M.Anbar, to be published. (3) hl. Anbar and H. Taube, J . A m . Chem. Soc., 7 6 , 6243 (1854). (4) L. Erdey and J. Inczedy, Acta Chim. Hung., 11, 125 (1956). (5) W.C. Schumb, C. N . Satterfield, R. L. Wentworth, “Hydrogen Peroxide,” A.C.S. Monograph Series No. 128, 1955, p. 665. Cf. H.H. Sisler, “Comprehensive Inorg. Chemistry,” Vol. V, Sneed and Rrasted, editors, D. Van Nostrand Co.,New York,N.Y., 1958,p. 101.
Experimental HzO1*was supplied from the distillation plant of the Weizmann Institute. Nitrite free nitric acid was prepared by vacuum distillation over urea. Sodium nitrate, perchloric acid and sulfuric acid were all of analytical purity. Hydrogen peroxide used was Merck Superoxol, stabilizer free reagent. The acid concentration was determined by acidimetric titration. The H202 was determined gasometrically after catalytic decomposition over platinum black. Procedure.-Solutions of H202 in HNOa were placed in a thermostat a t 60 f: 0.1” or a t 100 i 0.2”. Aliquots were taken a t intervals and the HzOz was decomposed on a vacuum line over pla.tinum black. I t mas found2 that platinum and water. black does not induce any exchange between HZOZ The gas was transferred into a glass tube with a break off seal and was submitted to the mass spectrometric analysis. The isotopic composition of the water was determined for each reaction mixture by equilibrating an aliquot with COz and analyzing the gas. I n experiments were the exchange reaction of nitrate ions was also followed, the nitrate was precipitated as Ba( NOa)z in an alcoholic solution. The dried Ba(NOl)z was heated with NHdCI in excess in a sealed glass tube for 45 minutes a t 400” to form NzO which was then analyzed. The mass spectrometric analysis was performed by a CEC model 21-401 Isotope Ratio Mass Spectrometer either by determining the ratio of masses 34/32 and 46/44 by the double collector or by determining the abundances of the individual masses. The rate constants were derived graphically by plotting A us. time, where A is atom per cent. the logarithm of A m excess of 0 ’ 8 in the sample.
-
Results and Discussion The isotopic exchange of oxygen between hydrogen peroxide and water was followed in perchloric, sulfuric and nitric acids: the results of these measurements a t 60’ are summarized in Table I. Considering first the results in perchloric acid i t seems that there exists no acid catalysed mechanism for the exchange. Such a mechanism might involve reaction l b as the rate determining step, reaction l a being a fast preequilibrium.6 PreHIO+ 4-HzOp HiOI
s HiOl+ + H20
+H2b
--t
HtO
+ Hpbz
(la) (Ib)
equilibrium l a would require the specific rate con(6) M. Anbar, A. Loewenstein and S. Meiboom, J . A m . Chem. Soc., 8 0 , 2630 (1858).
M. ANBARAND s. GUTTMANN
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TABLE I O N THE RATE OF THE EFFECTOF ACID CONCENTRATION H2OZ-Hz0ISOTOPIC EXCHANGE AT 60‘ k/ k/ (HzSOa), (HIJOa)iz,
Ho
0 . 1 A4 HClOi 4.OMHClO4 4 . 0 M HzSOc 6 . 0 MHpS04 7.0 M HpSOc 11.2 M HzSOI 1.0 M H N O ~ 2 . 1 MHhT03 3 . 0 MHNOs 4 . 0 MHN0.g 4.7 MHN08 5 . 6 MHNO3 6 . 0 hfHNO3 7 . 0 M Hh-03 8.0 MHNO, 10.0 &fHr\T03
k,
(”03),,
+
molel.
min.-I
1 -1.72 -1.85 -2.76 -3.32 -5.60
-0.18 -0.7 -1.02 -1.32 -1.50 -1.70 -1.79 -1.99 -2.10 -2.45
0.022 ,115 .285 ,600 ,850 1.43 1.71 2.52 3.52 5.90
mole -1 1. min.-’ X 108
2.2 10-8 2 , 8 10-8 4 . 5 10-8 5 . 5 10-8 6.810-8 1.1 10-7 1 . 010-7 1.510* 5 . 5 lo-’ 1.7 lo-’ 8 . 3 lo-’ 1.8 1 . 7 lo-‘ 2.i 6.9 2.9
mole - 2 1.2
min.-’ X 10’
1.1 0.9 1.0 1.0
2.06 1.14 0.68 0.47 1.15 0.88 .58 .43 .56 .83
stant to be dependent on the acidity €unction HO. Examination of the data in Table 1 indicates no such correlation (the values of Ho in Table I were taken or interpolated from Paul and Long’). Consequently, the exchange reaction in HC10, should be considered as an uncatalyzed interaction (IC), where water acts as a nucleophilic reagent.
+ Hz02
H*C*,
H20
+ Hz&
(IC)
The isotopic exchange in sulfuric acid is a slow process (the half-life in 11.2 molar acid being over ten years a t 60°), the rate constants are found proportional to sulfuric acid concentrations. The range of HzS04 concentration studies is not very extensive, as HzOz undergoes fast decomposition in more concentrated HzSOa solution under the experimental conditions. Because of this limited range of concentrations the first order dependence on HzS04 is not conclusive. If the rate of “spontaneous” exchange is taken into account, the rate of the H,SOI catalysed reaction is approximately proportional to the second power of H2S04concentration, which might then imply the participation of SO, in the exchange mechanism. This dependence on H2S04concentration may be explained by reaction 2a or by 2b. Qualitative observations of an exchange reaction of Caro’s
HG+ H ~ S O-+~ ~ 0 ~ +~H ~6O (2a) 6 ~ so3 + HZ HGSO~H (2b) H 0 3 S 6 8 H f H20 +HO3SbH + H6OH (2c) ----f
+ H20 +H03S6OH + Hz6 H O ~ S ~ O+ HH ~ O +HO~SOH+ HOOH
H03S66H
(2d) (2e)
acid (H2S05) with water in acid solutions have been reported,8 this exchange may proceed either eia reaction 2c or via reactions 2d and 2e. In order to extend Kolthoff’s experimentss into higher ( 7 ) ?VI. A. Paul and F. A. Long, Chem Rets., 57, 1 (1957). ( 8 ) I . A T . Kolthoff and I. K. Rfiller, J. A m . C h e w . Soc., 7 3 , 3055 (,l!l.jH
Vol. 83
acid concentrations, where the SO4 radical mechanism is unlikely, potassium persulfate was hydrolysed for 15 minutes in 14.4 N sulfuric acid in 82% H2OlS a t 100’; the reaction mixture was then decomposed over Pt black a t room temperature for 5 minutes and the evolved gas was analyzed; i t contained 13.82 atom per cent. Ols. This result suggests that reaction 2c is the exchange step on hydrolysis, because an exchange via reactions 2d and 2e would proceed a t a much slower rate considering the exchange kinetics presented in Table I . The isotopic exchange in nitric acid proceeds a t a much higher rate than in perchloric or in sulfuric acids. The dependence of the rate of exchange on H202 concentration showed that i t remains constant within f 12% when H2Oz concentration was changed by a factor of ten. Moreover hydrogen peroxide concentration was decreasing constantly during the exchange experiments owing to its decomposition, still no deviation from the first order dependence was observed. These results imply a first order dependence on HzOz concentration R = k [HtOz]. The dependence of the rate of exchange on the stoichiometrical concentration of HNO, shows approximately a fourth power dependence; i t is hard however to suggest a mechanism with four molecules of nitric acid involved. Plotting log k vs. Ho in nitric acid yields a second order dependence, which impliesg a rather improbable species (H3N03++)as intermediate. The next parameter examined was the concentration of undissociatednitric acidlo (HN03)r ; values reported for 25’ were used because i t was hard to estimate dissociation constants a t higher temperatures by extrapolation. 1 1 , 1 2 I t was found that the specific rate constants are proportional to the second power of the concentration of undissociated nitric acid, R = k [H2O:+][HN03]2. 9 s may be seen from Table I, this dependence holds throughout the whole range investigated, with rate constants changing by four orders of magnitude. The exchange mechanism suggested is therefore H30f “03
+ SO3+ HT\:03 **
+ HOOH OpKOOH + HzO ** 02XOOH + H 2 0 ** 0zNOOH + Hp0 KpO5
**
+ H?O KpOj + H20 ** “01 + 0pNOOH 02NbH + HO6H ** * O Z S 0 0 H + HpO * OzNOH + HOOH HN03
(3a)
(3b)
(3~) (3d) (3e) (3f)
The fast preequilibria assumed are (3a), (3b) and (3c) and the rate determining step are either reaction 3d or reaction 3e followed by (3fj. I n order to test the participation of preequilibrium 3a in the exchange mechanism, the effects of free nitrate ions as well as of a strong acid on the rate of exchange were examined. These results are presented in Table 11. The concentration of un(9) F. A. Long and hf.-4.Paul, Ciiem. Revs., ET, 983 (1957). (10) 0. Redlich and J. Bigeieisen, J. A m . Chem. SOL.,65, 1833
(1943). (11) 0. Redlich, Chem. Reus., 39, 333 (1946). (12) N R . Rao, Indian J . Phys., 17, 2Y5 (1943).
May 5, 1961
OXYGEN
EXCHANGE I N NITRIC ACID SOLUTIONS
dissociated nitric acid in presence of HClOd and of NaN03 was calculated from the equilibrium constant for reaction 3a K = 2.77 X low2,which was calculated from the data of Redlich and BigeleiIt is evident therefore that sen for 2 M "03. equilibrium 3a is involved in the exchange mechanism. The participation of H2N03+which may also be formed from two molecules of nitric acid 2HN03 HzN03+ f NO3- is unlikely in view of the fact that nitrate ions increase the rate of exchange, whereas the concentration of HzN03+ is diminished in presence of Nos-.
-
TABLE 11 THEEFFECT OF NITRATEIONS AND OF PERCHLORIC ACID ON RATEOF EXCHANGE (100") k/ (HNOa)f* (HNOa)fi
mole 1.
1.0hfHNO3 2 . 0 i&f "03 2 . 0 M €€NO3 2.0 M HN03 3 . 0 MHNOI 6 . 0 M "01 6 . 6 M HNOi
+
+
0.022
k, mh-1
1.510-6 2 . 3 10-4 2 . 0 M NaN03 ,216 9 . 4 2 . 0 M HC104 ,216 1.1 ,285 1.710-3 1.71 4.5 2 . 1 8 7.510-2
.loo
mole2 1.8 min. -1 X 102
3.1 2.3 2.0 2.4 2.1 1.5 1.6
Changing the concentration of Hz02considerably did not change appreciably the rate of isotopic exchange, thus an exchange reaction zero order in Hs02 may be excluded. As the rate determining step involves a molecule of H20z,reaction 3b must also be a preequilibrium reaction. Reaction 3c, being an analogous step to the reverse of reaction 3b, may be considered & a fast reaction compared to reactions 3d or 3e; thus it is the stationary concentration of pernitric acid that determines the rate of exchange. Since there is no experimental value for the equilibrium constants of reactions 3b and 3c, no estimate can be made for the specific rate constant of the rate determining step. If the step involving isotopic exchange is reaction 3d, then an isotopic exchange between hydrogen peroxide and nitric acid would be expected. I n order to check on this, a kinetic experiment presented in Table I11 was carried out. Nitric acid a t an isotopic concentration lowe? than that of the solvent was mixed with hydrogen peroxide a t normal isotopic composition and the changes in the isotopic composition of nitric acid, hydrogen peroxide and water were followed. From these results i t is evident that as the H202 exchanges with the solvent, a decrease in the isotopic concentration of 01*in the nitric acid is observed, which implies an exchange reaction between H20z and HN03. At the final stage the isotopic concentration of 01*in "03 increases again, owing to its equilibration with the solvent. Reaction 3f, on the other hand, would require nitric acid to exchange with water a t least a t the rate of
2037
TABLE 111
SIMULTANEOUS EXCHANGE OF HYDROGEN PEROXIDE WITH WATERAND WITH NITRICACID = 6.6 N atom % 0 1 8 in
("02)
Time, min.
0 15 35 57 85
Hi0
38.2
37.7
(HzOP)= 1.5 M Temp 100' atom'% 0 1 8 in HzOz
Atpm % ' 018 ~n HNOp
0.21 3.66 8.60 13.8
36.8 34.5 33.0 32.3 34.2
the hydrogen peroxide exchange. It was found, however, that in nitric acid solutions above 4 M the rate of hydrogen peroxide exchange with water exceeds that of nitric acid with water. I n 6 molar nitric acid the rate of H2Oz exchange with water is 3 times faster than that of nitric acid with the ~ o l v e n t ' ~moreover, ; hydrogen peroxide was found to have no catalytic effect on the rate of HN03-H20 exchange. From the rate constants a t 60 and a 100' an activation energy of 34 kcal./mole for the exchange reaction may be derived. This value, however, has little meaning, because the temperature coefficient of the exchange reaction involves the temperature dependence of the preequilibria ; thus the measured activation energy is not necessarily the energy of activation of the rate determining step. The mechanism of formation of 02NOOH may be considered as a nitration reaction of HzOz, the nitrating agent in this case being not the nitronium ion but nitrogen p e n t 0 ~ i d e . I ~This reaction may just as well be looked a t as a nucleophilic substitution on Nz06, a reaction which is analogous to its hydro1ysis. Reactions 3d and its analogous reaction 2c are nucleophilic substitution reactions on the peroxide oxygen. This means that the 0-0bond undergoes polarization by the substituting group X forming 6-
6+
XO-OH. This type of polarization is favored by the negative inductive effect (-1) of the group X. The relative inductive effect of the SO3 group is smaller than that of N02,16 thus the nucleophilic substitution by water on OzNOOH proceeds a t a faster rate than that on 03SOOH. Attempts to detect a nucleophilic substitution by HS04- on 03SOOH have failed,16 probably because of the very low nucleophilic reactivity of the bisulfate ion. On the other hand induced isotopic exchange between hydrogen peroxide and water by hydroperoxides of transient nature2 demonstrates other cases analogous to reactions 2b and 3d. ( 1 3 ) M . Anbar and S. Guttmann, to be published. (14) C. K . Ingold, "Substitution a t Elements other than Carbon," Weizmann Science Press, 1959. (1.5) C. K. Ingold, "Structure and Mechanism in Organic Chemistry," Cornel1 Univ. Press, Ithaca. N. Y., 1953, Ch. 2,6. (16) G. Levey, D. R. Campbell, J. 0. Edwards and J. MacLachlan, J . A m . Chem. S O C 1, 9 , 1797 (1957).
