J. Phys. Chem. 1985,89, 3359-3363
3359
Reaction of Azide Radicals wlth Aromatic Compounds. Azide as a Selective Oxidant Zeev B. Alfassi and Robert H.Schuler* Radiation Laboratory' and Department of Chemistry, University of Notre Dame, Notre Dame, Indiana 46556, and Department of Nuclear Engineering, Ben Gurion University of the Negev, 841 21 Beer Sheva, Israel (Received: December 20, 1984)
In basic aqueous solution the N3.radical is found to oxidize aromatic systems such as aniline and phenoxide ions and their derivatives at rate constants of (3-5) X lo9 M-l s-l . In contrast to the reactions of OH radicals, where addition to the ring dominates, oxidation appears to be directly by electron transfer. Compounds such as benzene, toluene, and anisole are not observably oxidized by N3. Phenol, in its neutral form in acidic solution, is oxidized several orders of magnitude more slowly than is the phenoxide anion. The rate of oxidation of phenols is strongly dependent on substitution, with activating groups increasing the rate in the order para > ortho > meta. N3-can be readily prepared radiolytically by OH oxidation of azide. Being a neutral radical which does not absorb significantly above 300 nm it is very promising as a selective oxidant for pulse radiolysis studies.
Introduction Many radicals and radical ions prepared by radiolysis and photolysis can be used as oxidants to study the chemical and spectral properties of organic radicals and radical ions. Secondary oxidants commonly used are C12-., (CNS)2--, SO4-., Br2--, and Ag2+. Several investigators in recent years have used the N,. radical as an oxidant in radiation chemical research of interest to inorganic2v3and radiobiological system^.'^ The N3. radical is readily formed either by flash photolysis10or pulse radiolysis" of aqueous sodium azide solutions and has a very narrow absorption band a t -274 nm and little absorption above 300 nm. This low absorption provides an important advantage in optical pulse radiolysis studies since it readily allows measurements on many transients in a spectral region where radicals such as B r p , (SCN)2-., and Cl2; mask observations at short times. As a result one can directly examine the production of the oxidized product and study details of the oxidation process even while the initial radical is present. Another advantage of N3. is its nonionic character, whereas most of the more commonly used secondary oxidants are ionic. Effects of charge and dependence on ionic strength are, thus, not involved. While the hydroxyl radical is also neutral, it frequently reacts simultaneously via several competing pathways, e.g., abstraction of an H atom, addition to a double bond or to an aromatic ring, or by electron transfer, so that a number of intermediates may be initially present. In contrast N3. radicals appear to react primarily via electron thus simplifying the secondary chemistry. This latter&rLi been found to be particularly important in a recent study of the oxidation of aniline1*where the cyclohexadienyl radicals initially produced in the case of .OH addition have a relatively long lifetime. Because N3. reacts primarily by electron transfer it is also expected to be much more selective in its reactions than is .OH.We report here studies of the oxidation of aromatic systems chosen to illustrate the use of N3. as a selective secondary oxidant. (1) The research described herein was supported by the Office of Basic Energy Sciences of the Department of Energy. This is Document No. NDRL-2661 from the Notre Dame Radiation Laboratory. (2) Buxton, G. V.; Janovsky, I. J . Chem. SOC.,Faraday Trans. I , 1976, 72, 1884. (3) Eriksen, T. E.; Lind, J.; Merenyi, G. Radiochem. Radioanal. Lett. 1981, 48, 405.
(4) Land, E. J.; Prutz, W. A. Znt. J . Radiat. Biol. 1977, 32, 203. (5) Land, E. J.; Prutz, W. A. Znt. J. Radiat. Biol. 1979, 36, 75. (6) Chauvet, J. P.; Viovy, R.; Santus, R.; Land, E. J. J . Phys. Chem. 1981, 85, 3449. (7) Singh, A.; Koroll, G. W.; Cundall, R. C. Radiat. Phys. Chem. 1982, l . , 137. ( 8 ) Land, E. J.; Sloper, R. W.; Truscott, T. G. Radiat. Res. 1983,96,450. (9) Butler, J.; Land, E. J.; Swallow, A. J.; Prutz, W. Radiat. Phys. Chem. 1984, 23, 265. (10) Treinin, A.; Hayon, E. J. Chem. Phys. 1969, 50, 538. (11) Hayon, E.; Simic, M. J . Am. Chem. SOC.1970, 92, 7486. (12) Qin, L.; Tripathi, G. N. R.; Schuler, R. H. Z . Naturforsch., in press.
TABLE I: Rate Constants for Oxidation of N< by .OH PH 10-loko~ fOH' 10-'%,b
7.89 10.66 11.00
11.36 12.2 13.1 14.1
1.oo
1.25 1.19 1.05
0.95 0.89
0.84
0.78
0.34 0.076 0.033F
0.33 0.06
1.25
1.26 1.18 1.08
0.006
1.19 f 0.05 Fraction of .OH radical remaining, assuming pK, = 11.9. koM/ foH. eSlow oxidation in 1 M KOH appears to be attributable to reaction of 0-. with a rate constant 2 X lo8 M-' s-'.
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Experimental Section The N3- radicals were produced by pulse radiolysis of aqueous solutions of NaN3, purged of oxygen, and saturated with N,O to convert hydrated electrons to hydroxyl radicals. For measurement of the rate of formation of N3-the buildup of the radical was monitored directly at 274 nm. The rate of reaction of N3with the various reactants was measured by following the kinetics of the formation of the transient of interest in the latter's principal absorption band. All the organic chemicals used were from Aldrich, except p-methoxyphenol which was from Eastman, and were of the highest purity available. Sodium azide was from Alfa Products. The pH of the solutions was adjusted by using reagent grade HC104 or KOH. Solutions were prepared with water from a Millipore Milli-Q system. The solutions were irradiated with 5-11selectron pulses from an ARC0 LP-7 linear accelerator. The computer-controlled pulse radiolysis apparatus previously described13 allowed averaging of traces taken consecutively while the solution flowed through the irradiation cell. Digitization of the transient signal was usually with a Biomation 8 100. A Biomation 6500 was used for experiments below 1-ps full scale. Dosimetry was with a secondary emission monitor calibrated with N20-saturated 0.01 M SCNsolution, taking the yield of (SCN)2-. as 6.14 and its extinction coefficient at 472 nm as 7580 M-' cm-'.14 Most experiments were carried out at an initial radical concentration of -3 X lo6 M. Results and Discussion Production and Decay of N3-. In the pH range of 4-13 the N,. radical can be readily prepared by .OH oxidation of azide anion *OH
+ N3-
ki +
OH-
+ N3.
(1)
(13) Modified form of apparatus described by Patterson, L. K.; Lilie, J. Int. J. Radiat. Phys. Chem. 1974, 6, 129. (14) Schuler, R. H.; Patterson, L. K.; Janata, E. J . Phys. Chem. 19%0,84, 2088.
0022-3654/85/2089-3359.$01 .50/0 0 1985 American Chemical Society
3360 The Journal of Physical Chemistry, Vol. 89, No. 15, 1985
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Alfassi and Schuler
n
TIME
+ X w
0
1
1
260
280 300 WAVELENGTH
320
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340
I
nm
Figure 1. Absorption spectrum of N3. as determined 1 p s after pulse irradiating a N20-saturated 0.00077 M NaN3 solution at pH 1 1 . Extinction coefficients are based on SCN- dosimetry (ref 14) with the N3yield taken as 5.5. Absorption maximum is at 274 f 1 nm where the extinction coefficient is 2025 f 50 M-l cm-'.
The production of N3..by reaction 1 is demonstrated very nicely by the ESR spin-trapping experiments of Behar and Fessenden.I5 Its absorption spectrum, obtained 1 ps after pulse irradiating a N20-saturated 0.0008 M N a N 3 solution at p H 11.3, is given in Figure 1 . The extinction coefficient a t 274 nm, based on a N3- radiation chemical yield of 5.5 pertinent to the N3- concentration of this experiment,I6 is 2025 k 50 M-' cm-'.Though more detailed, this spectrum agrees with that previously reported from flash photolysisI0 and pulse radiolysis experiments7*' although the extinction coefficient is at considerable variance with the values of 14002,2300," and 1300 M-' cm-' reported by other investigators at -274 nm. Rate constants observed for N3.formation are given in Table I. After correction for the low rate of oxidation by 0-.,as is noted below, these values give an average of ( 1.19 f 0.05) X 1O'O M-' s-l for reaction 1, in complete accord with the rate constant of 1.2 X 1 O l o M-l s-l reported by Hayon and Simic" from competition studies. Measurements in the pH region of 12-13 show that for the most part the observed rate constant represents the contribution from reaction 1, with very little contribution from oxidation of N3- by
-
0 - e .
+ N3-
k2
02-+ N3* (2) A rate constant of 3 X IO8 M-l s-l was observed at pH 14 showing that k2 is at least two orders of magnitude lower than k l . After the small contribution from reaction 1 was corrected for, this measurement suggests that 0; does, in fact, oxidize N < at a rate constant -2 X IO8 M-' s-l. However, pending a more-detailed study at very high pH this value should be regarded only as an upper limit for k2. This low rate constant must, of course,be taken into account in designing experiments in strongly basic solutions. In highly acidic solution, reactions are complicated by the relatively high pK, of hydrazoic acid (4.74)17which is oxidized considerably more slowly by .OH than is the azide anion. While a spectrum similar to that of Figure 1, but slightly less intense, was observed a t pH 4 the N3-production period corresponds to a rate constant of only 1.9 X lo9 M-' s-l. Since hydrazoic acid is 15% ionized at this lower pH, this value essentially represents reaction of -OH with the residual N;. The rate constant for .OH oxidation of hydrazoic acid is considerably less than lo9 M-I s-l and as a result N3.is not a practical oxidant much below pH 4. Absorption of HN3 in the near-ultraviolet interferes with direct 0--
+
(15) D. Behar, R. W. Fessenden, J . Phys. Chem. 1972, 76, 1710. (16) Schuler, R. H.; Hartzell, A. L.; Behar, B. J . Phys. Chem. 1981, 85,
192.
(17) Cotton, F. A,; Wilkinson, G . "Advanced Inorganic Chemistry"; Wiley-Interscience: New York, 1966; 2nd ed, p 339.
- ps
F m e 2. Decay of N3. as measured at 274 nm for solutions initially (0) 5 X lo6, (X) 3 X IO6, and (A) 1 X 10" M in N,.. Data represent averages of 100 experiments. The latter two sets of data are displayed, as described in the text, with delays of 15 and 110 ps indicated by the arrows to take into account the lower initial radical concentration. The solid curve corresponds to second-order decay with a rate constant (2k3) of 8.8 X IO9 M-' s-I.
measurements of .OH attack in more acidic media where relatively high H N 3 concentrations are required. In the absence of other solutes N3-decays considerably more rapidly than many radicals so that the mechanism for this decay is of some interest. The decay, as observed in a 1.3 X lo4 M azide solution at initial radical concentrations of 0.9 X IO", 2.9 X lo4, and 4.7 X 10" M, is illustrated in Figure 2. In these cases, in order to assure reliability at long times, the analyzing lamp was not pulsed and as a result the signal-to-noise ratio was relatively poor (
