T H E
J O U R N A L
OF
PHYSICAL CHEMISTRY Registered in US.Patent Office 0 Copyright, 1978, by the American Chemical Society
VOLUME 82, NUMBER 14
JULY 13, 1978
Rate Constants for the Reaction of the OH Radical with Selected Alkanes at 300 K Karen R. Darnall, Roger Atkinson," and James N. Pitts, Jr. State wide Air Poiiution Research Center and Department of Chemistry, University of California, Riverside, California 9252 1 (Received March 1, 1978) Publication costs assisted by the California Air Resources Board and the National Science Foundation
Rate constants for the reaction of OH radicals with propane, isobutane, neopentane, n-pentane, isopentane, cyclopentane, and 2,3-dimethylbutane have been obtained from their rates of disappearance, relative to that for n-butane, in irradiated NO,-hydrocarbon-air mixtures at atmospheric pressure and 300 f 1 K. The rate constants obtained are as follows (10l2 k2 cm3 molecule-' s-l): propane, 1.59 f 0.22; isobutane, 2.52 f 0.05; neopentane, 1.04 f 0.17; n-pentane, 3.74 f 0.13; isopentane, 3.78 f 0.07; cyclopentane, 4.72 f 0.28; and 2,3-dimethylbutane, 5.67 f 0.29. From an analysis of these and the available literature rate constant data for r C 3 alkanes, revised rate constants for the reaction of OH radicals per primary, secondary, and tertiary C-H + Nkrtx 2.1 x cm3molecule-' bond were obtained, yielding k2 = (Nprim X 6.5 x W4+ N,,, x 5.8 x s-l at 300 K , where Nprim, N,,,, and Ntertare the numbers of primary, secondary, and tertiary C-H bonds in the alkanes. This equation can then be used to calculate values of k2 for alkanes for which no rate constant data presently exist.
Introduction The hydroxyl radical is known to play a dominant role in the chemistry of the atmosphere and in combustion processes. For the alkanes, which comprise a significant portion of the hydrocarbons emitted into polluted urban areas, reaction with the OH radical is essentially their sole loss process in the troposphere. With the increasing emphasis toward lowering the reactivity of the organics emitted into the troposphere and the development of a reactivity scale based upon the rates of reaction of these organics with the OH radical,'-3 rate qonstant data for the reaction of the OH radical with the alkanes have become necessary. While numerous a b ~ o I u t e ~ -and ' ~ relative rate19-28constant studies have been carried out for the alkanes to date, there are still in many cases significant discrepancies. Greiner6 derived a formula enabling rate constants to be calculated from the number of primary, secondary, and tertiary C-H bonds and the rate constants for the reaction of the OH radical per primary, secondary, and tertiary C-H bond. Such a formula in principle allows rate constants to be derived for alkanes for which rate data 0022-3654/78/2082-1581$01 .OO/O
are not available. However, because of the scatter in the literature data, especially for alkanes containing tertiary C-H bond^,^^^^,^^^^^ use of this formula to calculate rate constants for such alkanes may lead to large uncertainties. In order to resolve certain of the discrepancies between the literature data, rate constants for the reaction of OH radicals with selected alkanes have been determined in this work. As a result, it is possible t o calculate a revised set of rate constants for the primary, secondary, and tertiary C-H bonds for use in the Greiner formula.
Experimental Section The experimental techniques used to determine OH radical rate constants have been described in detail p r e v i o ~ s l yhence , ~ ~ ~only ~ ~ a brief summary will be given here with emphasis on the modifications to the techniques pertinent to this study. Irradiations of the NO,-hydrocarbon-air mixtures were carried out in a nonrigid Teflon bag of -5500-L volume which was placed inside an 6400-L Pyrex environmental chamber. The chamber temperature was maintained a t 300 f 1 K during the
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0 1978 American Chemical Society
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irradiations. The alkane concentrations were monitored as a function of irradiation time by gas chromatography (FID) using a 36 ft X 1/8 in. stainless steel column of 10% 2,4-dimethylsulfolane on C-22 Firebrick (60-80 mesh) followed by a 18 in. X 1/8 in. stainless steel column of 10% Carbowax 600 on C-22 Firebrick (60-80 mesh) operated a t 273 K. The initial concentrations of the alkanes were as follows: propane, 0.047-0.049 ppm (1 ppm = 2.37 X 1013molecule a t 300 K and 735 Torr total pressure); n-butane, 0.073-0.075 ppm; isobutane, 0.055-0.058 ppm; neopentane, 0.053-0.059 ppm; n-pentane, 0.128-0.137 ppm; isopentane, 0.101-0.119 ppm; cyclopentane, 0.155-0.174 ppm; and 2,3-dimethylbutane, 0.105-0.114 ppm. In addition to these alkanes, propene, isobutene, and formaldehyde were also initially present in order to obtain a reactive hydrocarbon-NO, mixture. Four 4-h irradiations were carried out, each with two completely independent analyses of the alkanes every 0.5 h. In a nonirradiated experiment the alkane concentrations as measured by gas chromatography were invariant over the 4-h sampling period to better than f1.5% (corresponding to 1 4 % of the typical reactive losses during the irradiation for the case of n-butane), thus verifying that the alkanes showed no dilution losses due to sampling.
Results Rates of disappearance of the alkanes propane, n-butane, isobutane, neopentane, n-pentane, isopentane, cyclopentane, and 2,3-dimethylbutane were obtained from four separate irradiations of NO,-hydrocarbon-air mixtures a t atmospheric pressure and 300 f 1 K. With the present experimental system, dilution losses of the reactants due to sampling were eliminated, as confirmed by the nonirradiated control experiment. Furthermore it has been shown from chemical kinetic computer modeling studies3'J2 that the sole reactive loss process for the alkanes under conditions such as those employed in the present study is via reaction with the OH radical. For instance, ~ negligible reaction with O3 can be readily ~ a l c u l a t e dto~ be a t the O3 concentrations observed (C0.3 ppm), while reaction with O(3P)atoms can be ~ a l c u l a t e d , from ~ ~ J ~the maximum NO2 concentrations and the light intensity, to contribute 50.6% of the loss due to reaction with the OH radical. Hence the observed disappearance of the alkanes during the irradiations is solely due to reaction with the OH radical: -d In [alkane]/dt = k[OHl (1) where h is the rate constant for the reaction OH + alkane H 2 0 + products
K. R.
Darnall, R. Atkinson, J. N. Pitts
TABLE I: Rate Constants k , for the Reaction of OH Radicals with Selected Alkanes at 300 K alkane propane is0 bu tane neopentane n-pentane isopen tane cyclopen tane 2,3-dimethylbutane
lO'*k, cm3 molecule-' s-lavb 1.59 t 0.22 2 - 5 2 ? 0.05 1.04 t 0.17 3.74 t 0.13 3.78 * 0.07 4.72 t 0.28 5.67 t 0.29
The indicated error limits are one standard deviation. Values of k , calculated using k , = 2.73 X lo-'' cm3 molecule-' $ - ' for n-butane at 300 K,'*
inated by ratioing the gas chromatographic response to those for a reference alkane, n-butane in this case. Thus
where [A,],,, [A2],,are the concentrations of alkanes Al and A2 a t time to, [A,], and [A,], are the corresponding concentrations a t time t , and kl and h2 are the rate constants for the reactions of alkanes Al and A2 with the OH radical, n-Butane was used as the reference alkane A,, as the three absolute rate constant determinations using flash photolysis techniques agree well.639Js A rate constant a t 300 K of kl = 2.73 X lo-'' cm3 molecule-l s-l was used, as calculated from the Arrhenius expression of Perry, Atkinson, and Pitts.ls From least-squares analyses of plots of In ([A2],/[Al],) against time, and the OH radical concentrations as obtained from least-squares analyses of the n-butane disappearance rates, a total of eight sets of rate constants h2 were obtained. These rate data were then weighted39using their least-squares standard deviations, and the weighted mean values, together with the standard deviations, are given in Table I. (The unit weighted mean values were within 4% of the weighted mean values, though with somewhat larger standard deviations.)
Discussion The rate constants h2 obtained in this work are compared with literature values in Table 11,which also includes all the available data for 2C3 alkanes (for many of the relative rate studies the reference reaction rate constant has been reevaluated to take into account more recent rate constant data). In many cases it can be seen that the present rate constants are in agreement with the literature data within the cumulative error limits, but there are obviously significant discrepancies, especially for n-pentane and 2,3-dimethylbutane. For 2,3-dimethylbutane, the present value of k2 is substantially lower than the absolute In all the irradiations it was observed that, within the rate constant determined by Greiner,6while it is somewhat analytical errors, plots of In [alkane] against irradiation higher than, though in agreement within the cumulative time were linear, indicating that in each case the OH error limits with, the two previous rate constant radical concentration was essentially constant over the determination^^^,^^ from this laboratory. The present rate duration of the irradiation. This has been noted in constant for propane (k2 = (1.59 f 0.22) X cm3 previous OH radical rate studies from this laboratory using molecule-l s-l) is seen to be approximately in the middle Using the literature rate constant this of the room temperature literature though for the reaction of OH radicals with n-butane,ls the OH it is in excellent agreement with k2 = 1.62 X cm3 radical concentrations during the irradiations were in the molecule-l s-l, calculated using the Arrhenius activation ~, to the concenrange (2-3) X IO6 molecule ~ m - similar trations observed in the previous rate s t ~ d i e s . ~ ~ , ~ ~energy , ~ ~ of, ~Greiner6 ~ - ~ ~to convert the rate constant of Harker and Burton,13 obtained a t 329 K, to 300 K. Equation 1 could hence be used in the integrated form Using the formula due to Greiner6 In ([alkane],,/[alkane],) = k[OHl(t - to) (2) (4) h2 = Nprimhprim + N sech sec + Nterthtert where [alkane],, and [alkane], are the alkane concentrawhere Nprim, N,,,, and Ntertare the numbers of primary, tions a t time to and t , respectively. Furthermore, errors secondary, and tertiary C-H bonds in the alkane, and hph, due to small differences in sampling volumes were elim-
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Reaction of the OH Radical with Selected Alkanes
The Journal of Physical Chemistty, Vol. 82, No. 14, 1978 1583
TABLE 11: Comparison of the Literature Room Temperature Rate Constants k, for >C, Alkanes with the Values Obtained in This Study and with Those Calculated from Eq 5 10'*k, 10'Zk,cm3 molecule-'s-' 300 K alkane cm3 molecule-' s-* at T K ref calcd 1.20 296-299 6 1.55 0.83 f 0.17 300 8 20 298 2.2 f 0.6" 295f 2 14 2.02f 0.10 1.98 f 0.08 13 3 2300 9* 1.62b 300 This work 1 . 5 9 f 0.22 298-301 2.71 n-butane 6 2.57 7 298 4.1 9 298 2.35 f 0.35 2.9 f 0.7a 20 298 12 4.22 298 22 2.59 t 0.16c 292 18 2.72 f 0.27 298 297-305 2.69 2.46 6 is0bu tane 3.5 t 0.9a 298 20 2.ld 303 27 28 1.86 f O.2lc 305 This work 3 00 2.52 f 0.05 4.64 20 298 1.2 f 0.3a cyclobu tane 0.78 6 292-298 neopentane 0.825 This work 300 1.04 f 0.17 3.87 27 303 6.2d n-pentane This work 300 3.74 f 0.13 3.84 23 305 3.1 f 0.6e isopentane This work 300 3.78 f 0.07 21 5.80 298 6.1a cy clopentane This work 300 4.12 f 0.28 305 23 5.01 5.0 f l.Oe 2-methylpentane 5.01 23 305 3-methylpentane 6.8 f 1.4e 5.03 305 5.9 f 1.2e 23 n-hexane 303 27 5.7d 292 25 6.1 f 0.4e 6.96 6 295 7.95 cyclohexane 20 298 6.7 f 1.5= 303 27 6.2d 4.98 300 6 2,3-dimethylbutane 7.45 26 305 4.8 t l . O d 303 4.3 f 1.4f 24 This work 300 5.67 f 0.29 6 296-303 3.08 5.05 2,2,34rimethylbutane 26 305 3 . 6 f 0.7d 1.17 1.12 2,2,3,3-tetramethylbutane 6 294-301 4.24 2,2,4-trimethylpentane 3.73 6 298-305 7.35 295 n-octane 8.42 6 a Relative to OH t CO, placed on an absolute basis using k(OH t CO) = 1.5 X lo-'' cm3 molecule-' s-'.~' Converted t o Relative to OH t CO; carried out at 100 Torr total pressure of CO" 300 K using the activation energy obtained by ref 6. or N, + 0,.z8It is assumed that CO is analogous t o N, t 0, as a third body in the OH t CO pressure d e p e n d e n ~ e , ~and ~ ? a~ ' value of k(OH t CO) = 1.75 X cm3 molecule-' s-' has been used.29 Relative to OH t isobutenei6 or OH t cis-2butene," placed on an absolute basis using k(OH + isobutene) or k(OH t cis-2-butene) calculated from the Arrhenius expression of ref 41. e Relative to OH t n-butane, placed on an absolute basis using k(OH + n-butane) calculated from the Arrhenius expression of ref 18. f Relative to OH t ethane, placed on an absolute basis using k(OH + ethane) calculated from the Arrhenius expression of ref 6. propane
I
h,,,, and htertare the corresponding rate constants for H atom abstraction per C-H bond, the rate constant data in Table I1 are judged to be best fit (based mainly on the present study and those of Greiner,6 Perry, Atkinson, and Pitts,18 and Darnall et a1.26)a t 300 K by
lzz = N~~~~x 6.5 x 10-14 + N,,, x 5.8 x 10-13 + Ntert X 2.1 x cm3 molecule-' s-l (5) (Greiner6derived hprh = 6.6 X cm3molecule-I s-l; k,, = 5.6 X cm3 molecule-' s-l, and h,,,, = 2.9 X molecule-' s-l, the major difference obviously being the value for Equation 5 fits the available rate constant data to better than approximately f20% in most cases (Table II), the exception being cyclobutane, probably due to the effects of ring strain. Because of the closeness of the present values of kprim and k,,, with those obtained by Greiner6 and because the
only temperature dependent rate constant studies for 2 C 3 alkanes are those of Greiner6 and Perry, Atkinson, and Pitts18(which agree very well for n-butane, the only alkane studied by Perry, Atkinson, and Pitts18), the Arrhenius activation energies of Greiner6for hpim and h,, can be used to derive h2 as a function of temperature. For the alkanes containing tertiary C-H bonds, an activation energy of E,, N 0 kcal mol-l gives a reasonably good fit to the high temperature (-493-498 K) rate data of Greinera6Hence the equation k 2 = 1.01 x 10-'2e-'"5/RrN prim . + 2.41 x 10-12e-850/~~~,,~ + 2.10 X 10-12Nt,rtcm3 molecule-l s-l (6) fits the available literature data over the temperature range 300-500 K within acceptable limits (-*20% or better) and can hopefully be used to calculate rate constants for
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G. Levin
the reaction of OH radicals with ZC, alkanes, apart from cyclobutane (and by analogy, cyclopropane) for which no data presently exist.
(15) (16) (17) (18)
Acknowledgment. The authors gratefully acknowledge the assistance of F. R. Burleson and G. C. Vogelaar for carrying out the gas chromatographic analyses, and W. D. Long for valuable assistance in conducting the chamber experiments. This work was supported by the California Air Resources Board (ARB Contract No. A6-172-30) and by t h e National Science Foundation (Grant No. CHE76-10447).
(19) (20) (21) (22) (23) (24) (25)
References and Notes
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(1) K. R. Darnall, A. C. Lloyd, A. M. Winer, and J. N. Pitts, Jr., Environ. Sci. Techno/., 10, 692 (1976). (2) J. N. Pitts, Jr., A. C. Lloyd, A. M. Winer, K. R. Darnall, and G. J. Doyle, 69th Annual Air Pollution Control Association Meeting, Portland, Oreg., June 27-July 1, 1976, Paper No. 76-31.1. (3) J. N. Pitts, Jr., A. M. Winer, K. R. Darnall, A. C. Lloyd, and G. J. Doyle, "International Conference on Photochemical Oxidation Pollution and Its Control, Proceedings", Vol. 11, B. Dimitrlades, Ed., EPA-600/ 3-77-001b, p 687, Jan 1977. (4) W. E. Wilson and A. A. Westenberg, Symp. Int. Combust. [ f r o c . ] , I l t h , 1966, 1143 (1967). (5) N. R. Greiner, J. Chem. fhys., 46, 3389 (1967). (6) N. R. Greiner, J. Chem. fhys., 53, 1070 (1970). (7) E. D. Morris, Jr., and H. Niki, J . fhys. Chem., 75, 3640 (1971). (8) J. N. Bradley, W. Hack, K. Hoyermann, and H. Gg. Wagner, J. Chem. Soc., Faraday Trans. 1 , 69, 1889 (1973). (9) F. Stuhl, Z . Naturforsch. A , 28, 1383 (1973). (10) D. D. Davis, S. Fischer, and R. Schiff, J. Chem. fhys., 61, 2213 (1974). (1 1) J. J. Margitan, F. Kaufman, and J. G. Anderson, Geophys. Res. Lett ., 1, 80 (1974). (12) S. Gordon and W. A. Mulac, Int. J . Chem. Kinet., Symp. 1, 289 (1975). (13) A. B. Harker and C. S. Burton, Int. J . Chem. Kinet., 7, 907 (1975). (14) R. P. Overend, G. Paraskevopoulos, and R. J. Cvetanovic, Can. J. Chem., 53, 3374 (1975).
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C. J. Howard and K. M. Evenson, J . Chem. fhys., 64, 197 (1976). R. Zellner and W. Steinert, Int. J. Chem. Kinet., 6, 397 (1976). C. J. Howard and K. M. Evenson, J . Chem. fhys., 64, 4303 (1976). R. A. Perry, R. Atkinson, and J. N. Pitts, Jr., J . Chem. fhys., 64, 5314 (1976). R. A. Gorse and D. H. Volman, J . fhotochem., 1, 1 (1972). R. A. Gorse and D. H. Volman, J . fhotochem., 3, 115 (1974). D. H. Volman, Int. J. Chem. Kinet., Symp. 1, 358 (1975). I. M. Campbell, B. J. Handy, and R. M. Kirby, J. Chem. SOC.,Faraday Trans. 1 . 71. 867 11975). A. C. Lloyd, K. R. Darnall,'A. M. Winer, and J. N. Pitts, Jr., J . fhys. Chem., 80, 789 (1976). R. Atkinson, K. R. Darnall, A. M. Wlner, and J. N. Pitts, Jr.. Final ReDort to E. I. duPont de Nemours and Co., Inc., Feb 1, 1976. I. M. Campbell, D. F. McLaughlin, and B. J. Handy, Chem. phys. Lett., 36, 362 (1976). K. R. Darnall, A. M. Wlner, A. C. Lloyd, and J. N. Pitts, Jr., Chem. Phys. Lett., 44, 415 (1976). C. H. Wu, S. M. Japar, and H. Niki, J . Environ. Sci. Health, A l l , 191 (1976). R. Butler, I. J. Solomon, and A. Snelson, Chem. fhys. Lett., 54, 19 (1978). R. Atkinson, K. R. Darnall, A. C. Lloyd, A. M. Winer, and J. N. Pitts, Jr., Adv. fhotochem., in press. G. J. Doyle, A. C. Lloyd, K. R. Darnall, A. M. Winer, and J. N. Pitts, Jr., Environ. Sci. Techno/.,9, 237 (1975). A. C. Baldwin, J. R. Barker, D. M. Golden, and D. 0. Hendry, J. fhys. Chem., 81, 2483 (1977). W. P. L. Carter, A. C. Lloyd, J. L. Sprung, and J. N. Pitts, Jr., Int. J. Chem. Kinet., in press. C. C. Schubert and R. N. Pease, J . Chem. fhys., 24, 919 (1956). J. T. Herron and R. E. Huie, J. Fhys. Chem. Ref. Data, 2, 467 (1973). F. Kaufman and J. R. Kelso. J . Chem. fhvs.. 46. 4541 (1967). A. M. Winer, A. C. Lloyd, K. d. Darnall, and j .N: Pitts, Jr., J : fhys. Chem., 80, 1635 (1976). A. C. Lloyd, K. R. Darnall, A. M. Winer, and J. N. Pitts, Jr., Chem. fhys. Lett., 42, 205 (1976). A. M. Winer, A. C. Lloyd, K. R. Darnall, R. Atkinson, and J. N. Pitts, Jr., Chem. fhys. Lett., 51, 221 (1977). L. G. Parratt, "Probability and Experimental Errors in Science", Wiley, New York, N.Y., 1961. R. A. Perry, R. Atkinson, and J. N. Pitts, Jr., J . Chem. fhys., 67, 5577 (1977), and references therein. R. Atkinson and J. N. Pitts, Jr., J . Chem. fhys., 63, 3591 (1975).
Photooxidation of Aromatic Hydrocarbons by Europium(111) Salts G. Levin Department of Chemistry, State University of New York, College of Environmental Science and Forestry, Syracuse, New York 13210 (Received January 12, 1978) Publication costs assisted by the State University of New York
Aromatic hydrocarbons such as tetracene (T),perylene (P),coronene (C), and naphthalene (N) are photooxidized in acetonitrile by Eu(II1) perchlorate or nitrate to their respective radical cations. The reversible reaction, , place in the dark period. The rate constants being 2.8 X lo6, 1.6 AH.+ + Eu(I1) AH Eu(III), M 0takes x lo7,1.3 x lo9, and 3.9 x IO9M-ls-l for perylene, tetracene, coronene, and naphthalene, respectively. These satisfactorily correlate with the reduction potential llEo. In addition, the rate constant of oxidation by Eu3+ of the excited singlets of tetracene and coronene were determined. The latter are diffusion controlled being 4.5 x 1Olo and 4.0 x 1O1O M-l s-l for tetracene and coronene, respectively. The triplets of tetracene and coronene are not oxidized by Eu3+ implying that the rate constant of these reactions are smaller than lo6 M-' s-'.
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Introduction Photochemical reactions, and especially photooxidations and photoreductions, are attracting much attention as possible routes for harnessing solar energy. Since photooxidation or photoreduction resulting from electron transfer induced by light proceed in homogeneous solutions simultaneously with the reverse thermal electron transfer, the resulting products are rapidly destroyed. It is our 0022-3654/78/2082-1584$01.00/0
intention to control and, if possible, to slow down the undesirable thermal processes. With this in mind, we investigated some photooxidations and photoreductions and studied the kinetics of the respective reverse thermal electron-transfer processes using aromatic hydrocarbons as the reducing agents and europium salts as the oxidants. The results of our studies are reported in this communication. 0 1978 American
Chemical Society
