T H E
J O U R N A L
O F
PHYSICAL CHEMISTRY Registered in U.S. Patent Office
0 Copyright, 2978, by t h e American Chemical Society
VOLUME 82, NUMBER 12
JUNE 15, 1978
Rates of OH Radical Reactions. 4. Reactions with Methanol, Ethanol, I-Propanol, and 2-Propanol at 296 K' Ralph Overend and George Paraskevopoulos' Division of Chemistry, National Research Council of Canada, Ottawa, Canada, K1A OR9 (Received December 28, 1977) Publication costs assisted by the National Research Council of Canada
Absolute rate constants of hydrogen abstraction from CH,OH, C2HSOH, 1-C3H70H,and 2-C3H70H by OH radicals at 296 f 2 K have been measured in the gas phase using the flash-photolysis resonance-absorption technique. The influence of secondary reactions on the measured rates has been discussed and assessed. The H f 0.22 rate constants in units of cm3 mol-'^-^ were found to be: ~ C H ~ O=H 0.64 f 0.06 x lo1', ~ C ~ H ~=O2.25 x lo1', hl.CsH,OH = 3.21 f 0.32 X 10", and k 2 . c 3 ~ ?=0 ~3.30 f 0.33 x 10".
Introduction There is considerable interest in the reactions of the hydroxyl radical with organic molecules because of the important role that this radical plays in the chemistry of the atmosphere and in combustion processes. Indeed, Pith and co-workers2have formulated recently a reactivity scale for hydrocarbons based on their rate of reaction with the hydroxyl radical. As a result, a large effort has been directed recently toward obtaining accurate rate constants for OH radical reactions. However, little has been reported to date on the rates of hydrogen abstraction by OH from alcohols in the gas phase. To our knowledge there has been no measurements of absolute rate constants; all previous consist of competitive determinations of relative rates. The two values reported for the reaction with m e t h a t ~ o lvary ~ , ~ by a factor of 5-10 depending on the value of the rate constant of the reaction OH + CO, which is used as a standard in one of the studies3 and which has been found recently to be pressure dependent.6 The reactions of OH with alcohols should be of atmospheric interest, as alcohols are used in commercial solvents and are seriously considered as transport fuels in mixtures with gasoline. The reactions are also of interest to chemical kinetics; a large collection of hydrogen abstraction rate data should enable better understanding and eventual prediction of the rates of such reactions. In addition, the abstraction of hydrogen from an alcohol by OH would be expected to 0022-3654/78/2082-1329$01 .OO/O
result in the same or very similar radical product as is formed by addition of OH to the corresponding olefin. It should be interesting, therefore, to test whether these radical products are involved in secondary reactions with OH and thus gain a better understanding of possible similar secondary reactions in the case of the faster OH + olefin r e a ~ t i o n . ~ In the present paper we report values of the absolute rate constants of the reactions of OH radicals with CH30H, CzH50H,1-C3H70H,and 2-C3H70H;these were measured using the flash-photolysis resonance-absorption technique described earlier.s We have also assessed the effect of secondary radical-radical reactions involving OH on the measured rate constants and established the approximate magnitude of such an effect for our experimental conditions.
Experimental Section The apparatus and experimental technique were described in detail before8 and will be mentioned here briefly. The apparatus consists of a fast flash-photolysis system coupled to a spectrophotometric detection system. The flash lamp is coaxial with the reaction vessel8 and was operated at 160 J. The OH concentration was monitored by following the time resolved attenuation of the OH resonance radiation (Q13or Q14rotational line of the (0,O) band of the A2Z+ X 2 n transition) produced by a mi-
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0 1978 American Chemical Society
1330
The Journal of Physical Chemistry, Vol. 82, No. 72, 1978
crowave discharge in a low pressure Ar/HzO mixture. The radiation was detected with a photomultiplier (EM1 9783B) mounted a t the exit slit of a 1-m Czerny-Turner monochromator (Jarrell-Ash). The signal was amplified with a current-to-voltage converter employing a fast-settling operational amplifier and was then filtered with a variable low-pass electronic filter which has a four-pole Bessel response (Ithaco 4210). The cutoff frequency of the filter was set a t 3 to 4 times the rate to be measured thereby providing adequate filtering of high-frequency noise with tolerable distortion of the rise to maximum OH absorption. The filtered transient signal was recorded in digital form by a fast transient recorder-analog to digital converter (Biomation 610B transient recorder) interfaced with a signal averager (Nicolet Model 1074),where the signal was accumulated. The transient recorder has an amplitude resolution of 1 part in 64 and a time resolution of 256 channels. Signal averaging improved the signal-to-noise ratio;8 rate measurements in the range 0.5-8 X lo4 s-l were made employing 2-64 flashes. The contents of the signal averager memory, consisting of values of Al = Io - I at various time intervals, were transferred through an interface to an on line P D P 11/10 minicomputer for further processing. In the majority of the experiments, the source of hydroxyl radicals was photolysis of small pressures of water vapor in the vacuum UV (A >160 nm, Le., the Suprasil window cutoff). Experiments were also made using the alternative source of OH radicals, i.e., photolysis of N20/H2 mixtures. More details on this technique will be given below. The materials were Fisher Spectranalyzed methanol and 2-propanol, Fisher certified 1-propanol, and absolute ethanol, all thoroughly degassed before use. Analysis of the ethanol by gas chromatography and UV spectrophotometry showed no detectable impurity; Matheson Ultrahigh Purity Hz was used as is; Matheson N20,stated purity 9890 minimum was used after degassing and three bulb-to-bulb distillations in vacuo in which only the middle half was retained; Canadian Oxygen He, stated purity 99.9990 minimum, was used directly from the cylinder.
R. Overend and G. Paraskevopoulos
CHANNEL Figure 1. Oscilloscope tracing of the digital data accumulated in the signal averager after 32 flashes for OH 1-C3H70H. The experimental conditions were as follows: sweep rate = 1 jdchannel; H,O = 0.8 Torr, [l-C3H,0H] = 9.97 X 10'' mol ~ m - ~ .
+
I.00-
0 a
0%
o""8,
'
Results Various aspects of the experimental conditions procedure and treatment of the data are given in ref 8. Mixtures of HzO/alcohol/He were prepared in large mixing bulbs from which several aliquots could be transferred to the reaction vessel. The effect of multiple flashes (Le., accumulation of reaction products) on the measured rates was tested a t the outset, for every alcohol, by flashing a given mixture 1, 2, 4, 8, and 16 times and comparing the resulting rates. No significant difference was detected; nevertheless, a maximum of 8 flashes, and usually only 2-4 flashes per mixture were used. In addition, the condition of minimum OH concentration (Le., low flash energy of 160 J and low H 2 0 pressure -0.8 Torr) were used in this work. Under these conditions undesirable processes such as secondary reactions of free radicals with OH are kept a t a minimum. The OH concentration was estimated approximately, as described before,8 from the initial absorption of the resonance radiation in experiments in which only H20 in He was flashed. In the absence of another reactant the decay of OH, due to the reaction OH + OH, is very slow and the initial absorption was estimated without a long extrapolation of the decay curve. It was found that the initial absorption, In (Zo/Ztr), so obtained, varied linearly with the pressure of H 2 0 and it was quite reproducible. The OH concentration was estimated to be 5 to 10 x lo-" mol cm-3
-
0
20
40
TiME
60
80
100
(PSI
Figure 2. Example of computer plots of (a) normalized [OH], and (b) In [OH] against time for OH 4- l-C3H70H. The experimental conditions were the same as in Figure 1.
depending on the type of experiment; the actual values will be given together with the corresponding experiments. The concentration of the alcohol was always much greater than the initial OH concentration, so that the reaction was pseudo-first order. Experiments were made a t 150 Torr pressure. An oscilloscope tracing of the raw data (i.e., values of hZ = I , - Z a t various time intervals) after 32 flashes is shown in Figure 1. As previously,Rthe data were smoothed using a nine-point quadratic approximation suggested by Savitzky and G01ay;~the OH concentration [OH] = In (Z0/ZtJ expressed in normalized units (i.e., as fraction of the concentration a t maximum absorption), and the In [OH] = In (In (ZJIJ) were calculated and plotted against the
Rates of
OH
Radical Reactions
The Journal of Physical Chemistry, Vol. 82, No. 12, 1978 1331
Y
7
* X
0 246810 8,Of
6.0
20
-
5.0 -
4.0 3.0
-
i'
-
2.0-
/+
'.$i
0 2 4 6 810 20 ALCOHOL CONCENTRATION
ALCOHOL CONCENTRATION x lo9 (mol cm-3)
Flgure 3. Plots of
pseudefirst-order rate constant k , against the alcohol concentration for the reactions of OH with CH30H and C2H,0H. TABLE I ? Measured Values of the Second-Order Rate Constant, k11, for t h e Reaction OH + ROH a t 296 t 2 K hII x l o - ' * (cm' mol" s-I) Reactant R O H
?lo
CH,OH C,H,OH l-C,H,OH 2-C ,H,OH N,O/H,/C,H,OH
0.64 t 0.05 2.40 i 0.08 3.21 r 0.18 3.30 ? 0.15 2.25 t 0.12
No. of experiments 11
14 14 10 8
a The experimental conditions were as follows: Flash photolysis of H,O; H,O = 0 . 8 Torr, the approximate OH concentration estimated from the initial absorption of the resonance radiation* was [ O H ] = 9-10 X lo-''mol cm-,. Flash photolysis of N,O/H,; the ratio N,O/H, was constant and equal t o 0.1 (Le., [ N , O ] = 2.8 X lo-' mol ~ m - ~ , [ H , ] = 2.6 X l o u 6mol cm-'), [ O H ] = 5 x lo-''mol ~ m - ~ .
time by the computer. Such plots are shown in Figure 2. The pseudo-first-order plots (Figure 2b) were linear over 2-4 half-lives depending on the conditions. The pseudo-first-order rate constant for the OH decay, kI, was obtained from the slopes of the linear plots as described before;8 the values of hI obtained over a range of alcohol concentrations are related to the second-order rate constant, hU,by the equation k1= LY + hIIIROH]. Plots of kIagainst the alcohol concentration are shown in Figures 3 and 4,for CH30H, C2H,0H, 1-C3H70H,and 2-C3H70H, respectively. The second-order rate constants, kII, were evaluated from the slopes of these plots by least-squares fits of the equation kI = a kIIIROH] to the points with weights of 1 / 2 , where r is the error in hI evaluated as described before.sJO The measured values of hn from these experiments together with the experimental conditions are given in Table I (first four rows). The possibility of some unspecified influence of H 2 0 on the measured rates of C2H50Hwas examined in experi-
+
OH+ 1 -C3H70H
x109 (mol cm-3) pseudo-first-orderrate constant k , against the alcohol concentratlon for the reactions of OH with 1-C3H,0H and 2-C3H,0H. Figure 4. Plots of
ments in which OH was produced in the absence of H 2 0 by flashing N20/H2/C2H50Hmixtures as mentioned earlier. The following reactions take place:
-
+ hu O ( ' D ) + N, O ( ' D ) + H, O H + H O H + H, H,O + H OH + C,H,OH products N,O
+
-t
+
From the relative rates of the reactions of O('D) with H211and N20,12and taking the rate constant of O(lD) + C2H50H to be equal to that of O('D) C2H613J4(Le., ~ H ~ : ~ N ~ ~ := ~ 1.00:0.571:2.02), c ~ H ~ o H and the concentrations used (Le., in mol cm-3, H, = 2.60 X lo-", N20 = 2.8 X C2H50H= 3.4-16 X it was estimated that virtually all the O(lD) atoms will react with H, and not with N 2 0 or C2H50H. The pseudo-first-order rate constant kI, measured a t constant H2 concentration and over a range of C2H50Hconcentrations, is related to the second-order rate constants hII of H2 and C2H,0H by the equation
+
kI = kII"Z)[ H2 3
+ kII(C,H50H)[C2 H, OH]
For a constant H2 concentration and a range of C2H50H concentrations, a plot of hI against [C2H50H]should be linear with a slope equal to hII(C2H50H) and an intercept equal to ~ I I " ~ [ H ~Such ] . a plot is shown in Figure 5. The data were fitted to the equation above by least squares with so obtained, weights = l/a2. The value of hIl(C2H50H), together with the experimental conditions, is given in Table I (last row). From the intercept, which may be subject to a larger error than that indicated by the standard deviation, we calculated kII(H2)= 3.85 f 0.60 X lo9 cm3 mol-' s-' which agrees well with 3.5 f 0.16 X lo9 cm3 mol-' s-l obtained in our previous direct determinations
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The Journal of Physical Chemistry, Vol. 82, No. 12, 1978
7
R. Overend and G. Paraskevopoulos
-0.8 T
2
0
4
6
8
13
12
C;?H~OHCONCENTRATION x
14
16
18
io9(mol cm-3)
Figure 5. Plot of pseudefirst-order rate constant, k,,against the C2H50H concentration in the absence of H20 (OH produced by flashing N20/H2/C2H50Hmixtures).
Discussion Consider the following reactions: H,O + hu OH + H --+
+ R O H - .R’OH + H,O OH + .R‘OH products OH + OH + M - H,O, + M OH + H + M - H,O + M OH + O H - H,O + 0 OH
+
where .R’OH is the radical resulting when hydrogen is abstracted from the alcohol ROH by OH. In our measurements of the pseudo-first-order rate constants, hl, of the OH decay, possible sources of systematic error are the radical-radical reactions involving OH. We will first examine if (and to what extent) such reactions affect our measured values of hI and then we will compare our values with values in the literature. Effect of Radical-Radicab Reactions Involving OH. In our earlier work8 the differential equations describing the system were integrated numerically for various initial conditions and it was shown that for measured pseudofirst-order rate constants, hI, larger than 5 X lo3 s-l the influence of reactions 3 , 4 and 5 is negligible. We are left then with the influence of reaction 2 on reaction 1; for pseudo-first-order conditions [ROH] >> [OH] and hI = hlIIROH], the differential equations for reactions 1 and 2 are - d [ O H ] / d t = kI[OH]
+ kz [ O H ] [ * R ‘ O H ]
d [ * R ’ O H ] / d t= kI[OH] - hz [ O H ] [’R’OH] Let hI(apparent) be the measured overall rate of the OH decay (including a possible contribution from reaction 2), then, it may be seen that the value of hl(apparent) lies between hI and 2121 depending on the parameter k2[OHIo/hI ([OH], = the initial concentration of OH); when the parameter is very small hI(apparent) = hl, whereas when it is large the radical .R’OH is in steady state and hl(apparent) = 2hI. The ratio hI(apparent)/hl may be estimated by numerical integration of the equations above for various initial conditions8 (Le., values of [OH],, h2, and kI). We have calculated values of the ratio for each one of our experiments for the following conditions: (1) The experimental values of [OH], and hI were used. (2) The rate of reaction 2 was taken initially to be the collisional rate h2 = 2 x 1014cm3 mol”’ s-l. (3) The calculation was made
CH, OH CONCENTRATION x lo9(mol cm-3) Figure 6. Effect of reaction 2 on k,,for methanol: lines a and b result from least-squares fits of (a), k,(apparent), and (b), k,(corrected) for k, = 2 X l O I 4 cm3 mol-’ s-‘.
for OH decays over the range 0.1 < [OH],/[OH] < 0.9, these limits correspond to realistic experimental conditions, Le., missing record between t = 0 and some later time due to electrical noise or “blinding” of the photomultiplier and limited range of observations of 2-4 half-lives. From the calculated values of the ratio, hl(apparent)/hI, and the measured values of hl(apparent) we estimated corrected values of hl which were then used to obtain corrected values of kII by least-squares fits of the equation hl = a + hl,[ROH] with the same weights as those used in fitting the experimental points. This is illustrated for methanol in Figure 6 in which lines a and b result from the least-squares fits of the measured and corrected values of hl, respectively. The following comments may be made: For h2 = collisional rate and measured values of hI larger than about lo4 SI, the resulting line a is nearly parallel to line b, the difference in the slopes is 11% . On the other hand for small values of hl (dashed line) the slope is significantly larger than that of line b. In view of the comments above, the small intercept resulting from the fit has no obvious physical significance; in the two limiting cases (i.e., hI(apparent) = h I and hl(apparent) = 2hI) the intercept should be zero. The corrected values of hU,so obtained, are (in cm3mol-’ s l): CH30H, 0.57 f 0.05 X 10l2;C2H50H,2.13 f 0.08 X 10l2;and 2.21 f 0.12 X 10l2for OH from H 2 0 and N20/H2 photolysis, respectively; 1-C3H70H,2.85 f 0.16 X 10l2; 2-C3H70H,3.03 f 0.10 X 10l2. The systematic error that corresponds to these values is 11, 12, 11, and 8% for CH30H, C2H50H,1-C3H70H,and 2-C3H70H,respectively. Since these values were calculated using h2 = collisional rate, the error above should represent the maximum systematic error due to reaction 2. A similar calculation using hz = 1014cm3 mol-’ s-l, Le., half the collisional rate, gave systematic errors that were smaller than the random errors shown in Table I. Comparison of Rate Constant Values. Table I1 lists the values of hll obtained in this and other work in the literature. The error limits of all our measurements were taken to be &lo%; these limits include a possible systematic error due to reaction 2, as discussed above. The two values of the rate constant of C2H50H(i.e., 2.25 f 0.08 x 10l2and 2.40 f 0.08 x 10l2),obtained by using the two sources of OH, are not different statistically. We have taken the former in Table 11,because in this determination, in which the photolysis of N 2 0 / H 2was the source of OH radicals, the initial OH concentration was half that used in the other determination and, therefore, the resulting value of the rate constant is subject to smaller systematic
The Journal of Physical Chemistry, Vol. 82, No. 12, 1978
Reactions of 0 Atoms with Sulfides and Thiols
TABLE 11: Comparison of t h e Room T e m p e r a t u r e R a t e Constants for t h e R e a c t i o n OH t ROH D e t e r m i n e d by D i f f e r e n t Techniques
hII x Reactant CH,OH
( c m 3 mol-' s-I)
This w o r k
Lit, value
Ref
0.64f 0.06
0.57 f 0.06 0.12 f 0.02
4a 3b 16' 4a
0.90 C,H,OH l-C,H,OH
2-C3H,0H
2.25
t
0.22
* 0.32 3.30* 0 . 3 3 3.21
1.8 t 0 . 2 1.85 2.3 * 0 . 2 2.8
16'
4a 16'
4.3 f 1.3
5d
2.15
16'
a Relative to t h e r a t e constant OH + n-butane, for w h i c h a value of 1.4 X 10" c m 3 mol-' s-' has been used. Relative to t h e r a t e constant of OH w i t h CO, for w h i c h w e used t h e h i g h pressure value of 1.9 X 10" c m 3 mol-' s-lu6 Liquid phase d e t e r m i n a t i o n ; OH was generated by pulse radiolysis o f H,O. R a t e constants are m e a n values r e p o r t e d in ref 16. T h e r a t e constant was d e t e r m i n e d a t 305 K,relative to t h e r a t e constant of OH isobutene a t 306 K,for w h i c h a value of 3.05 X l O I 3 c m ' mol-' s-' has been used.
+
error from the secondary reaction 2. As mentioned in the Introduction all the other studies are relative determinations. Our values agree well with those of Campbell et ala4for CH30H and C2H50H,while for 1-C3H70Htheir value is -30% lower than ours; they used a rate constant for the reaction of OH + n-butane of 1.4 X 10l2cm3 mol-' s-l, if the more recent value of 1.65 X 1OI2 cm3 mol-' s-l is used,15 very good agreement is obtained for all alcohols. Our value for 2-C3H70H is somewhat lower than, but within the wide error limits of
1333
the value reported by Lloyd et aL5 On the other hand the value of Osif et al.3 for CH30H is lower than all the other determinations by a factor of about 5. It is interesting that the values of the gas-phase rate constants are, with the exception of 2-propanol, very close to those determined in the liquid phase,16 which are also given in Table 11.
References and Notes (1) Issued as N.R.C.C. No. 16591.
(2j K. R. Darnall, A. C. Lloyd, A. M. Winer, and J. N. Pitts, Jr., Envifon. Sci. Techno/., 10, 692 (1976). (3) T. L. Osif, R. Sirmaitis. and J. Heicklen, J. photochem., 4, 233 (1975). (4) I. M. Campbell, D. F. Mdaughlln, and B. J. Handy, Chem. phys. Left., 38. ... 362 ..-(19761. (5) A. C. Lloyd, K,'R. Darnall, A. M. Winer, and J. N. Pins, Jr., Chem. f h y s . Left.,42, 205 (1976). (6) R. Overend and G. Paraskevopoulos, Chem. fhys. Lett.,49, 109 (1977). (7) R. Overend and G. Paraskevopoulos, J. Chem. phys., 67,674 (1977). (8) R. P. Overend, G. Paraskevopoulos, and R. J. Cvetanovic, Can. J . Chem., 53, 3374 (1975). (9) A. Savltzky and M. J. E. Golay, Anal. Chem., 36, 1627 (1964). (10) R. J. Cvetanovlc, R. P. Overend, and G. Paraskevopoulos,Proceedings of the Svrnwsium on Chemlcal Kinetics Data for the Lower and UDW Atmosher'e, Warrenton, Va., Sept 15-18, 1974; Int. J. Chem. klriet. Symp., 1, 249 (1975). (1 . 1). G. ParaskevoDoulos and R. J. Cvetanovic, J . Am. Chem. Soc., 91, 7572 (1969): (12) G. Paraskevopoulos, V. 8. Symonds, and R. J. Cvetanovic, Can. J. Chem., 50, 1838 (1972). (13) P. Mlchaud, G. Paraskevopoulos, and R. J. Cvetanovlc, J . Phys. Chem., 78, 1457 (1974). (14) The assumption that the rate constants of the reactions of O('D) with C&OH and CPHBare approximately equal is Justifiedfor the purpose of this estimate because kCehis close to the collision frequency; the estimate is stili valid for a value of the rate twice as large as that used. (15) R. A. Perry, R. Atkinson, and J. N. Pitts, Jr., J . Chem. fhys., 64, 5314 (1976). (16) R. L. Wilson, C. L. Greenstock, G. E. Adams, R. Wageman, and L. M. Dorfman, Int. J . Radiat. fhys. Chem., 3, 211 (1971).
Study of the Reactions of Oxygen Atoms with Hydrogen Sulfide, Methanethiol, Ethanethiol, and Methyl Sulfide Irene R. Slagle, Frank Balocchl, and David Gutman' Department of Chemistry, Iiiinois Instifufe of Technology, Chicago, Illinois 606 16 (Received November 17, 1977; Revised Manuscript Received March IO, 1978) Publication costs assisted by the I//inoisInstitute of Technology
Rate constants were measured as a function of temperature between 250 and 500 K for the reactions of oxygen atoms with four divalent sulfur compounds: hydrogen sulfide, methanethiol, ethanethiol, and methyl sulfide. The purpose of the study was to gain additional information on the mechanisms of these reactions. The trend in the rate constants obtained for the 0 + CH3SH, C2H5SH,and CH3SCH3reactions support an addition mechanism analogous to that for 0 + olefin reactions. T h e 0 + H2S rate constants give no clear indication of the reaction mechanism.
Introduction In an earlier study of the reactions of oxygen atoms with organic compounds containing reduced sulfur (mercaptans and sulfides), we suggested that this class of reactions proceeds by the electrophilic addition of the oxygen atom t o the sulfur contained in the compound to form an energy-rich complex which could either stabilize through subsequent collisions or unimolecularly decompose.' The evidence cited to support the proposed mechanism included the direct detection and identification of the adducts formed by some of these reactions and a noted linear 0022-3654/78/2082-1333$01 .OO/O
relationship between the log of the room temperature rate constants for this set of reactions and the ionization potential of the sulfur-bearing reactant. We have sought to extend our observations on the reactivity of oxygen atoms with sulfides and thiols in order to obtain additional information on the mechanisms of these reactions. Specifically we have now measured the temperature dependencies of the rate constants of the 0-atom reactions with H2S, CH3SH, C2H5SH, and CHBSCHBin order to obtain the Arrhenius parameters of the respective rate constants. The data and the infor-
0 1978 American
Chemical Society
