Radiation Chemistry

15 Pulse Radiolysis Studies. XIII. Rate Constants for the Reaction of Hydroxyl

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Radicals with Aromatic Compounds in Aqueous Solutions P. N E T A and LEON M . D O R F M A N The Ohio State University, Columbus, Ohio 43210

Absolute rate constants have been determined for the reac tion of the hydroxyl radical with a variety of aromatic com pounds in aqueous solution. The rate constants obtained are significantly higher than values previously reported. Rate constants for the reaction of the hydroxyl radical with methyl alcohol and ethyl alcohol have also been determined by competition kinetics using three of these absolute rate constants as reference values. Comparison of our results with the published values from competition kinetics suggests that the rate constants for the reaction of hydroxyl radicals with iodide ion and thiocyanate ion are significantly higher than reported in earlier work. The ultraviolet absorption bands of the various substituted hydroxycyclohexadienyl radicals formed have been observed.

Absolute rate constants for the reaction of hydroxyl radicals with aromatic compounds in aqueous solution have been determined i n a number of pulse radiolysis investigations (6, 10, 11, 15). The rate con stants were obtained directly from the formation curves of the hydroxycyclohexadienyl adduct free radical. Recent observations i n this (22) and other laboratories (6, 15, 18) have shown that the cyclohexadienyl radical formed by Η-atom addition has an ultraviolet absorption band which overlaps that of the hydroxycyclohexadienyl radical and is equally intense. The rate constants for the addition of hydrogen atoms to benzene and toluene (18), to phenol (15) and to benzoic acid (16) i n water 222 In Radiation Chemistry; Hart, E.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

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Hydroxyl Radicals


appear to be significantly lower than the rate constants for hydroxyl radical addition. It appeared worthwhile, therefore, to determine absolute rate constants for OH-addition for benzoate ions, benzene and a number of other aromatic compounds from the formation curves in systems in which the hydrogen atoms have been essentially eliminated or at least sharply reduced in concentration. Solutions containing a sufficient con centration of nitrous oxide to scavenge the hydrated electron and convert it to hydroxyl radical at p H 6 to 9 have been used for this purpose. Relative reactivities of organic compounds toward hydroxyl radicals have also been determined by competition kinetics using carbonate ions ( I ) , thiocyanate ions (J, 2, 19), and iodide ions (21) as the reference reactant. However, the values reported for carbonate differ by a factor of two (1, 21, 23) and the values for both iodide and thiocyanate have been questioned (7,12) since there is a possibility that diiodide formation rather than iodine atom formation may be rate determining (12) and the analogous possibility has been suggested (7) for thiocyanate. There appears to be a need for further work to establish accurate values for a number of rate constants which may serve as reference values. Relative reactivities of substituted benzenes and benzoate ions toward hydroxyl radicals have been correlated (4) with Hammett's equation. In the present work such a correlation has also been carried out using absolute rather than relative values. In the course of this work the absorption spectra for the various substituted hydroxycyclohexadienyl radicals have been determined. Experimental The detailed experimental technique, using a Varian V-7715A linear accelerator has been outlined (13). 3.5 to 4 Mev. electrons were used with pulse duration of 50 to 100 nsec. during most of the work, although occasionally pulses as long as 0.5 /xsec. were used. The pulse current was in the range 310 to 330 ma. The dose with the 50 nsec. pulse at maximum current is 3 X 10 e.v./gram as determined with the modified Fricke dosimeter previously discussed (9). Taking G H + G - = 5.2 molecule/ 100 e.v., this gives an initial O H concentration of 3 X 10" M. The detec tion system used was for the most part as described (13), but a few runs in cases which exhibited higher rate constants were done with a nano second detection system (20) similar to that designated by Hunt and Thomas (14). The irradiation cell was 2 cm. long, using a double pass of the analyzing light beam and 0.8 cm. deep in the direction of the electron beam. The concentration range of the aromatic compound in each case was at least threefold and at least four runs were done at each concentration. Depending upon the rate constant the concentrations varied from 5 X 10~ to 5 X 10" M. The carboxylic acids were neutralized with sodium 1G

0

c

a q

6

δ

4

In Radiation Chemistry; Hart, E.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

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1

hydroxide to p H 6 to 8, and some were also irradiated at p H 9.4 in borax buffered solution. In the case of volatile compounds of low solubility such as benzene and anisole, the solution was made up as follows. The stock solution was made up by pipetting the solute into a large volumetric flask in which the remaining gaseous volume was very small compared with the liquid volume. The water was first deaerated by bubbling N 0 through it and was thus also saturated with N 0 at atmospheric pressure. The stock solution was then pipetted in and the final reaction solution transferred to the cell using a large syringe. It would appear that any non-random uncertainty in the solute concentrations would tend to give a rate constant on the low side. The nitrous oxide was passed through two successive alkaline pyrogallol solutions and finally through triply distilled water. The aromatic compounds were either Baker Analyzed reagents or Baker Grade reagents. 2


2

Results and Discussion Under our experimental conditions the hydrated electron is effectively scavenged by nitrous oxide in competition with the reactions with hydrogen ion and with the aromatic molecule: e-aq

+ N 0 = N 02

(1)

2

A t one atmosphere pressure the N 0 concentration exceeds ! 0 ~ M and & i [ N 0 ] is thus much greater than fc - +H [ H ] at p H 6 to 9 and ^e- f aromatic [aromatic] at the solute concentrations used. There is ample evidence (3, 13, 17) that the lifetime of N 0 " is very short and that O H is rapidly formed (17) either through O " formation followed by protonation involving water or by direct protonation of the nitrous oxide anion involving water. The yield of hydroxyl radical is approximately doubled, and the yield of hydrogen atom is thus about one-tenth the total hydroxyl radical yield under these conditions. The system is then very nearly a one-radical system in which the hydroxycyclohexadienyl radical is the observed species formed in the reaction: 2

2

2

E

+

NQ

ag

2

OH + C H COO = (OH)C H COO6

5

0

5

(2)

The nitrous oxide does not interfere with the hydroxyl radical kinetics as is shown in this and earlier work (17). Absorption Spectra. The absorption spectra of the various hydroxycyclohexadienyl radicals, in accord with earlier work (6, 8, 10, 11, 15), show a strong band in the ultraviolet. The absorption bands for eleven different compounds have been determined in this investigation and are shown in Figure 1. The wavelength corresponding to the maximum for each band is also given in Table I. The contribution of the hydrogen atom adduct to these absorption bands is very small, and its effect may be neglected in view of the low yield of hydrogen atoms although the


15.

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AND

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225


hydrogen adduct spectrum is similar i n its wavelength maximum as well as extinction coefficient (6, 15, 22). Calculation of the bathochromic shifts showed a fair agreement with the results of Chutny (8).

Figure 1. Ultraviolet absorption bands of various hydroxycyclohexadienyl radicals in aqueous solution containing nitrous oxide at neutral pH. The ordinate gives the relative optical density and the abscissa the wavelength in n.m.

Absolute Rate Constants. Absolute rate constants for the hydroxyl radical reactions, as determined from the formation curves of the hydroxycyclohexadienyl radicals, are summarized i n Table I. Detailed data for benzoate ion are shown in Table II. In all cases the rate curves fit closely to a first order rate law. A detailed examination of this case seems warranted not only as an example of the data, but because of the possible use of this reaction as a reference reaction i n competition kinetics.


226

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As mentioned in the experimental section, the initial O H concentration with the 50 nsec. pulse is 3 X 10" M. Thus, the ratio [ O H ] / [ C H C O O ~ ] is 0.03 and 0.06 for the 50 nsec. and 100 nsec. pulse respectively at 1 X 10" M benzoate ion, the lowest concentration used. Since k o n o H = 1-2 X 10 M s e c . (17), the initial rate for O H recombination is 0.07 and 0.13 of the OH-addition rate for this initial OH-concentration, and falls to considerably less over the region of rate curve analysis. The effect on the observed rate constant is thus on the order of 1 to 2 % for the lowest benzoate concentration. 6

6

0

5

4

10


+

1

1

Table I. Absolute Rate Constants for the Reactions of Hydroxyl Radical with Aromatic Compounds

Compound

Wavelength Maximum of OH-Adduct Radical (nm.)

Absolute Rate Constant (M' sec.' X 10~ ) 0

pH

1

1

ko

l

H

+

PhZ

OH + PhH

9

K

C H N0 C H,SO C H CN C H,COOH C H,COO C H COCH C H CH COOC H C H,CH OH CoH.OCH^ C H OH

410 315 348 347 330 372 325 313 320 330 330

7 7 7 3 6-9.4 7 6-8 7 7 7 7.4-7.7

3.2 4.7 4.9 4.3 6.0 6.5 7.9 7.8 8.4 12 14

± ± ± ± ± ± ± ± ± ± ±

0.4 0.6 0.6 0.8 0.7 0.7 1.1 1.1 1.2 3 3

-0.39 -0.22 -0.20 -0.26 -0.11 -0.08 -0.005 0.00 0.03 0.19 0.25

C H COOp-N0 C H COOp-ClC H COOp-OHC H COOp-OC H COO-

330 420 345 375 425

6-9.4 6-9.4 6-9.4 7 9.4

6.0 2.6 5.0 9

± ± ± ± —

0.7 0.4 0.8 2

0.00 -0.36 -0.08 -0.17 —

6

5

2

6

G

s

5

e

6

6

6

5

6

5

6

s

2

6

6

6

G

2

b

5

5

2

6

6

4

4

6

f)

4

4

Taken from Ref. 22. Taken from Ref. 15. The correction for the effect upon the value for the OH-addition rate constant ( as determined graphically ) of the two reactions O H + aromatic and O H + O H has been estimated using an Electronic Associates Incorporated, Model TR-20 analog computer, and is included in this table. It ranges from +10% to + 8 % with decreasing concentration of the aromatic compound, the correction for O H + O H ranging from 0 to -2%. a

ft c

No p H effect was found for the reaction with benzoate ion i n the region of p H 6-9.4, as may be seen in Table II. This was also found to be the case for the other compounds listed. It should be noted that the rate constants for benzoic acid (22) and benzoate ion show a slight


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227

Hydroxyl Radicals

difference. In the case of benzoate ion the same rate constant was obtained when chloroform (at a concentration five times that of benzoate) was used as a scavenger for the hydrated electron instead of nitrous oxide. This demonstrates that the presence of N 0 has no effect upon the rate constant as has been shown in another reaction investigated by Rabani and Matheson (17). 2


Table II.

Absolute Rate Constant for the Reaction of Hydroxyl Radicals with Benzoate Ion

Benzoate Ion Concentration (M X 10 )

Scavenger

h

1.0 1.0 2.0 2.0 2.0 2.5 2.5 3.0 3.0 3.0 4.0 5.0 5.0 5.0 average

N 0 (1 N 0 N 0 N 0 CHC1 N 0 N 0 N 0 N 0 CHC1 N 0 N 0 N 0 CHC1 2

atm.)

2 2

2

3

(ImM)

H

(1.5mM)

H

(2.5mM)

2

2

2 2

2 2

2

pH 9.4 6 6 9.4 6 6 9.4 6 9.4 6 6 6 9.4 6 6-9.4

Rate Constant (M- seer X 10~ ) 1

1

9

6.6 6.9 7.2 6.6 6.6 5.2 6.1 6.5 5.9 5.8 5.3 4.8 5.4 5.6 6.0 ± 0.7

The absolute rate constants determined in this investigation are significantly higher ( see Reference 5 for comparison ) than those determined by competition kinetics with iodide ion ( 21 ) and thiocyanate ion ( 2 ) as well as those determined absolutely by formation kinetics (10, 11). In the latter case this may be understood from the fact that the addition of hydrogen atoms to the aromatic ring, as has been shown for benzene and toluene (18) and benzoic acid (16) has a much lower rate constant than the addition of hydroxyl radicals. Since the H-adduct free radical has a similar absorption, and is formed in comparable amount in acidic solution (11) in the absence of an electron scavenger, its observation concurrently with the OH-adduct would give an apparent lower rate constant. Our value for nitrobenzene is, however, lower than the recent value of Asmus, et al. (6). W i t h respect to results from competition kinetics, the difference shown in the results of this investigation indicates that there may still be some uncertainty in the value of the rate constants of the reference reactions used. A plot of log * ° + * from these H

P h

*OH + PhH


228

RADIATION CHEMISTRY 1

absolute values for monosubstituted benzene and p-substituted benzoate ions as well as the value for phenol (15) and for benzoic acid (22) against the Hammett σ values is shown in Figure 2. This correlation has been discussed (4). The ρ value obtained from the slope is —0.5, in good agreement with th évalue —0.4 previously obtained (4). The contribution of the aliphatic side chain to the overall rate appears to be very


Small, Comparing "OH+acetone (5).

&OH+aromatic

Figure 2.

With

&0H+methyl alcohol,

^OH+acetate

ΟΓ

Plot of correlation with Hammett's equation

para, #, and meta, O, are used as discussed in Ref. 4

a

a

Competition Kinetics. Since the foregoing absolute rate constants have been determined independently, and appear to be accurate values, it is of interest to carry out competition kinetics with some reactants for which relative rate constants have been determined by competition kinetics with iodide ion and thiocyanate ion as the reference reactants. In this way, some assessment may be made of the validity of the existing values for the rate constants of the reactions: and

O H + I" = O H " + I

(3)

O H + SCN- = O H " + S C N

(4)



15.

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229

Hydroxyl Radicals

Rate constants for methanol and ethyl alcohol relative to those for ben zoate ion, phenylacetate ion and p-nitrobenzoate ion are shown in Table III. Each value in the table consists of experiments at five separate concentration ratios. The random uncertainty in each value is less than ± 1 0 % . In determining these rate constants from optical density ratios it was necessary to make a small correction for the contribution to the optical density by the H-adduct free radical. The molar extinction co efficients at 340-350 π\μ for the H-adduct and OH-adduct are similar for benzoic acid (22) and were assumed to be comparable for the other two aromatic ions in the table. The correction is necessary since the rate constants for the reaction of hydrogen atoms with the alcohols used are two orders of magnitude lower than the rate constants for hydrogen atom addition to the aromatic ring, while the analogous hydroxyl rate constants are roughly comparable. The data i n Table III show excellent self-consistency for the three aromatic rate constants, the individual values of which are nevertheless quite different. Comparison with rate constants determined by iodide ion competition kinetics (21) indicates that the absolute values deter mined in this work are higher for benzene, ethyl alcohol, and methyl alco hol. Comparison with the rate constants determined by thiocyanate ion competition kinetics (1,2) indicates that our absolute values are higher for benzene, benzoate ion, methyl alcohol and ethyl alcohol. This comparison indicates that the actual rate constants for Reactions 3 and 4 may be higher than the values which have been determined (J, 2, 21) from the formation curves of the optically absorbing product by as much as a factor of 1.6 to 2.3 in the former case and 1.7 in the latter case. That is, the true values may be nearer h = 2 Χ 10 and k = 1.2 Χ 10 M " sec." at 25 °C. This same qualitative conclusion for the case of thio cyanate has recently been reached by a different method by Baxendale and Stott ( 7 ) whose suggested value for fc is approximately 2 X 1 0 M sec." . The difficulties i n the direct absolute determination lie in the complexities of the kinetics and hence i n the interpretation of the forma tion rate curves. 10

A

10

4

1

1

1 0

4

_ 1

1

Table III. Rate Constants for the Reaction of Hydroxyl Radicals with Methanol and Ethyl Alcohol Determined by Competition Kinetics Reference Aromatic Compound C H COO" C H CH COOp-N0 C H COO average 6

5

e

5

2

2

6

4

koH+CHsOH (X lO-'M'

1

seer )

8.5 8.3 8.4 8.4 ± 1

1

18.5 18.3 18.1 18.3 ± 2


230

RADIATION

CHEMISTRY

1

Acknowledgment


This report is based upon work supported by the U.S. Atomic Energy Commission. The assistance obtained from the Graduate School of the Ohio State University i n supporting P. N . as a visiting scientist under the University Postdoctoral Program is gratefully acknowledged. W e are indebted to E . G . Wendell for his operation and maintenance of the linear accelerator and the electronic detection equipment. W e are grate ful to C . J. Geankoplis for his helpful advice concerning the use of the E A I analog computer. Literature Cited (1) Adams, G . E., Boag, J. W., Michael, B. D., Trans. Faraday Soc. 61, 1417 (1965). (2) Adams, G . E., Boag, J. W., Currat, J., Michael, B. D., "Pulse Radiolysis," p. 131, Baxendale, Ebert, Keene, Swallow, eds., Academic Press, New York, Ν. Y., 1965. (3) Adams, G . E., Boag, J. W., Michael, B. D., Proc. Royal Soc. (London) A289, 321 (1966). (4) Anbar, M . Meyerstein, D., Neta, P., J. Phys. Chem. 70, 2660 (1966). (5) Anbar, M . , Neta, P., Intern. J. Appl. Radiation Isotopes 18, 493 (1967). (6) Asmus, K. D., Cercek, B., Ebert M . , Henglein, Α., Wigger, Α., Trans. Faraday Soc. 63, 2435 (1967). (7) Baxendale, J. H., Stott, D . Α., Chem. Commun. 1967, 699. (8) Chutny, B., Nature 213, 593 (1967). (9) Dorfman, L . M . , Taub, I. Α., J. Am. Chem. Soc. 85, 2370 (1963). (10) Dorfman, L . M . , Taub, I. Α., Buhler, R. E., J. Chem. Phys. 36, 3051 (1962). (11) Dorfman, L . M . , Taub, I. Α., Harter, D . Α., J. Chem. Phys. 41, 2954 (1964). (12) Dorfman, L . M . , Firestone, R. F., Ann. Rev. Phys. Chem. 18, 177 (1967). (13) Felix, W . D., Gall, B. L., Dorfman, L . M . , J. Phys. Chem. 71, 384 (1967). (14) Hunt, J. W., Thomas, J. K., Radiation Res. 32, 149 (1967). (15) Land, E . J., Ebert, M., Trans. Faraday Soc. 63, 1181 (1967). (16) Neta, P., Dorfman, L . M. (to be published). (17) Rabani, J., Matheson, M . S., J. Phys. Chem. 70, 761 (1966). (18) Sauer, M . C., Jr., Ward, B., J. Phys. Chem. 71, 3971 (1967). (19) Scholes, G., Shaw, P., Willson, R. L., Ebert M., "Pulse Radiolysis," p. 151, Baxendale, Ebert, Keene, Swallow, eds., Academic Press, New York, Ν. Y., 1965. (20) Shank, N., Dorfman, L . M . (to be published). (21) Thomas, J. K., Trans. Faraday Soc. 61, 702 (1965). (22) Wander, R., Neta, P., Dorfman, L . M . , J. Phys. Chem. 72 (in press). (23) Weeks, J. L., Rabani, J., J. Phys. Chem. 70, 2100 (1966). RECEIVED

January 2, 1968.


Radiation Chemistry

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