OH-Initiated Oxidation of Imidazoles in Tropospheric Aqueous-Phase

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A: Kinetics, Dynamics, Photochemistry, and Excited States

OH-Initiated Oxidation of Imidazoles in Tropospheric Aqueous-Phase Chemistry Tamara Felber, Thomas Schaefer, and Hartmut Herrmann J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.8b11636 • Publication Date (Web): 30 Jan 2019 Downloaded from http://pubs.acs.org on February 2, 2019

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The Journal of Physical Chemistry

OH-Initiated Oxidation of Imidazoles in Tropospheric Aqueous-Phase Chemistry Tamara Felber, Thomas Schaefer, and Hartmut Herrmann* Atmospheric Chemistry Department (ACD), Leibniz-Institute for Tropospheric Research (TROPOS), Permoserstrasse 15, 04318 Leipzig, Germany

Abstract Imidazoles formed via the reaction of dicarbonyls with nitrogen containing compounds in the atmospheric particle-phase can be expected to initiate secondary organic aerosol (SOA) growth due to their potential to act as photosensitizers. Recent field studies quantified imidazoles for the first time in ambient aerosol samples from Europe and China. However, the knowledge about the kinetics and mechanisms of particle-phase reactions involving imidazoles is still very limited. In the present study, the radical-driven aqueous-phase oxidative degradation reactions of the hydroxyl radical (OH) with imidazoles were investigated. For the imidazoles following rate constants at 298 K and acidic conditions are obtained as: k(imidazole-2-carboxaldehyde) = (1.8 ± 0.1) × 109 L mol-1 s-1, k(1methylimidazole) = (2.3 ± 0.1) × 109 L mol-1 s-1, k(2-methylimidazole) = (3.9 ± 0.1) × 109 L mol-1 s-1, k(4-methylimidazole) = (3.2 ± 0.2) × 109 L mol-1 s-1, k(1-ethylimidazole) = (2.6 ± 0.1) × 109 L mol-1 s1, k(2-ethylimidazole) = (3.9 ± 0.1) × 109 L mol-1 s-1. Temperature and pH dependencies of the reactions as well as the activation parameters have been determined and discussed. The possible implications and restrictions for imidazoles acting as photosensitizers in tropospheric particles have been considered.

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Introduction Imidazoles in tropospheric aerosol particles are widely discussed in literature. Recent studies suggest that atmospheric particle-phase imidazoles might act as photosensitizers similar to what is known for surface water chemistry.1-3 Moreover, they might affect aerosol optical properties due to their contribution to Brown Carbon (BrC).4-6 Therefore, the interest in imidazoles and their role in atmospheric chemistry increased over the last years.4-7 Laboratory studies have shown that imidazoles can be formed as secondary products via the reaction of α,β-dicarbonyls such as glyoxal with nitrogen containing compounds such as ammonium ions, amino acids or amines8-16 thus contributing to SOA formation. Galloway et al.8 identified first the formation of imidazoles on ammonium sulfate seeds after the uptake of glyoxal and classified imidazole-2-carboxaldehyde (2-IC) as one of the products formed. Relating to the light absorption properties of glyoxal and ammonium sulfate mixtures, an absorption band at a wavelength of 280 nm was observed.9 Small imidazoles absorb at short wavelength (λ < 280 nm) and thus, in the near UV range.13,14 Kampf et al.14 further identified biimidazole (BI) as a major contributor to the 280 nm absorbance band and 2-IC as a minor contributor. Hence, imidazoles contribute to atmospheric BrC,17,18 which is defined as organic carbon which absorbs light in the near UV to visible spectrum of the solar spectrum.6 Imidazoles are also suggested to be potential photosensitizers.19-21 Photosensitizers absorb light in the range of solar actinic radiation. The excited photosensitizer might directly react in its excited state or transfers its excess energy to other molecules and thus, through either pathway, induces further reactions resulting in the oxidation of substrates22 and initiate SOA growth.19,23 2-IC has been shown to act as an efficient photosensitizer19 with regard to SOA growth,19,24 hydroperoxyl radical (HO2) formation,25,26 and the reaction with halide anions20 and isoprene.21 Separate from these laboratory-based studies, Teich et al. recently identified


2

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and quantified five imidazoles in ambient aerosol samples from different environment in Europe and China in a concentration range between 0.2 and 14 ng m-3.27 According to this study, imidazoles occur in particulate matter (PM) at sites with strong biomass burning influence and in more polluted air regions. It is important to emphasize that besides the secondary formation of imidazoles from the above-mentioned precursors, it also might be possible that imidazoles are directly emitted from sources such as biomass burning and thus be of primary origin. The identification of imidazoles in ambient aerosol samples motivates further studies relating to the chemistry of these compounds. Accordingly, temperature-dependent kinetic investigations

of

hydroxyl

radical

(OH)

reactions

regarding

the

oxidation

of

imidazole-2-carboxaldehyde (2-IC), 1-methylimidazole (1-MI), 2-methylimidazole (2-MI), 4-methylimidazole (4-MI), 1-ethylimidazole (1-EI), and 2-ethylimidazole (2-EI) were performed to study the degradation of imidazoles by particle-phase radical chemistry. The measurements of the present study intend to better characterize the atmospheric lifetime of imidazoles in atmospheric aqueous particles.

Experimental Setup Laser flash photolysis is an established method to measure the rate constants of reactions with transient species in aqueous solution.28-30 In the present study a thermostated laser flash photolysis-laser long path absorption (LFP-LLPA) setup was used, which has been described in more detail previously.28-35 The rate constants were measured by competition kinetic method using thiocyanate anion (KSCN; 2 × 10-5 mol L-1) as reference system.31,35-37 OH radicals were generated by the photolysis of hydrogen peroxide (H2O2; 2 × 10-4 mol L-1) with an excimer


3


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laser (COMPex 201, Lambda-Physik) at λ = 248 nm. The OH radical concentrations were in a range of 2 to 5 × 10-7 mol L-1 depending on the excimer laser energy. The concentration of the H2O2 (c = 0.2 mol L-1) and KSCN (c = 0.018/ 0.02 mol L-1) stock solutions were determined by UV/VIS spectrometry (Lambda 900 double beam spectrometer, Perkin Elmer). All kinetic experiments were carried out under first order conditions adding the current alkyl-substituted imidazole in excess to the formed radical with concentrations from 0 to 2 × 10-4 mol L-1 at pH = 2 or from 0 to 1 × 10-5 mol L-1 at pH = 9. For the measurements with 2-IC following concentrations were used: 0 to 4 × 10-5 mol L-1 at pH = 0, 0 to 1 × 10-4 mol L-1 at pH = 4.2, and 0 to 8 × 10-6 mol L-1 at pH = 9. The measurements were performed at temperatures in the range of 278 K ≤ T ≤ 318 K. Perchloric acid and potassium hydroxide were used for pH adjustment. The measurements were done in a high-purity silica cell, which is continuously flushed by the solution, with the dimensions: width = 4.0 cm, height = 2.0 cm, length = 3.5 cm, volume = 28 mL. A White cell mirror configuration was adjusted for 12 passes giving an absorption path length of 48 cm. The temporal change of the radical concentration was measured by two different continuous wave lasers at the wavelengths of λ = 407 nm (Laser LSR 407 nm Model 0222-583-00, Coherent) under acidic conditions or λ = 594 nm (LHYR-0200, Laser 2000) under alkaline conditions and was detected by a photodiode (S1336-44BQ, Hamamatsu). The signals were recorded by an oscilloscope (Data Sys 944, Gould) and transferred to a computer for further data processing.

Competition kinetic method The investigation of OH-initiated oxidation reactions of imidazoles has been performed using the well-established competition kinetic method with thiocyanate anion SCN- as reference compound.31,35-37


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H2O2 + hν(λ = 248 nm)  2OH

(R-1)

OH + reactant  R + H2O

(R-2)

OH + SCN -  SCNOH -

(R-3)

 SCN + OH -

(R-4)

 SCN + SCN -  (SCN)2 -

(R-5)

(SCN)2 - + (SCN)2 -  products

(R-6)

SCNOH

-

Hydrogen peroxide (H2O2) was used as a radical precursor and by its photolysis at λ = 248 nm the OH radicals were generated (R-1). Afterwards, the OH radicals react in competition with both the organic reactant (R-2) and the reference compound (R-3). The reaction of OH radical with thiocyanate anion SCN- leads to the formation of the absorbing transient species dithiocyanate radical anion (SCN)2- (R-3)-(R-5), whose yield can be used for determining the rate constant of the OH reaction with the organic reactant. The OH radical rate constant was calculated using following equation:35 gA  SCN2  k2 R  0,x =1+ gk3 SCN-  A  SCN2   x

(1)

The term A[(SCN)2-]0,x corresponds to the maximum absorbance of (SCN)2- in the absence of organic reactant. By addition of the organic reactant the yield and absorbance maximum of (SCN)2- decreases due to the competition reaction, which is described by A[(SCN)2-]x. The investigated organic reactant can absorb in the range of the photolysis wavelength (λ = 248 nm) of the radical precursor. An occurring absorbance has to be considered because the internal absorption effect35 leads to an decrease of the yield of OH radicals, resulting in an overestimation of the rate constant. To avoid this, the rate constant has to be corrected by the absorption contribution of the organic reactant. Therefore, the term A0 of equation (eq.


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1) has to be corrected for the particular reactant concentration by calculating first the initial OH radical concentration [OH]0 in the absence of the reactant, then the initial OH radical concentration in the presence of the reactant. The initial OH radical concentration [OH]0 was estimated using the molar absorption coefficients of H2O2 (ε248nm = 25.65 ± 1 L mol-1 cm-1) and KSCN (ε248nm = 60 ± 3 L mol-1 cm-1) as well as the quantum yield of H2O2 (ɸ248nm = 1.02 ± 0.1) at λ = 248 nm.30,36 The resulting percentage change of the OH radical concentrations in the presence of different reactant concentrations, Δ[OH]0,x, leads to the relative changed value of A[(SCN)

2

-

]0,x (cf. “Internal absorption effect” at SI). The strength of correction by the

internal absorption effect depends on the molar absorption coefficient and the particular concentration of the reactant at λ = 248 nm. 2-IC can be excited into its singlet state by the excimer laser pulse at λ = 248 nm used for the photolysis of H2O2 followed by intersystem crossing leading to triplet state 2-IC. In order to evaluate the impact of the triplet state of 2-IC on the OH radical kinetics the following conditions were considered. For a possible reaction of the triplet state of 2-IC with thiocyanate resulting in the formation of (SCN)2- a rate constant was estimated. To do so, the one-electron reduction potentials of the halides iodide and bromide as well as of the pseudohalogen SCN38

were correlated with the rate constants of iodide and bromide.20 This way, a rate constant

of kq = 1.6 × 108 L mol-1 s-1 for the reaction of the triplet state of 2-IC with SCN- was obtained as an estimation. The quenching rate constant of the triplet state of 2-IC with water is given as 1 × 105 s-1. Furthermore, the triplet state of 2-IC is supposed to react rapidly with molecular oxygen forming singlet oxygen39 under the used experimental conditions. The formed singlet oxygen is expected to be quenched by water. In the present study, the concentration of dissolved molecular oxygen in aqueous solution is 2.6 × 10-4 mol L-1 at 298 K. The rate constants for the triplet state quenching of 2-IC by oxygen are given as kq = 3.9 × 109 L mol-1


6

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s-1 20 and kq = (3.1 ± 0.1) × 109 L mol-1 s-1.21 From the given concentrations and rate constants a ratio of the effectivity of the sink processes for the 2-IC triplet state can be estimated as: O2:H2O:SCN- = 10:1:0.03,20 and O2:H2O:SCN- = 8:1:0.0321 depending on the rate constant used. This result shows that under the present experimental conditions the reaction of the triplet state of 2-IC with molecular oxygen establishes by far the highest sink. The formation of (SCN)2- from the triplet state of 2-IC shows the lowest contribution. To characterize the additional (SCN)2- concentration from the reaction of SCN- with triplet state 2-IC a quantum yield ɸ = 1 as an upper limit for the triplet state formation was assumed. Under the used conditions, the additional contribution was calculated to be smaller than 1% to the total (SCN)2- concentration. Thus, the photo-induced reaction of the organic reactant (2-IC) has a negligible impact on the OH radical kinetics. Each OH radical rate constant was determined by measuring five solutions increasing the concentration of the organic reactant stepwise with each solution resulting in a reduced yield of (SCN)2-. Thus, five ratios of A[(SCN)2-]0,x/A[(SCN)2-]x were obtained and are plotted against the corresponding concentration ratios [R]/[SCN-], leading to a linear relationship with a slope of m = k2/k3 (eq 1). The term k2 is the unknown OH radical rate constant and k3 is the established temperature-dependent reference rate constant, which is expressed by the equation of Zhu et al. (eq 2).37 For each measurement, 256 single traces were measured and averaged. For these five measurement data and their 1σ, 5% confidence interval errors a linear regression was performed and the statistical error of the linear regression line was calculated. The errors of the rate constants are statistical errors with 95% confidence interval considering Student’s t-factor. The linear regression was forced through the origin of 1 in the competition kinetic plots.

 1690 K  -1 -1 ln k(T) = 28.87 -   L mol s  T  ACS Paragon Plus Environment

(2) 7


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Results and Discussion Reaction mechanism Regarding the reaction mechanism (Scheme 1), Samuni and Neta40 suggested an OH addition to a double bond of the heterocyclic aromatic rather than a hydrogen abstraction. These authors supposed that the OH addition would take place at the positions 2 and 5 of the heteroaromatic ring system resulting in an allylic type radical. Quantum chemistry and density functional theory (DFT) calculations showed that the addition at the 5-position is energetically favored.41

Scheme 1. Supposed reaction mechanism of the OH radical reaction with imidazoles.

Bansal et al.42 investigated the OH reaction of imidazole, 1-methylimidazole, and 2methylimidazole via pulse radiolysis. These authors obtained a first-order decay of the OH adduct radicals from imidazole and 2-methylimidazole at alkaline conditions (pH ≥ 11) and described this decay (kfirst,

max

= 104 s-1) as a water elimination reaction (Scheme 1, blue

pathway). It is supposed that the formed OH adducts might undergo a base-catalyzed water elimination, which includes the OH group and the hydrogen atom at the N1 position.40-42 Only ACS Paragon Plus Environment

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1-methylimidazole did not show a first-order decay, which leads to the assumption that the water elimination cannot occur for substituted 1-imidazoles.42 This can be explained by the replacement of the H atom at the N1 position with the substituent. Thus, a rearrangement of the aromatic system via a water elimination is not possible. However, the base-catalyzed water elimination only occurs at high pH values (pH ≥ 11) and thus this pathway is irrelevant for the present study. In any case, for all imidazoles, an OH addition to the aromatic ring system takes place leading to the formation of an alkyl radical, which might likely react with oxygen to generate a peroxyl radical, undergoes recombination reactions or decompose to ring opening products. Unfortunately, the fate of the imidazole oxidation by OH radicals in aqueous solution is rarely investigated.

pH dependency The pH value was varied to study the effect of the acid-base equilibrium of the imidazoles (Scheme 2) on the reactivity. At acidic conditions, the tertiary nitrogen atom is protonated and the system is positively charged. The state of protonation of the nitrogen atom affects the rate constant.

Scheme 2. Acid-base equilibrium of imidazoles.

The results from the next section (Kinetics) show that the deprotonated forms react faster with the OH radicals than the protonated forms, which can be explained by the different electron density within the aromatic ring system due to the electron-withdrawing property of the protonated nitrogen. ACS Paragon Plus Environment

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Kinetics Imidazole-2-carboxaldehyde (2-IC) Imidazole-2-carboxaldehyde (2-IC) reacts with OH radicals by OH radical addition to the imidazole ring (Scheme 1). To describe the acid-base equilibrium of 2-IC (Scheme 3), pKa values of 1.3 ± 0.1 and 5.25 ± 0.05 for the ring system and aldehyde group were determined by using the acid-base titration method (Figure S5 and Figure S6). Recently, Ackendorf et al.43 investigated the acid-base equilibrium of 2-IC and determined pKa values of 2.5 ± 0.4 and 5.94 ± 0.05 by NMR studies. The pKa values obtained in the present study are lower than those determined by Ackendorf et al.43 The acid-base titration method applied in our study only determines the free acid concentration at a given pH. This is in contrast to measure the compound itself, as it was done in the cited the NMR study.43 Thus, the discrepancy of the pKa values might be due to the two different experimental methods. Compared to the other imidazoles 2-IC has a second acid-base equilibrium (Scheme 3) due to its aldehyde group. At acidic conditions, the hydrated form of 2-IC is preferred (cf. Figure S1 and Figure S2).

Scheme 3. Acid-base and hydration equilibrium of imidazole-2-carboxaldehyde (2-IC).

According to the determined pKa values the pH values 0, 4.2, and 9.1 were chosen to measure OH radical reactivity of the protonated diol (PD), the mixture of the deprotonated diol (DD) and protonated aldehyde (Ald+), and the aldehyde (Ald) form of 2-IC. The internal absorption effect35 has to be considered especially for this reaction due to the high molar absorption


10

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coefficient at λ = 248 nm (cf. SI). If the internal absorption effect would be not considered the rate constant would be overestimated by 9-25% as shown in Figure 1 for the OH radical reaction of 2-IC at pH = 9.1. Furthermore, the contributions of the protonated diol (PD) and the aldehyde (Ald) form to the reaction rate constant of the deprotonated diol (DD)/protonated aldehyde (Ald+) form of 2-IC has to be considered due to their molar fractions at pH = 4.2 (χDD/Ald+ = 92.21%, χPD = 0.12%, χAld = 7.67%; c.f. Figure S7). The reaction rate constant of DD/Ald+ is calculated by eq 3. It should be noted that at pH = 4.2 the deprotonated diol (DD) and the protonated aldehyde (Ald+) are present at the same time. About 7-8% of Ald+ are present at pH = 4.2 calculated from the UV measurements (cf. SI Figure S1) and the hydration constant (Khyd ≈ 6) from Ackendorf et al.43 Thus, the reaction of OH radicals with the protonated aldehyde (Ald+) form contributes about 10% to the obtained rate constant.

kDD/Ald = 

kobserved - kAld × χ Ald + kPD × χPD  L mol-1 s-1 χDD/Ald

(3)



The corrected rate constants for the OH radical addition were obtained as k298 K (pH = 0) = (1.8 ± 0.1) × 109 L mol-1 s-1, k298 K (pH = 4.2) = (1.8 ± 0.2) × 109 L mol-1 s-1 and k298 K (pH = 9.1) = (4.4 ± 0.6) × 109 L mol-1 s-1.

Figure 1. Uncorrected and corrected second order rate constants of the OH radical reaction with 2-IC at 298 K and pH = 9.1.


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Published rate constants for the OH reactions with imidazoles in aqueous solution are rare. Rao et al.44 determined the OH reaction rate constants of imidazole and 1-methylimidazole at different pH values using pulse radiolysis to form the OH radicals. These authors obtained rate constants of kimidazole = 5.5 × 109 L mol-1 s-1 at pH = 3.4, kimidazole = 8.7 × 109 L mol-1 s-1 at pH = 6.8, kimidazole = 1.2 × 1010 L mol-1 s-1 at pH = 10.2, k1-methylimidazole = 5.0 × 109 L mol-1 s-1 at pH = 5.4, and k1-methylimidazole = 8.1 × 109 L mol-1 s-1 at pH = 9.4 by the direct observation of the formed OH-imidazole adduct. Aruoma et al.45 and Ching et al.46 determined for the oxidation of imidazole by OH radicals the following rate constants kimidazole = 3.9 × 109 L mol-1 s-1 and kimidazole = (5.2 ± 0.7) × 109 L mol-1 s-1 at pH = 7. The OH radicals were generated using the Fenton reaction and deoxyribose was used as the reference reactant (kdeoxyribose + OH = 2.5 × 109 L mol-1 s-1).47,48 In both studies, the imidazole oxidation has been measured at pH = 7, due to the pKa value of imidazoles (pKa = 6.95)49 both the protonated and deprotonated form contributes to the reported rate constants.45,46 However, both values indicate that the pH dependent rate constants for imidazole appears to be lower than in the study of Rao et al.44 The obtained rate constants in the present study are in the same range compared to the literature values from Aruoma et al.45 and Ching et al.46. For the first time, the T-dependent rate constants of the OH radical reaction with imidazole-2-carboxaldehyde (2-IC) have been determined and is depicted in Figure 2.


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Figure 2. Arrhenius plot of the temperature-dependent measurement of 2-IC with OH radicals at pH = 0 (black points), pH = 4.2 (blue points), and pH = 9.1 (green points).

The derived Arrhenius expressions are shown in eq (4) for pH = 0, in eq (5) for pH = 4.2, and in eq (6) for pH =9.1.

k(T, pH = 0) = (7.5 ± 0.1) × 109 × exp  - 430 ± 80 K  L mol-1 s-1 T  

(4)

k(T, pH = 4.2) = (9.0 ± 0.4) × 1010 × exp  - 1140 ± 300 K  L mol-1 s-1 T  

(5)

k(T, pH = 9.1) = (3.7 ± 0.2) × 1011 × exp  - 1320 ± 360 K  L mol-1 s-1 T  

(6)













The temperature dependency of the OH radical reaction with 2-IC can be described in the following order: pH = 0 < pH = 4.2 < pH = 9.1. Under alkaline conditions, the reaction is more T-dependent than under acidic conditions as it can be seen in the Arrhenius plot and the derived Arrhenius expressions.

1-Methylimidazole (1-MI) The pKa value of 1-methylimidazole (1-MI) was taken from Lenarcik and Ojczenasz49 (pKa = 7.21) to describe its acid-base equilibrium. The absorbance was characterized at pH = 2.0 and 9.1 to clarify the possible influence of the internal absorption effect.35 The molar absorption coefficients of 1-MI ε248 nm = 0.3 ± 0.2 L mol-1 cm-1 at pH = 2.0 and ε248 nm = 1.2 ± 0.2 L mol-1 cm-1 at pH = 9.1 (cf. Figure S3) obtained from UV-measurements are low in comparison to 2-IC. These values indicate that under the used conditions the internal absorption effect is negligible. A resulting change in the initial OH radical concentrations of 0.02% at pH = 2.0 and 0.01% at pH = 9.1 was calculated. The rate constants for the OH radical addition were obtained


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as k298 K (pH = 2.0) = (2.3 ± 0.1) × 109 L mol-1 s-1 and k298 K (pH = 9.1) = (6.1 ± 0.6) × 109 L mol-1 s1. The comparison with the values from the study of Rao et al. (k1-methylimidazole = 5.0 × 109 L mol1 s1

at pH = 5.4, and k1-methylimidazole = 8.1 × 109 L mol-1 s-1 at pH = 9.4) shows that under acidic

conditions the rate constant is a factor of two higher and under alkaline conditions it is a factor of one higher than the obtained value from this study. This is in contrast to the reported values from the imidazole oxidation from Aruoma et al.45 Ching et al46 as described in the upper text area (see 2-IC section), which is in good agreement with the received rate constants in this study. The T-dependency of the OH radical rate constants is shown in Figure 3.

Figure 3. Arrhenius plot of the temperature-dependent measurement of 1-MI with OH radicals at pH = 2.0 (black points) and at pH = 9.1 (green points).

In the equations (eq 7) and (eq 8) the Arrhenius expressions obtained from the T-dependent measurement for pH = 2.0 and pH = 9.1 are shown.

k(T, pH = 2.0) = (1.1 ± 0.1) × 1010 × exp  - 480 ± 120 K  L mol-1 s-1 T  

(7)

k(T, pH = 9.1) = (1.8 ± 0.1) × 1012 × exp  - 1660 ± 400 K  L mol-1 s-1 T  

(8)









Similar to the above-mentioned compound imidazole-2-carboxaldehyde the T-dependency of the OH radical reaction of 1-MI at pH = 9.1 has a strong character and the activation energy is around a factor of seven higher in comparison to pH = 2.0. ACS Paragon Plus Environment

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2-Methylimidazole (2-MI) The pH-dependent rate constants of the OH radical oxidation reaction of 2-methylimidazole (pKa = 7.85)49 were determined as k298 K (pH = 2.0) = (3.9 ± 0.1) × 109 L mol-1 s-1 and k298 K (pH = 9.2) = (5.7 ± 0.4) × 109 L mol-1 s-1. The derived molar absorption coefficients can be given by ε248 nm = 1.6 ± 0.2 L mol-1 cm-1 at pH = 2.0 and ε248 nm = 2.9 ± 0.3 L mol-1 cm-1 at pH = 9.2 (cf. Figure S3). According to these results, the internal absorption effect is negligible under the used conditions. The calculated change in the OH radical concentrations is lower than 0.13% at pH = 2.0 and 0.01% at pH = 9.2. The Arrhenius expressions of the temperature-dependent measurement of the reaction of OH radicals with 2-MI at pH = 2.0 and 9.2 are displayed in eq (9) and eq (10) and the Arrhenius plot is shown in Figure S10.

k(T, pH = 2.0) = (1.2 ± 0.1) × 1011 × exp  - 1080 ± 120 K  L mol-1 s-1 T  

(9)

k(T, pH = 9.2) = (2.1 ± 0.1) × 1013 × exp  - 2410 ± 360 K  L mol-1 s-1 T  

(10)









The T-dependency pattern of the OH radical reaction with 2-MI shows the general behavior as 2-IC and 1-MI. The activation energy is around factor two higher under alkaline conditions (pH = 9.2) in comparison to pH = 2.0.

4-Methylimidazole (4-MI) The pH-dependent OH radical reactivity was investigated at pH =2.0 and pH = 9.1 according to the pKa value of 7.69.49 Furthermore, the molar absorption coefficients were determined by UV measurements, yielding: ε248 nm = 11.6 ± 1.2 L mol-1 cm-1 at pH = 2.0 and ε248 nm = 15.1 ± 0.5 L mol-1 cm-1 at pH = 9.1 (cf. Figure S3). The resulting change of the initial OH radical concentrations of 0.91% at pH = 2.0 and 0.06% at pH = 9.1 is negligible. The rate constants


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observed are given by k298 K (pH = 2.0) = (3.2 ± 0.2) × 109 L mol-1 s-1 and k298 K (pH = 9.1) = (6.9 ± 1.2) × 109 L mol-1 s-1. The obtained Arrhenius plot of 4-MI oxidation by OH radicals at pH = 2.0 and 9.2 is shown in Figure S11. The derived Arrhenius expressions at pH = 2.0 and pH = 9.1 are described by eq (11) and eq (12).

k(T, pH = 2.0) = (5.7 ± 0.1) × 1010 × exp  - 840 ± 120 K  L mol-1 s-1 T  

(11)

k(T, pH = 9.1) = (7.4 ± 0.3) × 1012 × exp  - 2170 ± 360 K  L mol-1 s-1 T  

(12)









Under acidic conditions, the 4-MI oxidation by OH radicals is less T-dependent than under alkaline conditions due to the protonation of the imidazole ring system.

1-Ethylimidazole (1-EI) The properties of the compound 1-ethylimidazole are given as follows: pKa = 7.2649 and the molar absorption coefficients of ε248 nm = 2.4 ± 0.1 L mol-1 cm-1 at pH = 2.0 and ε248 nm = 2.7 ± 0.1 L mol-1 cm-1 at pH = 9.1 (cf. Figure S4). The OH radical rate constants k298 K (pH = 2.0) = (2.6 ± 0.1) × 109 L mol-1 s-1 and k298 K (pH = 9.1) = (5.7 ± 0.5) × 109 L mol-1 s-1 obtained in this study were not influenced by the internal UV absorption effect. The corresponding Arrhenius plot is shown in Figure S12. The obtained data of the OH radical reaction of 1-EI at pH = 2.0 is represented by the Arrhenius expression eq (13) corresponding to an activation energy of EA = 5 ± 1 kJ mol-1.

k(T, pH = 2.0) = (2.0 ± 0.1) × 1010 × exp  - 600 ± 120 K  L mol-1 s-1 T   



(13)

The same reaction under pH = 9.1 yields the Arrhenius expression in eq (14).

k(T, pH = 9.1) = (1.6 ± 0.1) × 1012 × exp  - 1680 ± 120 K  L mol-1 s-1 T   




(14)

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The corresponding activation energy (EA = 14 ± 1 kJ mol-1) is a factor of around three higher than the activation energy at pH = 2.0.

2-Ethylimidazole (2-EI) The oxidation of 2-ethylimidazole by OH radicals was conducted at pH = 2.0 and pH = 9.2 regarding to its pKa = 7.99 reported by Lenarcik and Ojczenasz49. The received rate constants k298 K (pH = 2.0) = (3.9 ± 0.1) × 109 L mol-1 s-1 and k298 K (pH = 9.2) = (6.4 ± 1.2) × 109 L mol-1 s-1 are not influenced by the internal UV absorption of 2-EI (ε248 nm = 1.3 ± 0.2 L mol-1 cm-1 at pH = 2.0 and ε248 nm = 2.1 ± 0.4 L mol-1 cm-1 at pH = 9.2; cf. Figure S4). From the Arrhenius plot (Figure S13) the following T-dependencies are obtained:

k(T, pH = 2.0) = (2.3 ± 0.1) × 1011 × exp  - 1200 ± 120 K  L mol-1 s-1 T  

(15)

k(T, pH = 9.2) = (6.3 ± 0.1) × 1012 × exp  - 2050 ± 120 K  L mol-1 s-1 T  

(16)









The derived activation energies depict the same pattern as the aforementioned substituted imidazoles from this study. In this case the activation energy at pH = 9.2 is a factor of around two higher as at pH = 2.0. The reactivity toward OH radicals of the investigated substituted imidazoles under acidic conditions can be described as: k(2-IC) < k(1-MI) < k(1-EI) < k(4-MI) < k(2-EI) ≈ k(2-MI). In contrast to the rate constants under acidic conditions the sequence for the rate constants at pH = 9 is: k(2-IC) < k(2-MI) ≈ k(1-EI) < k(1-MI) < k(2-EI) < k(4-MI). The calculated ratios of the rate constants at alkaline conditions (pH = 9) and at acidic conditions (pH = 2, exception for 2IC: pH = 0) at different temperatures are presented in Table 1.

Table 1. Rate constant ratios k(pH = alkaline)/ k(pH = acidic) at different temperatures.


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k(pH = alkaline)/ k(pH = acidic) Compound T = 278 K

T = 288 K

T = 298 K

T = 308 K

T = 318 K

2-IC

1.7

2.4

2.4

2.2

2.9

1-MI

2.6

2.4

2.7

3.5

4.4

2-MI

1.0

1.5

1.5

1.7

2.2

4-MI

1.3

1.8

2.2

2.3

2.3

1-EI

1.9

2.1

2.2

2.8

3.0

2-EI

1.4

1.4

1.6

1.7

2.0

In conclusion, the obtained ratios in Table 1 indicate that the investigated compounds can be divided into two groups. The compounds in the first group 2-MI, 4-MI and 2-EI exhibit electron-releasing alkyl groups -CH3 and -C2H5 at the carbon atoms of the imidazole ring, which results in smaller T-dependency ratios. As well as the second group: 2-IC, 1-MI and 1-EI, while the carboxylate group of the IC has an electron-withdrawing inductive effect and the alkyl groups of 1-MI and 1-EI are bound to the heteroatom of the imidazole ring.

Diffusion limit Because the reactions of imidazoles with OH radicals occur very fast especially under alkaline conditions and higher temperatures it might be possible that these reactions are diffusion controlled. Thus, the diffusion rate constants considering the temperature dependency for each imidazole were calculated (cf. SI). Although the imidazoles exhibit different molar volumes, radii, and diffusion coefficients (cf. Table S6), they all show the same value for the diffusion limit rate constants (in unit of L mol-1 s-1): kD (278 K) = 0.7 × 1010, kD (288 K) = 1.0 × 1010, kD (298 K) = 1.4 × 1010, kD (308 K) = 1.7 × 1010 and kD (318 K) = 2.2 × 1010. In Table S7 the calculated ratios of the measured rate constant (ksecond) against the diffusion limit rate ACS Paragon Plus Environment

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constant (kD) are shown. The highest ratios up to 74% have been obtained from the Tdependent measurements at pH = 9, especially for the lower temperatures at T = 278 K and 288 K. For the measurements at pH = 2 the highest ratio was obtained at T = 278 K from the oxidation of 2-EI by OH radicals. However, these obtained values indicate that the rate constants of the OH radical reactions with imidazoles are not diffusion controlled.

Activation parameters The Arrhenius equation shows a linear relationship between the logarithm of the rate constant and the inverse temperature and can be used to calculate the pre-exponential factor A and the activation energy EA. The activation enthalpy ΔHǂ, the activation entropy ΔSǂ, and the free Gibbs enthalpy ΔGǂ were calculated via the transition state theory.50 Table 4 shows all calculated activation parameters. The activation enthalpy ΔHǂ is associated with EA and can be derived by following equation:

ΔH‡ = EA - R × T

(17)

Moreover, the pre-exponential factor A is a proportion for the change of the activation entropy ΔSǂ, which is described by following equation:

 k × T   ΔS‡ = R ×  ln A - ln  B  - 1  h   

(18)

where kB is the Boltzmann constant and h is the Planck constant. The free Gibbs enthalpy of activation ΔGǂ can be obtained from ΔHǂ and ΔSǂ using the Gibbs-Helmholtz equation (eq 19).

ΔG‡ = ΔH‡ - T × ΔS‡

(19)

ΔGǂ describes the difference in the potential free energy between the educts and the transition state.


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Table 4. Calculated activation parameters of the temperature-dependent OH reactions with imidazoles in aqueous solution. Compound

2-IC

pH

A (L mol-1 s-1)

EA (kJ mol-1)

ΔHǂ (kJ mol-1)

ΔSǂ (J K-1 mol-1)

ΔGǂ (kJ mol-1)

0

(7.5 ± 0.1) ×109

4±1

1±1

-(64 ± 1)

20 ± 4

4.2

(9.0 ± 0.4) ×1010

10 ± 3

7±2

-(44 ± 2)

20 ± 6

9.1

(3.7 ± 0.2) ×1011

11 ± 3

9±3

-(32 ± 2)

18 ± 6

2.0

(1.1 ± 0.1) ×1010

4±1

1±1

-(61 ± 1)

20 ± 4

9.1

(1.8 ± 0.1) ×1012

14 ± 3

11 ± 3

-(19 ± 1)

17 ± 5

2.0

(1.2 ± 0.1) ×1011

9±1

6±1

-(41 ± 1)

18 ± 3

9.2

(2.1 ± 0.1) ×1013

20 ± 3

18 ± 3

2±1

17 ± 4

2.0

(5.7 ± 0.1) ×1010

7±1

5±1

-(47 ± 1)

19 ± 2

9.1

(7.4 ± 0.3) ×1012

18 ± 3

15 ± 3

-(7 ± 1)

17 ± 4

2.0

(2.0 ± 0.1) ×1010

5±1

3±1

-(56 ± 1)

19 ± 5

9.1

(1.6 ± 0.1) ×1012

14 ± 1

11 ± 1

-(20 ± 1)

17 ± 2

2.0

(2.3 ± 0.1) ×1011

10 ± 1

8±1

-(36 ± 1)

18 ± 2

9.2

(6.3 ± 0.1) ×1012

17 ± 1

15 ± 1

-(8 ± 1)

17 ± 2

1-MI

2-MI

4-MI

1-EI

2-EI

The value of the activation energy EA change by a factor of one to three by changing the pH value. In the same manner the derived pre-exponential factors A increases by two orders of magnitude, except for 2-EI. Conclusively, the OH radical reaction with substituted imidazoles at acidic conditions is less T-dependent than at pH ≈ 9. In this study, positive values for ∆H‡ in the range from 1 to 8 kJ mol-1 for acidic conditions and 9 to 18 kJ mol-1 for alkaline conditions were obtained. In contrast to ∆H‡, the calculated values for the free Gibbs enthalpy ΔGǂ do not show that strong pH-dependency. Mean values of the free Gibbs enthalpies ΔGǂ = 19 kJ mol-1 under acidic conditions and ΔGǂ = 17 kJ mol-1 under alkaline conditions were received. By comparison with the literature of OH radical reactions with hydroxylated compounds


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similar values of ΔGǂ were derived.29,32,33 For almost all reactions negative values of ΔSǂ were obtained, the only exceptions is 2-MI at pH = 9.2. In addition, the positon of the alkyl group at the imidazole ring influences the reactivity.

Atmospheric Aqueous-Phase Lifetimes One main object of this present study was the clarification of the stability of imidazole-based photosensitizers. Hence, the atmospheric aqueous-phase lifetimes (τ) of the investigated imidazoles were calculated based on equation (eq 20).

τ = k × OH radicals 

-1

(20)

The reported mean tropospheric OH radical concentrations for an urban cloud case of 3.5 × 10-15 mol L-1, a remote cloud case of 2.2 × 10-14 mol L-1, an urban aerosol case 4.4 × 10-13 mol L-1, and a remote aerosol case 3.0 × 10-12 mol L-1 were taken from CAPRAM 3.0i multiphase mechanism from Herrmann et al.36

Table 5. Atmospheric aqueous-phase lifetimes calculated for urban and remote areas. Compound

2-IC

τcloud droplet

pH

τaerosol

urban

remote

urban

remote

0

1.8 d

7h

21 min

3 min

4.2

1.8 d

7h

21 min

3 min

9.1

18.1 h

2.9 h

9 min

1 min

2.0

1.4 d

5.5 h

16 min

2 min

9.1

13.1 h

2.1 h

6 min

1 min

2.0

20.4 h

3.2 h

10 min

1 min

9.2

13.9 h

2.2 h

7 min

1 min

2.0

1.1 d

3.9 h

12 min

2 min

1-MI

2-MI

4-MI


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9.1

11.5 h

1.8 h

5 min

1 min

2.0

1.3 d

4.9 h

15 min

2 min

9.1

13.9 h

2.2 h

7 min

1 min

2.0

20.4 h

3.2 h

10 min

1 min

9.1

12.4 h

2h

6 min

1 min

1-EI

2-EI

Due to the difference of the OH radical concentration, the calculated lifetime of imidazoles in remote areas is smaller than in urban areas. The usual pH for aerosols in urban areas ranges from -2 to 5 and for clouds from 4 to 64. Thus, the oxidation of the imidazoles is dominated by their protonated form. Under these conditions 2-MI and 2-EI have the shortest and 2-IC the longest lifetime. In the case of the remote areas 2-IC has the longest lifetime, whereas 2-MI and 2-EI has the shortest lifetime. However, a much shorter lifetime of the imidazoles was calculated for the remote case. Due to the atmospheric lifetime of around 20 minutes in aerosols and two days in clouds of 2IC photosensitized reactions might possibly occur in the aqueous phase subject to the actual OH radical concentration encountered in the atmospheric aqueous system of particles or clouds. Though, the hydration equilibrium of 2-IC (Scheme 3) has to be considered within the urban aerosol and cloud pH range (pHaerosol = -2 to 5, pHcloud = 4 to 6).4 Under these conditions the deprotonated diol as well as the aldehyde form of 2-IC are existent. The aldehyde form of 2-IC can act as a photosensitizer19,20,24 due to its carbonyl group whereas the formation of the diol form results in the loss of this reaction pathway. Thus, under atmospheric conditions photosensitized reactions including 2-IC will be partly inhibited and this should also be considered in treatments of this chemistry for atmospheric aqueous phase systems. Furthermore, the comparison of the atmospheric lifetimes of 2-IC indicates that the contribution from photosensitized reactions under atmospheric conditions is expected to be ACS Paragon Plus Environment

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higher in clouds than in aerosols. The present study supports the conclusion of Tsui et al.51 These authors performed model studies regarding to the SOA formation from photosensitized reactions of 2-IC with volatile organic compounds (VOC) and determined a contribution of less than 0.3% to the total aqueous SOA growth under remote ambient conditions. A treatment of the pH dependency of the molecular form of 2-IC and its impact on photosensitization is missing in this modelling study.

Conclusion In this study, the aqueous-phase reactivity of imidazoles toward the most atmospherically relevant oxidant, the OH radical, was investigated. Due to the common occurrence of OH radicals in the atmosphere, they play an important role for the chemistry of organic compounds. These OH-initiated oxidation reactions represent a sink for organic compounds and are a source for new species, which may have different properties from their precursor compounds or affect other important atmospheric processes (e.g. particle mass production, atmospheric oxidation potential). The temperature-dependent kinetic investigation shows that the oxidation of imidazoles with OH radicals occurs very fast like other aromatic compounds.36,52-54 The calculated atmospheric lifetimes refer to an efficient oxidation of the imidazoles under tropospheric aqueous phase conditions. Thus, atmospheric aqueous-phase oxidation processes should be considered in the study of the chemistry of imidazoles. Product studies are necessary to identify the oxidation products as well as the reaction mechanism. Several studies predict that the OH radical add to a double bond of the heteroaromatic compound leading to an alkyl radical,40-42 which can further undergo different reactions such as a base-catalyzed water elimination,40,42 formation of peroxyl radicals via the reaction with oxygen, recombination reactions, or ring opening reactions. Nevertheless, the


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understanding of the degradation pathway of imidazoles is still lacking. Due to this, investigations concerning the product formation of the OH-initiated oxidation of imidazoles are necessary to get a better understanding of the reaction mechanism. The calculated atmospheric lifetimes allow for the occurrence of photosensitization but this can be limited by OH oxidation and hydration of the studied compounds.

Associated Content Supporting Information Materials, absorbance spectra, acid-base titration, internal absorption effect, temperaturedependent second order rate constants

Author Information Corresponding Author *H. Herrmann. Phone: +49 341 2717 7024. Fax: +49 341 2717 99 7024. E-mail: [email protected] ORCID Hartmut Herrmann: 0000-0001-7044-2101 Notes The authors declare no competing financial interest.

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23. Monge, M. E.; Rosenørn, T.; Favez, O.; Müller, M.; Adler, G.; Abo Riziq, A.; Rudich, Y.; Herrmann, H.; George, C.; D’Anna, B. Alternative Pathway for Atmospheric Particles Growth. PNAS 2012, 109, 6840-6844. 24. Rossignol, S.; Aregahegn, K. Z.; Tinel, L.; Fine, L.; Nozière, B.; George, C. Glyoxal Induced Atmospheric Photosensitized Chemistry Leading to Organic Aerosol Growth. Environ. Sci. Technol. 2014, 48, 3218-3227. 25. González Palacios, L.; Corral Arroyo, P.; Aregahegn, K. Z.; Steimer, S. S.; Bartels-Rausch, T.; Nozière, B.; George, C.; Ammann, M.; Volkamer, R. Heterogeneous Photochemistry of Imidazole-2carboxaldehyde: HO2 Radical Formation and Aerosol Growth. Atmos. Chem. Phys. 2016, 16, 1182311836. 26. Corral Arroyo, P.; Bartels-Rausch, T.; Alpert, P. A.; Dumas, S.; Perrier, S.; George, C.; Ammann, M. Particle-Phase Photosensitized Radical Production and Aerosol Aging. Environ. Sci. Technol. 2018, 52, 7680-7688. 27. Teich, M.; van Pinxteren, D.; Kecorius, S.; Wang, Z.; Herrmann, H. First Quantification of Imidazoles in Ambient Aerosol Particles: Potential Photosensitizers, Brown Carbon Constituents, and Hazardous Components. Environ. Sci. Technol. 2016, 50, 1166-1173. 28. Herrmann, H. Kinetics of Aqueous Phase Reactions Relevant for Atmospheric Chemistry. Chem. Rev. 2003, 103, 4691-4716. 29. Hoffmann, D.; Weigert, B.; Barzaghi, P.; Herrmann, H. Reactivity of Poly-Alcohols towards OH, NO3 and SO4- in Aqueous Solution. Phys. Chem. Chem. Phys. 2009, 11, 9351-9363. 30. Schaefer, T.; Schindelka, J.; Hoffmann, D.; Herrmann, H. Laboratory Kinetic and Mechanistic Studies on the OH-Initiated Oxidation of Acetone in Aqueous Solution. J. Phys. Chem. A 2012, 116, 6317-6326. 31. Ervens, B.; Gligorovski, S.; Herrmann, H. Temperature-Dependent Rate Constants for Hydroxyl Radical Reactions with Organic Compounds in Aqueous Solutions. Phys. Chem. Chem. Phys. 2003, 5, 1811-1824. 32. Schöne, L.; Schindelka, J.; Szeremeta, E.; Schaefer, T.; Hoffmann, D.; Rudzinski, K. J.; Szmigielski, R.; Herrmann, H. Atmospheric Aqueous Phase Radical Chemistry of the Isoprene Oxidation Products Methacrolein, Methyl Vinyl Ketone, Methacrylic Acid and Acrylic Acid – Kinetics and Product Studies. Phys. Chem. Chem. Phys. 2014, 16, 6257-6272. 33. Schaefer, T.; van Pinxteren, D.; Herrmann, H. Multiphase Chemistry of Glyoxal: Revised Kinetics of the Alkyl Radical Reaction with Molecular Oxygen and the Reaction of Glyoxal with OH, NO3, and SO4– in Aqueous Solution. Environ. Sci. & Technol. 2015, 49, 343-350. 34. Otto, T.; Stieger, B.; Mettke, P.; Herrmann, H. Tropospheric Aqueous-Phase Oxidation of Isoprene-Derived Dihydroxycarbonyl Compounds. J. Phys. Chem. A 2017, 121, 6460-6470. 35. Schaefer, T.; Herrmann, H. Competition Kinetics of OH Radical Reactions with Oxygenated Organic Compounds in Aqueous Solution: Rate Constants and Internal Optical Absorption Effects. Phys. Chem. Chem. Phys. 2018, 20, 10939-10948. 36. Herrmann, H.; Hoffmann, D.; Schaefer, T.; Bräuer, P.; Tilgner, A. Tropospheric Aqueous-Phase Free-Radical Chemistry: Radical Sources, Spectra, Reaction Kinetics and Prediction Tools. ChemPhysChem 2010, 11, 3796-3822. 37. Zhu, L.; Nicovich, J. M.; Wine, P. H. Temperature-Dependent Kinetics Studies of Aqueous Phase Reactions of Hydroxyl Radicals with Dimethylsulfoxide, Dimethylsulfone, and Methanesulfonate. Aquat. Sci. 2003, 65, 425-435. 38. Stanbury, D. M. Reduction Potentials Involving Inorganic Free Radicals in Aqueous Solution. In Advances in Inorganic Chemistry, Sykes, A. G., Ed. Academic Press: 1989; Vol. 33, pp 69-138. 39. Wilkinson, F.; Helman, W. P.; Ross, A. B. Rate Constants for the Decay and Reactions of the Lowest Electronically Excited Singlet State of Molecular Oxygen in Solution. An Expanded and Revised Compilation. Journal of Physical and Chemical Reference Data 1995, 24, 663-677. 40. Samunl, A.; Neta, P. Electron Spin Resonance Study of the Reaction of Hydroxyl Radicals with Pyrrole, Imidazole, and Related Compounds. J. Phys. Chem. 1973, 77, 1629-1635. 41. Llano, J.; Eriksson, L. A. Mechanism of Hydroxyl Radical Addition to Imidazole and Subsequent Water Elimination. J. Phys. Chem. B 1999, 103, 5598-5607. ACS Paragon Plus Environment

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42. Bansal, K. M.; Sellers, R. M. Polarographic and Optical Pulse Radiolysis Study of the Radicals Formed by Hydroxyl Radical Attack on Imidazole and Related Compounds in Aqueous Solutions. J. Phys. Chem. 1975, 79, 1775-1780. 43. Ackendorf, J. M.; Ippolito, M. G.; Galloway, M. M. pH Dependence of the Imidazole-2carboxaldehyde Hydration Equilibrium: Implications for Atmospheric Light Absorbance. Environ. Sci. Technol. Lett. 2017, 4, 551-555. 44. Rao, P. S.; Simic, M.; Hayon, E. Pulse Radiolysis Study of Imidazole and Histidine in Water. J. Phys. Chem. 1975, 79, 1260-1263. 45. Aruoma, O. I.; Laughton, M. J.; Halliwell, B. Carnosine, Homocarnosine and Anserine: Could they act as Antioxidants in vivo? Biochem. J. 1989, 264, 863-869. 46. Ching, T.-L.; Haenen, G. R. M. M.; Bast, A. Cimetidine and other H2 Receptor Antagonists as Powerful Hydroxyl Radical Scavengers. Chemico-Biological Interactions 1993, 86, 119-127. 47. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical Review of Rate Constants for Reactions of Hydrated Electrons, Hydrogen Atoms and Hydroxyl Radicals (⋅OH/⋅O−) in Aqueous Solution. J. Phys. Chem. Ref. Data 1988, 17, 513-886. 48. NIST Solution Kinetics Database Version 3.0, National Institute of Science and Technology, Gaithersburg, MD. 1998. 49. Lenarcik, B.; Ojczenasz, P. The Influence of the Size and Position of the Alkyl Groups in Alkylimidazole Molecules on their Acid-Base Properties. J. Heterocycl. Chem. 2009, 39, 287-290. 50. Laidler, K. J.; King, M. C. Development of Transition-State Theory. J. Phys. Chem. 1983, 87, 2657-2664. 51. Tsui, W. G.; Rao, Y.; Dai, H.-L.; McNeill, V. F. Modeling Photosensitized Secondary Organic Aerosol Formation in Laboratory and Ambient Aerosols. Environmental Science & Technology 2017, 51, 7496-7501. 52. Anbar, M.; Meyerstein, D.; Neta, P. The Reactivity of Aromatic Compounds toward Hydroxyl Radicals. J. Phys. Chem. 1966, 70, 2660-2662. 53. Neta, P.; Hoffman, M. Z.; Simic, M. Electron Spin Resonance and Pulse Radiolysis Studies of the Reactions of OH and O- Radicals with Aromatic and Olefinic Compounds. J. Phys. Chem. A 1972, 76, 847-853. 54. Ashton, L.; Buxton, G. V.; Stuart, C. R. Temperature Dependence of the Rate of Reaction of OH with some Aromatic Compounds in Aqueous Solution. Evidence for the Formation of a π-Complex Intermediate? J. Chem. Soc. Faraday Trans. 1995, 91, 1631-1633.

Table of Contents (TOC) Graphic


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Scheme 1. Supposed reaction mechanism of the OH radical reaction with imidazoles. 127x64mm (600 x 600 DPI)


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Scheme 2. Acid-base equilibrium of imidazoles. 66x17mm (600 x 600 DPI)



Scheme 3. Acid-base and hydration equilibrium of imidazole 2 carboxaldehyde (2-IC). 132x22mm (600 x 600 DPI)


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Figure 1. Uncorrected and corrected second order rate constants of the OH radical reaction with 2-IC at 298 K and pH = 9.1. 79x59mm (600 x 600 DPI)



Figure 2. Arrhenius plot of the temperature-dependent measurement of 2-IC with OH radicals at pH = 0 (black points), pH = 4.2 (blue points), and pH = 9.1 (green points). 80x53mm (600 x 600 DPI)


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Figure 3. Arrhenius plot of the temperature-dependent measurement of 1-MI with OH radicals at pH = 2.0 (black points) and at pH = 9.1 (green points). 79x51mm (600 x 600 DPI)


OH-Initiated Oxidation of Imidazoles in Tropospheric Aqueous-Phase

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