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. 2024 Mar 1;128(10):1880–1891. doi: 10.1021/acs.jpca.4c00156

Temperature-Dependent Kinetics of the Reactions of the Criegee Intermediate CH2OO with Hydroxyketones

Zachary A Cornwell †, Jonas J Enders †, Aaron W Harrison ‡, Craig Murray †,*
PMCID: PMC10945482  PMID: 38428028

Abstract

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Though there is a growing body of literature on the kinetics of CIs with simple carbonyls, CI reactions with functionalized carbonyls such as hydroxyketones remain unexplored. In this work, the temperature-dependent kinetics of the reactions of CH2OO with two hydroxyketones, hydroxyacetone (AcOH) and 4-hydroxy-2-butanone (4H2B), have been studied using a laser flash photolysis transient absorption spectroscopy technique and complementary quantum chemistry calculations. Bimolecular rate constants were determined from CH2OO loss rates observed under pseudo-first-order conditions across the temperature range 275–335 K. Arrhenius plots were linear and yielded T-dependent bimolecular rate constants: kAcOH(T) = (4.3 ± 1.7) × 10–15 exp[(1630 ± 120)/T] and k4H2B(T) = (3.5 ± 2.6) × 10–15 exp[(1700 ± 200)/T]. Both reactions show negative temperature dependences and overall very similar rate constants. Stationary points on the reaction energy surfaces were characterized using the composite CBS-QB3 method. Transition states were identified for both 1,3-dipolar cycloaddition reactions across the carbonyl and 1,2-insertion/addition at the hydroxyl group. The free-energy barriers for the latter reaction pathways are higher by ∼4–5 kcal mol–1, and their contributions are presumed to be negligible for both AcOH and 4H2B. The cycloaddition reactions are highly exothermic and form cyclic secondary ozonides that are the typical primary products of Criegee intermediate reactions with carbonyl compounds. The reactivity of the hydroxyketones toward CH2OO appears to be similar to that of acetaldehyde, which can be rationalized by consideration of the energies of the frontier molecular orbitals involved in the cycloaddition. The CH2OO + hydroxyketone reactions are likely too slow to be of significance in the atmosphere, except at very low temperatures.

Introduction

Criegee intermediates (CIs) are a zwitterionic species formed in alkene ozonolysis that can impact the oxidizing capacity of the troposphere.1−4 Alkene ozonolysis proceeds via a 1,3-dipolar cycloaddition reaction to produce a cyclic 1,2,3-trioxolane, or primary ozonide (POZ), that promptly decomposes to form a CI and a carbonyl compound.1,5 While larger CIs tend to undergo unimolecular decomposition,6,7 generating OH radicals, the smallest CI, CH2OO, has a longer lifetime and may undergo bimolecular reaction with trace atmospheric gases after collisional stabilization.8,9 Reaction with water vapor, primarily in the form of water dimer, is the major reactive sink for CH2OO,4 although bimolecular reactions with other trace atmospheric gases, such as SO2 and organic acids, are occasionally competitive in specific environments and under favorable conditions.4,10

CIs react with carbonyl species in a concerted 1,3-dipolar cycloaddition reaction to form a cyclic 1,2,4-trioxolane, or secondary ozonide (SOZ). Recently, work in our laboratory has explored the effect of varying the carbonyl substituents (R1R2CO, where R1 and R2 are alkyl or acyl groups) on the gas-phase reactivity of a series of ketones and diketones with CH2OO.11,12 Reactivity trends can be rationalized using concepts from frontier molecular orbital (FMO) theory.13−16 The cycloaddition mechanism arises primarily from π–π* interactions between the occupied n(pC–pO) orbital of the electron-rich species (CH2OO) and the unoccupied π* orbital of the electron-deficient species (R1R2CO). Electron-withdrawing groups (EWGs) on the carbonyl lower the energy of the carbonyl π* orbital, which reduces the energy gap with the CI orbital, stabilizes the transition state (TS), and increases the reactivity. Electron-donating groups (EDGs) have the opposite effect and ultimately decrease reactivity. Hammett substituent constants provide a useful qualitative proxy for the electron-donating or withdrawing-character of the substituents on the carbonyl.17,18

The kinetics of the reaction of acetone (Ac, R1 = R2 = CH3), a representative model ketone, with CH2OO has been thoroughly investigated experimentally,11,12,19−22 with recent measurements converging on a 298 K rate constant of ∼5 × 10–13 cm3 s–1. Aldehydes (R1 = H, R2 = H or alkyl) react faster, with rate constants in the range (1–4) × 10–12 cm3 s –1.19,20,23−25 The presence of strongly EWGs increases rate constants further. For example, hexafluoroacetone (HFA, R1 = R2 = CF3) has a rate constant of ∼3 × 10–11 cm3 s–1.19,26 The α-diketones, biacetyl (BiAc, R1 = CH3, R2 = CH3CO), and acetyl propionyl (AcPr, R1 = CH3/C2H5, R2 = C2H5CO/CH3CO) have both electron-donating alkyl and electron-withdrawing acyl substituents, and the reactions with CH2OO have rate constants of ∼1 × 10–11 cm3 s–1.11,12 Where examined, all reactions of carbonyls with CH2OO show a negative temperature dependence.

Our focus in this study is the reactions of CH2OO with two hydroxyketones: hydroxyacetone (acetol, AcOH, R1 = CH3, R2 = CH2OH) and 4-hydroxy-2-butanone (4H2B, R1 = CH3, R2 = CH2CH2OH). Hydroxyketones are multifunctional VOCs, which present multiple reactive sites and may exhibit cooperative effects.27 The reactions of AcOH and 4H2B with CH2OO can occur at the carbonyl or the hydroxyl moieties. The latter pathway is expected to be minor, as aliphatic alcohols react with CH2OO to form alkoxymethyl hydroperoxides, with rate constants in the range (1–2) × 10–13 cm3 s–1 at 298 K,28,29 smaller than those of most carbonyls. The hydroxymethyl (CH2OH) group has Hammett substituent constants of zero and is neither electron-donating nor withdrawing. Consequently, the rate constant for the cycloaddition reaction is anticipated to be comparable to that of acetaldehyde.

AcOH and 4H2B have also been identified as trace gases in the troposphere,30,31 where they are formed as secondary oxidation products of isoprene and other atmospheric hydrocarbons.32−35 AcOH is also produced directly from biomass burning.36,37 The major reactive sink for hydroxyketones is reaction with OH radicals, which results in lifetimes of a few days. Kinetics studies of the OH + AcOH reaction have produced surprisingly inconsistent results; the IUPAC recommendation for the 298 K rate constant is 5.9 × 10–12 cm3 s–1.38−47 The OH + 4H2B kinetics measurements show similar inconsistencies, but the 298 K rate constant appears to be similar.40,48−52 Photolysis of both species is far slower than reaction with OH, particularly for AcOH where the presence of the α-hydroxyl group causes a ∼10 nm blue-shift of the first absorption band that reduces the absorption at actinic wavelengths (see Figure S2 in the Supporting Information).53

The results of laser flash photolysis transient absorption spectroscopy measurements quantifying the temperature-dependent kinetics of the reactions of CH2OO with AcOH and 4H2B across the range 275–335 K are reported here. Complementary ab initio calculations map out the reaction energy profiles and provide a basis for explaining reactivity trends using FMO theory. The potential implications of the title reactions in the atmosphere are briefly discussed.

Methods

The temperature-controlled flash photolysis, transient absorption spectroscopy apparatus has been described in detail previously12 and will be summarized briefly here.

CH2OO was produced in the flow reactor by the photolysis of diiodomethane (CH2I2) in the presence of excess O2 using the 355 nm output of an Nd:YAG laser (Continuum Surelite II-10). Typical pulse energies were ∼10 mJ, resulting in fluences of ∼28 mJ cm–2. Absorption spectra were obtained by dispersing the output of pulsed LEDs (LightSpeed Technologies) in a spectrograph (Andor Shamrock 303i with iDus 420 CCD camera). Kinetics measurements used an LED nominally centered at 365 nm to obtain transient absorption spectra of CH2OO (and IO) in the range 360–395 nm at various time delays after photolysis. Independent absorption measurements were performed in the range 270–295 nm using an LED centered at 280 nm to quantify reactant concentrations. A digital delay generator (Quantum Composers, 9528) synchronized the photolysis laser, LED driver, and CCD camera.

The flow reactor itself comprises a jacketed quartz tube, with an effective path length of 90 cm. A unistat (Huber Tango) precisely controlled the reactor temperature (within <1 K) over the range 275–335 K. Gas flows into the reactor were controlled using a range of choked-flow orifices (O’Keefe). Gases (O2 and N2) were used directly from the cylinders, while liquids (CH2I2, AcOH, and 4H2B) were placed in smog bubblers and carried into the cell by a flow of N2. The smog bubblers were held in a water bath maintained at 295 K to prevent evaporative cooling and vapor pressure drop off. Measurements with AcOH used a total flow rate of 3.8 sLpm, resulting in a pressure of 78 Torr in the reactor. The lower vapor pressure of 4H2B (1.2 Torr versus 3.5 Torr at 295 K)49,54 required a larger total flow rate of 4.9 sLpm and a reactor pressure of 100 Torr. Typically, the concentrations in the reactor were [CH2I2] = 1.1 × 1015 cm–3, [O2] = 2.1 × 1017 cm–3, [AcOH] = (1–5) × 1015 cm–3, and [4H2B] = (0.5–1.5) × 1015 cm–3, with N2 balance. All chemicals were used as supplied: O2 (Airgas, UHP 4.4), N2 (Airgas, industrial grade), CH2I2 (Sigma-Aldrich, 99%), AcOH (Acros Organics, 90%), and 4H2B (Tokyo Chemical Industry, 95%). FT-IR spectra of the headspace above samples of liquid AcOH and 4H2B were recorded (see Figure S1 in Supporting Information) using a JASCO 4700 spectrometer to identify any impurities. No bands associated with any other organic species were identified, consistent with previous suggestions that the likely impurity in AcOH is residual H2O,55 although at levels too low to affect the kinetics measurements.

Electronic structure calculations were performed with the GAMESS and Gaussian 16 programs.56−59 Geometries of reactants, products, entrance channel complexes, and TS structures were initially optimized using the B3LYP functional with the Dunning-type cc-pVDZ basis set and the harmonic frequencies subsequently calculated. The presence of zero or one imaginary frequency confirmed that the optimized geometries were true minima or TSs, respectively. Intrinsic reaction coordinate (IRC) calculations were performed to verify that the expected reactants and products were reached on either side of the TS. Reaction thermochemistry was determined using rigid-rotor harmonic-oscillator (RRHO) partition functions. Additional calculations were performed using the composite CBS-QB3 method to refine the calculated energies.60 The CBS-QB3 method provides reliable thermochemistry at modest computational cost.61−63 In addition, the FMO energies are obtained from the optimization output at the B3LYP/cc-pVDZ level of theory. Previous work has shown the frontier orbital energies calculated at a similar level of theory to provide linear correlation with molecular properties such as ionization potential, electron affinity, and excitation energy.64

Results

Reactant Concentration Measurements

As in our previous work,11,12,28 concentrations of the hydroxyketone reactants in the flow cell during kinetics measurements can be estimated using reported vapor pressures (3.50 ± 0.17 Torr for AcOH, and 1.24 ± 0.04 Torr for 4H2B)49,54 and fractional flow rates. Since the first UV absorption bands of both hydroxyketones can be observed using an LED centered at 280 nm,53 their concentrations can also be measured directly. Absorption spectra of the hydroxyketones are recorded in the wavelength range 275–290 nm under conditions that are otherwise identical to those used in the kinetics measurements across the 275–335 K temperature range. Absolute hydroxyketone number densities are determined using previously reported absorption cross sections. The first UV absorption bands of AcOH and 4H2B have peak cross sections of ∼6 × 10–20 cm2 at ∼270 and 280 nm, respectively, as is typical of the excitation to the S1(nπ*) state in carbonyls.53 The only reported measurement for 4H2B is by Messaadia et al.,65 while various measurements exist for AcOH.39,42,53,65 There is a discrepancy between the AcOH cross sections recommended by the IUPAC Task Group on Atmospheric Chemical Kinetic Data Evaluation and JPL Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies Evaluation.47,66 While IUPAC prefers the values of Orlando et al.,39 JPL uses an average of that and lower values measured by Butkovskaya et al.39,42 Since the other reported measurements agree well with the JPL recommendation,53,65 we elect to use it to quantify [AcOH]exp and determine a concentration scale factor. The UV absorption spectra of both hydroxyketones are shown in Figure S2 in the Supporting Information.

The gradients of plots of measured against estimated concentrations ([X]exp vs [X]est) (Figures S3 and S4 in the Supporting Information) provide a scaling factor that can be used to correct the estimated hydroxyketone concentrations. Previously,11,12 we have found concentration scaling factor values within 10% of unity for acetone and diketones, indicating that the estimates give values close to the actual concentrations in the flow cell. The hydroxyketone concentration scaling factors deviate from unity, however, as can be seen from Figure 1. For AcOH, the scale factor is independent of temperature with an average value of 1.23 ± 0.07, indicating that the actual concentration is slightly higher than estimated. In contrast, the values determined for 4H2B suggest that its concentration is overestimated, with an average of 0.59 ± 0.08. The overall uncertainties are estimated from the spread in values obtained in multiple measurements across the temperature range.

Figure 1.

Figure 1

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Scale factors determined from gradient of [X]exp versus [X]est calibration plots as a function of temperature. Solid lines are the T-independent average, while shaded areas represent the estimated experimental uncertainty, based on the variability of the measurements. Values >1 or <1 indicate that reactant concentration is underestimated (AcOH) or overestimated (4H2B), respectively.

Deviations from unity for the concentration scaling factors derive primarily from systematic errors in [X]est and/or [X]exp, which depend respectively on the hydroxyketone vapor pressures Pvap,X and absorption cross sections σX(λ). The T-dependent vapor pressures of AcOH and 4H2B have been measured to a high degree of precision,49,54 although several groups have discussed evidence of hydroxyketone “stickiness” in the course of kinetics measurements.43,44,65 Wall losses between the bubbler and the flow reactor would lead to concentration overestimates and scaling factors <1, which is consistent with the observed value for 4H2B but not for AcOH. Based on the FT-IR spectrum of AcOH shown in Figure S1 in the Supporting Information and reported band intensities,36 we estimate a vapor pressure of 3.9 ± 0.4 Torr. If the C=O stretch bands of AcOH and 4H2B are assumed to have the same intensity (as supported by ab initio calculations), we estimate the vapor pressure of 4H2B to be 0.76 ± 0.11 Torr. The ratios of these estimated vapor pressures to the literature values49,54 are 1.11 ± 0.12 for AcOH and 0.61 ± 0.09, in good agreement with the measured scaling factors and suggesting that the systematic error originates in the reported values of Pvap,X. Another possibility for values <1 is that the smog bubbler headspace is not saturated with the organic vapor, although the use of the experimental scale factor corrects for this effect.

For the purposes of the kinetics measurements, the absorption cross sections are ultimately of most significance as they are used to determine [X]exp directly. As noted above, σAcOH(λ) values recommended by IUPAC are 10% greater than the JPL values, which would produce lower values of [AcOH]exp and bring the concentration scaling factor value closer to unity (1.12 ± 0.06) and closer to the vapor pressure ratio estimated from the FT-IR spectra. Such a change would also require an increase in the bimolecular rate constant for reaction with CH2OO proportionally. We proceed on the basis that the σX(λ) values used are accurate. The rate constants determined in the experiments discussed below are inversely proportional to σX(λ) and can be adjusted appropriately if improved values become available in the future.

Kinetics Measurements

The kinetics of the reactions of CH2OO with AcOH and 4H2B were studied under pseudo-first-order conditions of excess hydroxyketone at four temperatures in the range 275–335 K. The lowest hydroxyketone concentrations were approximately 2 orders of magnitude greater than [CH2OO]0. Transient absorption spectra obtained in the range 363–395 nm at various photolysis-probe delay times were decomposed into contributions from CH2OO and IO using known absorption spectra67,68 to generate [CH2OO]t and [IO]t concentration–time profiles. Examples of typical experimental transient spectra recorded with and without AcOH and the resulting [CH2OO]t profiles are shown in Figure S5 in the Supporting Information. As expected, the CH2OO concentrations are observed to decrease more rapidly with increasing hydroxyketone concentration, while the IO concentration profiles remain unaffected. Peak CI concentrations of [CH2OO]0 = (6–8) × 1012 cm–3 are significantly smaller than those of the hydroxyketone reactants, ensuring pseudo-first-order conditions.

Analysis of the [CH2OO]t profiles used the same kinetic model as described previously and summarized in Supporting Information.11,12 The differential rate law for CH2OO loss includes a quadratic term for bimolecular self-reaction with rate constant kself and a linear term for pseudo-first-order reactions with rate constant kloss. The pseudo-first-order rate constants kloss at each hydroxyketone concentration are derived from fits of the [CH2OO]t profiles to the integrated rate law, with kself fixed to a T-independent value of 7.8 × 10–11 cm3 s–1.12 Plots of kloss against [hydroxyketone] are linear, and a least-squares fit (weighted by the uncertainties in both kloss and [hydroxyketone]) returns a background loss rate kbgd and the bimolecular rate constant khydroxyketone as the gradient. Examples are shown in Figure 2 for the reactions of CH2OO with AcOH and 4H2B at 295 K, and the full set of measurements at all four temperatures is shown in Figures S6 and S7 in the Supporting Information. At 295 K, the bimolecular rate constants for the CH2OO + AcOH and CH2OO + 4H2B reactions are the same within the statistical uncertainties of the fits: kAcOH = (1.09 ± 0.15) × 10–12 cm3 s–1 and k4H2B = (1.11 ± 0.26) × 10–12 cm3 s–1. As has been observed for other reactions of CH2OO with carbonyl species,12,20,22,25 the rate constants for the AcOH and 4H2B reactions also decrease with increasing temperature. The complete set of measured T-dependent rate constants is summarized in Table 1. Background loss rates, attributed to reaction with I atoms or other species present in the flow reactor, are typically between ∼1000 and 1500 s–1 and decrease slightly with increasing temperature. The bimolecular rate constants for the CH2OO + hydroxyketone reactions were obtained by averaging three individual kinetic runs at each temperature, with errors representing the statistical uncertainty (1σ) in the fit. The same values, although with smaller uncertainties, are obtained from global fits to the complete data sets at each temperature.

Figure 2.

Figure 2

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Pseudo-first-order plots for the reactions of CH2OO with (a) AcOH and (b) 4H2B at 295 K. Vertical and horizontal error bars represent statistical uncertainties (1σ) in the loss rates determined from fitting [CH2OO] time profiles and uncertainties in the hydroxyketone concentration calibration measurements, respectively. Weighted linear fits to the experimental data are also shown, with shaded area representing 1σ prediction bands.

Table 1. T-dependent Bimolecular Rate Constants, Arrhenius Parameters, and Standard Enthalpies, Entropies, and Gibbs Energies of Activation at 298 K for the Reactions of CH2OO with Hydroxyacetone and 4H2Ba.

T/K kAcOH/10–12 cm3 s–1 k4H2B/10–12 cm3 s–1
275 1.63 ± 0.07 1.70 ± 0.04
295 1.09 ± 0.15 1.11 ± 0.26
315 0.75 ± 0.03 0.75 ± 0.09
335 0.62 ± 0.09 0.60 ± 0.12
A/10–15 cm3 s–1 4.3 ± 1.7 3.5 ± 2.6
(Ea/R)/K –1630 ± 120 –1700 ± 200
Ea/kcal mol–1 –3.24 ± 0.23 –3.38 ± 0.40
Δ⧧H°/kcal mol–1 –4.43 ± 0.23 –4.56 ± 0.40
Δ⧧S°/cal K–1 mol–1 –39.5 ± 0.4 –39.9 ± 0.7
Δ⧧G°/kcal mol–1 +7.36 ± 0.26 +7.34 ± 0.46
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a

Uncertainties are 1σ statistical uncertainties from the fits.

Arrhenius plots for both reactions are shown in Figure 3. The plots are linear over the temperature range explored in the experiments, and the T-dependent bimolecular rate constants can be expressed in Arrhenius form as kAcOH(T) = (4.3 ± 1.7) × 10–15 exp[(1630 ± 120)/T] and k4H2B(T) = (3.5 ± 2.6) × 10–15 exp[(1700 ± 200)/T]. The similarity of the −Ea/R values indicates that both reactions are equally sensitive to temperature, within the 1σ experimental uncertainties, with negative energies of activation of Ea ≈ −3.3 kcal mol–1.

Figure 3.

Figure 3

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Arrhenius plots for the reaction of CH2OO with AcOH (red) and 4H2B (blue), with 1σ statistical uncertainties. The AcOH data have been offset vertically for clarity. Solid black lines are linear fits with shaded areas representing 1σ prediction bands.

The thermodynamic formulation of canonical transition state theory (CTST) for a bimolecular gas-phase reaction can be used to extract the standard entropy, enthalpy, and free energy of activation

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The factor k2B/hp° is evaluated to be 2.87 × 10–12 cm3 s–1 K–2 using p° = 105 Pa. Δ⧧S° and Δ⧧H° are the standard entropy and enthalpy of activation, respectively. The latter is related to the activation energy by Δ⧧H° = Ea – 2RT for a bimolecular reaction.69 Unsurprisingly, given the similarities in the T-dependent rate constants, the thermodynamic parameters derived for the AcOH and 4H2B reactions are also similar, with Δ⧧S° ≈ −40 cal K–1 mol–1, Δ⧧H° ≈ −4.5 kcal mol–1, and Δ⧧G° ≈ +7.3 kcal mol–1 at 298 K. The pre-exponential factors, activation energies, and thermodynamics of activation for AcOH and 4H2B are compiled in Table 1. An alternative analysis using linear least-squares fits of ln (k/T2) vs 1/T, as shown in Figure S8 and Table S1 in Supporting Information, yields identical thermodynamic parameters within the experimental uncertainties.

The pressure dependence of the CH2OO + hydroxyketone reactions was also investigated by measuring kloss across the 80–120 Torr range at 295 K and using a single concentration of each hydroxyketone ([AcOH] = (5.1 ± 0.3) × 1015 cm–3 and [4H2B] = (8.0 ± 1.1) × 1014 cm–3). The total pressure was changed by varying the flow rate of the N2 buffer gas, while leaving all other flow rates unchanged. The average kloss values as a function of total pressure are shown in Figure S9 in Supporting Information. No change in kloss is observed as the total pressure is varied, and the average values are the same as those obtained in the kinetics measurements. Based on similar measurements for the CH2OO + Ac reaction,12,22 we conclude that the rate constants measured in 80–100 Torr of N2 likely represent the high-pressure limit.

Computational Results

We have performed ab initio calculations to characterize the reactions of CH2OO with AcOH and 4H2B. Geometries of the reactants, entrance channel complexes, TSs, and primary products were optimized, initially at the B3LYP/cc-pVDZ level of theory, and harmonic frequency analysis performed to confirm minima and saddle points. IRC calculations were performed to ensure that the TSs connected the reactant and product minima. Subsequent calculations were performed at the CBS-QB3 level of theory to provide improved thermochemistry. The Cartesian coordinates and energies for all species are compiled in the Supporting Information, and the CBS-QB3 thermochemistry data, Δ(E + ZPE) at 0 K and ΔH°, ΔG° at 298 K, are summarized in Table 2.

Table 2. CBS-QB3 Thermochemistry for the Reactions of CH2OO with AcOH and 4H2B: ΔE + ZPE @ 0 K (ΔH° @ 298 K) [ΔG° @ 298 K] in kcal mol–1.

    cycloaddition A cycloaddition B 1,2-addition
AcOH vdW –8.7 (−8.3) [+1.3] –8.7 (−8.5) [+2.1] –4.0 (−3.7) [+5.1]
  TS –5.6 (−6.5) [+6.9] –6.2 (−7.2) [+6.7] +2.2 (+1.4) [+13.6]
  product –48.1 (−49.5) [−34.9] –48.3 (−49.8) [−34.8] –48.0 (−48.9) [−35.9]
4H2B vdW –8.3 (−8.0) [+1.8] –8.9 (−8.6) [+1.3] –4.9(−4.6) [+4.1]
  TS –5.7 (−6.5) [+6.4] –6.7 (−7.6) [+5.6] –0.5 (−1.1) [+10.3]
  product –49.2 (−50.6) [−35.8] –49.6 (−51.0) [−36.2] –46.6 (−46.9) [−36.1]
Ac TS –6.1 (−6.4) [+5.5]    
MeOH TS     –2.0 (−3.1) [+9.1]
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Preliminary calculations at the B3LYP/cc-pVDZ level of theory were performed to assess the possible relevance of different hydroxyketone conformers. Four low-energy conformers were identified for each molecule, distinguished primarily by rotations about the α C–C bond and the C–O bond. The optimized geometries are shown in Figure 4. The most stable hydroxyketone conformers are distinguished by the presence of an intramolecular hydrogen bond, resulting in cyclic five and six membered structures for AcOH and 4H2B, respectively. For both hydroxyketones, the Boltzmann population distribution is dominated at all temperatures (>98%) by the intramolecular H-bonded conformers, which are at least 3 kcal mol–1 lower in energy than the next lowest conformer. Higher energy conformers were assumed to play no significant role in the reactions with CH2OO and were neglected in subsequent calculations.

Figure 4.

Figure 4

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Low-energy conformers of AcOH (left) and 4H2B (right), calculated at the B3LYP/cc-pVDZ level. The zero-point corrected energy (kcal mol–1) of each conformer relative to the most stable cyclic H-bonded structures is indicated.

Hydroxyketones can react with CH2OO as either carbonyls or alcohols. Reaction at the carbonyl site occurs via a 1,3-dipolar cycloaddition, leading to the formation of a five membered cyclic trioxolane, or SOZ. Reaction at the hydroxyl group can occur via a 1,2-addition (or insertion) mechanism, leading to a substituted hydroperoxide. The overall zero-point corrected energy [Δ(E + ZPE) at 0 K] and Gibbs free energy (ΔG° at 298 K) profiles are shown in Figure 5. The reactions are similarly exoergic with products lying at values of Δ(E + ZPE) ≈ −50 kcal mol–1 (ΔG° ≈ −35 kcal mol–1). The entrance channels of both reactions support van der Waals complexes that are bound by 4–9 kcal mol–1, while the TSs have energies that are generally below the separated reactants. On the free energy surface, all entrance channel complexes and TSs are higher in energy than the reactants.

Figure 5.

Figure 5

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CBS-QB3 energy and free-energy profiles for the CH2OO + AcOH (red) and 4H2B (blue). (a) Δ(E + ZPE) at 0 K. (b) ΔG at 298 K for both reactions. Solid lines connect stationary points for cycloaddition A, dotted lines for cycloaddition B, and dashed lines for 1,2-addition at the OH group. The inset shows a magnified view of the entrance channel van der Waals complex (vdW) and the transition state (TS) for the cycloaddition pathways.

Cycloaddition at the carbonyl site of each hydroxyketone can occur via two near-equivalent pathways that differ in the orientation of the CH2OO with respect to the OH group of the hydroxyketone. In pathway A, the central O atom of CH2OO is oriented toward the hydroxyalkyl side at TSA, while pathway B has the central O oriented toward the methyl side of the hydroxyketone at TSB. The optimized TSA and TSB geometries are shown in Figure 6. Unsurprisingly, the energies and free energies are similar. The free-energy barriers for the A and B pathways of the AcOH reaction differ by only 0.2 kcal mol–1, with a difference of 0.8 kcal is found for the 4H2B reaction, slightly favoring TSB. The larger energy difference for 4H2B arises because of a geometrical distortion of the hydroxyethyl group out of the plane, away from the attacking CH2OO, in TSA while the carbon backbone maintains planarity in TSB. The TS free energies for reaction at the OH site of the hydroxyketones are significantly higher than those for the cycloaddition reactions (almost 7 kcal mol–1 for the AcOH reaction and ∼4 kcal mol–1 higher for the 4H2B reaction), as shown in Figure 5 and Table 2.

Figure 6.

Figure 6

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Optimized TS geometries for the 1,3-dipolar cycloaddition reactions of CH2OO with AcOH (a,b) and 4H2B (c,d). The orientation of the central O atom of CH2OO oriented toward (a,c) or away (b,d) from the hydroxyl group of the hydroxyketone identifies cycloaddition A or B, respectively. Red arrows indicate displacement vectors along the reaction coordinate.

Discussion

The CH2OO + AcOH and CH2OO + 4H2B reactions are found to have almost identical rate constants across the 275–335 K temperature range studied experimentally, as summarized in Table 1. At room temperature (∼295 K), the rate constants for both reactions are 1.1 × 10–12 cm3 s–1, which is greater than the reactions of simple ketones by around a factor of 2. The rate constant for the CH2OO + acetone (Ac) reaction has been measured by various groups, with values in the range kAc = (2.3–4.8) × 10–13 cm3 s–1 reported,11,12,19−22 where the range is likely a consequence of the reaction pressure dependence. Measurements at the high-P limit have a weighted average value kAc = 4.3 × 10–13 cm3 s–1.11,12,21,22 The rate constant for the CH2OO + methylethylketone (MEK) reaction is slightly larger than that of Ac with a value kMEK = 6.4 × 10–13 cm3 s–1.70 The room temperature rate constants for the hydroxyketone reactions are much closer to that of acetaldehyde (MeCHO), for which rate constants in the range kMeCHO = (1.0–1.7) × 10–12 cm3 s–1 have been reported.19,20,23 The CH2OO + R1R2CO reactions also all show similar negative temperature dependences.

The reactivity of the hydroxyketones toward CH2OO is increased relative to aliphatic ketones, although the effect is largely insensitive to whether the OH is at the α or β position. The OH group provides an additional site for reaction via the 1,2-addition mechanism that has been characterized for alcohols.28,29,71,72 In general, alcohols tend to react with CH2OO much more slowly than carbonyls. For example, the rate constant for the CH2OO + methanol (MeOH) reaction at room temperature is kMeOH = 1.2 × 10–13 cm3 s–1.28,29 The increased rate constant for the hydroxyketones relative to acetone is likely due to enhancement of the 1,3-dipolar cycloaddition mechanism rather than being due to the presence of an additional reaction pathway. We note that synergistic rate constant increases arising from the presence of different functional groups have been demonstrated in other CI reactions. For example, 3-aminopropanol reacts with acetaldehyde oxide, CH3CHOO, in a concerted double hydrogen atom transfer step, where both the amine and hydroxyl functional groups interact with the CI simultaneously, significantly faster than simple amines or alcohols.27,73 Additionally, acetylacetone (AcAc), which exists predominantly as its enolone tautomer, has C=O, OH, and C=C sites for reaction with CH2OO. The CH2OO + AcAc reaction is twice as fast as CH2OO + Ac at room temperature (kAcAc = 8.0 × 10–13 cm3 s–1) and shows a weak temperature dependence (−Ea/R = 460 K) which can be explained in part by the existence of competitive pathways for reaction at both the C=O and C=C sites, where ab initio calculations find similar Δ⧧G° values at 298 K.11,12 Rate constants for CH2OO + alkene reactions are generally much smaller than that for carbonyls74 and show a positive temperature dependence. For the AcAc reaction, it appears that the adjacent carbonyl group may enhance the reactivity at the C=C site.

The experimental kinetics observations are supported by the ab initio calculations, which show the presence of relatively stable entrance channel complexes followed by TS barriers that are in most cases submerged relative to the reactants, consistent with the negative temperature dependences. Additionally, no dependence on total pressure was observed in the range 80–120 Torr of N2. A two-step mechanism for either the cycloaddition or 1,2-addition reactions can be written as

graphic file with name jp4c00156_m002.jpg R1
graphic file with name jp4c00156_m003.jpg R2

where the products are either SOZs or hydroperoxide species (see the Supporting Information). In the high-P limit, equilibrium is established for reaction R1, and the overall rate constant can be represented as k = k2K1, where K1 = k1/k–1. From the perspective of transition-state theory, the magnitude of the overall experimental rate constant is largely determined by the standard free energy of activation Δ⧧G° at the TS. The results of the CBS-QB3 ab initio calculations are summarized in Table 2 and Figure 5. Free-energy barriers for the cycloaddition reactions at the carbonyl are calculated to be broadly similar (∼+6 kcal mol–1), and analysis of the temperature dependence of the rate constants results in values that are in good agreement (see Table 1). In contrast, the calculated free-energy barriers for reaction at the hydroxyl group are significantly higher (>10 kcal mol–1). For comparison, the free energy of activation for CH2OO reacting with methanol is calculated to be ΔG° = +9.1 kcal mol–1 at the same CBS-QB3 level of theory (see Table 2). That is, the energy barrier for reaction at the OH position is even higher in the hydroxyketones than that in a simple alcohol like methanol, for which the observed rate constant is an order of magnitude lower. The higher energy barrier is consistent with the disruption of the intramolecular H-bond in the minimum energy structure of hydroxyketones required for reaction at the hydroxyl group. It is only the slightly higher-energy conformers, as shown in Figure 4, that can take part in the 1,2-addition with CH2OO. Consequently, we conclude that the dominant reaction between CH2OO and the hydroxyketones is the 1,3-dipolar cycloaddition at the C=O reaction site. The presence of the hydroxyl group on either the α or β carbon appears to have no significant effect on the observed rate constants, consistent with the similarity of the calculated free-energy barriers.

Previously,11 we have attempted to rationalize the reactivity trends in 1,3-dipolar cycloaddition reactions between CH2OO and carbonyl compounds (R1R2CO, where R1 and R2 are alkyl or acyl substituents) using a FMO theory approach.13−16 Starting with the model of symmetry-allowed orbital interactions developed by Sustmann,13 the cycloaddition between carbonyl compounds and CH2OO primarily involves interactions between the out-of-plane π and π* orbitals of CH2OO and R1R2CO. The dominant interaction is between the occupied nonbonding n(pC–pO) molecular orbital of the electron-rich species CH2OO (the 1,3-dipole) and the lowest unoccupied π* molecular orbital of the electron-deficient carbonyl, R1R2CO (the dipolarophile), with energy gap |ΔEA|. The energy gap |ΔES| between the occupied carbonyl π orbital and the unoccupied π* orbital of CH2OO is, in general, larger and makes a smaller contribution to the reactivity. The energies of the carbonyl FMOs are affected by the electron-donating or electron-withdrawing nature of the substituents R1 and R2. EDGs on the carbonyl raise the energy of the orbitals, increasing the magnitude of |ΔEA|, leading to a decreased reactivity. In contrast, EWGs lower the orbital energies and have the opposite effect on reactivity. The orbital interactions are illustrated in Figure 7, which shows the FMOs of CH2OO, and the carbonyls formaldehyde (HCHO), AcOH, 4H2B, and acetone (Ac), calculated at the B3LYP/cc-pVDZ level of theory.

Figure 7.

Figure 7

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Frontier orbital energies calculated at the B3LYP/cc-pVDZ level for CH2OO and the carbonyls formaldehyde, acetone, and the hydroxyketones AcOH and 4H2B. Red: A orbitals (nonbonding n(pC–pO) for CH2OO, π* LUMO for carbonyls), black: S orbitals (π bonding orbitals of carbonyls, π* antibonding for CH2OO), and gray: nonbonding (nO orbitals).

The EWG or EDG character of the R1 and R2 substituents can be represented by Hammett substituent constants, σm and σp, where the subscripts refer to substitution at the meta/para position of benzoic acid.18 Positive values indicate EWG character and increased reactivity, while negative values indicate EDG character and reduced reactivity. Among the carbonyls HCHO, AcOH, 4H2B, and Ac, the substituents and their Hammett constants are H (σm = σp = 0), CH3 (σm = −0.07, σp = −0.17), and CH2OH (σm = σp = 0), while values for CH2CH2OH are unknown. Methyl is electron-donating, raising the energies of the frontier orbitals while the others have no significant effect. The effect of the electron-donating methyl groups on the calculated FMO energies is evident in Figure 7 and Table 3. Relative to H, each methyl group substituent increases the carbonyl FMO energies by ∼0.4 eV. Based on the calculated orbital energies in Table 3, it is likely that CH2CH2OH is marginally more electron-donating than CH2OH due to the presence of the additional methylene group between the hydroxyl and the carbonyl. Although the experimental rate constants were found to be indistinguishable, accounting for the possible systematic error in the UV absorption cross section of AcOH described above would lead to a 10% increase in kAcOH, in line with the expectations of the orbital analysis.

Table 3. B3LYP/cc-pVDZ Calculated Orbital Energies (eV) of HCHO, AcOH, 4H2B, and Aca.

  HCHO AcOH 4H2B Ac
π* –1.116 –0.762 –0.735 –0.299
|ΔEA| 5.578 5.932 5.959 6.395
π –10.803 –9.959 –9.714 –9.388
|ΔES| 6.980 7.538 7.293 6.966
Open in a new tab
a

CH2OO π* E = −2.422 eV, n(pC–pO) E = −6.694 eV. |ΔEA| is the magnitude of the energy difference between the occupied n(pC–pO) orbital of CH2OO and unoccupied π* of the carbonyl (the dominant interaction). |ΔES| is the magnitude of the difference between the unoccupied π* orbital of CH2OO and the occupied π orbital of the carbonyl. The |ΔEA| interaction is always smaller.

Previously, we showed that the relationship between observed rate constants and calculated orbital energy gaps could be used quantitatively.11Figure 8 shows an updated plot of ln k against the magnitude of the orbital energy gap |ΔEA|, for a range of CH2OO + R1R2CO reactions. Experimental rate constants at room temperature have been obtained from various sources11,12,19,22,25,26,70,75,76 and are compiled in Table S2 of Supporting Information. The data set includes reactions of CH2OO with ketones, α-diketones, aldehydes, and α,β-unsaturated enones and enals. Where more than one experimental value is available, the weighted average rate constant is used. Orbital energies have been calculated at the B3LYP/cc-pVDZ level, rather than M06-2X/aug-cc-pVTZ as used previously.11 The effect of the electron-donating or electron-withdrawing character of the R1 and R2 substituents on the orbital energy gaps and rate constants leads to a strong negative linear correlation. The fastest reaction is with HFA, which has strongly electron-withdrawing substituents (R1 = R2 = CF3, σm = 0.43, σp = 0.54), while the slowest reaction is with acetone, which has electron-donating substituents (R1 = R2 = CH3). For species with R1 ≠ R2, such as acetaldehyde (R1 = H, R2 = CH3) or the hydroxyketones, the effect is additive. The α,β-unsaturated carbonyls such as MVK, MACR, ACR, and the enolone form of AcAc are shown in Figure 8 but clearly deviate from the trend and are not included in the fit. The |ΔEA| values for these species imply that the reactions with CH2OO should be much faster than observed experimentally,11,12,75,76 suggesting that delocalization of the π system reduces the reactivity of the carbonyl. Further work is required to explain the reactivity of α,β-unsaturated carbonyls toward CH2OO, which deviate from the trend.

Figure 8.

Figure 8

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Inverse correlation between reported experimental rate constants at room temperature for the reactions of CH2OO with a series of carbonyl compounds and the energy gap |ΔEA| between the π* orbital of the carbonyl and the n(pC–pO) orbital of CH2OO. Orbital energies were calculated at the B3LYP/cc-pVDZ level. A linear fit is shown, where the shaded area represents 1σ prediction bands. Experimental rate constants are drawn from various sources, see the text for details. Reactions involving α,β-unsaturated carbonyls (gray) were excluded from the fit.

Atmospheric Implications

The major reactive sink for hydroxyketones in the atmosphere is reaction with OH radicals. Rate constants for reaction with OH are 2.0 × 10–12 exp(−320/T) cm3 s–1 and 1.3 × 10–12 exp(−400/T) cm3 s–1 for AcOH and 4H2B, respectively.47,51 Typical lifetimes for both hydroxyketones reacting with OH radical are ∼3 days in the troposphere.38,39,43,44,49,51 Comparatively, photolysis has a minor contribution with lifetimes of ∼12–14 days for AcOH and 26 days for 4H2B.43,51 Temperature-dependent lifetimes for AcOH and 4H2B were estimated using typical average tropospheric concentrations for CH2OO and OH. The average CH2OO concentration was assumed to be 2 × 104 cm–3, although concentrations as high as 1 × 105 cm–3 have been reported.3,77 The average concentration of OH was estimated to be 5 × 106 cm–3 during the day and 2 × 105 cm–3 at night.78,79 At 295 K, the hydroxyketone reactions with CH2OO are insignificant, with estimated lifetimes >500 days. However, the CH2OO reactions show a strong negative T dependence and may become relatively more important at lower temperatures. While the hydroxyketone loss rates due to CH2OO increase at lower temperature, with lifetimes of ∼80 days at 220 K, OH remains the most important reactive sink for both hydroxyketones across the temperature range, even at night when OH concentrations are markedly lower.

Conclusions

The kinetics of the CH2OO + AcOH and CH2OO + 4H2B reactions were measured across the temperature range 275–335 K using a flash photolysis, transient absorption spectroscopy technique. The temperature-dependent bimolecular rate constants are kAcOH = (4.3 ± 1.7) × 10–15 exp[(1630 ± 120)/T] and k4H2B = (3.5 ± 2.6) × 10–15 exp[(1700 ± 200)/T]. Complementary ab initio calculations confirm that both reactions proceed via 1,3-dipolar cycloaddition at the carbonyl to form cyclic SOZs, while reaction at the hydroxyl group via 1,2-addition is insignificant. The increased reactivity of hydroxyketones relative to acetone can be understood from a FMO theory approach, wherein the cycloaddition involves an interaction between the occupied n(pC–pO) orbital of the CH2OO and the unoccupied π* orbital of the carbonyl. Electron-donating or electron-withdrawing substituents increase or decrease the π* orbital energy. Alkyl substituents are electron-donating, which decreases reactivity, while hydroxyalkyl substituents are, like H, neither electron-donating nor electron-withdrawing. A strong inverse correlation is found between the logarithm of the rate constants and the orbital energy gap for a range of R1R2CO species. The reactions of CH2OO with AcOH and 4H2B are unlikely to be significant in the troposphere, where reaction with hydroxyl radicals and photolysis control the hydroxyketone lifetimes.

Acknowledgments

This material is based upon work supported by the National Science Foundation under grant no. ECS-1905364.

Supporting Information Available

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acs.jpca.4c00156.

  • Description of the calibration procedure; FT-IR spectra of AcOH and 4H2B; UV absorption spectra of AcOH and 4H2B; concentration calibration plots; description of the kinetic model; example transient absorption spectra and [CH2OO] time profile; global pseudo-first-order plots; plot of ln(k/T2) against 1/T; experimental values of enthalpy, entropy, and Gibbs free energy of activation; pressure dependence of CH2OO loss rates; compilation of room temperature rate constants for CH2OO + R1R2CO reactions; and CBS-QB3 Cartesian coordinates and energies of the reactants, complexes, TSs, and products (PDF)

The authors declare no competing financial interest.

Supplementary Material

jp4c00156_si_001.pdf (1.7MB, pdf)

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