Kinetics of the reactions of the OH radical with 2-methyl-1-propanol, 3-methyl-1-butanol and 3-methyl-2-butanol between 241 and 373 K

Absolute rate constants have been measured for the gas-phase reactions of hydroxyl radicals with 2-methyl-1-propanol (k1), 3-methyl-1-butanol (k2) and 3-methyl-2-butanol (k3).

Experiments were carried out using two different techniques, the relative rate method and the pulsed laser photolysis-laser induced fluorescence technique.

The kinetic data were used to derive the following Arrhenius expressions in the temperature range 241–373 K(in units of cm3 molecule−1 s−1):k1 = (3.1±0.9) × 10−12exp[(352±82)/T]k2 = (2.8±0.9) × 10−12exp[(503±98)/T]k3 = (2.6±0.6) × 10−12exp[(456±65)/T]At 298 K, the reaction rate constants obtained by the two methods were in very good agreement.

The results are presented, discussed and used to estimate the atmospheric lifetimes for the studied alcohols.


Alcohols are emitted into the atmosphere by a wide number of anthropogenic and natural processes.1

The main sources are due to their use as solvent in different industries.

For example, 2-methyl-1-propanol (isobutanol) is commonly used in coatings, as an additive in gasoline, an intermediate for glycol ethers and agricultural chemicals, and in anti-corrosion and anti-wear additives in engine oils.

As a chemical intermediate, 2-methyl-1-propanol is used to make esters and ethers.

3-methyl-1-butanol (isopentanol) is used as a chemical intermediate and solvent, and in pharmaceutical products and medicines.

The release of these oxygenated volatile organic compounds is likely to contribute to the formation of ozone and other components of photochemical smog such as aldehydes, ketones and PAN-type compounds in urban areas.

To assess the impact of these chemical species on the environment, a detailed understanding of the kinetics and mechanisms of their atmospheric degradation is required.

The gas-phase atmospheric oxidation of alcohols is primary initiated by reaction with OH radicals.2,3

Several kinetic studies of the OH reaction with alcohols have been reported previously.2,3

However, only a limited number of studies have been conducted on the absolute measurements of the OH reaction rate constant with alcohols with more than four carbon atoms.

In this article, we report the rate coefficient data for the reactions of OH radicals with 2-methyl-1-propanol (k1), 3-methyl-1-butanol (k2) and 3-methyl-2-butanol (k3): OH + (CH3)2CHCH2OH→products: k1OH + (CH3)2CHCH2CH2OH→products: k2OH + (CH3)2CHCH(OH)CH3→products: k3Experiments were carried out using two different techniques.

The pulsed laser photolysis–laser induced fluorescence (PLP-LIF) technique was used to determine the absolute rate coefficients for the above reactions in the temperature range 241–373 K at around 100 Torr while the relative method enabled us to conduct the measurements at room temperature and atmospheric pressure.

As far as we know, this work provides the first temperature dependence studies for the three reactions.

The only existing data reported in the literature are those of Hai et al. for 2-methyl-1-propanol4 Saunders et al. for 3-methyl-1-butanol5 and Wallington et al. for 3-methyl-2-butanol.6

The present work also aims at extending the kinetic data base to be used to develop structure-reactivity relationships for OH reaction with oxygenated VOCs.

Experimental section

Relative measurements

The experimental set-up used to measure the relative rate constant consisted of a 120 L Teflon bag and lamps positioned in a wooden box with the internal faces covered with aluminium foil.

Measured amounts of reagents were flushed from calibrated bulbs into the Teflon bag through a stream of ultra pure air.

The photoreactor was then filled to its full capacity at atmospheric pressure with ultra-pure air.

OH radicals were produced by photolysing H2O2 using lamps emitting at 254 nm (Sylvania G30W) and illuminating the hole gas sample.

A gas chromatograph equipped with flame ionisation detector (GC-FID, CP-3800 or Star 3600 CX, Varian) was used for the quantitative analysis of the reactants.

Chromatographic separation was achieved by using a DB-1 capillary column (J&W Scientific, 30 m, 0.25 nm id, 5 μm film).

The column was operated at the temperature of 318 K, and helium was used as the carrier gas.

Absolute measurements

The pulsed laser photolysis–laser induced fluorescence (PLP-LIF) technique and the methodology used have been described in detail previously [e.g. in ref. 7], only a brief description is given here.

Hydroxyl radicals were produced by photolysis of H2O2 at λ = 248 nm (KrF excimer laser) and their concentration was monitored at various reaction times ranging from about 10 μs to 10 ms by pulsed laser induced fluorescence.

A Nd:YAG-pumped frequency-doubled dye laser was used to excite the OH radicals at λ = 282 nm.

Fluorescence from the OH radicals was detected by a photomultiplier, fitted with a 309 nm narrow bandpass filter.

The output pulse from the photomultiplier was integrated for a preset period by a gated charge integrator.

Typically the fluorescence signal from 10 different delay times and 100 probe laser shots were averaged to generate OH concentration-time profiles over at least three lifetimes.

H2O2 was introduced into the reaction cell by passing a small flow of helium through a glass bubbler containing H2O2 solution.

Alcohols were premixed with helium in a 10 L glass bulb to form 0.2–0.6% mixture at a total pressure of about 850 Torr.

The gas mixture, the photolytic precursor (H2O2), and the bath gas (helium) flowed through the cell with a linear velocity ranging from 3 to 10 cm s−1.

The gas mixture illuminated was then renewed every two laser shots.

The concentrations of the reactants were calculated from their mass flow rates, the temperature, and the pressure in the reaction cell.

All flow rates were measured with mass flow meters calibrated by measuring the rate of pressure increase in a known volume.

The pressure in the cell was measured with a capacitance manometer connected at the entrance of the cell.

The helium carrier gas (UHP certified to >99.9995% (Alphagas)) was used without purification.

The 50 wt % H2O2 solution, from Prolabo, was concentrated by bubbling helium through the solution to remove water for several days prior to use and constantly during the experiment.

2-methyl-1-propanol (≥99.8 %), 3-methyl-1-butanol (≥99.5) and 3-methyl-2-butanol (≥99.5%) were from Fluka, they were further purified by repeated freeze, pump, and thaw cycles and fractional distillation before use.

Results and discussion

Relative rate measurements

Relative rate constants were measured by comparing the OH reaction rate with alcohols to that with reference compounds:OH + alcohol → products: kalcoholOH + reference → products: krefAssuming that the compounds are only consumed by reaction with OH, it can be shown that:ln([alcohol]0/[alcohol]t) = (kalcohol/kref)ln([ref]0/[ref]t)where the subscripts 0 and t indicate concentrations before irradiation and at time t, respectively.

The rate constant for the reaction of OH with the three alcohols were measured at (298 ± 3) K and 760 Torr, relatively to the rate constant of OH with 1-butanol and 1,3-dioxolane.

The concentrations of alcohols and references were in the range 30–100 ppm.

The compounds showed a good stability when they were mixed in the chamber with H2O2 and kept in the dark for about 1 h.

Moreover, in absence of H2O2, the photolysis of the mixtures (alcohols and references in air) for more than one hour did not show any decrease in the concentrations.

Fig. 1 shows an example of the obtained relative loss of alcohols versus that of the reference compounds in the presence of OH.

In Table 1, the obtained rate constant ratios are given along with the derived rate constants for the OH reactions with the three alcohols.

The rate constants of the reaction of OH with the references used in this work were taken as (8.47 ± 0.34) × 10−12 for 1-butanol,8 and (11.1 ± 0.9) × 10−12 cm3 molecule−1 s−1 for 1,3-dioxolane.9

While averaging the values from different experiments, we obtained:k1 = (0.92 ± 0.15) × 10−11 cm3 molecule−1 s−1.k2 = (1.4 ± 0.2) × 10−11 cm3 molecule−1 s−1.k3 = (1.25 ± 0.2) × 10−11 cm3 molecule−1 s−1.The quoted errors represent twice the standard deviation, to which we have added an estimated error of 10% on the reference rate constants.

Absolute measurements

Kinetic experiments have been performed under pseudo-first-order conditions with the concentration of alcohols in large excess over that of OH concentration ([alcohol]0 > 100[OH]0).

Typically, the initial OH concentration, [OH]0, was around 2×1011 molecule cm−3.

The rate of disappearance of the OH radical followed a simple exponential rate law:[OH]t = [OH]0ekt where k′ = ki [alcohol] + k0ki represents the rate constants for the reaction of OH with the three alcohols.

The decay rate k′ is the first-order OH disappearance rate in the presence of the alcohol and k′0 is the first-order rate constant for OH removal in absence of alcohol (attributed to the diffusion of OH radicals out of the detection zone and to their reaction with H2O2).

Typically, alcohol and H2O2 concentrations were in the ranges (1.6–92)×1013 and (1–10)×1013 molecule cm−3, respectively. k′0 and k′ were in the ranges 60–140 s−1 and 200–13000 s−1, respectively.

Experiments were conducted in the temperature range 241–373 K and a total pressure of (106 ± 5) Torr of helium.

In all conditions, the OH decays were found to be exponential over at least three lifetimes.

The high [alcohol]/[OH]0 ratios and low OH concentrations made negligible contribution from secondary reactions involving the products of reactions to the measured rate constants.

The contribution of the reaction of OH with photofragments of alcohol was negligible since alcohols are not photolysed at 248 nm,10 the wavelength used to generate OH radicals.

As expected, variation in the photolysis fluence (3–26 mJ cm−2) had no effect on the determined rate constants.

The three alcohols were purified to better than 99.5% and hence loss of OH radicals by reaction with impurities in the gas mixtures is expected to be insignificant.

Fig. 2 shows examples of the plots of k′ versus the alcohol concentrations obtained at 298 K.

The obtained values of k1k3 and the experimental conditions are listed in table 2.

The room temperature rate constant taken as the average of all values obtained at (298 ± 2) K are:k1 = (1.0 ± 0.1)×10−11 cm3 molecule−1 s−1.k2 = (1.5 ± 0.1)×10−11 cm3 molecule−1 s−1.k3 = (1.2 ± 0.1)×10−11 cm3 molecule−1 s−1.The quoted errors for k1k3 include 2σ from the least-squares analysis and the systematic error (5%, due to uncertainties in measured concentrations).

These values are in good agreement with those obtained by the relative method.

The measured values of k1k3 shown in Table 2 are plotted in the Arrhenius form in Fig. 3.

An un-weighted least squares analysis of the ln kivs. 1/T plot yields the following expressions for the temperature dependence of k1k3 in the temperature range 241–373 K (in units of cm3 molecule−1 s−1 ):k1 = (3.1±0.9) × 10−12exp[(352±82)/T]k2 = (2.8±0.9) × 10−12exp[(503±98)/T]k3 = (2.6±0.6) × 10−12exp[(456±65)/T]Uncertainties are 2lnA and 2σE/R, for A and E/R, respectively.

The rate constants obtained in this work are compared with the previously reported data in Table 3.

For isobutanol, the measured value of k1 at room temperature is in good agreement with that obtained by Wu et al. using the relative rate method.4

These authors measured k1 relative to the rate constant of OH reactions with propane and cyclohexane and obtained (9.08±0.35) × 10−12 and (9.94±0.47) × 10−12 cm3 molecule−1 s−1, respectively.

Wu et al. measured also the rate constant for the reaction of OH with 3-methyl-1-butanol using the same references, they obtained (13.8±0.5) × 10−12 and (13.7±1.1) × 10−12 cm3 molecule−1 s−1 with propane and cyclohexane, respectively.

Our measured value of k2 and that of Wu et al4. are in agreement with that previously reported by Saunders et al. using the discharge flow–laser induced fluorescence (DF-LIF) technique at 1 Torr total pressure and 298 K (k2 = (13.1±2.6) × 10−12 cm3 molecule−1 s−1).5

For 3-methyl-2-butanol, our value of k3 at room temperature agrees with that reported earlier by Wallington et al. obtained using the flash photolysis–resonance fluorescence technique (k3 = (12.4±0.7) × 10−12 cm3 molecule−1 s−1).6

This work provides the first temperature dependence parameters for the OH reaction with 2-methyl-1-propanol, 3-methyl 1-butanol and 3-methyl 2-butanol.

The results show a negative temperature dependence for k1, k2 and k3 similarly to other alcohols for which temperature dependence data are available such as n-propanol, i-propanol and n-butanol.8

The present rate constant values k1k3 can be compared with those calculated from the structure–activity relationship (SAR) of Atkinson based on group reactivity using the revised substituent factors F(–OH) = 2.9 and F(–CH2OH) = F(CHOH) = F(COH) = 2..611,12

To calculate k1k3 we used also the following parameters as defined by Kwok and Atkinson: F(–CH3) = 1, F(–CH2–) = F(CH–) = F(C) = 1.23, kprim = 0.136, ksec = 0.934, ktert = 1.94 and kOH = 0.14 (k are in units of 10−12 cm3 molecule−1 s−1) where F(–CH3), F(–CH2–), F(CH–), and F(C) are the substituent factors and kprim, ksec, ktert and kOH represent the rate constants for the H-atom abstraction from –CH3, –CH2–, CH–, and –OH, respectively.11

The calculated k1k3 values compared to the experimental ones (in brackets) are k1 = 0.89 (0.96) ; k2 = 0.92 (1.45) and k3 = 1.28 (1.23) (units of 10−11 cm3 molecule−1 s−1).

While the experimental values of k1 and k3 are in good agreement with the calculated ones, that of k2 is a factor of 1.6 higher than that calculated using the SAR.

It has to be noticed that this kind of discrepancy between the experimental and calculated rate constants values for the reaction of OH with hydroxyl-containing organic compounds has already been mentioned earlier.12,13

The measured k2 value supports the statement of Bethel et al. indicating that accurate estimation of OH radical rate constants for all hydroxyl-containing compounds (to within better than a factor of two) does not appear possible using the general approach of the SAR developed by Atkinson.12

One of the possible reasons for this discrepancy is that the SAR of Atkinson considers only effects of the OH substituent on H-atom abstraction at the α- and the β- positions for hydroxyl-containing compounds while the available data shows that the oxygenated functional groups have long-range effects with respect to H-atom abstraction at sites remote from the substituent groups.3

This long-range effect indicates an alternative pathway to the direct concerted hydrogen abstraction process such as the formation of a hydrogen-bonded complex in which, first, a hydrogen bond is formed between the H atom of the OH radical and the O atom of the oxygenated compound then a second hydrogen bond is formed between the O atom of the OH radical and a H atom in the hydrocarbon chain, resulting in intermolecular H-atom transfer via a cyclic transition state.3,14

In all cases, the resulting reaction is an abstraction of a hydrogen atom from the oxygenated substrate.

At room temperature, reaction of OH radicals with alcohols proceeds mainly by an H-atom abstraction process from the C–H bonds of the CH3–, –CH2- and CH– groups with a minor contribution of the H-atom abstraction from the -OH group.2,15–20

Product distribution studies on the reaction of OH radicals with long chain linear alcohols in the presence of NOx in air have shown that H-atom abstraction from different positions (α, β, or γ) to the –OH group makes a significant contribution to the total reaction.2,11–16

Under atmospheric conditions, the resulting hydroxy alkyl radicals react with oxygen.

The α-hydroxy alkyl radicals formed through H-atom abstraction from the CHx group bonded to OH react solely with O2 to lead to carbonyls:2,20RR′CHOH + OH → RR′C˙OH + H2ORR′C˙OH + O2 → RR′C(O) + HO2The other hydroxy alkyl radicals (β, γ, etc.) react with O2 and NO to lead to different alkoxy hydroxy radicals which then react with O2 or decompose.

In addition, recent mechanistic studies have suggested that isomerization of this type of intermediate oxy radicals may also be important in the oxidation of long chain alcohols, through a six-membered transition state, leading to the formation of hydroxyketones and hydroxyaldehydes.16–19

In the gas phase, the primary degradation step of alcohols in the troposphere is reaction with OH radicals since other processes such as reactions with O3 and NO3, and photolysis are of minor importance.2,3

The measured OH reaction rate constants (k) along with a typical OH concentration of [OH] = 2 × 106 molecule cm−3,21 can be used to estimate the tropospheric lifetimes (τ = 1/k[OH]) of the studied alcohols.

The derived lifetimes are around 10 h for the three alcohols which indicate that these compounds are quickly removed from the atmosphere and give rise to other stable products that have longer lifetimes such as acetone.