Carboxylic Acid Catalyzed Hydration of Acetaldehyde - The Journal of

Article pubs.acs.org/JPCA

Carboxylic Acid Catalyzed Hydration of Acetaldehyde Heather A. Rypkema,*,† Amitabha Sinha,‡ and Joseph S. Francisco§ †

Department of Atmospheric, Oceanic, and Space Sciences, University of Michigan, Ann Arbor, Michigan 48109, United States Department of Chemistry and Biochemistry, University of California San Diego, La Jolla, California 92093, United States § Department of Chemistry, H.C. Brown Building, Purdue University, West Lafayette, Indiana 47906, United States ‡

S Supporting Information *

ABSTRACT: Electronic structure calculations of the pertinent stationary points on the potential energy surface show that carboxylic acids can act effectively as catalysts in the hydration of acetaldehyde. Barriers to this catalyzed process correlate strongly with the pKa of the acid, providing the potential to provide the predictive capacity of the effectiveness of carboxylic acid catalysts. Transition states for the acid-catalyzed systems take the form of pseudo-six-membered rings through the linear nature of their hydrogen bonds, which accounts for their relative stability compared to the more strained direct and water-catalyzed systems. When considered as a stepwise reaction of a dimerization followed by reaction/complexation, it is likely that collisional stabilization of the prereactive complex is more likely than reaction in the free gas phase, although the catalyzed hydration does retain the potential to proceed on water surfaces or in droplets. Lastly, it is observed that postreactive diol−acid complexes are significantly stable (∼12−17 kcal/mol) relative to isolated products, suggesting the possibility of long-lived hygroscopic species that could act as a seed molecule for condensation of secondary organic aerosols.

■

INTRODUCTION Acetaldehyde (AA) is a prominent carbonyl compound in atmospheric chemistry, and the second most abundant aldehyde after formaldehyde. It has both biogenic1−4 and anthropogenic sources, including combustion and biomass burning,5,6 although overall global sources are not fully characterized.7 Acetaldehyde may also be produced in situ via the oxidation of multiple common alkyl species (i.e., methane and some others) and has been predicted to be on the order of 128 Tg/year,8 photochemical production from VOCs,9 and as a minor product in isoprene oxidation. Reactions of acetaldehyde in the atmosphere contribute to the budgets of ozone, HOx,10 and peroxyacetyl nitrate (PAN)11 among other species, and it is readily photolyzed. The prevalence of acetaldehyde and other oxygenated volatile organic compounds (OVOCs) is expected to be significant in controlling the OH radical budget in remote marine environments8,12 as well as those of OH and O3 in polar atmospheres.13 Combined with acetone and methanol, it has been measured to constitute up to 85% of the nonmethane volatile organic compounds (NMVOC) in North Atlantic marine air.14 Example rural site measurements in the United States and Europe measured a mean acetaldehyde concentration of 0.70 ppb,15,16 whereas urban measurements have found values up to 15.9 ppbv in China,17 and as high as 45.60 ppb in Brazil.18,19 The atmospheric lifetime of acetaldehyde has been estimated at 0.8 days.8 Its chemistry is of particular interest due to the recent demonstration of its photochemical tautomerization to vinyl alcohol, which could provide a © XXXX American Chemical Society

mechanism for the production of organic acids, whose in situ production is not currently fully accounted for.20 Read et al. found that acetaldehyde in marine air was consistently underpredicted by the global CAM-Chem model, suggesting at least one missing source.12 Because of acetaldehyde’s ubiquity in the atmospheric environment, and the many remaining questions surrounding sources, sinks, and reactivity of acetaldehyde, a comprehensive study of its chemistry is necessary to gain a full understanding of its complete role in atmospheric chemistry. Additionally, acetaldehyde can act as a hydrogen bond acceptor and is infinitely soluble in water, which suggests that its molecular interactions at high humidity, at surfaces, and in bulk-phase water droplets could contribute substantively to its ultimate chemical impact. It has already been shown that inorganic acids can catalyze its tautomerization to vinyl alcohol,21 which has significant atmospheric implications as mentioned above. To further the general understanding of water’s impact on acetaldehyde reactivity, we have undertaken to study the quantum chemistry of the hydration reaction. In the presence of water, acetaldehyde can undergo hydration to produce 1,1-ethanediol Special Issue: Mario Molina Festschrift Received: October 24, 2014 Revised: February 28, 2015

A

DOI: 10.1021/jp510704j J. Phys. Chem. A XXXX, XXX, XXX−XXX

Article

The Journal of Physical Chemistry A

In this work, we present the results of quantum chemical calculations on reactions 1A−1C for the purpose of evaluating the feasibility and implications of the hydration of acetaldehyde in the atmosphere.

■

COMPUTATIONAL METHODS The Gaussian 09 suite of programs51 was used to perform the quantum chemical calculations reported herein. Isolated molecules and pre- and postreactive complexes were optimized using multiple theoretical methodologies. These include the full and direct MP2 method using a 6-311++G(2df,2p) basis set and the hybrid M06-2X52 method with a MG3S basis set. The latter basis set was selected due to its consistent performance in predicting transition state geometries with the M06 method. Calculations for all stationary points were also performed using M06-2X with an aug-cc-pvtz basis set, but these have been omitted from our results for clarity, as the values they generated were virtually identical to those produced by the MG3S basis set. Following optimization of stationary points, frequency calculations were performed at each level of theory to obtain zero point energy corrections. Single point energies were also calculated using CCSD(T) with the 6-311++G(2df,2p) basis set at the optimized MP2 geometry of each species. Reported CCSD(T) energies have been zero-point corrected using the ZPE determined at the MP2 level of theory. Transition states were optimized to ensure a single imaginary frequency, which was confirmed to correspond to the desired reaction via visual inspection of the vibration, and intrinsic reaction coordinate (IRC) calculations were performed to ensure that the appropriate pre- and postreactive complexes corresponding to the reaction of interest had been found.

The product species is a geminal diol, which is not photolabile in the atmosphere22 and is expected to be more hygroscopic than acetaldehyde due to its increased oxygen functionalization.23 This species can furthermore act as either a donor or acceptor of hydrogen bonds, making it more versatile for forming complexes not only with water but also with other oxygenated organic compounds as a preliminary seed for oligomers capable of condensing into SOA. Hydration of acetaldehyde, as well as the reverse dehydration reaction, has been studied in the aqueous phase24−27 in synthetic and biochemical contexts, but to our knowledge the gas-phase reaction has not been directly examined, particularly in the context of atmospheric implications. Acid-catalyzed hydration of carbonyl compounds has been shown to conform to multidimensional Marcus Theory in the solvent phase,28 but the equivalent gas-phase study has not been reported due to the absence of experimental data. The dehydration of geminal diols in aqueous solution was found to have an activation barrier unaffected by electron withdrawing substituents, but the barrier was lowered for electron donating substituents.29 To evaluate the potential atmospheric impact of acetaldehyde hydration in the gas phase, we have undertaken to examine the energetics of the reaction under three catalytic conditions, as given by reactions 1A−1C. CH3CHO + H 2O → CH3(OH)2

(1A)

CH3CHO + 2H 2O → CH3(OH)2 + H 2O

(1B)

■

RESULTS AND DISCUSSION Analysis of the gas-phase hydration of acetaldehyde was conducted for the direct reaction as well as that catalyzed by water and a suite of eight carboxylic acids: formic, acetic, propionic, butyric, chloroacetic, fluoroacetic, oxalic, and glyoxylic monohydrate. Figure 1 depicts the shifts in the potential energy surface for the hydration reaction in the uncatalyzed, water catalyzed, and formic acid catalyzed cases.

CH3CHO + H 2O + RCOOH → CH3(OH)2 + RCOOH (1C)

where reaction 1A is uncatalyzed, reaction 1B is watercatalyzed, and reaction 1C is catalyzed by a carboxylic acid. Water catalysis of a broad range of reactions has been demonstrated both experimentally and theoretically. A classic example of water-catalyzed hydration is the addition of a water molecule to SO3 to form sulfuric acid.30−32 Water has been found to catalyze abstraction of hydrogen by OH at low temperatures, although its effect was determined not to have a significant atmospheric impact in the gas phase.33,34 More atmospherically significant is the role of water in mediating gasphase reactions with ozone,35 the OH + HCl reaction,36 and many others.37,38 Although the direct and water-catalyzed reactions are clear choices, we elected to also and specifically examine the catalytic effect of carboxylic acids because of their relatively high abundance in the atmosphere. Formic acid (FA), the most prevalent species, has been measured at levels as high as 20 ppbv,39 and FA has been shown to be an effective catalyst in the hydration and hydrolysis of several species, including SO340,41 (which has also been found to be catalyzed by sulfuric acid42), formaldehyde,43 glyoxal,44 ketene,45 and alkyl formate.46,47 FA has also been shown, along with water and the water dimer, to catalyze the tautomerization of nitroguanidine.48 Carboxylic acids have further been demonstrated to catalyze the decomposition of α-hydroxyalkyl hydroperoxides49 and carbonic acid.50

Figure 1. Potential energy profiles of the hydration of acetaldehyde in the presence and absence of catalysts. Reported energies are at the MP2/6-311++G(2df,2p) level of theory. B


Article


Figure 2. Structures of the complexes and transition state structures for reactions 1A−1C.

The touchstone of this understanding resides in the consideration of the very nature of hydrogen bonds and their origins. Traditionally, we are trained to think of hydrogen bonds as primarily electrostatic in nature; electronegative atoms such as O, N, and F withdraw electron density from adjacent protons, resulting in partial charge characteristics on each atom, which can attract an oppositely signed charge on another molecule to form a loose intermolecular association. However, there is also a covalent component to hydrogen bonds54 such that in many systems the most favorable geometry involves an 180° angle between the donor atom, the hydrogen, and the acceptor atom.55 Although the molecular environment does not always allow for this configuration, deviation from this linear arrangement comes with a cost of decreased stabilization due to poorer favorability for π delocalization. With this in mind, let us inspect the electrostatic and covalent properties of the transition state. Figure 3 provides the charge distribution on the isolated catalysts where relevant (a) and the transition states (b) in the uncatalyzed, water-catalyzed and FA-catalyzed systems. The structure of the direct hydration transition state is obviously strained, with an HBA OHO bond angle of 123.4°, and the destabilization of the hydrogen bond manifests in a longer hydrogen bond length between the hydrogen and the carbonyl oxygen: 1.33 Å compared with 1.28 and 1.19 Å for water and formic acid, respectively. Bond angles for HBA in the two catalyzed systems are 156.0° and 178.4° for water and formic acid, respectively, showing that the formic acid system is better able to exploit the semicovalent character of its hydrogen bond. Mulliken analysis at different theory levels predict different partial charges on the pertinent hydrogen of the

In each case, the system forms a prereactive complex via hydrogen bonding, whereas both catalyzed reactions exhibit a postreactive complex as well. No postreactive complex is observed in the uncatalyzed case, as the product is a single molecular species. The direct hydration barrier is too high (36.9 kcal/mol) for the reaction to be kinetically competitive under atmospheric conditions, but it is substantially decreased in both catalyzed cases. Although the water-catalyzed barrier remains relatively high (15.5 kcal/mol), the formic acid catalyzed barrier drops below the energy of the reactant entrance channel. This suggests that formic acid and other carboxylic acids could potentially facilitate the hydration of acetaldehyde to yield a geminal diol in the atmosphere. In all three cases the prereactive complexes and transition structure exhibit ring-like properties via their hydrogen bonding networks. The geometries of these stationary points are shown in Figure 2. It should be noted that in a previous study of formic acid catalyzed hydration of formaldehyde,53 several preliminary complexes and associated transition states were found. However, the ring-shaped transition state analogous to the structure depicted below was the lowest in energy at every level of theory. We have therefore confined our study to this family of transition state structures as the most kinetically favorable. It is clear from this and previously mentioned studies that formic acid is a superior catalyst to water in this type of hydration reaction, but the question of why has not been directly addressed. The structure of these species, the transition states in particular, provides an immediate qualitative explanation simply through the steric stability of the hydrogen bonding network. C


Article


these systems, with the varying substituents nonparticipatory except via the tuning of the catalysts’ molecular properties. This tuning manifests in a range of catalytic effectiveness such that the barrier height relative to reactants scales linearly with the pKa of the catalytic acid, as shown in Figure 4a. Although there is the expected deviation in barrier height values among theory levels, the trend prediction is consistent within each series. Furthermore, the difference in absolute values appears to be systematic for each methodology, such that the stabilization of the acid-catalyzed barrier relative to water catalysis shows highly consistent results, as shown in Figure 4b. Standard deviation among the predicted stabilization for the three methods ranges from 0.4 to 0.8 kcal/mol over the series. This trend is in accordance with the Brønsted acid catalysis law, which is well-studied in the solution phase, although it has not, to our knowledge, been reported specifically for systems in the gas phase. This result suggests a powerful predictive tool for estimating the effectiveness of additional carboxylic acid catalysts not addressed in this study. Energies of pertinent stationary points for all the systems studied are summarized in Table 1. Stability of pre- and postreactive complexes follows the same trend as the transition states, with superior catalysts producing more stable complexes. This is likely due to the subtle tuning of hydrogen bond strength imposed by the molecular properties of the molecular catalysts. The overall reaction is predicted to be exothermic at all levels of theory, with values comparable to the experimental exothermicity measured by aqueous-phase experiments. The overall hydration reaction is therefore thermodynamically viable in the atmosphere. Using even the most conservative estimatethe highest CCSD(T) valuesall of the acidcatalyzed barrier heights are predicted to be near zero or below.

Figure 3. Catalyst Mulliken charge distribution (a) and transition state hydrogen bonding (b) in the direct, water-catalyzed, and formic acidcatalyzed acetaldehyde hydration systems.

catalysts: 0.26 for water and 0.29 for formic acid (MP2/6-311+ +G(3df,3p), 0.24 and 0.28 (M06-2X/MG3S), and 0.30 and 0.32 (B3LYP/6-31G**). In each case, the slightly higher partial charge on the donating hydrogen of formic acid supports the more energetically favorable hydrogen bond. The angles for HBB are 158.7° and 174.7° for water and formic acid, whereas the bond length is significantly smaller for water (1.25 Å vs 1.46 Å). A similar analysis for the prereactive complexes shows a difference in bond stability with lengths decreasing from 1.93 Å to 1.88 Å to 1.67 Å as the effectiveness of the catalyst increases, whereas the bond angles increase from 123.4° to 155.6° to 176.8° as we proceed to more effective catalytic systems. In terms of complex stability, the catalyzed systems obviously have energetic advantages in that they each contain two hydrogen bonds compared with only one in the direct case. But the substantially greater difference in stabilization of the catalyzed transition structures compared to the stabilization of the complexes may be attributed to the disfavorable hydrogen bond angle as well as the steric strain of the four-membered ring in the direct hydration transition state. It is clear from Figures 2 and 3 that the enhanced stability of the fomic acid catalyzed system can be largely attributed to its more favorable geometric properties. The transition state is an eight-membered pseudoring in which the nearly linear hydrogen bonds allow the OHO moieties to be treated as a single pseudobond, giving it characteristics approximating a sixmembered ring. Hydrogen bond pseudorings are well-known and have a broad prevalence in the literature.55−60 Beyond formic acid, seven additional carboxylic acids were studied. All exhibited transition state and complex structures that were qualitatively identical to that of formic acid. The acid group is the structural driving force behind catalytic activity in

■

ATMOSPHERIC IMPLICATIONS A spontaneous termolecular collision among three specific molecules not including nitrogen or oxygen is obviously a lowprobability phenomenon. Therefore, it is worthwhile to consider the more realistic mechanism by which a preliminary dimer is formed, followed by the addition of the third species to the complex. This process can proceed through two distinct pathways: (1) dimerization of water and acetaldehyde, followed by complexation with the acid, and (2) dimerization of water and the acid, followed by complexation with acetaldehyde. Although a third process, in which the acetaldehyde forms a preliminary complex with the carboxylic acid, is theoretically

Figure 4. (a) Hydration barrier height vs pKa of the catalyst for the suite of reactions studied at multiple levels of theory. (b) Stabilization of carboxylic acid catalyzed barrier height values relative to water catalysis. D


Article

The Journal of Physical Chemistry A Table 1. Energies (kcal/mol) of Stationary Points Relative to Reactants for All Catalytic Systems Studieda pKa

stationary point

MP2/6-311++G(2df,2p)

M063X/3GS

CCSD(T)/6-311++G(2df,2p)

none

catalyst

N/A

water

15.7

propionic acid

4.87

butyric acid

4.82

acetic acid

4.75

formic acid

3.75

glyoxylic acid

3.31

chloroacetic acid

2.86

fluoroacetic acid

2.59

oxalic acid

1.23

pre TS pre TS post pre TS post pre TS post pre TS post pre TS post pre TS post pre TS post pre TS post pre TS post reactants products

−4.7 36.9 −11.0 15.5 −12.2 −15.8 −2.0 −18.2 −15.8 −2.0 −18.5 −15.8 −2.1 −19.0 −15.8 −2.6 −18.9 −17.3 −4.8 −19.5 −17.5 −4.7 −19.7 −17.2 −4.4 −19.7 −18.6 −8.0 −23.0 0.0 −6.1

−4.6 35.5 −11.3 12.0 −15.9 −16.1 −5.1 −20.6 −15.8 −4.5 −20.4 −16.4 −5.3 −21.9 −16.3 −5.8 −21.7 −17.6 −7.4 −21.9 −17.7 −7.1 −22.3 −17.6 −7.3 −22.3 −19.1 −10.9 −25.9 0.0 −9.1

−4.6 38.7 −10.6 18.2 −11.5 −15.5 0.19 −17.3 −15.5 0.14 −17.6 −15.5 0.02 −18.2 −15.4 −0.4 −18.0 −17.0 −2.8 −18.6 −17.0 −2.5 −18.8 −16.8 −2.4 −18.8 −18.1 −6.3 −22.0 0.0 −5.6

all a

Acid pKa values provided for reference.

Table 2. Gibbs Free Energy Changes and Calculated Kp (atm−1) at 298 K for the Complexation Processes Leading to Acetaldehyde Hydration species

water dimer ΔG (J)

acetaldehyde formic acid butyric acid glyoxylic acetic propionic oxalic water chloroacetic fluoroacetic

14124 8208 9938 14790 8004 10391 1021 13147 6679 6890

Kp (dimer) 3.3 3.6 1.8 2.6 4.0 1.5 6.6 5.0 6.7 6.2

× × × × × × × × × ×

10−03 10−02 10−02 10−03 10−02 10−02 10−01 10−03 10−02 10−02

AA-water → trimer ΔG (J) N/A −514.6 241.5 −4545 −777.1 280.9 −10090 16270 −6979 −4910

Kp (trimer) pathway 1 CA-water → trimer ΔG (J) N/A 1.2 × 9.1 × 6.3 × 1.4 × 8.9 × 5.9 × 1.4 × 1.7 × 7.3 ×

1000 10−01 1000 1000 10−01 1001 10−03 1001 1000

N/A 8168 6385 −5527 8205 5760 9147 17710 3949 5708

Kp (trimer) pathway 2 N/A 3.7 × 7.6 × 9.3 × 3.6 × 9.8 × 2.5 × 7.9 × 2.0 × 1.0 ×

10−02 10−02 1000 10−02 10−02 10−02 10−04 10−01 10−01

same order of magnitude, with pathway 1 being larger by a factor of 1−3. Given the value of Kp for the water− acetaldehyde dimer, and pressure values for urban acetaldehyde and water at high humidity, the largest concentration expected for the AA−water complex is roughly 5 pptv. Put another way, the ratio of free to water-complexed AA is expected to be just under 10 000 at 100% humidity. These values suggest that formic acid, which has been measured at levels as high as 20 ppb,39 as an example abundance values for the dimer could reach up to 25 ppt. However, an earlier study suggests that up to 12% of formic acid could exist in its water-complexed form.61

possible, we did not consider it due to the extremely low number density of these species relative to water. Free energy changes and equilibrium constants at 298 K for pathways 1 and 2 are presented in Table 2. Although the entropic disadvantage leads to low equilibrium constants, formation of the prereactive complex from the dimer is largely spontaneous or neutral in the case of water complexing with acetaldehyde followed by addition of a carboxylic acid (pathway 1). This process is less favorable in pathway 2, but because the preliminary water dimerization is more favorable for the acids than for acetaldehyde, the product of the two pertinent equilibrium constants remains within the E


Article


conditions. Stable precomplexes are formed through a network of hydrogen bonds, leading to a ring-shaped transition state in which the linear H-bonds may be approximated as pseudobonds to form the effective geometry of a six-membered ring. The overall reaction is exothermic, and the barrier heights are predicted to be near or below the energy of the entrance channel for all acids studied. The barrier heights to these catalyzed reactions scale linearly with pH, providing a predictive tool for the likely effectiveness of additional carboxylic acid catalysts, as well as their applicability to similar reactions. In the gas phase, formation of the prereactive complexes is likely to dominate over the full reaction, whereas both complexation and reaction have the potential to impact the chemistry on surfaces and in droplets. Both the prereactive and subsequent postreactive complexes are highly stable relative to both reactants and products and are likely to remain in this associated form, which is highly hygroscopic and could act as a seed for cloud or SOA formation.

Following both pathways, we can calculate the bimolecular rate constants of the preliminarily formed dimers with the pertinent using canonical transition state theory: k=

kBT Q⧧ e−Ea / kBT h Q dimerQ monomer

where the values of the partition functions (Q) are taken directly from the Gaussian 09 frequency and optimization calculations of the relevant species. The results of these calculations, performed at the mp2-2df level of theory and evaluated at 298 K, are summarized in Table S1 of the supplementary materials. We choose to include this in supplementary data because the rate constants were calculated for the purposes of rule-of-thumb estimation and not an elaborate rate model, which is beyond the scope of this work. Their values are not intended to constitute a recommendation for comprehensive kinetic modeling. Scaled relative to the preceding dimers, the bimolecular barrier heights for carboxylic acids in pathway 1 become slightly positive, on the order of 11 kJ or fewer. Both glyoxylic and oxalic acids maintain negative barrier heights; these values were replaced by a barrier value of zero for the purposes of this relatively simple calculation. Acid catalysis barrier heights for pathway 2 range from 0.55 kJ (oxalic acid) to 26 kJ (acetic acid). For reference, we have also included the stabilization energy of the postreactive complex relative to separated products. Though these values can easily be calculated from the values in Table 1, it may benefit the reader to obtain a sense of the degree of stabilization experienced by these postreactive complexes. Again taking formic acid as an example with a number density of 5 × 109 mol/cm3, our calculation estimates a lifetime of years for the acetaldehyde−water complex against this reaction in the free gas phase. These values suggest that this family of reactions is unlikely to have any significant atmospheric impact in the free gas phase, although the prereactive complexes could be potentially significant. An encounter between the AA−water dimer and an isolated acid favors formation of the trimer and is only enhanced by the stabilization provided by additional collisions. Although the hydration reaction is unlikely to have significance in the gas phase, there is better potential for reactions to occur on surfaces and in the bulk-phase droplets. The acid-catalysis rate constants are consistently 10−12 orders of magnitude faster than the water-catalyzed hydration of acetaldehyde. As a comparison, the direct bimolecular rate constant calculated at this same level of theory has a value of 6 × 10−40, which is roughly an additional 10 orders of magnitude slower than even the water-catalyzed case. Furthermore, in an aqueous or surface environment, the extreme stability of both pre- and postreactive complexes would be enhanced by additional stabilization through hydrogen bonding from the solvent cage. These complexes possess a wealth of oxygenated functionalization capable of forming hydrogen bonds not just with water but other oxygenated VOCs. These stable complexes provide an excellent seed for cloud formation as well as providing a mechanism for the association of molecules into oligomers or condense into SOA.

■

ASSOCIATED CONTENT

S Supporting Information *

Bimolecular rate constants for pathways 1 and 2 and postcomplex stabilization energies relative to products. This material is available free of charge via the Internet at http:// pubs.acs.org.

■ ■

AUTHOR INFORMATION

Notes

The authors declare no competing financial interest.

REFERENCES

(1) Fall, R.; Karl, T.; Hansel, A.; Jordan, A.; Lindinger, W. Volatile Organic Compounds Emitted after Leaf Wounding: On-Line Analysis by Proton-Transfer-Reaction Mass Spectrometry. J. Geophys. Res.Atmos. 1999, 104, 15963−15974. (2) Guenther, A.; Hewitt, C. N.; Erickson, D.; Fall, R.; Geron, C.; Graedel, T.; Harley, P.; Klinger, L.; Lerdau, M.; McKay, W. A.; et al. A Global Model of Natural Volatile Organic Compound Emissions. J. Geophys. Res.-Atmos. 1995, 100, 8873−8892. (3) Warneke, C.; Karl, T.; Judmaier, H.; Hansel, A.; Jordan, A.; Lindinger, W.; Crutzen, P. J. Acetone, Methanol, and Other Partially Oxidized Volatile Organic Emissions from Dead Plant Matter by Abiological Processes: Significance for Atmospheric Hox Chemistry. Global Biogeochem. Cycles 1999, 13, 9−17. (4) Karl, T.; Potosnak, M.; Guenther, A.; Clark, D.; Walker, J.; Herrick, J. D.; Geron, C. Exchange Processes of Volatile Organic Compounds above a Tropical Rain Forest: Implications for Modeling Tropospheric Chemistry above Dense Vegetation. J. Geophys. Res.Atmos. 2004, 109, D18306. (5) Yamada, K.; Hattori, R.; Ito, Y.; Shibata, H.; Yoshida, N. Carbon Isotopic Signatures of Methanol and Acetaldehyde Emitted from Biomass Burning Source. Geophys. Res. Lett. 2009, 36, L18807. (6) Holzinger, R.; Warneke, C.; Hansel, A.; Jordan, A.; Lindinger, W.; Scharffe, D. H.; Schade, G.; Crutzen, P. J. Biomass Burning as a Source of Formaldehyde, Acetaldehyde, Methanol, Acetone, Acetonitrile, and Hydrogen Cyanide. Geophys. Res. Lett. 1999, 26, 1161−1164. (7) Singh, H. B.; Salas, L. J.; Chatfield, R. B.; Czech, E.; Fried, A.; Walega, J.; Evans, M. J.; Field, B. D.; Jacob, D. J.; Blake, D.; et al. Analysis of the Atmospheric Distribution, Sources, and Sinks of Oxygenated Volatile Organic Chemicals Based on Measurements over the Pacific During Trace-P. J. Geophys. Res.-Atmos. 2004, 109, D15S07. (8) Millet, D. B.; Guenther, A.; Siegel, D. A.; Nelson, N. B.; Singh, H. B.; de Gouw, J. A.; Warneke, C.; Williams, J.; Eerdekens, G.; Sinha, V.; et al. Global Atmospheric Budget of Acetaldehyde: 3-D Model Analysis and Constraints from in-Situ and Satellite Observations. Atmos. Chem. Phys. 2010, 10, 3405−3425.

■

CONCLUSIONS We have demonstrated that carboxylic acids can successfully catalyze the hydration of acetaldehyde under atmospheric F


Article


Dehydration Reactions of Carbon and Silicon Geminal Diols. Phys. Chem. Chem. Phys. 2011, 13, 18507−18515. (30) Kolb, C. E.; Jayne, J. T.; Worsnop, D. R.; Molina, M. J.; Meads, R. F.; Viggiano, A. A. Gas Phase Reaction of Sulfur Trioxide with Water Vapor. J. Am. Chem. Soc. 1994, 116, 10314−10315. (31) Gonzalez, J.; Torrent-Sucarrat, M.; Anglada, J. M. The Reactions of SO3 with HO2 Radical and H2OHO2 Radical Complex. Theoretical Study on the Atmospheric Formation of HSO5 and H2SO4. Phys. Chem. Chem. Phys. 2010, 12, 2116−2125. (32) Morokuma, K.; Muguruma, C. Ab Initio Molecular Orbital Study of the Mechanism of the Gas Phase Reaction SO3 + H2O: Importance of the Second Water Molecule. J. Am. Chem. Soc. 1994, 116, 10316−10315. (33) Thomsen, D. L.; Kurten, T.; Jorgensen, S.; Wallington, T. J.; Baggesen, S. B.; Aalling, C.; Kjaergaard, H. G. On the Possible Catalysis by Single Water Molecules of Gas-Phase Hydrogen Abstraction Reactions by OH Radicals. Phys. Chem. Chem. Phys. 2012, 14, 12992−12999. (34) Elm, J.; Bilde, M.; Mikkelsen, K. V. Influence of Nucleation Precursors on the Reaction Kinetics of Methanol with the Oh Radical. J. Phys. Chem. A 2013, 117, 6695−6701. (35) Varandas, A. J. C. Odd-Hydrogen: An Account on Electronic Structure, Kinetics, and Role of Water in Mediating Reactions with Atmospheric Ozone. Just a Catalyst or Far Beyond? Int. J. Quantum Chem. 2014, 114, 1327−1349. (36) Buszek, R. J.; Barker, J. R.; Francisco, J. S. Water Effect on the OH + HCl Reaction. J. Phys. Chem. A 2012, 116, 4712−4719. (37) Buszek, R. J.; Francisco, J. S.; Anglada, J. M. Water Effects on Atmospheric Reactions. Int. Rev. Phys. Chem. 2011, 30, 335−369. (38) Vaida, V. Perspective: Water Cluster Mediated Atmospheric Chemistry. J. Chem. Phys. 2011, 135, 020901. (39) Usher, C. R.; Blatrusaitis, J.; Grassian, V. H. Spatially Resolved Product Formation in the Reaction of Formic Acid with Calcium Carbonate (1014) the Role of Step Density and Absorbed WaterAssisted Ion Mobility. Langmuir 2007, 23, 7039−7045. (40) Hazra, M. K.; Sinha, A. Formic Acid Catalyzed Hydrolysis of SO3 in the Gas Phase: A Barrierless Mechanism for Sulfuric Acid Production of Potential Atmospheric Importance. J. Am. Chem. Soc. 2011, 133, 17444−17453. (41) Long, B.; Long, Z.-w.; Wang, Y.-b.; Tan, X.-f.; Han, Y.-h.; Long, C.-y.; Qin, S.-j.; Zhang, W.-j. Formic Acid Catalyzed Gas-Phase Reaction of H2O with SO3 and the Reverse Reaction: A Theoretical Study. ChemPhysChem 2012, 13, 323−329. (42) Torrent-Sucarrat, M.; Francisco, J. S.; Anglada, J. M. Sulfuric Acid as Autocatalyst in the Formation of Sulfuric Acid. J. Am. Chem. Soc. 2012, 134, 20632−20644. (43) Hazra, M. K.; Francisco, J. S.; Sinha, A. Gas Phase Hydrolysis of Formaldehyde to Form Methanediol: Impact of Formic Acid Catalysis. J. Phys. Chem. A 2013, 117, 11704−11710. (44) Hazra, M. K.; Francisco, J. S.; Sinha, A. Hydrolysis of Glyoxal in Water-Restricted Environments: Formation of Organic Aerosol Precursors through Formic Acid Catalysis. J. Phys. Chem. A 2014, 118, 4095−4105. (45) Louie, M. K.; Francisco, J. S.; Verdicchio, M.; Klippenstein, S. J.; Sinha, A. Hydrolysis of Ketene Catalyzed by Formic Acid: Modification of Reaction Mechanism; Energetics, and Kinetics with Organic Acid Catalysis. J. Phys. Chem. A 2015, 10.1021/jp5076725. (46) Jogunola, O.; Salmi, T.; Eränen, K.; Mikkola, J.-P. Qualitative Treatment of Catalytic Hydrolysis of Alkyl Formates. Appl. Catal. AGen. 2010, 384, 36−44. (47) Jogunola, O.; Salmi, T.; Wärnå, J.; Mikkola, J.-P. Kinetic Studies of Alkyl Formate Hydrolysis Using Formic Acid as a Catalyst. J. Chem. Technol. Biotechnol. 2012, 87, 286−293. (48) Du, B.; Zhang, W. Catalytic Effect of Water, Water Dimer, or Formic Acid on the Tautomerization of Nitroguanidine. Comput. Theor. Chem. 2014, 1049, 90−96. (49) Kumar, M.; Busch, D. H.; Subramaniam, B.; Thompson, W. H. Role of Tunable Acid Catalysis in Decomposition of a-Hydroxyalkyl

(9) Millet, D. B.; Guenther, A.; Siegel, D. A.; Nelson, N. B.; Singh, H. B.; de Gouw, J. A.; Warneke, C.; Williams, J.; Eerdekens, G.; Sinha, V. Global Atmospheric Budget of Acetaldehyde: 3-D Model Analysis and Constraints from in-Situ and Satellite Observations. Atmos. Chem. Phys. 2010, 10, 3405−3425. (10) Muller, J.-F. Sources of Upper Tropospheric Hox: A ThreeDimensional Study. J. Geophys. Res. 1999, 104, 1705−1715. (11) Roberts, J. M. The Atmospheric Chemistry of Organic Nitrates. Atmos. Environ. A-Gen. 1990, 24, 243−287. (12) Read, K. A.; Carpenter, L. J.; Arnold, S. R.; Beale, R.; Nightingale, P. D.; Hopkins, J. R.; Lewis, A. C.; Lee, J. D.; Mendes, L.; Pickering, S. J. Multiannual Observations of Acetone, Methanol, and Acetaldehyde in Remote Tropical Atlantic Air: Implications for Atmospheric Ovoc Budgets and Oxidative Capacity. Environ. Sci. Technol. 2012, 46, 11028−11039. (13) Kuo, M. H.; Moussa, S. G.; McNeill, V. F. Surface Disordering and Film Formation on Ice Induced by Formaldehyde and Acetaldehyde. J. Phys. Chem. C 2014, 118, 29108−29116. (14) Lewis, A. C.; Hopkins, J. R.; Carpenter, L. J.; Stanton, J.; Read, K. A.; Pilling, M. J. Sources and Sinks of Acetone, Methanol, and Acetaldehyde in North Atlantic Marine Air. Atmos. Chem. Phys. 2005, 5, 1963−1974. (15) Lee, Y.-N.; Zhou, X.; Hallock, K. Atmospheric Carbonyl Compounds at a Rural Southeastern United States Site. J. Geophys. Res.-Atmos. 1995, 100, 25933−25944. (16) Slemr, J.; Junkermann, W.; Volz-Thomas, A. Temporal Variations in Formaldehyde, Acetaldehyde and Acetone and Budget of Formaldehyde at a Rural Site in Southern Germany. Atmos. Environ. 1996, 30, 3667−3676. (17) Ho, K. F.; Ho, S.; Dai, W. T.; Cao, J. J.; Huang, R.-J.; Tian, L.; Deng, W. J. Seasonal Variations of Monocarbonyl and Dicarbonyl in Urban and Sub-Urban Sites of Xi’an, China. Environ. Monit. Assess. 2014, 186, 2835−2849. (18) Grosjean, D.; Miguel, A. H.; Tavares, T. M. Urban Air Pollution in Brazil: Acetaldehyde and Other Carbonyls. Atmos. Environ., Part B 1990, 24, 101−106. (19) Corrêa, S. M.; Martins, E. M.; Arbilla, G. Formaldehyde and Acetaldehyde in a High Traffic Street of Rio De Janeiro, Brazil. Atmos. Environ. 2003, 37, 23−29. (20) Andrews, D. U.; Heazlewood, B. R.; Maccarone, A. T.; Conroy, T.; Payne, R. J.; Jordan, M. J. T.; Kable, S. H. Photo-Tautomerization of Acetaldehyde to Vinyl Alcohol: A Potential Route to Tropospheric Acids. Science 2012, 337, 1203−1206. (21) Karton, A. Inorganic Acid-Catalyzed Tautomerization of Vinyl Alcohol to Acetaldehyde. Chem. Phys. Lett. 2014, 592, 330−333. (22) Epstein, S. A.; Nizkorodov, S. A. A Comparison of the Chemical Sinks of Atmospheric Organics in the Gas and Aqueous Phase. Atmos. Chem. Phys. 2012, 12, 8205−8222. (23) Hemming, B. L.; Seinfeld, J. H. On the Hygroscopic Behavior of Atmospheric Organic Aerosols. Ind. Eng. Chem. Res. 2001, 40, 4162− 4171. (24) Sorensen, P. E.; Jencks, W. P. Acid- and Base-Catalyzed Decomposition of Acetaldehyde Hydrate and Hemiacetals in Aqueous Solution. J. Am. Chem. Soc. 1987, 109, 4675−4690. (25) Bell, R. P.; Rand, M. H.; Wynne-Jones, K. M. A. Kinetics of the Hydration of Acetaldehyde. Trans. Faraday Soc. 1956, 52, 1093−1102. (26) Cheshnovsky, D.; Navon, G. Nuclear Magnetic Resonance Studies of Carbonic Anhydrase Catalyzed Reversible Hydration of Acetaldehyde by the Saturation Transfer Method. Biochemistry 1980, 19, 1866−1873. (27) Nakazawa, Y.; Takahashi, F. Thermodynamic and Magnetic Studies on the Formation of gem-Diol in an Aldehyde-Water Mixture. Bull. Chem. Soc. Jpn. 1999, 72, 975−980. (28) Guthrie, J. P. Hydration of Carbonyl Compounds, an Analysis in Terms of Multidimensional Marcus Theory. J. Am. Chem. Soc. 2000, 122, 5529−5538. (29) Ignatyev, I.; Montejo, M.; Ortega, P. G. R.; Gonzalez, J. J. L. Effect of Substituents and Hydrogen Bonding on Barrier Heights in G


Article

The Journal of Physical Chemistry A Hyddroperoxides and Mechanistic Implications for Tropospheric Chemistry. J. Phys. Chem. A 2014, 118, 9701−9711. (50) Kumar, M.; Busch, D. H.; Subramaniam, B.; Thompson, W. H. Organic Acids Tunably Catalyze Carbonic Acid Decomposition. J. Phys. Chem. A 2014, 118, 5020−5028. (51) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; et al. Gaussian 09, Revision D.01; Gaussian, Inc.: Wallingford, CT, 2009. (52) Zhao, Y.; Truhlar, D. The M06 Suite of Density Functionals for Main Group Thermochemistry, Thermochemical Kinetics, Noncovalent Interactions, Excited States, and Transition Elements: Two New Functionals and Systematic Testing of Four M06-Class Functionals and 12 Other Functionals. Theor. Chem. Acc. 2008, 120, 215−241. (53) Hazra, M. K.; Francisco, J. S.; Sinha, A. Gas Phase Hydrolysis of Formaldehyde To Form Methanediol: Impact of Formic Acid Catalysis. J. Phys. Chem. A 2013, 117, 11704−11710. (54) Alkorta, I.; Arunan, E.; Clary, D. C.; Crabtree, R. H.; Dannenberg, J. J.; Desiraju, G. R.; Hobza, P.; Kjaergaard, H. G.; Klein, R. A.; Legon, A. C.; et al. Defining the Hydrogen Bond: An Account (Iupac Technical Report). Pure Appl. Chem. 2011, 83, 1619. (55) Pakiari, A. H.; Eskandari, K. The Chemical Neture of Very Strong Hydrogen Bonds in Some Categories of Compounds. J. Mol. Struct.: THEOCHEM 2006, 759, 51−60. (56) Gilli, G.; Gilli, P. Towards a Unified Hydrogen Bond Theory. J. Mol. Struct. 2000, 552, 1−15. (57) Sakamoto, T.; Koga, Y.; Hikota, M.; Matsuki, K.; Murakami, M.; Kikkawa, K.; Fujishige, K.; Kotera, J.; Omori, K.; Morimoto, H.; et al. Design and Synthesis of Novel 5-(3,4,5-Trimethoxybenzoyl)-4Aminopyrimidine Derivatives as Potent and Selective Phosphodiesterase 5 Inhibitors: Scaffold Hopping Using a Pseudo-Ring by Intramolecular Hydrogen Bond Formation. Bioorg. Med. Chem. Lett. 2014, 24, 5175−5180. (58) Asselin, P.; Soulard, P.; Madebene, B.; Goubet, M.; Huet, T. R.; Georges, R.; Pirali, O.; Roy, P. The Cyclic Ground State Structure of the HF Trimer Revealed by Far Infrared Jet-Cooled Fourier Transform Spectroscopy. Phys. Chem. Chem. Phys. 2014, 16, 4797− 4806. (59) Azofra, L. M.; Scheiner, S. Complexation of N SO2 Molecules (N = 1, 2, 3) with Formaldehyde and Thioformaldehyde. J. Chem. Phys. 2014, 140, 034302. (60) Jeffrey, G. A. Crystallographic Studies of Carbohydrates. Acta Crystallogr., Part B 1990, 46, 89−103. (61) Iuga, C.; Alvarez-Idaboy, J. R.; Vivier-Bunge, A. Mechanism and Kinetics of the Water-Assisted Formic Acid + Oh Reaction under Tropospheric Conditions. J. Phys. Chem. A 2011, 115, 5138−5146.

H


Carboxylic Acid Catalyzed Hydration of Acetaldehyde - The Journal of

Recommend Documents