Deprotonated Dicarboxylic Acid Homodimers: Hydrogen Bonds and

Article pubs.acs.org/JPCA

Deprotonated Dicarboxylic Acid Homodimers: Hydrogen Bonds and Atmospheric Implications Gao-Lei Hou,† Marat Valiev,*,‡ and Xue-Bin Wang*,† †

Physical Sciences Division and ‡Environmental Molecular Sciences Laboratory, Pacific Northwest National Laboratory, 902 Battelle Boulevard, P.O. Box 999, Richland, Washington 99352, United States S Supporting Information *

ABSTRACT: Dicarboxylic acids represent an important class of water-soluble organic compounds found in the atmosphere. In this work we are studying properties of dicarboxylic acid homodimer complexes (HO2C(CH2)nCO2−[HO2C(CH2)nCO2H], n = 0−12), as potentially important intermediates in aerosol formation processes. Our approach is based on experimental data from negative ion photoelectron spectra of the dimer complexes combined with updated measurements of the corresponding monomer species. These results are analyzed with quantum-mechanical calculations, which provide further information about equilibrium structures, thermochemical parameters associated with the complex formation, and evaporation rates. We find that upon formation of the dimer complexes the electron binding energies increase by 1.3−1.7 eV (30.0−39.2 kcal/ mol), indicating increased stability of the dimerized complexes. Calculations indicate that these dimer complexes are characterized by the presence of strong intermolecular hydrogen bonds with high binding energies and are thermodynamically favorable to form with low evaporation rates. Comparison with the previously studied HSO4−[HO2C(CH2)2CO2H] complex (J. Phys. Chem. Lett. 2013, 4, 779−785) shows that HO2C(CH2)2CO2−[HO2C(CH2)2CO2H] has very similar thermochemical properties. These results imply that dicarboxylic acids not only can contribute to the heterogeneous complexes formation involving sulfuric acid and dicarboxylic acids but also can promote the formation of homogeneous complexes by involving dicarboxylic acids themselves.

1. INTRODUCTION Organic acids represent an important component of atmospheric aerosols and have been found in abundance in a variety of urban,1−4 rural,4−6 and marine7 environments. While their significance in new particles formation and nucleation is supported by substantial experimental evidence,8−12 the precise mechanisms remain largely unclear. One of the particularly important questions is the role they play during the initial stage of aerosol particles formation. The relevant cluster structures, composition, and the nature of intermolecular interactions governing this process remain largely unknown.8 Such molecular-level information is not readily available from the current field measurements, which are limited to analysis of particle size larger than a few nanometers.8,13−16 A complementary way to investigate these issues, especially at the molecular scale, can be found in the analysis of clusters with tunable compositions and sizes.8,10−12,17,18 The development in this area has been so far mostly dominated by theoretical studies, which have provided useful information on the small clusters formed by atmospheric nucleation precursors. For example, Zhang and co-workers10,19 and Nadykto and coworkers20−22 as well as others11,23 have recently studied clusters consisting of a series of small organic species with sulfuric acid and/or H2O to investigate how organic acids enhance aerosol formation and growth. In addition to theoretical studies, several atmospherically relevant clusters consisting of sulfuric acid/ © XXXX American Chemical Society

bisulfate, nitric acid/nitrate, NH3/amine, and H2O have been investigated using infrared photodissociation spectroscopy,24−26 collision-induced dissociation,27 and mass spectrometry.28−30 Our approach to cluster characterizations, including those of atmospheric importance, consists of utilizing negative ion photoelectron spectroscopy (NIPES) measurements.12,18,31,32 These measurements are combined with accurate ab initio calculations to provide comprehensive and reliable molecular-level description of cluster properties. For example, we recently studied bisulfate−succinic acid (SUA) complexes and provided the first direct experimental evidence showing the thermodynamic advantage of organic (dicarboxylic) acid molecules in promoting bisulfate complexes/aerosols formation and growth.12 One of the important properties of organic acids during aerosol formation processes is their ability to form strong hydrogen bonds with nucleation precursors.10−12,19−22 This can provide significant stabilization energy and reduce the nucleation free energy barriers for particle growth. In addition, hydrogen bonds are ubiquitous in nature and are important in many other fields like supramolecular, biological, and medicinal systems, as well as enzymatic reactions.33−35 A good knowledge Received: February 2, 2016 Revised: March 22, 2016

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DOI: 10.1021/acs.jpca.6b01166 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

and small dimers HDCn−(H2DCn) with n ≤ 6, the aug-ccpVDZ basis set was used, while for large dimer complexes HDCn−(H2DCn) with n ≥ 7, the 6-31+G(d) basis set was used. We checked the performance of the two basis sets for n = 6 and found that both basis sets gave very similar structures with negligible differences (see Figure S1 in the Supporting Information). Harmonic vibrational frequencies analyses were performed to confirm that all optimized structures were true minima. Single-point energy calculations using M06-2X functional with the maug-cc-pVT(+d)Z basis set were performed. This combination of structure optimization using B3LYP functional followed by single-point energy calculations using M06-2X functional has been previously shown to give good agreement between calculations and experiments.12,18,47 Vertical detachment energy (VDE) was calculated as the total energy difference between the anionic and the neutral cluster both at the anion’s geometry. It should be pointed out that due to the flexibility of aliphatic chains and many possible ways to form hydrogen bonds, a plethora of close low-lying energy isomers may exist. This study does not aim for obtaining the global minimum for each structure; instead, we obtained low-energy structures by maximizing the number of hydrogen bonds and minimizing the steric strain energy of the aliphatic backbone. In addition, we also calculated the evaporation rates of these complexes by employing the method recently proposed by Ortega et al.48 (Details about evaporation rate calculations are described in the Supporting Information of this manuscript.) The calculations were conducted with the Gaussian0949 and NWChem50 suite of programs.

on the hydrogen bond strength will also be valuable to understand the reaction mechanisms and guide the drug and supramolecular design, but accurate measurement of hydrogen bond strengths is challenging. Previous studies12,36−40 have shown that NIPES can provide a useful estimate on the hydrogen bond strength, and the strengths of several lowbarrier hydrogen bonding interactions have been revealed in this way. In the present work, we are focusing on the singly deprotonated dicarboxylic acid anions HO2C(CH2)nCO2− (HDCn − for short hereafter) and their homogeneous complexes. These species constitute an important class of low-volatility compounds commonly found in atmospheric aerosols.41 Several studies have suggested that the formation of dimers and trimers may play a key role in the formation of initial nuclei and the subsequent growth.9,42−44 We have successfully used NIPES techniques in the past to characterize the chemical properties of dicarboxylic acid monomers and their complexes,12,18,45 and in the present work we extend the analysis to the homodimer complexes. Our investigation aims to provide detailed microscopic analysis of the process. Combining theoretical and experimental techniques data, we perform extensive structural characterization of the possible cluster systems and show the importance of hydrogen bonds in enhancing the stability of dimer complexes. We also explore the implications of these dimers in atmospheric particles formation.

2. EXPERIMENTAL METHODS The NIPES experiments were performed using a lowtemperature electrospray ionization (ESI) sourceNIPES apparatus, the details of which were previously described.46 The HDCn− and HDCn−(H2DCn) anions were produced via ESI of ∼0.1 mM H2DCn in acetonitrile/water (3:1 ratio) solvent with the solution pH tuned by adding ∼100 mM aqueous NaOH solution dropwise to pH ≈ 8. The produced anions were accumulated and collisionally cooled at 20 K by a cryogenically controlled three-dimensional ion trap to afford elimination of vibrational hot bands of the anions and to populate only the lowest energy minimum structures. The resultant cryogenic anions were transferred into the extraction zone of a time-of-flight (TOF) mass spectrometer. For each NIPES experiment, the HDCn− or HDCn−(H2DCn) anions were each mass-selected and decelerated before being photodetached by a 193 nm (6.424 eV) laser beam from an ArF excimer laser or a 157 nm (7.867 eV) laser beam from a F2 excimer laser in the interaction zone of a magnetic bottle photoelectron spectrometer. The lasers were operated at 20 Hz with the ion beam off at alternating laser shots, enabling shotby-shot background subtraction to be performed. Photoelectrons were collected with nearly 100% efficiency by the magnetic bottle and analyzed in a 5.2 m long calibrated electron flight tube. The electron binding energies presented in this work were obtained by subtracting the electron kinetic energies from the detachment photon energies. The energy resolution was ∼2%, that is, ∼20 meV for electrons with 1 eV of kinetic energy.

4. RESULTS AND DISCUSSION 4.1. Negative Ion Photoelectron Spectroscopy. Figure 1 presents the 20 K NIPE spectra of HDCn−(H2DCn) dimer anions at 157 nm (7.867 eV; red) and HDCn− monomer anions at 193 nm (6.424 eV; blue). The measured spectral features come from photodetachment of the oxygen lone pair electron on the carboxylate group.32,45 The 193 nm spectra of HDCn− anions at 20 K are consistent with the previously reported 193 nm spectra at 70 K.45 The spectral features for HDCn−(H2DCn) series show slowly rising features with onsets started from ∼6 eV, while for n = 0, a well-defined band centered at 6.7 eV is displayed. Since no vibrational progression is resolved in the onset region of the spectra, we estimated the electron binding energy of each complex anion using the threshold detachment energy (TDE)by adding the instrumental resolution to the electron binding energy at the crossing point of the spectral onset feature and the baseline. These are summarized in Table 1. For HDCn− (n = 1−10), the TDEs measured in this work are systematically slightly higher by 0.01 to 0.08 eV than the previously reported ones45 at 70 K. These differences are likely due to different temperatures (20 vs 70 K) under which the experiments were conducted, and they are within experimentally quoted uncertainties. As can be seen from the data, all the HDCn−(H2DCn) dimer anions have similar TDEs fluctuating around 6.0 eV. This indicates that aliphatic chain length has little effect on the electronic structure properties of the carboxylate group in these dimers, similar to the HDCn− cases.45 Compared to the monomer HDCn− anions,45 the TDEs of the dimer HDCn−(H2DCn) complexes are significantly higher, with differences ranging from 1.3 to 1.7 eV (Figure 1). This indicates that the formation of the dimer complex leads to the increased stabilization of the excess electron density. Given that the interaction between HDCn• and H2DCn is relatively

3. THEORETICAL DETAILS The equilibrium structures of HDCn− and HDCn−(H2DCn) anions, as well as H2DCn neutrals, were optimized with density functional theory using B3LYP functional without any symmetry constraints. For all monomers HDCn−, H2DCn, B


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Table 1. Experimental Threshold Detachment Energies of HDCn− and HDCn−(H2DCn) and Their Calculated Vertical Detachment Energies HDCn−(H2DCn) TDE (eV)

VDE (eV)

HDCn− TDE (eV)

VDE (eV)

calc

calc

n

exptla

B3LYPb

M06-2Xc

exptla

B3LYPd

M06-2Xe

0 1 2 3 4 5 6 7 8 9 10 11 12

5.85 6.25 6.35 6.25 6.00 6.15 6.05 6.10 6.15 6.10 6.10 6.05 6.15

5.76 5.92 5.86 5.83 5.84 5.63 5.67 (5.69) 5.61 5.67 5.71 5.70 5.69 5.67

6.60 6.89 6.93 6.79 6.77 6.55 6.55 (6.63) 6.37 6.43 6.53 6.52 6.52 6.51

4.35 4.70 4.70 4.60 4.52 4.49 4.58 4.63 4.70 4.67 4.70 4.75 4.75

4.64 4.67 4.82 4.76 4.68 4.67 4.70 4.57 4.54 4.71 4.72 4.66 4.66

5.08 5.25 5.30 5.22 5.05 5.00 5.07 4.92 4.95 5.09 5.09 5.08 5.05

The experimental uncertainty is ±0.1 eV. bB3LYP calculations with aug-cc-pVDZ for n = 0−5 and 6-31+G(d) for n = 7−12. For n = 6, results from both basis sets are shown with the number in parentheses using the smaller basis set, that is, 6-31+G(d). cSingle-point M06-2X/ maug-cc-pVT(+d)Z calculations based on B3LYP/aug-cc-pVDZ structures for n = 0−5 and B3LYP/6-31+G(d) structures for n = 7−12. For n = 6, results based on the structures optimized using the both basis sets are shown with the number in parentheses using the smaller basis set, that is, 6-31+G(d). dB3LYP/aug-cc-pVDZ (n = 0− 12) calculations. eSingle-point M06-2X/maug-cc-pVT(+d)Z // B3LYP/aug-cc-pVDZ (n = 0−12) calculations. a

sizes (n = 1−12) feature short O···O distances of ∼2.5 Å and nearly ideal ∠OHO angles of ∼180° (see Figure S3 for bond lengths and angles), indicating strong hydrogen-bond interactions.51 The protonated H2DCn neutral system retains folded configuration also stabilized by intramolecular H bonds (Figure 3), but the O−H···O hydrogen bonds are much weaker compared to their deprotonated counterparts. This weakening is because protonation of carboxylate group reduces negative electron density on the oxygen atoms participating in the hydrogen bond. The optimized structures of dimer complexes HDCn−(H2DCn) (n = 0−12) are shown in Figure 4. The formation of the dimers, for smaller sizes, n ≤ 6, retains folded structure of the constituents, HDCn− and H2DCn, which remain stabilized by intramolecular hydrogen bonding. Upon complexation, the intramolecular hydrogen bonds in HDCn− become slightly weaker, while the intramolecular hydrogen bonds in H2DCn simultaneously become slightly stronger. The dimer itself is kept intact by one strong intermolecular hydrogen bond with ∠OHO angle ranging from ∼170° to 180°. For large-size dimers, n ≥ 7, due to high flexibility of long aliphatic (CH2) chains, there emerges a second intermolecular hydrogen bond at the expense of the constituents’ intramolecular H bonds. 4.2.2. Energetic Analyses. The calculated VDEs for monomer, HDCn−, and dimer, HDCn−(H2DCn) (n = 0−12), anion species are listed in Table 1, and displayed in Figure 5 as a function of aliphatic chain length. It can be seen that the theoretical VDEs are systematically higher than the experimental TDEs for both monomer and dimer species. Such

Figure 1. The 20 K NIPE spectra of HDCn−(H2DCn) anions (red) recorded with 157 nm (7.867 eV) photons. For better comparison, the 20 K NIPE spectra of HDCn− (blue) recorded with 193 nm (6.424 eV) photons are also presented as insets. The electron binding energy positions in spectra from which TDEs are estimated are marked with green lines; the increase of TDE of each dimer anion with respect to that of HDCn− is also indicated.

weak, the increased electron stabilization energy can be related to the binding/interaction energy between HDCn− and H2DCn in the dimer systems (see Figure S2 in the Supporting Information for the schematic interpretation of the relationship between the electron binding energy difference and the binding energy difference). This will be analyzed in detail in the sections that follow. 4.2. Theoretical Calculations. 4.2.1. Structural Analyses. The optimized structures of HDCn− (n = 0−12) are shown in Figure 2. They appear similar to previously reported (n = 1− 10) structures45 and are characterized by folded configuration stabilized by intramolecular O−H···O hydrogen bonds. For n = 0, due to the limited aliphatic chain length, the O−H···O hydrogen bond is constrained with ∠OHO only ∼132°. Larger C


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Figure 2. Optimized B3LYP/aug-cc-pVDZ structures for singly deprotonated dicarboxylic acid anions HDCn− (n = 0−12). H bond lengths and angles are given in Figure S4 of the Supporting Information.

Figure 4. Optimized structures of HDCn−(H2DCn) at the B3LYP/augcc-pVDZ (n ≤ 6) and 6-31+G(d) (n ≥ 7) level of theory. H bond lengths and angles are given in Figure S6 of the Supporting Information.

Figure 5. Experimental TDEs (solid symbols) and calculated M06-2X VDEs (open symbols) for HDCn− (blue) and HDCn−(H2DCn) (red) as a function of the aliphatic chain length.

variation is expected, because experimentally measured TDEs include relaxation of the system after electron ejection, which is not present in VDEs. Overall the trends in calculated VDE match those of experimental TDEs reasonably well, particularly at smaller sizes (n < 4). For large sizes, the difference between the experimental TDEs and calculated VDEs is no longer

Figure 3. Optimized B3LYP/aug-cc-pVDZ dicarboxylic acid H2DCn (n = 0−12) structures. H bond lengths and angles are given in Figure S5 of the Supporting Information. D


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illustrated in Figure S2). Furthermore, since the interaction between HDCn− and H2DCn is mainly driven by strong intermolecular hydrogen bonding, ΔTDE can also serve as a measure of the strength of the all-formed intermolecular hydrogen bonds. Such experimental approach has been shown before to give reasonable estimates for the hydrogen bond strengths.36−39 To further validate this experimental scheme of estimating clustering binding energies based on the experimental ΔTDEs, we calculated the theoretical binding energies (BEs) between HDCn− and H2DCn for the n = 0−6 HDCn−(H2DCn) clusters. The M06-2X calculated BEs (Table 2) agree reasonably with the experimental values with the largest deviation of 0.4 eV and the average deviation of 0.25 eV. The qualitative agreement among experimental ΔTDEs and calculated ΔVDEs as well as theoretically computed BEs in these homodimer dicarboxylic acid hydrogen-bonded complexes indicates again that NIPES has the ability to provide useful estimates of hydrogen-bonding strengths in pure hydrogen-bonded complex systems. 4.3. Atmospheric Implications. 4.3.1. Thermochemical Analyses. The calculated thermochemical parameters (e.g., binding energy, enthalpy, entropy, and Gibbs free energy) for these HDCn−(H2DCn) complexes formation (n ≤ 6) at ambient conditions (1 atm and 298.15 K) are summarized in Table 3. (For the larger complexes HDCn−(H2DCn) n ≥ 7, we did not perform thermochemical calculations, since the basis sets used for HDCn−, H2DCn, and for HDCn−(H2DCn) are different. For details see Table S1 in the Supporting Information.) It can be seen that the absolute values obtained using the M06-2X functional are slightly larger (∼10% or less) than those obtained by the B3LYP functional. Since it was previously shown that the M06-2X energetic information was more reliable,12,47 we will employ the M06-2X results for discussion. The calculations show that all dimer complexes have large BE of ∼30 kcal/mol and that the Gibbs free energy changes (ΔG) for the complex formation are all ca. −20 kcal/mol. The large binding energies and negative Gibbs free energy changes indicate that these complexes are very stable and are thermodynamically favorable to form. Hence, they may survive for long times under atmospheric conditions (as will also be evidenced by the calculated evaporation rates, vide infra). Consequently, they may contribute to the reduction of the nucleation barriers and facilitate the growth of the initial nuclei

constant, most likely due to that not all of the optimized structures presented in Figure 4 are global minima. In addition, the long rising onsets in the experimental spectra may indicate large geometric changes upon removal of the excess electrons. Indeed, the optimized neutral structures for HDCn•(H2DCn) (n = 1−3; Figure S3 in the Supporting Information) show significant structural changes between the neutrals and their anionic counterparts. As shown in Figure 1, the formation of the dimer complex results in significant increase of electron binding energy. The latter can be described quantitatively by defining ΔTDE and ΔVDE as the differences between TDE and theoretical VDE values for monomer and dimer systems. As shown in Table 2, Table 2. Experimental Threshold Detachment Energies and Calculated Vertical Detachment Energies Increases of HDCn−(H2DCn) Relative to HDCn− as well as the Calculated Binding Energies ΔTDE/ΔVDE (eV)

BE (eV)

calc n

exptl

B3LYP

M06-2X

B3LYP

M06-2X

0 1 2 3 4 5 6 7 8 9 10 11 12

1.50 1.55 1.65 1.65 1.48 1.66 1.47 1.47 1.45 1.43 1.40 1.30 1.40

1.12 1.25 1.04 1.07 1.16 0.96 0.97 1.04 1.13 1.00 0.98 1.03 1.01

1.52 1.64 1.63 1.57 1.72 1.55 1.48 1.45 1.48 1.44 1.43 1.44 1.46

1.19 1.30 1.27 1.19 1.26 1.20 1.18

1.26 1.41 1.32 1.24 1.32 1.27 1.29

the experimental difference between monomer and dimer TDE (ΔTDE) spans the range between 1.30 and 1.66 eV, and similar behavior is observed theoretically in the differences of VDEs (ΔVDE). The agreement between experimental (ΔTDE) and calculated (ΔVDE) values is particularly good when using M06-2X exchange-correlation functional. The importance of ΔTDE is related to the fact that the increase in electron binding energy upon dimerization provides insight into the binding energies (interaction strength) between HDCn− and H2DCn (as

Table 3. Thermochemical Parameters for the Homodimer Dicarboxylic Acid Complex Formation and Evaporation Rates of the Neutral Dicarboxylic Acid Molecule Calculated at Ambient Conditions (298.15 K and 1 atm) BEa (kcal/mol)

ΔHa (kcal/mol)

ΔGa (kcal/mol)

γ (s−1)

ΔSa (cal/mol/K)

n

B3LYP

M06-2X

B3LYP

M06-2X

B3LYP

M06-2X

B3LYP

M06-2X

0 1 2 3 4 5 6

27.34 30.01 29.25 27.45 29.08 27.56 27.22

28.97 32.48 30.36 28.57 30.53 29.19 29.75

−26.85 −29.84 −28.90 −26.17 −28.69 −27.15 −26.83

−28.49 −32.31 −30.01 −27.29 −30.14 −28.78 −29.36

−17.97 −18.58 −19.14 −19.41 −18.35 −17.48 −16.82

−19.60 −21.05 −20.25 −20.53 −19.80 −19.11 −19.35

−29.78 −37.77 −32.74 −22.67 −34.68 −32.43 −33.57

−29.82 −37.77 −32.74 −22.67 −34.68 −32.43 −33.57

B3LYP 1.40 1.97 8.09 5.07 3.28 1.36 4.24

× × × × × × ×

10−3 10−3 10−4 10−4 10−3 10−2 10−2

M06-2X 8.64 3.07 1.25 7.69 2.85 8.76 5.98

× × × × × × ×

10−5 10−5 10−4 10−5 10−4 10−4 10−4

Binding energy, enthalpy, Gibbs free energy, and entropy changes are obtained by using the following equations: BE[HDCn−(H2DCn)] = E(HDCn−) + E(H2DCn) − E[HDCn−(H2DCn)] (with zero-point energy (ZPE) correction); ΔH[HDCn−(H2DCn)] = H[HDCn−(H2DCn)] − H(HDCn−) − H(H2DCn) (n = 0−6; with ZP enthalpy correction); ΔG[HDCn−(H2DCn)] = G[HDCn−(H2DCn)] − G(HDCn−) − G(H2DCn) (n = 0−6; with ZP Gibbs free energy correction); ΔS = (ΔH − ΔG)/T, T = 298.15 K. a

E


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The Journal of Physical Chemistry A into larger particles,8 consistent with the suggestions derived from mass spectrometric studies.42 Specifically, given our recent studies of bisulfate−SUA (succinic acid, H2DC2) complexes,12 it is instructive to compare the interaction of SUA molecule with succinate ion (SUA−, the singly deprotonated SUA anion) and SUA molecule with bisulfate ion. Previously12 we found that SUA molecule has significant thermodynamic advantage compared to H2O in promoting the bisulfate complexes formation and growth. This was attributed to high stability of bisulfate−SUA complexes induced by strong intermolecular hydrogen bonding. Upon the formation of HSO4−(SUA) complex, the electron binding energy increased by ∼1.7 eV (39.20 kcal/mol), which was reflected in calculated binding energy of ∼32 kcal/mol and ΔG of −20 kcal/mol. These thermochemical parameters for HSO4−(SUA) are comparable to those for SUA−(SUA) (Table S2), suggesting that SUA− is a potential ion as important as HSO4− when they interact with organic acids. Furthermore, it is interesting to note that the structure of SUA−(SUA) is very similar to the (SUA)2 part in the previously studied HSO4−(SUA)2 complex, which should more properly be recognized as (H2SO4)(SUA−)(SUA).12 The binding energy and ΔG for HSO4−(SUA)2 were calculated to be ∼50 and −27 kcal/mol, suggesting its high stability and thermodynamic propensity to form. We also calculated the binding energy of H2SO4 with SUA−(SUA) to be 31.38 kcal/mol and the ΔG for H2SO4 associating with SUA−(SUA) to form (H2SO4)(SUA−)(SUA) to be −19.77 kcal/mol. These results imply that SUA−(SUA) may bind H2SO4 molecules easily during the collision and increase growth rate. Considering the similar structures and thermochemical properties of all other size complexes HDCn−(H2DCn) with SUA−(SUA) (the n = 2 member of this series of homodimers), the conclusions drawn from succinic acid (n = 2) may be general and applicable to other dicarboxylic acid dimers. In the atmosphere, there are abundant organic dicarboxylic acids,6,9,41 and our study suggests that they may play important roles in stabilizing other ions or themselves by forming highly stable complexes and thus promoting organic aerosol nucleation processes. 4.3.2. Evaporation Rates. It is known that apart from thermodynamic issues, kinetics factors are also expected to play important roles in aerosol nucleation processes.8 In the atmosphere, the survival probability of initial aerosol nuclei, that is, a small cluster, depends on the competition between the evaporation rate and the collision rate of the cluster with a nucleation precursor molecule. In this context, a critical nuclei is referred to as a cluster with the size at which the growth and evaporation rates are equal. Hence, the knowledge of the evaporation rate of a specific cluster is critically important to understand its role in aerosol nucleation. The evaporation rates and vapor pressures of the H2DCn dicarboxylic acids have been studied extensively in the past,52 while those of the homogeneous complexes are not known. To address this issue, we estimated the evaporation rates of these complexes using recently developed theoretical arguments.48,53,54 Table S3 in the Supporting Information lists the calculated dipole moments (μ), polarizabilities (α), and the collision rates (β). The calculated evaporation rates for single neutral dicarboxylic acids in the HDCn−(H2DCn) complexes n = 0−6 are summarized in Table 3. It can be seen that the evaporation rates of the neutral dicarboxylic acid molecules are much smaller than 1 s−1, indicating that once the complexes form, they are highly stable and resistant to fragmentation due to

evaporation. It is also noted that the calculated evaporation rates are much smaller than the evaporations of sulfuric acid, ammonia, or dimethylamine from H2SO4−NH3 or H2SO4− (CH3)2NH complexes;48 both have been extensively studied and are believed to play key roles in the nucleation and growth of atmospheric aerosol particles. This is consistent with the suggestion from the thermochemical analyses, exhibiting the potentially important roles of organic dicarboxylic acids played in the initial nucleation of organic aerosols. In addition, it is also worth noting that the deviation of the calculated evaporation rates between M06-2X and B3LYP functionals are much larger than that of the calculated collision rates between these two functionals. This emphasizes the importance of experimental data to benchmark the reliability of the calculations on the thermochemical properties of small atmospherically relevant clusters.

5. CONCLUSIONS In this work, we investigated the singly deprotonated dicarboxylic acid homodimer complexes HDCn−(H2DCn) (n = 0−12) by NIPES and theoretical calculations. By comparing with the NIPE spectra of the corresponding singly deprotonated dicarboxylic acid anions HDCn−, the increases of electron binding energies due to dimerization were obtained, which amount to 1.3−1.7 eV. Theoretical calculations revealed that all of these homodimers are bonded together by strong intermolecular hydrogen bonds, and the calculated electron binding energy increases and binding energies (interaction strengths) agree reasonably with the experimentally obtained electron binding energy increases. This agreement indicates that the NIPES can potentially be used to quantitatively estimate hydrogen bond strengths for complicated hydrogenbonded systems. Thermochemical analyses and calculated evaporation rates indicate that these complexes are very stable and are thermodynamically favorable to form. Comparison of thermochemical properties of the SUA−(SUA) complex with the HSO4−(SUA) and HSO4−(SUA)2 complexes is also made, suggesting that singly deprotonated dicarboxylic acid anions, which are abundant in the atmosphere, are potentially important ions in promoting organic aerosol particle growth.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.6b01166. The optimized structures of HDC6−(H2DC6) at B3LYP/ aug-cc-pVDZ and B3LYP/6-31+G(d) levels of theory, respectively; schematic interpretation of the relationship between the electron binding energy difference and the binding energy difference; the optimized structures of HDCn•(H2DCn) (n = 1−3) neutrals at B3LYP/aug-ccpVDZ level of theory; the optimized structures of HDCn−, H2DCn, and HDCn−(H2DCn) (n = 0−12) with key bond lengths and angles labeled; details of the evaporation rates calculations; thermochemical parameters for the complex formation calculated at ambient conditions (298.15 K and 1 atm); comparison of the thermochemical parameters for HSO4−(SUA) and SUA−(SUA) (SUA = H2DC2); dipole moments (μ) and polarizabilities (α) as well as the collision rates (β) calculated at ambient conditions (298.15 K and 1 atm); F


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Cartesian coordinates for all structures in Figures 2−4. (PDF)

AUTHOR INFORMATION

Corresponding Authors

*Phone: +1 509-371-6459. E-mail: [email protected]. (M.V.) *Phone: +1 509-371-6132. E-mail: [email protected]. (X.B.W.) Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by EMSL Intramural Aerosol Science Theme Funding, and by U.S. Department of Energy (DOE), Office of Science, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences, and performed using EMSL, a national scientific user facility sponsored by DOE’s Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory, which is operated by Battelle Memorial Institute for the DOE. The theoretical calculations were conducted on the ScGrid and DeepComp 7000 of the Supercomputing Center, Computer Network Information Center of Chinese Academy of Sciences, as well as on the EMSL Cascade Supercomputer.

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