Side-by-Side Comparison of Hydroperoxide and Corresponding

Subscriber access provided by University of Newcastle, Australia

Article

Side by Side Comparison of Hydroperoxide and Corresponding Alcohol as Hydrogen Bond Donors Kristian H. Møller, Camilla Mia Tram, and Henrik G. Kjaergaard J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.7b01323 • Publication Date (Web): 01 Apr 2017 Downloaded from http://pubs.acs.org on April 4, 2017

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry A is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Side by Side Comparison of Hydroperoxide and Corresponding Alcohol as Hydrogen Bond Donors Kristian H. Møller, Camilla Mia Tram, and Henrik G. Kjaergaard

∗

Department of Chemistry, University of Copenhagen, Universitetsparken 5, DK-2100 Copenhagen Ø, Denmark

E-mail: [email protected]

Phone: +45-35320334. Fax: +45-35320322

Abstract Hydroperoxides are formed in signicant amounts in the atmosphere by oxidation of volatile organic compounds and are key in aerosol formation. In a room temperature experiment, we have detected the formation of bimolecular complexes of tert -butyl hydroperoxide (t -BuOOH) and the corresponding alcohol, tert -butanol (t -BuOH), with dimethyl ether (DME) as the hydrogen bond acceptor. Using a combination of Fourier transform infrared (FT-IR) spectroscopy and quantum chemical calculations, we compare the strength of the OH-O hydrogen bond and the total strength of complexation. We nd that both in terms of observed redshifts and determined equilibrium constants, t -BuOOH

is a signicantly better hydrogen bond donor than t -BuOH, a result which

is backed by a number of calculated parameters and can be explained by a weaker OHbond in the hydroperoxide. Based on combined experimental and theoretical results, we nd that the hydroperoxide complex is stabilized by about 4 kJ/mol (Gibbs free energy) more than the alcohol complex. Measured redshifts show the same trend in 1

ACS Paragon Plus Environment


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

hydrogen bond strength with trimethylamine (N acceptor atom) and dimethyl sulphide (S acceptor atom) as the hydrogen bond acceptors.

Introduction Hydroperoxides are formed in the atmosphere in large quantities under low NO x conditions from both conventional oxidation pathways and autoxidation. 1,2 Several studies, both modelling and experimental, have suggested that hydroperoxides represent a large constituent of secondary organic aerosols (SOA), some nding that it may be the major constituent. 35

Aerosol particles in the atmosphere have a large impact on the climate. 6 They aect the Earth's radiative balance due to their ability to scatter and absorb solar radiation (direct eect), but also due to their ability to act as cloud condensation nuclei and aect the radiative scattering of clouds (indirect eect). Aerosol particles still remain important sources of uncertainty in current climate models. 6

Models describing the formation of clusters at the molecular level are used in an attempt to achieve a better understanding of the formation and growth of SOA. 7 Networks of hydrogenbonded molecules have been discovered in these clusters and hydrogen bonds are believed to be a driving force in the formation and growth of clusters in the atmosphere, with the strength of the formed hydrogen bonds determining the thermodynamic stability of the clusters. 710

In spite of the importance of both hydroperoxides and hydrogen bonding for SOA formation and growth, the hydrogen bonding ability of organic hydroperoxides has, to the best of our knowledge, not been studied in the gas phase at conditions resembling the atmosphere. In the liquid phase, on the other hand, several studies of the hydrogen bonding of hydroperox-

2


Page 2 of 32

Page 3 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


ides have been carried out with the earlier results reviewed by Yabonskii et al. in 1972. 11 In a solvent of CCl4 , the presence of an intramolecular hydroperoxide OH- π hydrogen bond has been observed in isopropylbenzene hydroperoxide based on the observation of two peaks in the OH-stretching region, and similarly, an intramolecular OH-O hydrogen bond has been reported in 2,4-dimethyl-2-hydroperoxy-3-pentanon. 12,13 In various solvents, the dimerization and formation of higher order multimers of dierent hydroperoxides has been observed based on both IR, Raman and NMR spectroscopy. The hydroperoxides include tert -butyl hydroperoxide, deuterated tert -butyl hydroperoxide, isopropylbenzene hydroperoxide, tert pentyl hydroperoxide, phenyl cyclohexane hydroperoxide, 3-phenylmethyl hydroperoxide and two hydroperoxides from the oxidation of 2-methyltetrahydrofuran. 1420 Similarly, based on observed redshifts of the OH-stretching vibration in IR spectroscopy and changes in chemical shifts in NMR, complexes with various hydroperoxides as hydrogen bond donors have been reported in dierent solvents with a wide range of hydrogen bond acceptors: ketones, ethers, amines, alcohols and various aromatic compounds. 15,16,2123 One of these studies comparing a number of dierent hydrogen bond acceptors conclude that the hydrogen bonding ability of hydroperoxides in the liquid phase is between that of alcohols and phenol. 22

A few complexes of hydrogen peroxide as the hydrogen bond donor have also been reported in argon matrices including the hydrogen peroxide dimer and its complexes with water and dimetylether, all observing the hydroperoxide acting as a hydrogen bond donor. 2426 The complex with dimethyl ether is compared to the analogous complex of water nding that hydrogen peroxide is a better hydrogen bond donor. 24 Furthermore, theoretical studies have been carried out for the hydrogen peroxide-water complex and the complex of hydrogen peroxide with methyl hydroperoxide. 27,28

To resolve the relative hydrogen bonding ability of hydroperoxides to alcohols in the gas phase, we have used a combination of room temperature gas phase Fourier transform in-

3



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

frared (FT-IR) spectroscopy and quantum chemical calculations to study the hydrogen bond and complexation strength of tert -butanol (2-methylpropan-2-ol, t-BuOH) and tert -butyl hydroperoxide (2-methylpropane-2-peroxol, t -BuOOH) with dimethyl ether (DME) as hydrogen bond acceptor. DME is chosen as a simple model system for the wide range of organic atmospheric species with oxygen atoms. In addition, we have combined experiments and calculations to determine the equilibrium constant, or in other words the Gibbs free energy, ∆G , of formation for the two complexes. We characterize the interactions between the monomers in the two complexes with topological methods. Finally, we also measure the OH-stretching redshifts in the complexes of t -BuOH and t -BuOOH with trimethylamine (TMA) and dimethyl sulphide (DMS), which have nitrogen and sulphur as the acceptor atom, respectively, to assess the generality of the observed trend.

Experimental Dimethyl ether (Sigma-Aldrich, ≥99.9%) was used without further purication. tert-butanol (Sigma-Aldrich, anhydrous ≥99.5%) and tert-butyl hydroperoxide solution (Sigma-Aldrich, ∼5.5

M in decane over molecular sieves) were degassed by several freeze-pump-thaw cycles.

Due to the low vapor pressure of decane, the measured pressure of the t -BuOOH solution was not corrected for the presence of decane in the gas phase (see section S1.1). Attempts at distillations to separate the hydroperoxide from the decane proved unsuccessful, likely due to the instability of the hydroperoxide.

The FT-IR spectra were recorded at room temperature (295 K - 298 K) with a Bruker Vertex 70 or 80 FT-IR spectrometer equipped with an MIR light source, a KBr beam splitter and a liquid nitrogen cooled MCT detector. The spectra were recorded with a spectral resolution of 1 cm-1 and averaged over 500 scans. An aperture of 2-3.5 mm was employed.

4


Page 4 of 32

Page 5 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


The spectra were recorded with 2.4 m, 4.8 m, and 6.0 m multireection White cells (Infrared Analysis, Inc.) equipped with internal gold coated mirrors, external aluminium coated mirrors, and KCl windows. Background spectra were recorded with an evacuated cell and were subtracted from each of the sample spectra. The cells were evacuated on a glass vacuum line (J. Young) equipped with a rotary pump (Edwards RV3) and a turbo molecular high vacuum pump (Edwards E2M1.5 or Varian AX-65), creating a base pressure of ∼10−4 Torr. The pressure was measured using an Agilent PCG-750 Pirani capacitance diaphragm gauge and corrected for gas type dependence using an Agilent CDG-500 pressure gauge as described section S1.2. The pressure gauges were attached directly to the gas cells.

The t-BuOH·DME and t-BuOOH·DME mixtures were prepared in a glass mixing bowl which was attached to the vacuum line. The two components in each mixture were allowed to mix for more than 30 minutes to ensure that an equilibrium had been reached. The mixing bowl was covered by a piece of black cloth to block out light with samples containing t-BuOOH in order to minimize photodegradation of the hydroperoxy group.

Spectral subtractions and analyses were performed with the OPUS 6.5 (Bruker) and OriginPro 9.1 software. 29,30 The spectra of the complexes were obtained by subtracting reference spectra of the two monomers from a spectrum of the mixture of the two. The pressures of the monomer reference spectra were scaled in the subtraction to obtain at baselines in the areas where only the monomers absorb. This also corrects for possible condensation of the monomers during the recording of the mixture spectra.

The t -BuOH and t -BuOOH dimers were not observed within the applied monomer pressure ranges of 6-20 Torr and 0.2-4 Torr, respectively.

5



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 32

The equilibrium of the complexation can be expressed as:

A+B * ) Complex,

(1)

where A and B are the monomers and Complex is the hydrogen-bonded complex. The dimensionless equilibrium constant of the complexation, Kexpt , can be obtained from:

Kexpt =

PComplex · P , PA · PB

(2)

where PComplex is the pressure of the complex, PA and PB are the monomer pressures, and

P is the standard pressure (1 bar).

The pressure of each monomer, PA and PB was determined experimentally from the measured pressures of the reference monomer spectra and the associated scaling factors. The pressure of the complex formed, PComplex , was too low to be measured experimentally. Additionally, the rate of condensation of the complex is unknown, which further complicates an accurate determination of the complex pressure. Instead, the pressure of the complex was obtained from a combination of experimentally determined integrated absorbance and a calculated oscillator strength for the OH-stretching band of the complexes, using the expression:

PComplex = 2.6935 · 10−9 where T is the absolute temperature,

R

R T · A(˜ ν ) d˜ ν K−1 · Torr · m · cm , fcalc · l

(3)

A(˜ ν ) d˜ ν is the integrated absorbance of the OH-

stretching band, fcalc is the calculated dimensionless oscillator strength of the OH-stretch in the complex and l is the path length of the gas cell. Details of the specic experimental conditions for each of the measurements are given in section S1.3 and integration intervals are given in Table S3.

6


Page 7 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Theory and Calculations A systematic conformer search in Spartan '14 at the ω B97X-D/6-31+G(d) level showed that only one conformer exists for each of the monomers (see Figure S4). 3134 For t -BuOH and t -BuOOH,

this was conrmed by dihedral angle scans at the ω B97X-D/aug-cc-pVTZ level

in Gaussian 09, Revision D.01. 3537 The conformational space of the complexes was sampled by manually displacing the two monomer units relative to each other and optimizing the resulting structure at the ω B97X-D/6-31+G(d) level in Gaussian 09. All unique structures of monomers and complexes were subsequently optimized using ω B97X-D/aug-cc-pVTZ, with "verytight" convergence criteria and the ultrane integration grid, which has previously been found to yield good results for similar systems. 3840 The absence of imaginary frequencies in a subsequent frequency calculation conrmed that the stationary points obtained were minima on the potential energy surface. The binding energy, BE, of the complexes was calculated from the zero-point corrected electronic energies and corrected for basis-set superposition error using the counterpoise correction. 41

Vibrational frequencies of the fundamental OH-stretch in the monomers and complexes were calculated using both an anharmonic 1D local mode (LM) model assuming a Morse oscillator and an anharmonic 2D local mode model including the OH-stretch as well as the COH- or OOH-bending mode for the alcohol and hydroperoxide, respectively. For the complexes, the newly developed local mode perturbation theory (LMPT) model was also employed. These vibrational models are described elsewhere in the literature. 4248 In brief, the LMPT model employed uses an anharmonic 2D local mode model for the hydrogen bond donor and accounts for the eect of selected intermolecular modes by perturbation theory. It has been found to improve the LM results for a number of dierent bimolecular complexes. 48 The potential energy and dipole moment surfaces for all vibrational calculations were calculated at the ω B97X-D/aug-cc-pVTZ level in Gaussian 09.

7



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The theory of natural bond orbitals (NBO) was used to determine the character of the interaction between the two monomer units. 49 The NBO interaction energy was calculated using Gaussian 09 at the ω B97X-D/aug-cc-pVTZ level of theory with the keywords "pop=nbo" and "scf=verytight". The atoms in molecules (AIM) and non-covalent interactions (NCI) theories were used to visualize and characterize the interactions between the monomers in the complexes. 50,51 These were calculated using AIM2000 and NCIPLOT Version 3.0, respectively. 52,53 For NCI, the range of used for color coding the isosurfaces was

−0.015 a.u. < sign(λ2 ) · ρ(r) < 0.015 a.u. at s(r) = 0.5 a.u. Atomic polar tensor (APT) charges were calculated in Gaussian 09 to determine partial atomic charges.

8


Page 8 of 32

Page 9 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Results and Discussion While only one conformer is found for each of the monomers, calculations suggest the presence of several conformers for both complexes. For each complex, two hydrogen-bonded conformers and two non-hydrogen-bonded conformers are located (see Figure S5). In all cases, the non-hydrogen-bonded complexes are at least 12 kJ/mol higher in ω B97X-D/augcc-pVTZ zero-point corrected energy than the hydrogen-bonded conformers (see Table S4) and these will not be considered further. For each complex, the two hydrogen-bonded conformers are within 0.6 kJ/mol in zero-point corrected ω B97X-D/aug-cc-pVTZ energy of each other. Thus, at room temperature, both of these conformers will likely be present in similar amounts. As it has little impact on the conclusions drawn, we have focused on the lowest energy conformer (shown in Figure 1). For determination of the equilibrium constant, we have used an average oscillator strength rather than attempting to calculate a Boltzmann weight between the two conformers, as calculation of Gibbs free energies of hydrogen-bonded complexes has been found to be unreliable. Studies have found variations up to more than 10 kJ/mol between dierent computational approaches and equally large deviations from experimental values for some methods. 54,55 A signicant part of the challenge is the high sensitivity of the calculated entropy towards the low frequency intermolecular vibrational modes in these weakly bonded complexes.

Geometrical parameters of the two hydrogen-bonded conformers of each complex are given in Tables S5 and S6. The tables also include values obtained using M06-2X/aug-cc-pVTZ and MP2/aug-cc-pVTZ level yielding results comparable to the ω B97X-D/aug-cc-pVTZ results presented.

9



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

t -BuOH·DME

Page 10 of 32

t -BuOOH·DME

Figure 1: Structures of the lowest zero-point corrected energy conformer for each of the

two complexes at the ω B97X-D/aug-cc-pVTZ level of theory.

Spectra of the Complexes Upon spectral subtraction, a single, well-dened band is observed in the spectra of both t -BuOH·DME

and t -BuOOH·DME (see Figures 2 and S6) and these are assigned to the

fundamental OH-stretching bands ( νÕH ) of the complexes. The bands are observed with a maximum at νÕH = 3544 cm-1 and νÕH = 3428 cm-1 for t-BuOH·DME and t-BuOOH·DME, respectively. These values are in reasonable agreement with the wavenumbers calculated using the LMPT model, which predicts the bands to be found at 3568 cm -1 and 3394 cm-1 , respectively. For both complexes, the LMPT model predicts the two hydrogen-bonded conformers to be within 10 cm -1 of each other, which is in line with only one peak being observed in the spectra of the complexes. The 1D and 2D local mode approaches yield wavenumbers lower than the LMPT model by 20-50 cm -1 (see Table S9). This is in line with overestimation of the redshift in the local mode model found previously. 4648

In the spectrum of the t -BuOOH·DME mixture and complex (Figure 2, red and blue), an additional small sharp feature is observed at 3642 cm -1 . This correlates with the wavenumber of the t -BuOH monomer and is believed to arise due to photodegradation of the hydroperoxide group forming a small amount of t -BuOH. The shape and location of this feature matches that found for t -BuOH. As mentioned previously, the mixing bowl was covered to minimize 10


Page 11 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


this degradation. The small intensity of this band, suggests that only negligible amounts are present and we have not subtracted this band.

Figure 2: FT-IR spectra of a mixture of 6.8 Torr t-BuOOH and 95 Torr DME (red), 94.5 Torr DME (orange), 6.8 Torr t-BuOOH (green) and the spectrum of the complex (blue) obtained by spectral subtraction, all at T=295 K and with a path length of 2.4 m. The spectra have been oset for clarity.

Characteristics of Hydrogen Bonding For both complexes, the fundamental OH-stretching band is redshifted relative to the corresponding band in the monomer, which is a characteristic of hydrogen bonding. 56 The magnitude of the redshift of the OH-stretch in t-BuOH and t-BuOOH upon complexation with DME are visualized in Figure 3, by shifting the wavenumbers such that zero corre11



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

sponds to the wavenumber of the Q-branch of the fundamental OH-stretch of the monomer. The redshifts are calculated as the Q-branch of the OH-stretching transition of the monomer minus the band maximum of the complex. The redshift of 170 cm -1 for the t -BuOOH·DME complex is signicantly larger than the redshift of 98 cm -1 for the t -BuOH·DME complex, suggesting that the hydroperoxide is a signicantly better hydrogen bond donor than the alcohol. 56 This is in agreement with the observations in the liquid phase, which also reports a larger redshift for a hydroperoxide compared to the corresponding alcohol. 22 Furthermore, it is in line with the calculations suggesting that the elongation of the OH-bond upon hydrogen bond formation is larger in the hydroperoxide than the alcohol (see Table S11).

Figure 3: Spectra of the t -BuOH·DME and t -BuOOH·DME complexes in the fundamental OH-stretching region. The spectra have been shifted so zero corresponds to the Q-branch of the OH-stretching band for each of the monomers, t -BuOH and t -BuOOH, respectively.

Besides the redshift, the fundamental OH-stretches also display spectral broadening upon complex formation, which is another characteristics of hydrogen bond formation (full width at half maximum of the complexes are given in Table S12). 56 The broadening can likely be ascribed primarily to intramolecular vibrational energy redistribution (IVR) and predissociation eects. 5759

12


Page 12 of 32

Page 13 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Equilibrium Constants The equilibrium constant, which can easily be converted into ∆G , of complex formation has been determined to investigate the total interaction strength between the two monomers. As shown in the left panel of Figure 4, the intensity of the fundamental OH-stretching band in the t -BuOOH·DME complex increases with increasing monomer pressures. The same trend is observed for t -BuOH·DME, see Figure S7. We use the integrated absorbance of these bands with the calculated oscillator strengths of the band (see Table S10) and Equation 3 (taking into account the optical path length in the specic experiment) to obtain the pressure of the complex. As the two hydrogen-bonded conformers of each complex are calculated to be close in energy and have very similar calculated oscillator strengths (see Tables S4 and S10), we have used an average oscillator strength to determine the pressures of each of the complexes.

The fundamental OH-stretching transition is the only transition we can unambiguously assign to the complex and therefore provides the only band we can use for determining the equilibrium constant. While we cannot rule out coupling to overtones or combination bands, the smooth band shape does not suggest any important contributions from such eects. Furthermore, for the complex of methanol-dimetylamine, equilibrium constants obtained based on the fundamental OH-stretching transition and the second NH-stretching overtone transition yield similar results suggesting the fundamental OH-stretching transition is a suitable choice. 38,48 If, however, other transitions provide intensity to the observed band, the determined equilibrium constant would be overestimated.

The right panel of Figure 4 shows that for both the t -BuOH·DME and t -BuOOH·DME complexes, the pressure of the complex correlates linearly with the product of the pressures of the two monomers. The linearity conrms that a 1:1 complex is formed between the two monomers and the slope of this plot is equal to the equilibrium constant for complex 13



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

formation. The linear least square ts are forced through (0,0).

Figure 4: Left: Spectra of the t-BuOOH·DME complex in the fundamental OH-stretching

region with dierent combinations of monomer pressures with an optical path length of 2.4 m. Right: Pressure of the complex as a function of the product of monomer pressures. PComplex is calculated using the average calculated LMPT oscillator strengths of 7.57 × 10−5 and 1.02 × 10−4 for t-BuOH·DME and t-BuOOH·DME, respectively, and corrected for the dierent path lengths in the dierent experiments. All spectra are recorded at 295 K - 298 K. It is clear, that the slope for the t -BuOOH·DME complex is signicantly steeper than the slope for the t -BuOH·DME complex illustrating that the hydroperoxide forms a stronger complex with DME than the alcohol does. This is in line with the observed redshifts. Using the average LMPT calculated oscillator strengths ( 7.57 × 10−5 and 1.02 × 10−4 for t-BuOH·DME

and t-BuOOH·DME, respectively), we obtain equilibrium constants of 0.031

and 0.17 for the formation of t -BuOH·DME and t -BuOOH·DME, respectively, corresponding to ∆G values of 8.6 kJ/mol and 4.4 kJ/mol at room temperature. Due to the very similar oscillator strengths for the two conformers of each complex, using an average oscillator strength introduces a maximum uncertainty of less than 10 %. A Boltzmann weighted average of the two conformers would deviate even less from the average value, but would require calculation of Gibbs free energies which, as mentioned, are very dicult to calculate for these systems. The presence of small amounts of decane in the gas phase for the hydroperoxide 14


Page 14 of 32

Page 15 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


solution means that we will slightly underestimate the equilibrium constant. But as no exact determination of the contribution from each constituent to the total pressure can be made, it is not corrected for. However, this error is expected to be less than 10 % (see section S1.1). The uncertainty (95 % condence interval) derived from the slope of the linear t is 3 % for t -BuOH-DME and 7 % for t -BuOOH-DME. Despite recent eorts, specically the development of the LMPT model, the largest potential error in the determination is most likely still the calculated oscillator strength. 4648

The approach used here which combines experimentally determined integrated absorbances and calculated oscillator strengths to determine equilibrium constants has yielded good results for other similar systems. 60,61 The stronger total interaction in the hydroperoxide complex, shown by the larger equilibrium constant, is also reected in the calculated binding energies, which suggest that the hydroperoxide complex is stabilized by about 4.5 kJ/mol more than the alcohol complex (see Table S13).

Interactions NBO calculations (Table S14) show that for both the alcohol and hydroperoxide complexes, the hydrogen bond represents by far the largest interaction between the two monomer units and accounts for around 85 % of the total interaction. The strong hydrogen bond interaction is in line with the large redshifts observed experimentally. The larger observed redshift for the hydroperoxide complex correlates with the AIM calculations, which show an substantially increased charge density in the bond critical point of the hydrogen bond in t -BuOOH·DME compared to t -BuOH·DME (see Table S15 and Figure S8). 62,63

In Figure 5, the NCI isosurfaces of the two complexes show clearly both the strong hydrogen bond (blue) and weaker secondary interactions (green). The gure suggests that the

15



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 16 of 32

increased exibility of the hydroperoxide may allow for conformers with larger areas of secondary interactions between the monomers compared to the alcohol. For the hydroperoxide, the NCI calculation also reveals internal interactions between the oxygen atom and the two methyl groups close to it. These are both attractive (blue) and repulsive (red). As these are present in both the monomer (see Figure S10) and the complex, they are not expected to aect the strength of the binding between the monomers.

t-BuOH·DME (A)

t-BuOOH·DME (A)

Figure 5: The NCI isosurfaces of the ω B97X-D/aug-cc-pVTZ optimized

t-BuOH·DME

and t-BuOOH·DME complexes. The range of used for color coding the isosurfaces was −0.015 a.u. < sign(λ2 ) · ρ(r) < 0.015 a.u. at s(r) = 0.5 a.u. For the NCI isosurfaces, red corresponds to repulsive interactions, green to weak attractive interaction and blue to stronger attractive interactions.

Explanation for the Observed Hydrogen Bond Strength We attribute the observation that the hydroperoxide forms a stronger hydrogen bond than the alcohol to a weaker OH-bond in the hydroperoxide. This is evident from the wavenumber of the fundamental OH-stretching band, νÕH , in the monomers (see Figures 2 and S6). In good agreement with literature, we nd νÕH = 3642 cm-1 and 3598 cm-1 for t -BuOH and t -BuOOH, respectively. 6468 The lower wavenumber for the hydroperoxide indicates a weaker OH-bond. This experimental observation is backed by several calculated values. Our calculated (ω B97X-D/aug-cc-pVTZ) zero-point corrected bond dissociation energy of the OH-bond is 425 kJ/mol in t -BuOH compared to only 327 kJ/mol in t -BuOOH, in line with 16


Page 17 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


experimental values for other alcohols and hydroperoxides. 69 Furthermore, APT charges show a signicantly larger charge dierence between the O and H atoms of the bond in the alcohol, again indicating a stronger bond. Finally, it correlates with the acidity of t -BuOH and t -BuOOH, which have pKa values of 16.5 and 12.8, respectively. 7072 The weaker OHbond in the hydroperoxide means that the hydrogen atom is less stabilized making it more susceptible to forming a hydrogen bond and thereby increasing the hydrogen bond strength.

Comparison to Other Complexes The redshift of 98 cm -1 observed here for the t -BuOH·DME complex is very similar to that observed for other alkyl alcohol complexes with DME: The redshift for methanol-DME (MeOH·DME) is 103 cm-1 and the redshifts reported for ethanol-DME (EtOH ·DME) are 112 cm-1 and 94 cm-1 due to the two EtOH monomer conformers. 73,74 This similarity in the redshifts suggests, not surprisingly, similar hydrogen bond strength. For these complexes, equilibrium constants of complex formation have been obtained using the same combination of theory and experiment as in this study but with oscillator strengths calculated at a slightly dierent level. For all three alcohol-DME complexes, the equilibrium constants are around 0.01-0.03 (see Table S16). On the other hand, the corresponding values for the hydroperoxide complex, t -BuOOH·DME, are a redshift of 170 cm -1 and an equilibrium constant of 0.17. These values are closer to those obtained for the 2,2,2-triuoroethanol-DME complex (TFE-DME), which has a redshift of 198 cm -1 and an equilibrium constant of 0.26. 74 In constrast, a jet expansion experiment reports a signicantly larger redshift of 280 cm -1 for the phenolic OH-O hydrogen bond formed between 4-methylphenol ( p -cresol) and diethylether. 75 A hydrogen bond strength of the hydroperoxide between that of alcohols and phenolic compounds is in line with the results obtained in the liquid phase. 22 The similarity of the redshifts and equilibrium constants for the alkyl alcohols leads us to expect that the result obtained for t -BuOOH is similar to what would be observed with other simple hydroperoxides as hydrogen bond donors. Yet, the instability of most hydroperoxides renders 17



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

an investigation of dierent hydroperoxy hydrogen bond donors dicult.

We observe the superiority of t -BuOOH over t -BuOH as a hydrogen bond donor also with other hydrogen bond acceptors. Both in terms of experimentally observed redshifts of the fundamental OH-stretching and calculated binding energies, we nd the same trend with nitrogen and sulphur atoms as the hydrogen bond acceptors, exemplied by trimethylamine and dimethyl sulphide, respectively (see Tables 1 and S18 and Figures 6-7). This suggests that the trend is not limited to oxygen as the acceptor atom, but more general. In terms of hydrogen bond acceptor strength, we observe that nitrogen is a signicantly better acceptor than oxygen and sulphur in good agreement with results for other hydrogen bond donors in the literature. 40,60

Table 1: Experimentally observed redshifts a (in cm-1 ) of t -BuOH and t -BuOOH with various hydrogen bond acceptors.

t -BuOH

t -BuOOH a

DME

DMS

TMA

98 170

105 196

312b 433

Monomer Q-branch minus complex band maximum.

Monomer frequen-

cies are 3642 cm-1 and 3598 cm-1 for t -BuOH and t -BuOOH, respectively. b

To the most intense peak in the complex spectrum.

18


Page 18 of 32

Page 19 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 6: Spectra showing the magnitude of the redshift of the fundamental OH-stretching

band, ∆˜ νOH , in t-BuOH and t-BuOOH upon complexation with trimethylamine (TMA). The spectra have been shifted so zero corresponds to the Q-branch of the OH-stretching band of the monomer.

Figure 7: Spectra showing the magnitude of the redshift of the fundamental OH-stretching

band, ∆˜ νOH , in t-BuOH and t-BuOOH upon complexation with dimethyl sulphide (DMS). The spectra have been shifted so zero corresponds to the Q-branch of the OH-stretching band of the monomer.

The experimental and theoretical work in this study clearly shows, that t -BuOOH forms a signicantly stronger hydrogen-bonded complex with dimethyl ether (DME) and other hydro19



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

gen bond acceptors than the corresponding alcohol, t -BuOH. However, the hydroperoxides formed in the atmosphere are rarely simple, monofunctional species, as the one studied here. Multifunctional compounds have the ability to form intramolecular hydrogen bonds. 7680 The intramolecular hydrogen bonds compete with the formation of intermolecular interactions and may thereby increase the vapor pressure and aect the ability to partake in e.g. formation of molecular clusters and secondary organic aerosols. 78,79,81 The stabilization of the monomer (via intramolecular hydrogen bonds) with respect to its clusters has for instance been used to explain the limited interaction strength predicted for a diperoxy acid formed in the ozonolysis of cyclohexene with sulphuric acid. 82,83 However, the relative strength of complexation between alcohols and hydroperoxides in multifunctional systems is not expected to be simple and will likely depend on the specic system and its ability to form internal hydrogen bonds.

20


Page 20 of 32

Page 21 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Conclusion We have detected a hydrogen-bonded bimolecular gas phase complex between a hydroperoxide (hydrogen bond donor) and dimethyl ether (acceptor). The tert -butyl hydroperoxidedimethyl ether complex was observed using Fourier transform infrared (FT-IR) spectroscopy and compared to the corresponding alcohol complex; tert -butanol-dimethyl ether. Both in terms of measured redshift (170 cm -1 vs. 98 cm-1 ) and determined equilibrium constant (0.17 vs. 0.031), we nd that the hydroperoxide is a signicantly better hydrogen bond donor than the alcohol. The determined equilibrium constants correspond to the hydroperoxide complex being stabilized by about 4 kJ/mol (Gibbs free energy) more than the alcohol complex. The trend is explained by a weaker OH-bond in the hydroperoxide and correlates with the acidity of t -BuOH and t -BuOOH, which have pKa values of 16.5 and 12.8, respectively. Measured redshifts and calculated binding energies show the same trend in the complexes formed with trimethylamine and dimethyl sulphide as the hydrogen bond acceptors.

Acknowledgement We thank the Danish Center for Scientic Computing and the Center for Exploitation of Solar Energy, University of Copenhagen, Denmark for funding. We acknowledge the participants of the workshop "ELVOC in Dragør", June 2015, for fruitful discussions providing the inspiration for this study, in particular Anke Mutzel, Jonas Elm and Theo Kurtén.

Supporting Information Estimate of the eect of decane in the t -BuOOH solution, pressure gauge correction, experimental conditions, experimental integration intervals, relative conformer energies, conformer structures, geometrical parameters, spectral subtraction of t -BuOH·DME, experimental and calculated wavenumbers and intensities, calculated OH-bond elongation upon complexation, 21



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 32

experimental full width at half maximum of the complexes, spectra of the t-BuOH ·DME complex at dierent combinations of monomer pressures, calculated binding energies, NBO interaction energies, AIM charge densities and plots, NCI isosurfaces, experimental conditions and spectral subtractions for t -BuOH and t -BuOOH with TMA and DMS, binding energies of t -BuOH and t -BuOOH with TMA and DMS.

Gaussian 09 output les of the optimized structures of the hydrogen bonded complexes are available at http://www.erda.dk/public/archives/YXJjaGl2ZS1sTF9pZEw=/published-archive.

html

References (1) Hanst, P. L.; Gay, B. W. Atmospheric Oxidation of Hydrocarbons: Formation of Hydroperoxides and Peroxyacids. Atmos. Environ. (1967) 1983, 17, 2259 2265. (2) Crounse, J. D.; Nielsen, L. B.; Jørgensen, S.; Kjaergaard, H. G.; Wennberg, P. O. Autoxidation of Organic Compounds in the Atmosphere. J. Phys. Chem. Lett. 2013, 4,

35133520.

(3) Bonn, B.; von Kuhlmann, R.; Lawrence, M. G. High Contribution of Biogenic Hydroperoxides to Secondary Organic Aerosol Formation. Geophys. Res. Lett. 2004, 31, L10108. (4) Reinnig, M.-C.; Warnke, J.; Homann, T. Identication of Organic Hydroperoxides and Hydroperoxy Acids in Secondary Organic Aerosol Formed During the Ozonolysis of Dierent Monoterpenes and Sesquiterpenes by On-Line Analysis Using Atmospheric Pressure Chemical Ionization Ion Trap Mass Spectrometry. Rapid Commun. Mass Spectrom.

2009, 23, 17351741.

(5) Docherty, K. S.; Wu, W.; Lim, Y. B.; Ziemann, P. J. Contributions of Organic Peroxides 22


Page 23 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


to Secondary Aerosol Formed from Reactions of Monoterpenes with O 3 . Environ. Sci. Technol.

2005, 39, 40494059.

(6) Intergovernmental Panel on Climate Change, Climate Change 2013 - The Physical Science Basis ;

Cambridge University Press, 2014; Cambridge Books Online.

(7) Devarajan, A.; Windus, T. L.; Gordon, M. S. Implementation of Dynamical Nucleation Theory Eective Fragment Potentials Method for Modeling Aerosol Chemistry. J. Phys. Chem. A

2011, 115, 1398713996.

(8) Zhao, J.; Khalizov, A.; Zhang, R.; McGraw, R. Hydrogen-Bonding Interaction in Molecular Complexes and Clusters of Aerosol Nucleation Precursors. J. Phys. Chem. A 2009, 113,

680689.

(9) Feng, Y.-J.; Huang, T.; Wang, C.; Liu, Y.-R.; Jiang, S.; Miao, S.-K.; Chen, J.; Huang, W. π -Hydrogen Bonding of Aromatics on Surface of Aerosols: Insights from ab Initio and Molecular Dynamics Simulation. J. Phys. Chem. B 2016, 120, 6667 6673. (10) Zhang, R.; Khalizov, A.; Wang, L.; Hu, M.; Xu, W. Nucleation and Growth of Nanoparticles in the Atmosphere. Chem. Rev. 2012, 112, 19572011. (11) Yablonskii, O. P.; Belyaev, V. A.; Vinogradov, A. N. Association of Hydrocarbon Hydroperoxides. Russ. Chem. Rev. 1972, 41, 565. (12) Zharkov, V. V.; Rudnevskii, N. K. Intramolecular Hydrogen Bond in Isopropyl Benzene Hydroperoxide. Opt. Spectrosc. 1959, 7, 497. (13) Richardson, W. H.; Steed, R. F. Structure of α-Oxo Hydroperoxides. J. Org. Chem.

1967, 32, 771774. (14) Zharkov, V.; Rudnevsky, N. Spectroscopic Study of the Hydrogen Bond Between the

23



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Molecules of Tertiary Butyl Hydroperoxide in a Solution of Tetravalent Carbon. Opt. Spectrosc.

1962, 12, 479483.

(15) Walling, C.; Heaton, L. Hydrogen Bonding and Complex Formation in Solutions of t -Butyl

Hydroperoxide. J. Am. Chem. Soc. 1965, 87, 4851.

(16) Kamalova, D.; Kolyadko, I.; Remizov, A.; Semenov, M.; Skochilov, R. Hydrogen Bonds in Self- and Heteroassociates of Cumyl Hydroperoxide in Solutions and Polymeric Matrixes. J. Mol. Struct. 2010, 972, 104 110. (17) Yablonskii, O.; Lastochkina, N.; Belyayev, V. PMR Study of the Hydrogen Bond of Tertiary Hydroperoxides of Hydrocarbons. Petrol. Chem. USSR 1973, 13, 261 266. (18) Remizov, A.; Kamalova, D.; Skochilov, R.; Suvorova, I.; Batyrshin, N.; Kharlampidi, K. FT-IR Study of Self-Association of Some Hydroperoxides. J. of Mol. Struct. 2004, 700, 73 79. (19) Meshcheryakov, A. P.; Batuev, M. I.; Matveeva, A. D. Synthesis of tert-Butyl Hydroperoxide and tert-Butyl Peroxide and their Optical Investigation in Relation to the Problem of the Structure of Hydrogen Peroxide. Bull. Acad. Sci. USSR, Div. Chem. Sci.

1955, 4, 661666.

(20) Nikishin, G. I.; Glukhovtsev, V. G.; Nadtochii, M. A. Thermal Decomposition Reactions of 2-Methyltetrahydrofuran Hydroperoxide and Peroxide. Bull. Acad. Sci. USSR, Div. Chem. Sci.

1973, 22, 22242227.

(21) Shabalin, I.; Klimchuk, M.; Kiva, E. Hydrogen Bonding in Cumene Hydroperoxide Solutions. Opt. Spectrosc. 1968, 24, 294. (22) Shabalin, I.; Kiva, E. The Intermolecular Hydrogen Bond of Cumene Hydroperoxide. Opt. Spectrosc.

1968, 24, 377.

24


Page 24 of 32

Page 25 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(23) Yablonskii, O.; Bystrov, V.; Belyayev, V.; Vinogradov, A. Study by PMR of Complex Formation Between Hydrocarbon Hydroperoxides and Proton-Acceptor Compounds. Petrol. Chem. USSR

1971, 11, 277 282.

(24) Goebel, J.; Ault, B. S.; Del Bene, J. E. Matrix Isolation and ab Initio Study of the Hydrogen-Bonded Complex between H 2 O2 and (CH3 )2 O. J. Phys. Chem. A 2000, 104, 20332037. (25) Engdahl, A.; Nelander, B. The Structure of the Water-Hydrogen Peroxide Complex. A Matrix Isolation Study. Phys. Chem. Chem. Phys. 2000, 2, 39673970. (26) Engdahl, A.; Nelander, B.; Karlström, G. A Matrix Isolation and ab Initio Study of the Hydrogen Peroxide Dimer. J. Phys. Chem. A 2001, 105, 83938398. (27) Du, D.; Fu, A.; Zhou, Z. Theoretical Study of the Rotation Barrier of Hydrogen Peroxide in Hydrogen Bonded Structure of HOOHH 2 O Complexes in Gas and Solution Phase. J. Mol. Struct. THEOCHEM

2005, 717, 127 132.

(28) Du, D.; Zhou, Z. Chiral Discrimination in Hydrogen-Bonded Complexes of Hydrogen Peroxide with Methyl Hydroperoxide: Theoretical Study. Int. J. Quantum Chem. 2006, 106,

935942.

(29) Bruker Corporation, OPUS 6.5. (30) OriginPro 9.1. 2013; OriginLab, Northhampton, MA. (31) Spartan '14. Wavefunction Inc., Irvine, CA. (32) Chai, J.-D.; Head-Gordon, M. Long-Range Corrected Hybrid Density Functionals with Damped Atom-Atom Dispersion Corrections. Phys. Chem. Chem. Phys. 2008, 10, 66156620.

25



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(33) Hehre, W. J.; Ditcheld, R.; Pople, J. A. SelfConsistent Molecular Orbital Methods. XII. Further Extensions of GaussianType Basis Sets for Use in Molecular Orbital Studies of Organic Molecules. J. Chem. Phys. 1972, 56, 22572261. (34) Frisch, M. J.; Pople, J. A.; Binkley, J. S. Self-Consistent Molecular Orbital Methods 25. Supplementary Functions for Gaussian Basis Sets. J. Chem. Phys. 1984, 80, 32653269. (35) 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. (36) Dunning, T. H. Gaussian Basis Sets for Use in Correlated Molecular Calculations. I. The Atoms Boron Through Neon and Hydrogen. J. Chem. Phys. 1989, 90, 10071023. (37) Kendall, R. A.; Dunning, T. H.; Harrison, R. J. Electron Anities of the First-Row Atoms Revisited. Systematic Basis Sets and Wave Functions. J. Chem. Phys. 1992, 96,

67966806.

(38) Du, L.; Mackeprang, K.; Kjaergaard, H. G. Fundamental and Overtone Vibrational Spectroscopy, Enthalpy of Hydrogen Bond Formation and Equilibrium Constant Determination of the Methanol-Dimethylamine complex. Phys. Chem. Chem. Phys. 2013, 15,

1019410206.

(39) Du, L.; Lane, J. R.; Kjaergaard, H. G. Identication of the DimethylamineTrimethylamine Complex in the Gas Phase. J. Chem. Phys. 2012, 136, 184305. (40) Møller, K. H.; Hansen, A. S.; Kjaergaard, H. G. Gas Phase Detection of the NH-P Hydrogen Bond and Importance of Secondary Interactions. J. Phys. Chem. A 2015, 119,

1098810998.

(41) Boys, S.; Bernardi, F. The Calculation of Small Molecular Interactions by the Dier-

26


Page 26 of 32

Page 27 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


ences of Separate Total Energies. Some Procedures with Reduced Errors. Mol. Phys.

1970, 19, 553566. (42) Henry, B. R.; Kjaergaard, H. G. Local Modes. Can. J. Chem. 2002, 80, 16351642. (43) Henry, B. R. Use of Local Modes in the Description of Highly Vibrationally Excited Molecules. Acc. Chem. Res. 1977, 10, 207213. (44) Atkins, P. W. Molecular Quantum Mechanics , 4th ed.; Oxford University Press, 2008; pp 358359,538540. (45) Andersen, C. L.; Jensen, C. S.; Mackeprang, K.; Du, L.; Jørgensen, S.; Kjaergaard, H. G. Similar Strength of the NH ···O and NH···S Hydrogen Bonds in Binary Complexes. J. Phys. Chem. A 2014, 118, 1107411082. (46) Mackeprang, K.; Kjaergaard, H. G.; Salmi, T.; Hänninen, V.; Halonen, L. The Eect of Large Amplitude Motions on the Transition Frequency Redshift in Hydrogen Bonded Complexes: A Physical Picture. J. Chem. Phys. 2014, 140, 184309. (47) Mackeprang, K.; Hänninen, V.; Halonen, L.; Kjaergaard, H. G. The Eect of Large Amplitude Motions on the Vibrational Intensities in Hydrogen Bonded Complexes. J. Chem. Phys.

2015, 142, 094304.

(48) Mackeprang, K.; Kjaergaard, H. G. Vibrational Transitions in Hydrogen Bonded Bimolecular Complexes A Local Mode Perturbation Theory Approach to Transition Frequencies and Intensities. J. Mol. Spectrosc. 2017, 334, 1 9. (49) Foster, J. P.; Weinhold, F. Natural Hybrid Orbitals. J. Am. Chem. Soc. 1980, 102, 72117218. (50) Johnson, E. R.; Keinan, S.; Mori-Sánchez, P.; Contreras-García, J.; Cohen, A. J.; Yang, W. Revealing Noncovalent Interactions. J. Am. Chem. Soc. 2010, 132, 6498 6506. 27



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(51) Bader, R. F. W. Atoms in Molecules, A Quantum Theory ; Oxford University Press: New York, New York, 1990; pp 1424. (52) AIM2000. J. Comput. Chem. 2001, 22, 545559. (53) Contreras-García, J.; Johnson, E. R.; Keinan, S.; Chaudret, R.; Piquemal, J.-P.; Beratan, D. N.; Yang, W. NCIPLOT: A Program for Plotting Noncovalent Interaction Regions. J. Chem. Theory Comput. 2011, 7, 625632. (54) Elm, J.; Bilde, M.; Mikkelsen, K. V. Assessment of Density Functional Theory in Predicting Structures and Free Energies of Reaction of Atmospheric Prenucleation Clusters. Journal of Chemical Theory and Computation

2012, 8, 20712077.

(55) Bork, N.; Du, L.; Kjaergaard, H. G. Identication and Characterization of the HClDMS Gas Phase Molecular Complex via Infrared Spectroscopy and Electronic Structure Calculations. J. Phys. Chem. A 2014, 118, 13841389. (56) Arunan, E.; Desiraju, G. R.; Klein, R. A.; Sadlej, J.; Scheiner, S.; Alkorta, I.; Clary, D. C.; Crabtree, R. H.; Dannenberg, J. J.; Hobza, P. et al. Dening the Hydrogen Bond: An Account (IUPAC Technical Report). Pure Appl. Chem. 2011, 83, 1619 1636. (57) Huang, Z. S.; Miller, R. E. High-Resolution Near-Infrared Spectroscopy of Water Dimer. J. Chem. Phys.

1989, 91, 66136631.

(58) Nesbitt, D. J.; Field, R. W. Vibrational Energy Flow in Highly Excited Molecules: Role of Intramolecular Vibrational Redistribution. The Journal of Physical Chemistry 1996, 100,

1273512756.

(59) Pritchard, M.; Parr, J.; Li, G.; Reisler, H.; McCaery, A. J. The mechanism of H-bond rupture: the vibrational pre-dissociation of C2H2-HCl and C2H2-DCl. Phys. Chem. Chem. Phys.

2007, 9, 62416252. 28


Page 28 of 32

Page 29 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(60) Hansen, A. S.; Du, L.; Kjaergaard, H. G. Positively Charged Phosphorus as a Hydrogen Bond Acceptor. J. Phys. Chem. Lett. 2014, 5, 42254231. (61) Du, L.; Kjaergaard, H. G. Fourier Transform Infrared Spectroscopy and Theoretical Study of Dimethylamine Dimer in the Gas Phase. J. Phys. Chem. A 2011, 115, 12097 12104. (62) Espinosa, E.; Molins, E.; Lecomte, C. Hydrogen Bond Strengths Revealed by Topological Analyses of Experimentally Observed Electron Densities. Chem. Phys. Lett. 1998, 285,

170 173.

(63) Saleh, G.; Gatti, C.; Presti, L. L. Energetics of Non-Covalent Interactions from Electron and Energy Density Distributions. Comp. Theor. Chem. 2015, 1053, 53 59. (64) Kollipost, F.; Papendorf, K.; Lee, Y.-F.; Lee, Y.-P.; Suhm, M. A. Alcohol Dimers - How Much Diagonal OH Anharmonicity? Phys. Chem. Chem. Phys. 2014, 16, 1594815956. (65) Shreve, O. D.; Heether, M. R.; Knight, H. B.; Swern, D. Infrared Absorption Spectra of Some Hydroperoxides, Peroxides, and Related Compounds. Anal. Chem. 1951, 23, 282285. (66) Bernal, H. G.; Caero, L. C.; Finocchio, E.; Busca, G. An FT-IR Study of the Adsorption and Reactivity of tert-Butyl Hydroperoxide over Oxide Catalysts. Appl. Catal., A 2009, 369,

27 35.

(67) Fleming, P. R.; Rizzo, T. R. Infrared Spectrum of t -Butyl Hydroperoxide Excited to the 4vOH Vibrational Overtone Level. J. Chem. Phys. 1991, 95, 14611465. (68) Lange, K. R.; Wells, N. P.; Plegge, K. S.; Phillips, J. A. Integrated Intensities of O-H Stretching Bands: Fundamentals and Overtones in Vapor-Phase Alcohols and Acids. J. Phys. Chem. A

2001, 105, 34813486.

29



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(69) Blanksby, S. J.; Ellison, G. B. Bond Dissociation Energies of Organic Molecules. Acc. Chem. Res.

2003, 36, 255263.

(70) Shokri, A.; Abedin, A.; Fattahi, A.; Kass, S. R. Eect of Hydrogen Bonds on pKa Values: Importance of Networking. J. Am. Chem. Soc. 2012, 134, 1064610650. (71) Reeve, W.; Erikson, C. M.; Aluotto, P. F. A New Method for the Determination of the Relative Acidities of Alcohols in Alcoholic Solutions. The Nucleophilicities and Competitive Reactivities of Alkoxides and Phenoxides. Can. J. Chem. 1979, 57, 2747 2754. (72) Serjeant, E. P.; Dempsey, B.; IUPAC, Ionisation Constants of Organic Acids in Aqueous Solution ;

IUPAC Chemical Data Series 23; Pergamon Press, 1979; p 80.

(73) Howard, D. L.; Kjaergaard, H. G. Hydrogen Bonding to Divalent Sulfur. Phys. Chem. Chem. Phys.

2008, 10, 41134118.

(74) Du, L.; Tang, S.; Hansen, A. S.; Frandsen, B. N.; Maroun, Z.; Kjaergaard, H. G. Subtle Dierences in the Hydrogen Bonding of Alcohol to Divalent Oxygen and Sulfur. Chem. Phys. Lett.

2017, 667, 146 153.

(75) Biswal, H. S.; Wategaonkar, S. O-H · · · O versus O-H· · · S Hydrogen Bonding. 3. IRUV Double Resonance Study of Hydrogen Bonded Complexes of p -Cresol with Diethyl Ether and Its Sulfur Analog. J. Phys. Chem. A 2010, 114, 59475957. (76) Krueger, P.; Mettee, H. Spectroscopic Studies of Alcohols: Part VII. Intramolecular Hydrogen Bonds in Ethylene Glycol and 2-Methoxyethanol. J. Mol. Spectrosc. 1965, 18,

131 140.

(77) Buckley, P.; Brochu, M. Microwave Spectrum, Dipole Moment, and Intramolecular Hydrogen Bond of 2-Methoxyethanol. Can. J. Chem. 1972, 50, 11491156.

30


Page 30 of 32

Page 31 of 32

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(78) Graham, R. J.; Kroemer, R. T.; Mons, M.; Robertson, E. G.; Snoek, L. C.; Simons, J. P. Infrared Ion Dip Spectroscopy of a Noradrenaline Analogue: Hydrogen Bonding in 2Amino-1-phenylethanol and Its Singly Hydrated Complex. J. Phys. Chem. A 1999, 103,

97069711.

(79) Borho, N.; Suhm, M. A.; Le Barbu-Debus, K.; Zehnacker, A. Intra- vs. Intermolecular Hydrogen Bonding: Dimers of Alpha-Hydroxyesters with Methanol. Phys. Chem. Chem. Phys.

2006, 8, 44494460.

(80) Schrøder, S. D.; Wallberg, J. H.; Kroll, J. A.; Maroun, Z.; Vaida, V.; Kjaergaard, H. G. Intramolecular Hydrogen Bonding in Methyl Lactate. J. Phys. Chem. A 2015, 119, 96929702. (81) Kurtén, T.; Tiusanen, K.; Roldin, P.; Rissanen, M.; Luy, J.-N.; Boy, M.; Ehn, M.; Donahue, N. α-Pinene Autoxidation Products May Not Have Extremely Low Saturation Vapor Pressures Despite High O:C Ratios. J. Phys. Chem. A 2016, 120, 25692582. (82) Elm, J.; Myllys, N.; Hyttinen, N.; Kurtén, T. Computational Study of the Clustering of a Cyclohexene Autoxidation Product C 6 H8 O7 with Itself and Sulfuric Acid. J. Phys. Chem. A

2015, 119, 84148421.

(83) Elm, J.; Myllys, N.; Luy, J.-N.; Kurtén, T.; Vehkamäki, H. The Eect of Water and Bases on the Clustering of a Cyclohexene Autoxidation Product C 6 H8 O7 with Sulfuric Acid. J. Phys. Chem. A 2016, 120, 22402249.

31



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Table of Content Figure

TOC Graphic

32


Page 32 of 32

Side-by-Side Comparison of Hydroperoxide and Corresponding

Recommend Documents