Atmospheric Fate of Nitramines: An Experimental and Theoretical

Article pubs.acs.org/JPCA

Atmospheric Fate of Nitramines: An Experimental and Theoretical Study of the OH Reactions with CH3NHNO2 and (CH3)2NNO2 Mihayo Musabila Maguta,† Marius Aursnes,‡ Arne Joakim Coldevin Bunkan,† Katie Edelen,† Tomás ̌ Mikoviny,† Claus Jørgen Nielsen,*,† Yngve Stenstrøm,*,‡ Yizhen Tang,§ and Armin Wisthaler*,† †

Centre for Theoretical and Computational Chemistry, Department of Chemistry, University of Oslo, P. O. Box 1033, Blindern, 0315 Oslo, Norway ‡ Norwegian University of Life Sciences, IKBM, P.O. Box 5003, NO-1432 Aas, Norway § School of Environmental and Municipal Engineering, Qingdao Technological University, Fushun Road 11, 266033 Qingdao, Shandong P.R. China S Supporting Information *

ABSTRACT: The rates of CH3NHNO2 and (CH3)2NNO2 reaction with OH radicals were determined relative to CH3OCH3 and CH3OH at 298 ± 2 K and 1013 ± 10 hPa in purified air by long path FTIR spectroscopy, and the rate coefficients were determined to be kOH+CH3NHNO2 = (9.5 ± 1.9) × 10−13 and kOH+(CH3)2NNO2 = (3.5 ± 0.7) × 10−12 (2σ) cm3 molecule−1 s−1. Ozone was found to react very slowly with the two nitramines, kO3+nitramine < 10−21 cm3 molecule−1 s−1. Product formation in the photo-oxidation of CH3NHNO2 and (CH3)2NNO2 was studied by FTIR, PTR-ToF-MS, and quantum chemistry calculations; the major products in the OH-initiated degradation are the corresponding imines, CH2 NH and CH3NCH2, and N-nitro amides, CHONHNO2 and CHON(CH3)NO2. Atmospheric degradation mechanisms are presented.

1. INTRODUCTION Nitramines are formed in the atmospheric photo-oxidation of amines1−8 that, in turn, are emitted by a wide range of sources.9 Planned large scale implementation of amine-based technology for carbon capture may reduce CO2 emission from new and existing fossil fuel point sources, but the operation will likely result in small but still important discharges of solvent amines and other process degradation products, including simple aliphatic amines, to the atmosphere.10 Such installations may therefore present significant perturbations to the natural amine budget. For methylamine the reaction sequence leading to Nnitromethylamine is5 ̇ → CH ̇ 2NH 2 + H 2O CH3NH 2 + OH ̇ + H 2O → CH3NH ̇ + NO2 → CH3NHNO2 CH3NH → CH 2NH + HONO

There are only a few studies on health effects of aliphatic nitramines; they are suspected to be both mutagens and carcinogens although they seem to be less potent than the nitrosamines.11,12 Due to scarcity of toxicity data the Norwegian Institute of Public Health has suggested that the risk estimate for N-nitrosodimethylamine should be used also for exposure to nitramines and recommends that the total amount of nitrosamines and nitramines should not exceed 0.3 ng m−3 in air (corresponding to ca. 0.1 ppt dimethylnitrosamine) and 4 ng L−1 in drinking water to ensure minimal or negligible risk of cancer for the public from exposure to these substances.12 A recent health risk analysis for emissions to air from CO2 Technology Centre Mongstad (TCM) details how these conservative guideline values have been used to threshold acceptable emissions of amines and amine degradation products from the CO2-capture facility.13 There is only one previous study on the atmospheric chemistry of nitramines; the rate coefficient for OH radical reaction with N-nitrodimethylamine was reported by Tuazon and co-workers,14 who carried out relative rate studies employing long-path FTIR detection and dimethyl ether as reference compound. They reported kOH+(CH3)2NNO2 = (3.5 ± 0.7) × 10−12 cm3 molecule−1 s−1 at 296 K from experiments

(1a) (1b) (2a) (2b)

Although O2 reaction with amino radicals is 6 orders of magnitude slower than the corresponding NO and NO2 reactions,2 it is still dominating at atmospheric conditions, and N-nitromethylamine and N-nitrodimethylamine are only minor products in the photo-oxidation of respectively methylamine5 and dimethylamine.2,5 © 2014 American Chemical Society

Received: January 10, 2014 Revised: April 23, 2014 Published: April 25, 2014 3450

dx.doi.org/10.1021/jp500305w | J. Phys. Chem. A 2014, 118, 3450−3462

The Journal of Physical Chemistry A

Article

10−10 cm3 molecule−1 s−1,21 respectively (there are no rate coefficients available for the O(1D) reactions with CH3OCH3 and the nitramines). However, the mixing ratio of H2 is 3 orders of magnitude larger than those of the organics and the O(1D) reaction with these can safely be ignored. The OH radical concentration that can be achieved by this production method easily reaches 109 cm−3; in the present experiments the average OH radical concentration, including the waiting time between periods of photolysis and spectral recording, was around 108 cm−3. The kinetic study was carried out by the relative rate method in a static gas mixture, in which the removals of the reacting species are measured simultaneously as a function of reaction time. Assuming that the reactants react solely with the same radical species and that none of the reactants are re-formed in any side reactions, the relative rate coefficient, krel, is given according to the following expression:

using either CH3ONO or N2H4/O3 as OH source. This single rate coefficient was later employed as fix point by Atkinson in his general structure−activity-relationship for OH reactions with nitramines.15

2. EXPERIMENTAL SECTION 2.1. Relative Rate Experiments and Reference Spectra. The experiments were carried out in a 240 L electropolished stainless steel smog chamber (herein referred to as the Oslo chamber) equipped with a White type multiple reflection mirror system with a 120 m optical path length for rovibrationally resolved infrared spectroscopy. IR spectra were recorded with a Bruker IFS 66v FTIR instrument employing an LN2-cooled MCT detector. The reaction chamber was equipped with UV photolysis lamps mounted in a quartz tube inside the chamber, and all experiments were carried out in synthetic air (AGA 99.99%; CH4, CO and NOx < 100 ppbV) at 298 ± 2 K and 1013 ± 10 hPa. Initial partial pressures of the nitramines and reference compounds (methanol or dimethyl ether) were in the range 0.1−1.5 Pa, and the cell was filled to 1013 hPa with synthetic air. As a standard, 128 scans were coadded at a nominal resolution of 0.5 cm−1 and Fourier transformed using boxcar apodization. Hydroxyl radicals were generated by photolysis of O3 in the presence of H2 (99%, AGA). Ozone was produced from oxygen (99.995%, AGA) using a Model OZO1VVT generator (OZOMAX Inc.), having a conversion efficiency of approximately 5%, and collected in a trap filled with silica beads at 195 K. Typical partial pressures of ozone and hydrogen were 50 and 200 Pa, respectively. Photolysis of ozone was carried out at intervals of 2 min using a Philips TL 12 lamp (wavelength region 280−380 nm, λmax ∼ 315 nm): O3 + hν → O(1D) + O2

(3)

O(1D) + H 2 → OH + H

(4)

H + O3 → OH + O2

(5)

H + O2 + M → HO2 + M

(6)

HO2 + O3 → OH + 2O2

(7)

HO2 + HO2 → H 2O2 + O2

(8)

H 2O2 + hν → 2OH

(9)

OH + H 2 → H + H 2O

(10)

OH + O3 → HO2 + O2

(11)

OH + H 2O2 → HO2 + H 2O

(12)

OH + HO2 → H 2O + O2

(13)

⎧ [S] ⎫ ⎧ [R] ⎫ ln⎨ 0 ⎬ = k rel·ln⎨ 0 ⎬ ⎩ [S]t ⎭ ⎩ [R]t ⎭

k rel =

kS kR

(I)

where [S]0, [R]0, [S]t, and [R]t are concentrations of the substrate (nitramine) and the reference compound at start and at the time t, respectively, and kS and kR are the corresponding rate coefficients. A plot of ln{[S]0/[S]t} vs ln{[R]0/[R]t} will thus give the relative reaction rate coefficient krel = kS/kR as the slope. Control experiments were performed to check for loss of nitramines, dimethyl ether, and methanol via photolysis, dark chemistry, and heterogeneous reactions in the reactor. The lifetime of the nitramines in the reaction chamber was investigated with purified air as diluent and with the relevant radical precursor mixtures in purified air in experiments lasting around 2 h. The photostability of the nitramines (absorbance maximum around 235 nm22) toward the radiation used in generating the radicals (see above) was studied in a separate experiment with purified air as diluent: no direct photolysis in the reactor was detected. The experimental FTIR spectra were analyzed using a global nonlinear least-squares spectral fitting procedure.23 In this method, the spectrum of the mixture of absorbing species is first simulated by calculation from initial estimates of the absorber concentrations. The calculation is then iterated to minimize the residual between the measured and simulated spectrum to adjust the absorber concentrations, continuum level, and instrument line shape parameters. When possible, absorption coefficients are calculated from the HITRAN database (O3, H2O, and CO2);24 if HITRAN line parameter data are not available, high-resolution FTIR spectra are used to approximate the absorption coefficients. In the present study, we recorded reference spectra of CH3OH, CH3OCH3, CH3OCHO, CH3NHNO2, and (CH3)2NNO2; the latter two are shown in Figures S1 and S2, respectively, whereas Table S1 (Supporting Information) sums up the wavenumber regions employed and the chemical components included in the spectral analyses. The data from the independent experiments were analyzed jointly using a weighted least-squares procedure that includes uncertainties in both reactant concentrations.25 The estimated uncertainty in the concentration determination by FTIR was taken as either 1% of the initial concentration or 3σ from the least-squares spectral analysis, whichever is the largest.

This OH production method generates not only OH radicals in the ground state but also in excited vibrational states; reaction 4 results in vibrational levels with v ≤ 4 and reaction 5 in levels with v ≤ 9.16−18 The rate coefficients for vibrational relaxation of OH by N2 and O2 are on the order of 10−15 and 10−13 cm3 molecule−1 s−1, respectively,19 and as the mixing ratios of O2 and N2 are 4−5 orders of magnitude larger than those of the organic reactants, one may safely assume that these react exclusively with OH in the vibrational ground state. The rate coefficients for O(1D) reaction with CH3OH and H2 at 298 K are of comparable magnitude, 5.1 × 10−10 20 and 1.1 × 3451



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combined organic phase was dried (MgSO4), filtered, and evaporated in vacuo to yield dimethylnitramine as a white solid. Recrystallization (diethyl ether) gave 8.75 g (71%) of dimethylnitramine as large, white crystals. Mp: 55−56 °C (lit.28 54.0−56.0 °C). 1H NMR (300 MHz, CDCl3): δ 3.04 (s, 3H), 3.76 (s, 3H). 13C NMR (75 MHz, CDCl3): δ 40.48 ppm. 2.4. Electronic Structure Calculations. Reaction enthalpies were calculated from the G329 and G430 model chemistries, which are reported to reproduce the G3/05 test31 set with average absolute deviations of around 4.7 and 3.5 kJ mol−1, respectively. The G3 method is based on MP2/631G(d) structures and HF/6-31G(d) vibrational frequencies, whereas G4 model chemistry is based on B3LYP/6-31G(2df,p) structures and vibrational frequencies. G3 and G4 results are collected in Table S2 (Supporting Information). Theoretical IR spectra were obtained in B3LYP32−35 calculations employing the aug-cc-pVTZ basis set.36 Potential energy surfaces (PES) of reactions were explored in B3LYP and BHandHLYP calculations with the aug-cc-pVDZ basis set,36 and intrinsic reaction coordinate paths were followed to verify that the saddle points connect to the correct reactants and products. To obtain more reliable energies for the stationary points of the PES, single point energies were obtained in CCSD(T)/6-311+ +G(2d,2p)//BHandHLYP/aug-cc-pVDZ, G3 and G4 computations. Gaussian 0937 was used in all calculations. 2.5. Master Equation Modeling. The competition between collisional stabilization and unimolecular reaction of the chemically activated intermediates was studied using an energy grained master equation model as implemented in MESMER.38 Rate coefficients for reactions with a tight transition state were calculated using RRKM theory based on energies and ro-vibrational data from the G4 calculations, whereas the rate coefficients for the loose transition states were calculated assuming temperature independent rate coefficients of 1 × 10−11 cm3 molecule−1 s−1 for both reactions with both NO and O2. The applicability of this assumption was tested by varying the rate coefficients within a factor of 10 and the assumed NO concentration was varied between 1 and 100 ppm with any observed change in calculated yields. The fraction of vibrationally hot nitramino radicals that dissociate before being stabilized or before reacting with oxygen was estimated using a master equation as outlined above, but with the initial population set to have an excess of energy from the reaction of the parent nitramine with OH radicals. In addition to the limiting cases with all and none of the reaction energy deposited in the nitramino radical, the cases corresponding to one and two energy quanta deposited in the O−H vibrational modes of water were examined.

2.2. PTR-ToF-MS Product Study. Product studies were carried out in a 480 L Teflon coated reaction chamber (herein referred to as the Innsbruck chamber) made of glass and surrounded by 18 UV/vis lamps (λ ≥ 300 nm) to simulate sunlight. The experiments were performed in humidified synthetic air at 10−20% RH, an absolute pressure of 970 mbar, and absolute temperatures in the range between 298 K (start of experiment) and 308 K (end of experiment). A typical experiment lasted about 60 min. Starting conditions were 100− 150 ppb of the reagent nitramine and approximately 100 ppb of NO and NO2, respectively. The reagent nitramine was introduced by flushing the headspace of the pure sample into the chamber with synthetic air. OH radicals were produced by photolysis of nitrous acid (HONO), which was preinjected at triple-digit ppb levels before irradiation. HONO was synthesized from HCl vapor and NaNO2 according to Febo et al.26 NO2 was generated from NO and O3 with O3 being produced from pure O2 by UV-photolysis. Analytical instrumentation included a high-resolution proton-transfer-reaction time-of-flight mass spectrometer (PTRTOF 8000, Ionicon Analytik GmbH) for measurements of nitramines and nitramine degradation products, a chemiluminescence NO-detection instrument (CLD770 AL ppt, ECO PHYSICS) combined with a photolytic converter (PLC 760 MH, ECO PHYSICS) for NO2 detection and a temperature/ relative humidity sensor (UFT75-AT, MELTEC). 2.3. Organic Synthesis. Methylnitramine, CH3NHNO2. To 130 mL of 70% HNO3 cooled to 5 °C 23 g (124 mmol; MW 185.24) was added N-methyl-p-toluenesulfonamide portionwise with stirring. When all the sulfonamide had dissolved, the solution was cooled to 0 °C and 130 mL of 100% HNO3 was added. The solution was stirred at this temperature for another 10 min when this was poured carefully into a beaker with ice/ water ≈200 mL. The precipitated nitrosulfonamide (MW 230.24) was filtered off and recrystallized from ethanol (95%). Yield: 23.8 g (83%), Mp: 55−57 °C (lit.27 57 °C). 1H NMR (300 MHz, CDCl3): δ 2.44 (s, 3H), 3.68 (s, 3H), 7.35 (d, J 8.5 Hz, 2H), 7.87 (d, J 8.5 Hz, 2H). 13C NMR (75 MHz, CDCl3): δ 21.91 (CH3), 35.76 (CH3), 129.49 (CH × 2), 130.04 (CH × 2), 133.34 (C), 146.71 (C) ppm. To a solution of 40 mL of 1 M NaOH was added 4.6 g (20 mmol) of N-nitro-N-methyl-ptoluolsulfonamide, and the mixture was heated to reflux for 1 h or until a clear solution resulted. The solution was acidified with 6 M HCl(aq) by careful addition. Extraction with ether gave the methylnitramine (MW 76.05). Yield: 1.11g (73%). Mp: 34−36 °C (lit.27 32−36 °C). 1H NMR (300 MHz, CDCl3): δ 3.16 (s, 3H), 9.27 (bs, 1H). 13C NMR (75 MHz, CDCl3): δ 32 ppm. Dimethylnitramine, (CH3)2NNO2. Trifluoroacetic anhydride (91.7 mL, 0.663 mol) was placed in a 250 mL three-necked round-bottom flask equipped with a magnetic stirring bar and under nitrogen. The reaction mixture was cooled to −5 to 0 °C using a dry ice−ethanol cooling bath. Next, fuming nitric acid (30.8 mL, 0.733 mol) was added slowly by means of a dropping funnel. After complete addition of the nitric acid, the reaction mixture was cooled to −30 °C and dimethylformamide (10.6 mL, 0.137 mol) was added dropwise using a dropping funnel. After complete addition, the reaction was allowed to stand for 10 min before being concentrated to approximately half its volume in vacuo. Effervescence of brown gas was observed during this process. The resulting solution was poured on ice (∼150 g) and made alkaline (pH ∼10) using 50% NaOH. The aqueous solution was extracted six times with ether, and the

3. RESULTS AND DISCUSSION 3.1. Kinetic Study. The lifetimes of CH 3NHNO2 , (CH3)2NNO2, CH3OCH3, and CH3OH in the Oslo chamber when diluted in purified air were found to be on the order of days and dark-loss is therefore negligible in the present context. When mixed with H2 and O3 in purified air, the lifetimes were shorter; Figures S3−S6 (Supporting Information) show the observed decays. There is no report of H2 reacting with O3 in the gas phase, and the reaction of H2 with any O(3P) formed from ozone in the dark is slow: kO+H2 = 9 × 10−18 cm3 molecule−1 s−1.39 We therefore tentatively attribute the increased decays to bimolecular reactions with O3 and derive the following rate coefficients for O3 reaction: kO3+CH3NHNO2 = 3 3452



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× 10−22 cm3 molecule−1 s−1, kO3+(CH3)2NNO2 = 2 × 10−22 cm3 molecule−1 s−1, kO3+CH3OCH3 = 5 × 10−22 cm3 molecule−1 s−1, and kO3+CH3OH = 2 × 10−21 cm3 molecule−1 s−1, respectively. The kinetic analysis according to eq I assumes that only one loss process is taking place; other losses such as gas phase or surface reactions with the radical precursors will cause systematic errors in the relative rate determination. The OH experiments generally lasted for 1−2 h, including spectral recording and waiting time between photolysis periods, during which >90% of the initial nitramine reacted. The lifetimes of CH3NHNO2, (CH3)2NNO2, and CH3OCH3 with respect to reaction with O3 in the chamber in relation to the time span of the relative rate experiments are such that the systematic errors introduced by not considering the additional loss processes are almost negligible. For CH3OH, the additional loss due to reaction with O3 may be significant and the relative rate experiments involving CH3OH were therefore analyzed according to ⎧ [S] ⎫ ln⎨ 0 ⎬ − k O3 + S· ⎩ [S]t ⎭

∫0

t

[O3]· dt

⎧ ⎧ [R] ⎫ = k rel· ⎨ln⎨ 0 ⎬ − k O3 + R · ⎩ ⎩ [R]t ⎭ ⎪

⎪

∫0

t

⎫ [O3]· dt ⎬ ⎭

⎪

⎪

k rel =

kS kR (II)

Examples of the FTIR spectra obtained from the kinetic studies of the CH3NHNO2/CH3OCH3 reaction with OH radicals and the resulting residuals from the nonlinear leastsquares spectral analyses are shown in Figure 1. It can be seen that all spectral features are accounted for in the analysis. The only significant product bands appearing in the spectra during reaction in the CH3NHNO2/CH3OCH3 experiments stem from methyl formate originating in the oxidation of dimethyl ether. Spectra from the CH3NHNO2/CH3OH + OH experiments are illustrated in Figure S7 (Supporting Information); all spectral features are accounted for in the analyses. The decays of CH3NHNO2 and CH3OCH3 in the presence of OH radicals in three independent experiments are plotted as ln{[CH 3 NHNO 2 ] 0 /[CH 3 NHNO 2 ] t } vs ln{[CH3OCH3]0/[CH3OCH3]t} in Figure 2A, and Figure 2B shows a similar plot of the CH3NHNO2 and CH3OH decays in the presence of OH radicals in nine independent experiments. Least−squares fitting of the data resulted in the following relative rates (2σ statistical error limits): kOH+CH3NHNO2/ k O H + C H 3 O C H 3 = 0.304 ± 0.014 (16 data points), kOH+CH3NHNO2/kOH+CH3OH = 1.164 ± 0.016 (49 data points). Using today’s recommended values at 298 K for kOH+CH3OCH3 = 2.8 × 10−12 and kOH+CH3OH = 9.0 × 10−13 cm3 molecule−1 s−1 (both having uncertainty factors of 1.20), 40 places kOH+CH3NHNO2 at 0.85 ± 0.17 and 1.05 ± 0.21 × 10−12 cm3 molecule−1 s−1, respectively. Assuming that there are no additional molecule specific systematic errors, from the present experiments a best value of kOH+CH3NHNO2 = (9.5 ± 1.9) × 10−13 cm3 molecule−1 s−1 at 298 K can be derived. Examples of the FTIR spectra obtained from the kinetic studies of the (CH3)2NNO2/CH3OH reaction with Ȯ H radicals and the resulting residuals from the nonlinear leastsquares spectral analyses are shown in Figure 3. Representative spectra from the (CH3)2NNO2/CH3OCH3 + OH experiments are shown in Figure S8 (Supporting Information). The

Figure 1. Infrared transmission spectra of (A) the CH3NHNO2/ CH3OCH3/H2/O3 reaction mixture after 15 min photolysis, (B) CH3OCHO, (C) CH3NHNO2, (D) CH3OCH3, (E) O3, and (F) H2O. (B)−(F) have been shifted for the sake of clarity. Initial reactant mixing ratios: CH3NHNO2, 6 ppm; CH3OCH3, 4 ppm.

residuals show that all spectral features are accounted for in the analyses. The decays of (CH3)2NNO2 and CH3OCH3 in the presence of OH radicals in three independent experiments are plotted as ln{[(CH3)2NNO2]0/[(CH3)2NNO2]t} vs ln{[CH3OCH3]0/ [CH3OCH3]t} in Figure 4A, whereas Figure 4B shows a similar plot of the (CH3)2NNO2 and CH3OH decays in the presence of Ȯ H radicals in three independent experiment. Least−squares fitting of the data resulted in the following relative rates (2σ statistical error limits): kOH+(CH3)2NNO2/ k O H + C H 3 O C H 3 = 1.246 ± 0.015 (29 data points), kOH+(CH3)2NNO2/kOH+CH3OH = 3.83 ± 0.04 (39 data points). This places kOH+(CH3)2NNO2 at (3.49 ± 0.04) × 10−12 and (3.45 ± 0.04) × 10−12 cm3 molecule−1 s−1 at 298 K, respectively. Assuming no other molecule specific systematic errors and including the uncertainty factors in the reference rate constant values, a best value of kOH+(CH3)2NNO2 = (3.5 ± 0.7) × 10−12 cm3 molecule−1 s−1 at 298 K is derived from the present experiments which is in excellent agreement with the previous result of Tuazon et al., (3.6 ± 0.7) × 10−12 cm3 molecule−1 s−1 at 296 K.14 The available kinetic data show that reactivity of the two nitramines toward OH is around 20 times lower than that of their parent amine.41−44 The nitro group is electron withdrawing and its presence obviously changes the availability of the lone pair electrons on the amino-nitrogen to chemical reaction with the OH radicals. Electronic structure calculations were therefore undertaken to elucidate trends and differences in 3453



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Figure 3. Infrared transmission spectra of (A) the N-nitrodimethylamine/methanol/O3/H2 reaction mixture after 12 min photolysis, (B) N-methyl-N-nitroformamide, (C) N-nitrodimethylamine, (D) methanol, and the residual after spectral fitting. (B)−(D) have been shifted for the sake of clarity. Initial reactant mixing ratios: (CH3)2NNO2, 5 ppm; CH3OH, 6 ppm. Figure 2. (A) Relative rate plot showing the decays of CH3NHNO2 and CH3OCH3 at 1013 hPa and 298 K in the presence OH radicals; krel = 0.304 ± 0.014 (2σ) determined from 16 data points from 3 independent experiments. The data have been corrected for loss due to reaction with O3. (B) Relative rate plot showing the decay of CH3NHNO2 vs CH3OH at 1013 hPa and 298 K in the presence of Ȯ H radicals; krel = 0.823 ± 0.023 (2σ) determined from 49 data points from 9 independent experiments. The data have been corrected for loss due to reaction with O3, see text.

calculations are summarized in the Supporting Information (Tables S3−S6), structures of intermediates and saddle points are illustrated in Figure S9 (Supporting Information), and Cartesian coordinates are collected in Tables S7−S10 (Supporting Information). Results from G3, G4 and highlevel calculations on the parent amines44 have also been included in Tables S3−S6 for comparison. The potential energy surfaces of the OH reactions with amines and nitramines share similarities such as the existence of relatively strongly bonded prereaction (RMC) and postreaction (PMC) molecular complexes. The RMC structures are, however, different in the amines and nitramines; in the former, the Ȯ H radical is hydrogen bonded to the nitrogen lone-pair of the amino group, in the nitramines the bonding is to the oxygen of the nitro group (Figure S9 in the Supporting Information). The stabilization energies of the RMCs range from 10 to 16 kJ mol−1 for the nitramines, and from 25 to 28 kJ mol−1 for the amines. Following the PMCs, saddle points to hydrogen abstraction from either the amino or the methyl group are found with energies close to that of the reactant entrance energy. Finally, on the exit side there are relatively strongly bonded PMCs in which water is hydrogen bonded to the formed organic radical. The saddle point structures, Figure S9, show that the formed O−H bonds are relatively long, whereas the broken C−H and N−H bonds are fairly short. Thus, the transition states are reactant-like, and the reactions proceed via early barriers in accordance with the exothermic nature of these routes.

the potential energy hypersurfaces of amine and nitramine reactions with OH radicals and to assist in reaction mechanism development. High-level electron structure calculations are required to describe the kinetics of OH radical reactions with amines correctly.7,44 The nitro group introduces 26 additional electrons to a molecular system making similar high-level electron structure calculations very time-consuming for the nitramines. To uncover general trends resulting from nitration, we therefore settled to employ a modest calculation level in the present study (CCSD(T)/6-311++G(2d,2p)//BHandHLYP/ aug-cc-pVDZ), which is equivalent to the one used by Galano and Alvarez-Idaboy in their study of the OH reactions with CH3NH2, (CH3)2NH, and CH3CH2NH2.45 The stationary points on the potential energy surfaces of the OH radical reactions with the two nitramines CH 3 NHNO 2 and (CH 3 ) 2 NNO 2 and their corresponding parent amines, CH3NH2 and (CH3)2NH, are compared in Figures 5 and 6, respectively. Energies obtained in the electronic structure 3454



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Figure 5. Stationary points on the potential energy surfaces of the OH radical reactions with CH3NH2 (left) and CH3NHNO2 (right). Red: N−H abstraction routes. Black: C−H abstraction routes. Results from CCSD(T)/6-311++G(2d,2p)//BHandHLYP/aug-cc-pVDZ calculations.

Figure 4. (A) Relative rate plot showing the decay of (CH3)2NNO2 vs CH3OCH3 at 1013 hPa and 298 K in the presence of OH radicals; krel = 1.246 ± 0.015 (2σ) determined from 29 data points from 3 independent experiments. The data have been corrected for loss due to reaction with O3, see text. (B) Relative rate plot showing the decay of (CH3)2NNO2 vs CH3OH at 1013 hPa and 298 K in the presence of OH radicals; krel = 3.83 ± 0.04 (2σ) determined from 39 data points from 3 independent experiments. The data have been corrected for loss due to reaction with O3, see text.

Figure 6. Stationary points on the potential energy surfaces of the OH radical reactions with (CH3)2NH (left) and (CH3)2NNO2 (right). Red: N−H abstraction routes. Black: C−H abstraction routes. Results from CCSD(T)/6-311++G(2d,2p)//BHandHLYP/aug-cc-pVDZ calculations.

Under atmospheric conditions the Ċ H2NHNO2 radical is a priori expected to react with O2 following the same routes that were found for the Ċ H2NH2 radical,5,46 that is (1) via hydrogen abstraction leading to the corresponding N-nitro imine, CH2 NNO2, or (2) via the dioxy and oxy radicals leading to Nnitroformamide, CHONHNO2. The are two experimental results for the branching in the Ċ H2NH2 + O2 reaction; Nielsen et al.5 found around 85% imine formation (H-abstraction by O2) at atmospheric pressure from analyses of reactant and product time profiles in the CH3NH2 + OH reaction, whereas Onel et al.46 recently found close to 100% HO2 yield in the O2 reaction with Ċ H2NH2 radicals and no pressure dependency in the reaction at 25−75 mbar. The stationary points on the potential energy surfaces of the Ċ H2NH2 + O2 and Ċ H2NNO2 + O2 reactions are compared in Figure 7 (energies collected in Table S11 in the Supporting Information). Considering that the Ċ H2NHNO2 + O2 reaction is less exothermic than the Ċ H2NH2 + O2 reaction by ca. 15 kJ mol−1, that there are more internal degrees of freedom in Ȯ OCH2NHNO2 than in Ȯ OCH2NH2, and that the barrier heights from the ground state of the dioxy radical to HO2

In general, the barrier to C−H abstraction is found to be lower than that to N−H abstraction in the individual molecules. Further, the Born−Oppenheimer barrier heights evaluated at the saddle points increase from amine to nitramine by 5−6 kJ mol−1 for C−H abstraction and by 10−11 kJ mol−1 for N−H abstraction, reflecting the electron withdrawal of the nitro group. The calculations therefore suggest that C−H abstraction will be the dominant route in the OH radical reaction with CH3NHNO2. 3.2. Products in CH3NHNO2 Photo-oxidation. The reaction of CH3NHNO2 with OH radicals will proceed via hydrogen abstraction from either the −CH3 group or the >NH group: CH3NHNO2 + OH ̇ 2NHNO2 + H 2O ΔH298K ° = −97 kJ mol−1 → CH

(14a)

° = −81 kJ mol−1 ΔH298K

(14b)

̇ → CH3NNO 2 + H 2O

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Figure 7. Relative energies of stationary points on the potential energy surfaces of the Ċ H2NH2 + O2 and Ċ H2NHNO2 + O2 reactions. Results from G4 calculations.

elimination in the two systems are nearly the same intuitively suggests that the Ċ H2NHNO2 + O2 reaction will lead to less imine formation than the Ċ H2NH2 + O2 reaction. The competition between imine formation and stabilization of the chemically activated peroxy radical was investigated using a master equation model based on the potential energy surface sketched in Figure 7. In addition to the imine route, loss of the peroxy radical through reaction with NO was included. The calculations indicate the imine yield from the Ċ H2NH2 + O2 reaction to have a weak negative pressure dependence and an imine yield of 87% at 1 bar, whereas the imine yield from the Ċ H2NHNO2 was found to be strongly pressure dependent with an imine yield of only 4% at 1 bar. Variation of the barrier heights in the Ċ H2NH2 + O2 reaction system indicates that the large difference between the two systems is partially caused by the difference in barrier heights, as well as by the difference in number of degrees of freedom. The Ċ H 2NHNO2 radical, formed in (14a), will be vibrationally “hot” (ΔH°298K = −97 kJ mol−1) and may undergo N−N scission resulting in CH2NH and NO2 before reaction with O2; the dissociation is exothermic and with a barrier of only 30 kJ mol−1 (energies collected in Table S12 in the Supporting Information), a significant fraction of the vibrationally hot Ċ H2NHNO2 radical is therefore expected to follow this route. Master equation calculations show that direct dissociation of the Ċ H2NHNO2 radical to CH2NH and NO2 dominates completely in the relevant pressure range unless a large fraction of the CH3NHNO2 + OH reaction enthalpy (97 kJ mol−1) is deposited into the water molecule. It is not sufficient to have ν1 = 1 or ν3 = 1 in the product H2O (ν̃1 ≈ ν̃3 ≈ 3700 cm−1 corresponding to ∼44 kJ mol−1) and to have the remaining energy distributed according to the equipartition principle. Two quanta of the OH stretching vibrations need to be deposited in the product H2O to have the dissociation yield lowered significantly. With 2 quanta of the OH stretching vibration deposited in the product H2O, the direct dissociation at 1 bar is predicted to be around 7%. The CH3Ṅ NO2 radical reactions with O2, NO2, and NO are expected to resemble those established for the CH3Ṅ H radical.5,47−49 Figure 8a compares the lowest-energy stationary points on the PESs of CH3Ṅ H + O2 and CH3Ṅ NO2 + O2 reactions (energies are collected in Table S13 in the Supporting

Figure 8. Relative energies of stationary points on the potential energy surfaces of the (A) CH3NH + O2 and CH3NNO2 + O2 reactions, (B) CH3NH + NO2 and CH3NNO2 + NO2 reactions, and (C) CH3NH + NO and CH3NNO2 + NO reactions. Results from G4 calculations.

Information). As previously mentioned, the nitro group impacts the potential energy surfaces: the prereaction adduct between O2 and the amino radical, which is stabilized by 26 kJ mol−1 in the CH3Ṅ H + O2 system, is virtually nonexistent in the CH3Ṅ NO2 + O2 system, and the barrier to H-abstraction from the methyl group in CH3Ṅ NO2 by O2 leading to the corresponding imine increases dramatically from 22 kJ mol−1 above the entrance energy of the reactants in the CH3Ṅ H + O2 reaction48 to 64 kJ mol−1 in the CH3Ṅ NO2 + O2 reaction. A conventional transition state theory estimation of the rate coefficients suggests kCH3NH+O2 ∼ 10−18 cm3 molecule−1 s−1 and kCH3NNO2+O2 ∼10−26 cm3 molecule−1 s−1 for the two molecules; that is, the CH3Ṅ NO2 + O2 reaction will hardly be important at atmospheric conditions. In contrast, the NO and NO2 reactions 3456



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Scheme 1. Predicted Major Routes in the Atmospheric Photo-oxidation of CH3NHNO2a

a Reaction enthalpies (kJ mol−1) are from G4 calculations; values in italics refer to the corresponding reaction of the parent amine. The major product in atmospheric photo-oxidation of CH3NHNO2 is predicted to be CH2NH.

with CH3Ṅ H and CH3Ṅ NO2 proceed without electronic barriers. For CH3Ṅ H these reactions only constitute minor sinks at normal atmospheric conditions; for CH3Ṅ NO2, however, these reactions will constitute the dominant loss routes. The initial products in the CH3Ṅ NO2 + NO2 reaction are

Information) show that the n → π* transition of the nitroso group is red-shifted from 362(366) nm in syn(anti) CH3NHNO to 435 nm (275 kJ mol−1) in CH3N(NO2)NO with little change in its oscillator strength. The atmospheric lifetime of CH3N(NO2)NO with respect to photolysis will therefore be even shorter than that of other nitrosamines,7 i.e., only a few minutes. In summary, electronic structure calculations suggest that the by far dominant route in the OH radical reaction with CH3NHNO2 will be H-abstraction from the CH3 group, that CH2NH will be the major product in the OH-initiated atmospheric photo-oxidation of CH3NHNO2, and that CHONHNO 2 and CH 2 NNO 2 will constitute minor products. Hydrogen abstraction from the nitroamino group, constituting a minor route in the initial OH radical reaction with CH3NHNO2, will lead to formation of the dinitroamine, CH3N(NO2)2, and to nitrosomethane, CH3NO. Scheme 1 summarizes the major routes in atmospheric photo-oxidation of CH3NHNO2 according to quantum chemistry results; the scheme includes calculated reaction enthalpies (kJ mol−1, energies collected in Table S2 in the Supporting Information). FTIR spectra from OH initiated photo-oxidation experiments with CH3NHNO2 are congested and barely show any indications of products. CHONHNO2, CH2NH, and CH2 NNO2 are all likely to be more reactive than their parent precursor, which may explain why essentially no new bands attributable to organics are observed in the OH photooxidation experiments. Oxidation experiments with Cl atoms, however, reveal clear spectral features of the −NHNO2 moiety, Figure 8. The N−H stretching mode is red-shifted by around 5 cm−1, the two NO2 stretching bands are shifted from 1611 and 1331 cm−1 in CH3NHNO2 to 1631 and 1286 cm−1 in the product, and a carbonyl-stretching band appears at 1763 cm−1. These spectral features are consistent with N-nitroformamide, CHONHNO2, as a major product in the Cl-atom initiated photo-oxidation of N-nitromethylamine. Quantum chemistry calculations of vibrational wavenumbers of CH3NHNO2, CHONHNO2 and CH2NNO2, presented in Tables S17− S19 (Supporting Information), favor this interpretation. Figure 9 includes for comparison theoretical IR spectra of CHONHNO2 and CH2NNO2 (based on unscaled wavenumbers from B3LYP/aug-cc-pVTZ calculations). The spectral data do not allow an unambiguous identification of CH2 NNO2 as a product, and none of the rotational fine structure in the CH2NH spectrum (Figure S10 in the Supporting Information) could be identified during the Cl-initiated oxidation. It should be noted that the CH3NHNO2 + Cl atom reaction is around 65 kJ mol−1 less exothermic than the corresponding OH radical reaction. Consequently, the CH2NHNO2 radical formed in the CH3NHNO2 + Cl reaction will likely be less prone to dissociation than the CH2NHNO2 radical formed in the corresponding OH reaction (the relevant

̇ CH3NNO 2 + NO2 → CH 2N(NO2 )2‡

° = −125 kJ mol−1 ΔH298K

(15a)

° = −109 kJ mol−1 → CH 2N(NO2 )ONO‡ ΔH298K

(15b)

The vibrationally hot products/intermediates may undergo various isomerization and/or dissociation reactions; the lower energy stationary points of the CH3Ṅ NO2 + NO2 PES are compared to those of CH3Ṅ H + NO248 in Figure 8b (energies are collected in Table S14 in the Supporting Information). For the sake of simplicity, the figure does not include the high barrier routes from CH3N(NO2)ONO to CH2N(O)NO2 + HNO via a saddle point (SP) at ΔEElec+ZPE = 35.7 kJ mol−1, to CH2NNO2 + HONO via an SP at ΔEElec+ZPE = 131.0 kJ mol−1, and to CH2NONO + HONO via an SP at ΔEElec+ZPE = 53.5 kJ mol−1 above the entrance energy of the reactants. The CH3Ṅ H + NO2 → CH3NHONO → CH3NHȮ + NO route, which at atmospheric conditions will be followed by reaction with O2 to give CH3NO + HO2, was not confirmed in CH3NH2 photo-oxidation experiments,5 because its structural isomer, CHONH2 (formamide), is also among the products. It should be noted that the corresponding route in the (CH3)2Ṅ + NO2 reaction, leading to the (CH3)2NO radical, was confirmed in low pressure experiments by Lazarou et al.50 In summary, the quantum chemistry results show that the CH3Ṅ NO2 radical reaction with NO2 essentially leads to the dinitroamine, CH3N(NO2)2, and to nitrosomethane, CH3NO, and not to the N-nitro imine, CH2NNO2. Primary nitrosamines are thermally unstable and undergo isomerization reactions in the gas phase.47,49 Figure 8c compares stationary points of the CH3Ṅ NO2 + NO PES to the relevant ones of the CH3Ṅ H + NO PES (energies are collected in Table S15 in the Supporting Information). The calculations show barriers of CH3N(NO2)NO → CH2N(NO2)NOH isomerization of more the 50 kJ mol−1, and of Htransfer to the nitro group, leading to CH2NNO + OH + NO, of around 100 kJ mol−1 above the entrance energy of the reactants. According to the theoretical calculations, the mixed nitroso− nitramine will be thermally stable and its major atmospheric loss is consequently expected to be due to photolysis. The electronic absorption spectra of CH3NHNO, CH3N(NO2)NO, and CH3NHNO2 were mapped in TD-DFT calculations employing the B3LYP functional and the aug-cc-pVTZ basis set. The results (summarized in Table S16 in the Supporting 3457



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for the CH3NHĊ H2 radical,5,46 is blocked. Other routes that have been considered include (1) N−N scission in the CH 3 N(NO 2 )Ċ H 2 radical leading to NO 2 and methyl methanimine, CH3NCH2, (2) H-migration in the dioxy radical CH3N(NO2)CH2OȮ to Ċ H2N(NO2)CH2OOH, and (3) dissociation of the oxy radical CH3N(NO2)CH2Ȯ resulting in CH2O and the CH3Ṅ NO2 radical. The NN scission process CH3N(NO2)Ċ H2 → CH3N CH2 + NO2 is exothermic by 52 kJ mol−1 and involves a barrier of only 27 kJ mol−1 (energies of relevant species are collected in Table S20 in the Supporting Information). Considering that the CH3N(NO2)Ċ H2 radical is formed in a 107 kJ mol−1 exothermic reaction, it is likely that a significant fraction of the vibrationally hot CH3N(NO2)Ċ H2 will dissociate. The importance of this route was estimated in master equation calculations. Our results indicate that direct dissociation will dominate completely when the CH3N(NO2)Ċ H2 radical is formed with enough energy to overcome the barrier, and that 18% will dissociate directly at atmospheric pressures when energy corresponding to two quanta of the O−H vibrational modes are deposited in the product water. The second route, initiated by H-migration in the dioxy radical, was recently shown to be very efficient in OH regeneration in the (CH3)3N reaction with OH in the presence of O2.46 For the CH3NHCH2OȮ radical the barrier to HO2 elimination is around 20 kJ mol−1 lower than that to Hmigration resulting in negligible OH regeneration under similar conditions, Figure 10 (energies collected in Table S21 in the

Figure 9. Infrared absorption spectra from Cl atom reaction with Nnitromethylamine. (A) Spectrum of N-nitromethylamine. (B) Spectrum recorded after 5 min of reaction. (C) Predicted spectra of CH3NHNO2 (black curve), CHONHNO2 (red curve), and CH2 NNO2 (blue curve). Initial CH3NHNO2 mixing ratio: 5 ppm.

energies are included in Table S12 in the Supporting Information), and reservations are made with respect to the spectroscopic evidence for CHONHNO2 being a major product in the atmospheric photo-oxidation of CH3NHNO2. The PTR-TOF-MS study of the OH initiated CH3NHNO2 photo-oxidation was partly obstructed by the low reactivity and showed few ion signals that could be correlated to organic products in amounts above 1% of CH3NHNO2 reacted. However, significant amounts of HCN (ion signal at m/z 28.021, CH2N+) were detected. Of the products expected in CH3NHNO2 photo-oxidation, CHONHNO2 appears to be reasonably stable in the gas phase, as evidenced by the Cl experiments, whereas CH2NH has been shown to have very short lifetime in smog chamber experiments.5 The major product in the CH2NH + OH reaction is HCN.5 It is therefore concluded from the FTIR and PTR-ToF-MS results and from the theoretical studies that the major product in atmospheric CH3NHNO2 photo-oxidation is CH2NH, and that CHONHNO2 is only a minor product. 3.3. Products in (CH3)2NNO2 Photo-oxidation. The reaction of (CH3)2NNO2 with OH radicals proceeds via hydrogen abstraction from the −CH3 group:

Figure 10. Stationary points on the potential energy surfaces of the CH3NHĊ H2 + O2 and CH3N(NO2)Ċ H2 + O2 reaction surfaces. Results from G4 calculations.

Supporting Information). There are two reports on the branching in the CH3NHĊ H2 + O2 reaction; Nielsen et al.5 found around 55% CH3NCH2 + HO2 formation at atmospheric pressure from analyses of the time-profiles of reactants and products in the (CH3)2NH + OH reaction, whereas Onel et al.46 found close 100% HO2 formation in the CH3NHĊ H2 + O2 reaction and no pressure dependency at 25−100 mbar. The HO2 elimination route does not exist for CH3N(NO2)Ċ H2, and OH regeneration in the CH3N(NO2)Ċ H2 + O2 reaction may therefore be feasible. However, the barriers to Hmigration in CH3N(NO2)CH2OȮ and subsequent O−O scission in Ċ H2N(NO2)CH2O−OH are calculated to be considerably higher than the corresponding barriers in the parent amine system, Figure 10, suggesting that this route will not be important to CH3N(NO2)CH2OȮ .

̇ 2 + H 2O (CH3)2 NNO2 + OH → CH3N(NO2 )CH ° = −107 kJ mol−1 ΔH298K

(16)

Under atmospheric conditions the CH3N(NO2)Ċ H2 radical is expected to react with O2 via dioxy and oxy radicals leading to N-methyl-N-nitroformamide, CHON(CH3)NO2; the alternative route to imine formation, which was found to be dominant 3458



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Scheme 2. Major Routes in the Atmospheric Photo-oxidation of (CH3)2NNO2a

a Reaction enthalpies (kJ mol−1) are from G4 calculations; values in italics refer to the corresponding reaction for the parent amine. The major product in atmospheric photo-oxidation of (CH3)2NNO2 is predicted to be CH3NCH2.

The rate of CH3N(NO2)CH2Ȯ dissociation to CH3Ṅ (NO2) + CH2O, which is endothermic by 38 kJ mol−1, was estimated employing the Méreau et al.51 structure−activity relationship for alkoxyl radical decomposition reactions to be around 10% of the competing rate for H-abstraction by O2. Finally, the atmospheric fate of the CH3Ṅ (NO2) radical was discussed above; the major fate will be reaction with NO2 leading to the dinitro amine, CH3N(NO2)2. Scheme 2 summarizes the major routes in atmospheric (CH3)2NNO2 photo-oxidation according to results from quantum chemistry and includes as well the computed reaction enthalpies (/kJ mol−1). The FTIR spectra from the OH initiated photo-oxidation of (CH3)2NNO2 show product bands typical of the nitro group (Figure 3). Figure 11 shows the spectral time evolution during the (CH3)2NNO2 + Cl reaction, and a comparison with the theoretical, unscaled IR spectra of (CH3)2NNO2, CH3N(NO2)CHO, and CH2NNO2 (vibrational wavenumbers of

(CH3)2NNO2, CHON(NO2)CH3, and CH3NCH2 are presented in Tables S22−S24 in the Supporting Information) supports this interpretation. The Cl oxidation experiments confirm the spectral features attributed to N-methyl-Nnitroformamide; CH3NCH2 has a weak IR spectrum, and not even the characteristic C-type band at 1025 cm−1 (ν16(A″), Figure S11 in the Supporting Information) can be seen in the spectra. During the experiment a number of distinct bands having rotational fine structure can be seen to grow in HNO3 (1710, 1326, 879, 762 cm−1), ClNO2 (1685, 1319, 793 cm−1), CO2 (667 cm−1), and HCN (712 cm−1), the last which is obviously a secondary product in the (CH3)2NNO2 photooxidation. Again, it should be noted that the (CH3)2NNO2 + Cl atom reaction is around 65 kJ mol−1 less exothermic than the corresponding OH radical reaction, and less of the CH3N(NO2)Ċ H2 radical formed in the (CH3)2NNO2 + Cl reaction is expected to dissociate than in the corresponding OH reaction. There is, however, spectroscopic evidence for CH3N(NO2)CHO being a major product in the atmospheric photooxidation of (CH3)2NNO2. The time evolution/reproducibility of the PTR ion signals detected in two subsequent (CH3)2NNO2 photo-oxidation experiments is illustrated in Figure S12 (Supporting Information), whereas the time-integrated increase in signals of photooxidation products relative to the time-integrated decrease of the main reagent ion signal is shown in Figure S13 (Supporting Information). The largest ion signal in the PTR-TOF-MS study of the OH-initiated (CH3)2NNO2 degradation that can be attributed to a photo-oxidation product is m/z 105.030 (C2H5N2O3+, protonated N-methyl-N-nitroformamide). Additional, less intense ion signals, not due to chamber artifacts, were detected at m/z 77.033 (CH 5 N 2 O 2 + ), 75.033 (C2H5NO2+), 44.052 (C2H6N+, protonated CH3NCH2), 44.014 (CH2NO+, protonated HNCO), and 28.021 (CH2N+, protonated HCN).

4. CONCLUSIONS AND IMPLICATIONS The product studies show CH2NH and CH3NCH2 as the major products in the OH-initiated photo-oxidation of CH3NHNO2 and (CH3)2NNO2, respectively. N-Nitro amides, CHONHNO2 and CH3N(NO2)CHO, also form (the yields remain to be determined). CH2NH was reported to be the major product in CH3NH2 photo-oxidation experiments.5 It has been speculated that N2O could be a product in CH2NH photo-oxidation;52 this has, however, not been confirmed by experiments.5 CH3NCH2 was reported to be the major product in atmospheric photo-oxidation of both (CH3)2NH2,5 and (CH3)3N.5 In spite of their central role in amine chemistry, the atmospheric fate of imines is essentially still unmapped; clearly more research is needed. Concerning N-nitro amides,

Figure 11. Infrared absorption spectra from Cl atom reaction with Nnitrodimethylamine. (A) Spectrum of N-nitrodimethylamine. (B)− (D) Spectra recorded after 5, 10, and 15 min of reaction. (E) Predicted IR spectra of N-nitrodimethylamine (black curve), N-methyl-Nnitroformamide (red curve), and N-nitromethanimine (blue curve). Initial (CH3)2NNO2 mixing ratio: 5 ppm. 3459



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S10). FTIR spectrum of CH3NCH2 (Figure S11). Time variation of PTR-TOF-MS ion signals during (CH3)2NNO2 photo-oxidation experiments (Figure S12). Mass spectral difference observed during the photo-oxidation of (CH3 ) 2NNO2 (Figure S13). Wavenumber regions and chemical components included in spectral analyses of the different reaction mixtures (Table S1). Energies, enthalpies, and Gibbs free energies from G3 and G4 calculations (Table S2). Energies of stationary points of the potential energy surfaces of the OH radical reactions with CH3NH2, CH3NHNO2, (CH3)2NH, and (CH3)2NNO2 (Tables S3−S6). Cartesian coordinates of reactants, products and stationary points on the PES´s of the OH reactions with CH3NH2, CH3NHNO2, (CH3)2NH, and (CH3)2NNO2 (Tables S7−S10). Energies of stationary points on the potential energy surfaces of the CH2NH2 + O2 and CH2NHNO2 + O2 reactions (Table S11). Energies of stationary points on the potential energy surfaces of the CH3NHNO2 + OH and CH3NHNO2 + Cl reactions (Table S12). Energies of stationary points on the potential energy surfaces of the CH3NH + O2 and CH3NNO2 + O2 reactions (Table S13). Energies of stationary points on the potential energy surfaces of the CH3NH + NO2 and CH3NNO2 + NO2 reactions (Table S14). Energies of stationary points on the potential energy surfaces of the CH3NH + NO and CH3NNO2 + NO reactions (Table S15). Vertical excitation energies and oscillator strengths ( f) of singlet states in CH3NHNO, CH3N(NO2)NO, and CH3NHNO2 (Table S16). Cartesian coordinates, calculated vibrational wavenumbers, and integrated band intensities of CH3NHNO2, CHONHNO2, and CH2NNO2 (Tables S17−S19). Energies of stationary points on the potential energy surface of the (CH3)2NNO2 + OH and (CH3)2NNO2 + Cl reactions (Table S20). Energies of stationary points on the PESs of the CH3NHCH2 + O2 and CH3NNO2CH2 + O2 reactions (Table S21). Cartesian coordinates, calculated vibrational wavenumbers, and integrated band intensities of (CH3)2NNO2, CHON(NO2)CH3, and CH3NCH2 (Tables S22−S24). This material is available free of charge via the Internet at http:// pubs.acs.org.

very little experimental information is available. In particular, there are no toxicology data. The rate coefficients for gas phase reaction of OH radicals with the simplest primary and secondary aliphatic nitramines, CH3NHNO2 and (CH3)2NNO2, have been determined to be 0.95 and 3.5 × 10−12 cm3 molecule−1 s−1 at 298 K, respectively. Assuming a global average OH concentration of 106 cm−3,53 their estimated atmospheric lifetimes will be 12 and 3.3 days, respectively, with respect to reaction with OH radicals. The atmospheric lifetimes with respect to reaction with O3 (rate coefficients NNO2 = 9.3 (constrained to be the same as for −NH2, >NH, >N−, and >NNO). The parametrization is based on kinetic data from one single nitramine, (CH3)2NNO2, and is not applicable to primary nitramines. More data are clearly needed to establish a general OH-SAR for nitramines.

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AUTHOR INFORMATION

Corresponding Authors

*C. J. Nielsen: e-mail, [email protected]; fax, +47 2285 5441; tel, +47 2285 5680. *Y. Stenstrøm: e-mail, [email protected]; fax, +47 6496 5901; tel, +47 6496 5893. *A. Wisthaler: e-mail, [email protected]; fax, +47 2285 5441; tel, +47 2285 9181.

ASSOCIATED CONTENT

S Supporting Information *

FTIR spectrum of CH3NHNO2 vapor (Figure S1). FTIR spectrum of (CH3)2NNO2 vapor (Figure S2). Observed gas phase loss of CH3NHNO2 in the Oslo reaction chamber under dark conditions (Figure S3). Observed gas phase loss of (CH3)2NNO2 in the Oslo reaction chamber under dark conditions (Figure S4). Observed gas phase loss of CH3OH in the Oslo reaction chamber under dark conditions (Figure S5). Observed gas phase loss of CH3OCH3 in the Oslo reaction chamber under dark conditions (Figure S6). Examples of the FTIR spectra obtained during the kinetic studies of the CH3NHNO2/CH3OH reaction with OH radicals (Figure S7). Examples of the FTIR spectra obtained during the kinetic studies of the of a (CH3)2NNO2/CH3OCH3 reaction with OH radicals (Figure S8). Structures of stationary points on the potential energy surfaces of the OH radical reactions with CH 3 NH 2 , CH 3 NHNO 2 , (CH 3 ) 2 NH, and (CH 3 ) 2 NNO 2 (Figure S9). FTIR spectrum of CH2NH vapor (Figure

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work is part of the Atmospheric Degradation of Amines (ADA) project supported by Masdar, Statoil, Vattenfall, Shell, and the CLIMIT program under contracts 193438, 201604, and 208122 and has received additional support from the Norwegian Supercomputing Program (NOTUR) through a grant of computer time (Grant No. NN4654K), from the VISTA-program (project 6157 “Study of the formation and stability N-nitrosamines, N-nitramines, and N-nitroamides resulting from degradation of amines emitted to the atmosphere”), and from the Research Council of Norway 3460



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through a Centre of Excellence Grant (Grant No. 179568/ V30). The authors are indebted to Prof. Jean-Claude Guillemin, Ecole Nationale Supérieure de Chimie de Rennes, for help in synthesizing methanimine.

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