CH3ONO and CD3ONO Synthesis Method

DIRECT KINETIC STUDY OF METHOXY RADICAL REACTING WITH NO2 AND O2
AND DEUTERIUM KINETIC ISOTOPE EFFECTS
by
Jiajue Chai
A dissertation
submitted in partial fulfillment
of the requirements for the
Doctor of Philosophy Degree
State University of New York
College of Environmental Science and Forestry
Syracuse, New York
December 2013
Approved: Department of Chemistry
Theodore S. Dibble, Major Professor S. Scott Shannon, Chair

Examining Committee
Greg Boyer, Department Chair S. Scott Shannon, Dean

The Graduate School
© 2013
Copyright
Jiajue Chai
All rights reserved
Acknowledgements
A Ph.D. dissertation is never a one-person job. There are many people who deserve credit
for their contribution to the useful contents of this dissertation. I know I will not find the
appropriate amount of space here to properly thank all of those from whom I have benefited
while working on it.
First of all I would like to thank my advisor, Dr. Theodore Dibble. You have been very
supportive throughout my whole time here at ESF. Thank you for believing in me, for having
patience with my questions, and for leaving me enough liberty of working from entire project.
Your support went far beyond scientific, financial, and practical issues and enabled me to pursue
this dissertation. I also thank my steering committee members: Dr. Scott Shannon, Dr. Dave
Kieber, Dr. John Hassett, Dr. Joseph Chaiken, and Dr. Huiting Mao. In addition to taking time to
sit for my exams and provide me much needed feedback, each has made unique contributions to
my development I feel must be noted. I have been fortunate to get help and advice from
numerous people outside of ESF. Dr. Jeff Tyndall has given critical advice in experimental
design and writing manuscripts.
Next I must thank my lab mates who helped me during the experiment and dissertation
writing including Hongyi Hu, Feng Zhang, Yue Zeng, Karen Schmitt and Yuge Jiao. I would not
have been able to survive without my terrific group of friends behind me in Chemistry
Department. Thank you to the Jennifer Croskrey, Joanna Kinsey, Anna Flach, John Kelly,
Mingyu Li, Xin Liu, Adam Stringer, Ryan Tappel, Lili Wang, Xingfei Zhao, Inger Tyssebotn,
Alex Levine, Smriti Sharma, Sara Botton, Zhuyun Ye, Liang Chen, Caiping Lin, Josh Slocum,
Kyle Bell, Liyang Shao, Wenyang Pan, and so on.
Finally, I would like to thank my family. Thank you to my parents Aixiang Hou, Peihuai
Chai who helped me to the best of their ability at all the time especially after my daughter was
born; to my wife Xian Wang, having your support push me through my toughest time in the past
iii
Table of Contents
List of Tables .......................................................................................................... vi
List of Figures ....................................................................................................... viii
Abstract ................................................................................................................... xi
Chapter 1. Introduction...........................................................................................1
1.1 Tropospheric ozone ............................................................................................................... 2
1.2 Secondary organic aerosols (SOAs) ...................................................................................... 4
1.3 Research objectives ............................................................................................................... 6
References ................................................................................................................................. 10
Chapter 2. Background .........................................................................................13
2.1 Temperature profile of atmosphere ..................................................................................... 13
2.2 Atmospheric chemistry of alkoxy radicals .......................................................................... 14
2.2.1 RO• + O2 .................................................................................................................................... 17
2.2.2 Unimolecular isomerization ....................................................................................................... 24
2.2.2.1 1,5-H shift isomerization......................................................................................... 24
2.2.2.2 Other types of H-shift isomerization reactions ....................................................... 29
2.2.3 Unimolecular decomposition ..................................................................................................... 30
2.2.3.1 Decomposition of unsubstituted alkoxy radical ...................................................... 30
2.2.3.2 Structure activity relationship for decomposition of alkoxy radicals...................... 32
2.2.3.3 Decomposition of functionalized alkoxy radical .................................................... 34
2.3 Methoxy + NO2 ................................................................................................................... 40
2.4 Rate theory for recombination reaction (P-dependence) ..................................................... 46
2.5 Kinetics of CD3O• reactions ................................................................................................ 49
2.6 Principle of LFP/LIF method .............................................................................................. 51
2.6.1 CH3ONO photolysis................................................................................................................... 51
2.6.2 Principle of LIF spectroscopy .................................................................................................... 52
References ................................................................................................................................. 55
Chapter 3. Experimental Design ..........................................................................63
3.1 Preparation of gaseous reactant ........................................................................................... 63
iv
3.2 Absolute rate constant—LFP/LIF method .......................................................................... 66
3.2.1 Experimental apparatus.............................................................................................................. 66
3.2.2 Gas handling system .................................................................................................................. 71
3.2.3 Estimation of radical concentration ........................................................................................... 71
3.2.4 Kinetic experiment ..................................................................................................................... 74
References ................................................................................................................................. 79
Chapter 4. Rate constants and kinetic isotope effects for methoxy radical
reacting with NO2 and O2 ......................................................................................81
4.1 Introduction ......................................................................................................................... 81
4.2 Experiment .......................................................................................................................... 86
4.2.1 Preparation of gaseous reactants ................................................................................................ 86
4.2.2 Experimental setup for absolute rate constant measurement—LP/LIF method ........................ 87
4.3 Results and discussion ......................................................................................................... 89
4.3.1 Absolute rate constants for CH3O• + NO2 and CD3O• + NO2 ................................................... 89
4.3.2 Rate constant for CH3O• + O2 and CD3O• + O2 and tunneling effect........................................ 96
4.4 Conclusion ......................................................................................................................... 101
References ............................................................................................................................... 103
Chapter 5. Pressure dependence and kinetic isotope effects in the absolute
rate constant for methoxy radical reacting with NO2 ......................................107
5.1 Introduction ....................................................................................................................... 107
5.2 Experimental ..................................................................................................................... 111
5.3 Results and discussion ....................................................................................................... 113
5.4 Conclusion ......................................................................................................................... 131
References ............................................................................................................................... 133
Chapter 6. Conclusions ........................................................................................136
References ............................................................................................................................... 140
Appendix I. ...........................................................................................................141
Curriculum Vitae .................................................................................................152
v
List of Tables
Table 2.1 Absolute rate data in the form of Arrhenius expression for non-substituted alkoxy
radicals from C2 to C7. Both k and A are in units of cm3 molecule-1 s-1. Cited errors are
2σ. ..................................................................................................................................... 19
Table 2.2 Absolute rate constant Arrhenius expression for decomposition at high pressure limit.
........................................................................................................................................... 31
Table 2.3 The conditions of previous direct rate measurement on CH3O• + NO2. Note: DF-LIF is
discharge-flow/laser induced fluorescence; LFP-LIF is laser flash photolysis/laser
induced fluorescence. ........................................................................................................ 41
Table 4.1 Rate constants for CH3O + NO2 at 700 Torr. Cited errors are statistical 2σ, and the 5%
uncertainty for [NO2] measurement is not included. ........................................................ 95
Table 4.2 The ratios of the rate constants for the CH3ONO+O2/NO2 (i.e.: k1/k2) and CD3ONO +
O2/NO2 (i.e., k3/k4) experiments at all measured temperatures, 700 Torr. The error bars
for all numbers are 2σ. ...................................................................................................... 97
Table 4.3 Absolute rate constant for k1 and k3 (unit: cm3 molecule-1 s-1) and kinetic isotope effect
(KIE). The quoted errors (2σ) include statistical uncertainties from linear fitting of both
relative and absolute rate methods, uncertainty in the methyl nitrate concentration (10%),
plus uncertainties of concentration measurement of NO2 from absolute rate method. .... 97
Table 5.1 Pressure dependent rate constants k1 and at k2 different temperatures. Error bars
represent 2σ statistical error propagated with 5% uncertainty in [NO2] concentration
determination. ................................................................................................................. 117
Table 5.2 High pressure and low pressure limit rate constants resulted from fits of our data to
Troe expression (E5.4). Quoted errors are 2σ. Unit: k0—cm-6 molecule-2 s-1; k∞—cm-3
molecule-1 s-1 ................................................................................................................... 122
Table 5.3 KIE values (kH/kD) at different temperatures for k0, and k∞. Cited errors are 2σ. ...... 126
Table 5.4. Troe parameters resulting from different studies for the methoxy + NO2 reaction.
Cited errors are 2σ. Units for k0 and k∞ are cm6 molecule-2 s-1 and cm3 molecule-1 s-1
respectively. .................................................................................................................... 128
vi
Table I-1-1 CH3O + NO2 at 250 K ............................................................................................ 141
Table I-1-2 CH3O + NO2 at 265 K ............................................................................................ 142
Table I-1-3 CH3O + NO2 at 278 K ............................................................................................ 143
Table I-1-4 CH3O + NO2 at 295 K ............................................................................................ 144
Table I-1-5 CH3O + NO2 at 316 K ............................................................................................ 145
Table I-1-6 CH3O + NO2 at 333 K ............................................................................................ 146
Table I-2-1 CD3O + NO2 at 250 K ............................................................................................ 147
Table I-2-2 CD3O + NO2 at 277 K ............................................................................................ 148
Table I-2-3 CD3O + NO2 at 294 K ............................................................................................ 149
Table I-2-4 CD3O + NO2 at 319 K ............................................................................................ 150
Table I-2-5 CD3O + NO2 at 335 K ............................................................................................ 151
vii
List of Figures
Figure 2.1 Temperature structure of atmosphere below 50 km from U.S. Standard Atmosphere.
........................................................................................................................................... 14
Figure 2.2 Simplified energy diagram for reaction CH3O• + NO2. TS is transition state. ........... 44
Figure 2.3 Schematic fall-off curve for the apparent recombination rate constant krec implicitly
defined by E2.8. ................................................................................................................ 47
Figure 2.4 Schematic energy profile for two isotopic variants of the methoxy + O2 reaction. (TS
denotes transition state; Horizontal line represent zero-point energy level of
corresponding species) ...................................................................................................... 50
Figure 2.5 Principle of laser induced fluorescence (LIF) by simplified Jablonski diagram ......... 53
Figure 3.1 The experimental apparatus used in this study. ........................................................... 68
Figure 3.2 Timing sequence of an LIF experiment with a specific delay time of 5 μs between two
laser pulses preset in DDG. ............................................................................................... 70
Figure 3.3 UV spectra of CH3ONO and CD3ONO. Black line—CH3ONO; red line—CD3ONO.
........................................................................................................................................... 74
Figure 3.4 LIF spectrum of CH3O• in the range of 285.7-302.7 nm at 295 K. Details of excitation
transitions (𝐴2𝐴1 ← 𝑋2𝐸) in ν3 mode (C—O stretching) are denoted on top of each peak.
For example, in the case of 5’←0’’, 0’’ represents the first vibrational level of ground
electronic state, 5’ represents the sixth vibrational level of first excited electronic state.
Blue line—this work; black dot line—the work of Inoue et al.14 ..................................... 75
Figure 3.5 LIF spectrum of CD3O• in the range of 285.7-302.7 nm at 295 K. Details of excitation
transitions (𝐴2𝐴1 ← 𝑋2𝐸) in ν3 mode (C—O stretching) are denoted on top of each peak.
The meanings of the denotations are the same as Figure 3.4. Red line—this work; black
dot line—the work of Inoue et al.14 .................................................................................. 76
Figure 4.1 Typical linear decay of ln(LIF intensity) as a function of the delay time for
CH3O•+NO2 at total pressure 700 Torr and 295 K. NO2 concentrations in molecule cm-3
are: 9.2×1014, 1.84×1015, 3.06×1015, 4.58×1015, and 6.10×1015. Error bars are 2σ. ......... 91
viii
Figure 4.2 Plot of k’ versus [NO2] at 295 K under different pressures. Error bars are 2σ precision
of the fitted slope of ln(intensity) versus time. ................................................................. 91
Figure 4.3 Pressure dependence of k2 for CH3O• + NO2 at room temperature. Cited errors are 2σ
of precision in the fitted slopes of plots of k’ versus [NO2]. The black solid line is the fit
of the Troe expression to our results (see details below). All these data are listed in
Appendix I. ....................................................................................................................... 93
Figure 4.4 Comparison of the pressure dependent behavior for rate constants for CH3O + NO2
(squares) and CD3O + NO2 (triangles) at room temperature. Cited errors are 2σ statistical
errors in the fitted slopes of plots of k’ versus [NO2]. ...................................................... 94
Figure 4.5 Temperature dependence of methoxy + NO2 at 700 Torr. Error bars are statistical 2σ,
and the 5% uncertainty for [NO2] measurement is not included. ..................................... 95
Figure 4.6 Temperature dependence of methoxy + O2 at 700 Torr. The solid lines represent linear
least squares fits to the data. ............................................................................................. 98
Figure 4.7 Temperature dependent rate constant for CH3O+O2 in the range of 250 – 610 K. The
solid line represents the Arrhenius fit suggested in reference 20, Among the previous
experimental data, the ones from Wiebe et al.25 and Cox et al.24 at 298 K (denoted by*)
were derived by combining originally determined relative rate constant with the absolute
rate constant for the reference reaction measured in current work (R4.2) or elsewhere
(CH3O•+NO).46 ................................................................................................................. 99
Figure 5.1 Simplified energy diagram for reaction CH3O• + NO2. ............................................ 109
Figure 5.2 Typical linear decays of ln(LIF intensity) as a function of the delay time for
CH3O•+NO2 at total pressure 700 Torr and 250 K. NO2 concentrations in molecule cm-3
are: 6.3×1014, 1.19×1015, 1.89×1015, 2.73×1015, and 3.52×1015. .................................... 115
Figure 5.3 Plot of k’ for CH3O• + NO2 versus [NO2] at 250 K under different pressures. Error
bars are 2σ of the fitted slope of ln(intensity) versus time. ............................................ 116
Figure 5.4 Fall-off curve non-linear fitting of absolute rate constant k1a measured for CH3O• +
NO2 in this study to equation E5.4. T=250-333 K, P=30-700 Torr. Error bars represent 2σ
statistical error propagated with 5% uncertainty in [NO2] concentration determination.
The insertion demonstrates the magnification of the low pressure data. ........................ 120
ix
Figure 5.5 Fall-off curve non-linear fitting of absolute rate constant k2a measured for CD3O• +
NO2 in this study to equation E5.4. T=250-335 K, P=30-700 Torr. Error bars represent 2σ
The insertion demonstrates the magnification of the low pressure data. ........................ 121
Figure 5.6 Temperature dependence of low pressure limit rate constant k1a0 and k2a0 for
CH3O• + NO2 and CH3O• + NO2 respectively. The fitting is based on the equation k0 =
k298K0 × (T298) − n. Error bars are 2σ statistical error from non-linear fitting. ....... 123
Figure 5.7 Temperature dependence of high pressure limit rate constant k1a∞ and k2a∞ for
CH3O• + NO2 and CD3O• + NO2 respectively. The fitting is based on the equation k∞ =
k298K∞ × (T298) − m. Error bars are 2σ statistical error from non-linear fitting. ..... 124
Figure 5.8 Temperature dependent KIE for k0 and k∞. Cited errors are 2σ. ............................... 126
Figure 5.9 Comparison of fall-off curves (lines) for k1a calculated using parameters resulted from
current study together with the absolute T,P-dependent k1a values (symbols) determined
by Wollenhaupt et al. in Ar,13 except for results of Frost and Smith in Ar (noted in the
legend).11,12...................................................................................................................... 130
x
Abstract
J. Chai. Direct kinetic study of methoxy radical reacting with NO2 and O2 and deuterium kinetic isotope
effects, 153 pages, 10 tables, 26 figures, 2013.
Alkoxy radicals (RO•) are important intermediates in the photooxidation of volatile organic
compounds (VOCs) in the troposphere. The competition between their three fates (unimolecular
decomposition, unimolecular isomerization, and reaction with O2) greatly impacts the formation of
hazardous tropospheric ozone (O3) and secondary organic aerosols (SOAs). To date, direct kinetic studies
of RO• + O2 have been limited to alkoxy radicals derived from select C1-C7 alkanes and two halogenated
alkanes. The rate constants of O2 reactions are unknown for alkoxy radicals derived from oxygenated
VOCs or non-alkane hydrocarbons. This makes it difficult to build or validate structure-activity relations
(SARs) for the reactions of alkoxy radicals with O2.
The kinetics and mechanism of the methoxy + O2 reaction is the prototype for other RO• + O2
reactions. I investigated the temperature-dependent kinetics and deuterium isotope effects of the reaction
of methoxy + O2 at pressures up to 700 Torr of N2 and over 250 – 333 K. By combining my absolute rate
measurement for CH3O• + NO2 with our group’s relative rate measurement for CH3O• + O2/NO2, we
obtained kO2 as 1.3+0.9−0.5 × 10
−14
exp[−(663 ± 144)/𝑇] cm3 molecule-1 sec-1, corresponding to 1.4 × 10-15
3 -1 -1
cm molecule sec at 298 K. The rate constant at 298 K is in excellent agreement with previous work,
but the observed temperature dependence is less than previously reported. The deuterium isotope effect,
312±255
kCH3O•+O2 / kCD3O•+O2, can be expressed in Arrhenius form as (1.6+2.1−0.9 ) × exp( T
), which provides
insights into the effect of tunneling on the CH3O + O2 rate constant.
The reaction of methoxy + NO2 is a typical radical-radical recombination reaction in the
atmosphere. Pressure dependent kinetic data for this reaction provides information to investigate
collisional energy transfer. I studied the pressure dependent (30-700 Torr of N2) rate constant for CH3O•
(CD3O•) + NO2 over 250-333 K. The low pressure limiting rate constants are k 0CH3O+NO2 = 4.29+0.40 −0.37 ×
−29 −(1.65±1.11) 0 +1.00 −29 −(4.79±0.92)
10 (T/298) and k CD3O+NO2 = 9.97−0.91 × 10 (T/298) cm molecule-2 s-1
6
∞
respectively. The high pressure limiting rate constants are given by k CH3O+NO2 = (1.95 ± 0.03) ×
10−11 (T/298)−(1.13±0.18) and k ∞ CD3O+NO2 = (1.91 ± 0.02) × 10
−11
(T/298)−(1.11±0.09) cm3 molecule-1
-1
s respectively. The rate constants for the two isotopologues track each other closely as the high pressure
limit is approached.
Key Words: alkoxy radical, ozone, secondary organic aerosol, structure–reactivity relationship, tunneling,
pressure dependence
J. Chai
Candidate for the degree of Philosophy of Science, December 2013
Theodore S. Dibble, Ph.D.
Department of Chemistry
State University of New York College of Environmental Science and Forestry
Syracuse, New York
Theodore S. Dibble, Ph.D. ______________________
xi
Chapter 1. Introduction
Among all the planets in the solar system, only the planet Earth supports life, and does so
primarily by virtue of its unique atmospheric composition. The Earth’s atmosphere consists of
the gases N2 (78%), O2 (20.9%), Ar (0.9%), water vapor and smaller amounts of what are
referred to as trace gases (<0.1%).1 The trace gases CO2 and H2O serve as the materials for
photosynthesis, while O2 is indispensable for respiration. Trace gases like H2O, CO2, CH4, N2O
and O3 make the earth warmer than it would be in their absence because they absorb infrared
radiation emitted by the Earth’s surface and convert that radiation to heat. Ozone in the
stratosphere (formed following photolysis of O2 by short-wavelength ultraviolet light) filters out
solar radiation of wavelengths shorter than 290 nm. This means that stratospheric ozone protects
life on earth from the harmful ultraviolet light that is not absorbed by O2. While playing an
essential role in the Earth’s radiative balance, trace gases also greatly affect the chemical
properties of the atmosphere. The Earth’s atmosphere is a tremendously active photochemical
reactor driven by the solar energy, where photochemistry produces and removes trace gases. In
preindustrial times, natural emissions were the major source of trace gases. Since the Industrial
Revolution, anthropogenic emissions of trace species such as NOx, SOx and halocarbons have
increased drastically.2 Severe air pollution episodes and the Antarctic ozone hole drew increased
attention to the role of trace gases because these phenomena threaten the health of people and/or
ecosystems. Due to the slow air exchange between troposphere and stratosphere, ground level
emissions are mostly relevant to troposphere, except for long-lived gases like methane and
halocarbons.3 With regard to air pollution, the two major concerns are ozone (O3) and secondary
organic aerosols (SOAs). Both result from photo-oxidation of volatile organic compounds
(VOCs) in the presence of NOx.
1
1.1 Tropospheric ozone
The impact of VOCs on air quality was first brought into the world’s view in the 1950s,
due to their contribution along with the anthropogenically emitted NOx (NOx=NO+NO2) to the
formation of Los Angeles’ urban photochemical smog.4 O3 and SOAs are produced from the
photooxidation of VOCs, and are two of the major components of photochemical smog. The
trace gas O3 acts as a pollutant at ground-level that can cause respiratory problems in humans
and deterioration of vegetation, in contrast to the beneficial UV-blocking effect of stratospheric
O3. Compared to the maximum ozone mole fraction of about 10 × 10-6 (mixing ratio 10 ppmv) in
the stratosphere, ground-level ozone concentrations generally ranges from 20 to 60 ppbv, with
values exceeding 100 ppbv in urban areas and broad regions downwind of polluted urban areas.5
As of 2008, EPA’s National Ambient Air Quality Standard (NAAQS) for ozone is an average of
75 ppbv over 8 hours. Ozone is primarily generated from the NO-NO2 cycle, which is facilitated
by HOx cycling and reactions of organic peroxy radicals with NO,5 as demonstrated in Scheme
1.1. In areas with low VOC emissions, such as the remote marine atmosphere, O3 formation is
sustained by OH-oxidation of long-lived carbon monoxide (CO) and methane (CH4), and has low
concentrations due to the low mixing ratio of NOx (0.01-0.1 ppb).5 Ozone formation in the urban
and regional atmosphere is driven by much shorter-lived anthropogenic and biogenic VOCs,
including alkanes, alkenes, aromatics and oxygenated organic compounds. Moreover, because of
biomass burning and biogenic VOC emissions, it has been recognized that the impact of VOC
photooxidation is not just confined to the urban ground level atmosphere, but also extends to
rural regions and the upper troposphere.
2
Scheme 1.1 Photochemical cycle of NO2-NO
Scheme 1.2 VOC oxidation mechanism. (VOC is represented here by RH)
Alkoxy radicals (RO•) are important intermediates during the photooxidation of VOC. As
shown in Scheme 1.2, their formation is primarily initiated by •OH radicals reacting with VOCs
(forming an alkyl radical R•), followed by reaction of the organic radicals with O2 (forming
3
RO2•) and then NO (forming RO•). As illustrated in Scheme 1.1, both HO2• (from RO• + O2)
and RO2• act to convert NO to NO2, a conversion which drives ozone formation. As more NO2 is
produced, the O3 level is increased.6
Once formed, alkoxy radicals mostly undergo one of three competing reaction pathways:
unimolecular decomposition, unimolecular isomerization, or reaction with O2.6 Reaction with O2
leads to the formation of hydroperoxy radicals (HO2•); and decomposition and isomerization of
alkoxy radicals produce a second generation of carbon-centered radicals (R’•), which are
subsequently transformed to second generation alkylperoxy radicals (R’O2•).6 As shown in
Scheme 1.2, the R’O2• resulting from unimolecular decomposition and unimolecular
isomerization can further convert NO to NO2 while producing an alkoxy radical that
subsequently produces HO2. Therefore, when the concentration of NO is not too low, the two
unimolecular reaction pathways of alkoxy radicals lead to more O3 production than the reaction
of alkoxy radicals with O2.7 Competition between the three pathways strongly influences ozone
formation under polluted conditions and the identity of the stable products of VOC oxidation.
The physical and chemical properties of these oxidation products influences their potential for
gas-particle partitioning, and, therefore, the extent of their contribution to secondary organic
aerosol.8,9
1.2 Secondary organic aerosols (SOAs)
The atmospheric oxidation of VOCs and semivolatile organic compounds (SVOCs) can
lead to the formation of secondary organic aerosols (SOAs), which comprise a large fraction
(65%-95%) of atmospheric organic aerosol matter in the northern hemisphere.10,11 Aerosols are
defined as relatively stable suspensions of solid or liquid particles in a gas.12 Among a number of
4
properties of particles, size matters the most due to its impact on human health.12 Particles with
diameters, D, less than 10 µm can penetrate into lungs where they are then deposited in the
respiratory bronchioles or alveolar sacs. The size range of health concern for particulate matter
(PM) has shifted from D < 10 to D < 2.5 µm, as a consequence of our increasing knowledge of
the effect of fine particles on health and our increasing ability to measure these particles.13
Coarse particles (defined a D 1µm) usually originate from seasalt aerosols, pollen and
mechanical processes including human-made and natural dust. “Fine” particles (D < 1 µm),
which are so small that they can only be seen with an electron microscope, arise primarily from
combustion processes, and the oxidation of volatile compounds such as SO2, NOx, and VOCs to
less volatile compounds. Organic compounds account for a substantial proportion (~50% by
mass) of fine particles in the urban-affected northern hemisphere and as high as 90% in the
tropical forested area.14
Gas-phase oxidation reactions can reduce the volatility of organic compounds (by adding
polar functional groups) or increase their volatility (by the cleavage of carbon–carbon bonds).11
Key branching points that control the volatility of the stable products are the site of initial attack
of the oxidant, and subsequent reactions of the alkylperoxy and alkoxy radicals. In this thesis I
will only focus on the chemistry of the alkoxy radical. As was already shown in Scheme 1.2,
there are three major competing pathways for alkoxy reactions. Reaction with O2 introduces a
carbonyl group, and isomerization leads to the addition of both hydroxyl and carbonyl groups to
the original carbon chain; both reactions result in the reduction of volatility in the reaction
products. On the other hand, decomposition breaks down the carbon chain, leading to more
volatile products than the products from competing O2 reaction and isomerization. The more
volatile fragmentation products are less likely to partition to an organic aerosol phase, but some
5
(such as formaldehyde, acetic acid, diacids and hydroxy acids) may have sufficiently high water
solubility to partition into an aqueous phase.10 Reaction products that remain in the gas phase can
undergo further oxidation and/or oligomerization to form second- and higher-generation reaction
products.
1.3 Research objectives
The fates of alkoxy radicals significantly impact ozone formation and also affect gas-
particle partitioning of (aerosol formation and growth from) the eventual stable products of VOC
oxidation.6,15 Therefore, elucidation of the chemistry of alkoxy radicals is essential. To date,
direct kinetic studies of RO• + O2 have been limited to alkoxy radicals derived from select C1-C7
alkanes16,17,18,19,20,21,22,23,24,25,26,27,28 and two halogenated alkanes.29,30,31 With these exceptions, the
rate constants of O2 reactions are unknown for alkoxy radicals derived from non-alkane VOCs.
Many previous studies of alkoxy radical kinetics have only determined the rate constant ratio
kunimolecular/kO2, and rely on estimates of kO2 32 to determine kunimolecular. Previous estimation of kO2
was based on only the experimental determinations for C ≤ 4 alkane derived alkoxy radicals. kO2
for 4 ≤ C ≤ 7 alkane derived alkoxy radicals shows varying temperature dependencies that differ
significantly from those exhibited by smaller alkoxy radicals.32 Consequently, the estimated kO2
cannot represent larger alkane-derived alkoxy radicals, let alone those derived from alkenes or
oxygenated VOCs. The lack of absolute rate constants kRO•+O2 limits the determination of
kunimolecular, thus preventing us from establishing structure-reactivity relations (SARs) for the
unimolecular reactions. SARs are needed to predict the tropospheric fate of larger and
6
functionalized alkoxy radicals, for which experimental data are largely absent and not easily
obtained from experiment.6, 32
The methoxy radical (CH3O•) is the prototype for all alkoxy radicals. Similarly, the
reaction methoxy + O2 (R1.1):
k1
R1.1 CH3O + O2 HCHO + HO2
is the prototype for studying the kinetics and mechanism for other RO• + O2 reactions. However,
the rate constant k1 has never been determined experimentally at temperatures below 296 K.
Also, the previous experimental results over the temperature range of 298–973 K showed
significant non-Arrhenius behavior.33 Theoretical calculations by Bofill et al.34 suggested that
tunneling dominated the rate constant at room temperature. In contrast, Setokuchi and Sato35
computed that tunneling was significantly less important at room temperature than found by
Bofill et al.. Measurement of k1 below room temperature may help determine the extent of
tunneling and validate computational efforts to understand the kinetics of the methoxy + O2
reaction. A reliable theoretical approach for computing rate constants for alkoxy + O2 reactions
would be valuable for calculating kO2 for larger or functionalized alkoxy radical, since absolute
rate measurement does not apply.
It has been realized that the impact of alkoxy radicals is not just confined to the ground
level, but also the upper troposphere.36,37,38 Hence it is necessary to study alkoxy chemistry not
only at conditions common to ozone pollution episodes at the earth’s surface, but also at the
much lower temperatures and pressures relevant to the upper troposphere. In several field studies
of upper tropospheric chemistry, photolysis of various oxygenated organic compounds such as
7
acetone (either generated in situ or transported via convection) have been shown to be a major
source of unexpectedly high concentrations of OH and HO2 radicals. This higher-than-expected
concentration of HOx radicals would drive faster ozone production than expected. This is of
concern, because ozone in this region of the troposphere acts effectively as a greenhouse gas.
Also, these measurements suggested that in the upper troposphere, the production rate of ozone
increases rapidly with the concentration of NO from aircraft emission or biomass burning. For
these reasons, the chemistry of alkoxy radicals is also important in the upper troposphere.36,37,38
Consequently, the first goal of my Ph.D. research was to determine k1 (T) over a
temperature range extended below room temperature, which is presented in Chapter 4. It is
difficult to measure k1 directly below room temperature using the typical method (laser flash
photolysis-laser coupled to induced fluorescence: LFP-LIF). This is because the reaction CH3O•
+ O2 has a low rate constant which decreases as temperature decreases so that the use of high
concentration of O2 (>50 Torr) is required. Unfortunately, O2 efficiently quenches fluorescence
of excited CH3O• (Ã2A1),39 so that the even higher [O2] required at temperatures below 293 K
greatly reduce signal:noise. To overcome this difficulty, I combined relative rate measurements
of the ratio k1/k2 with absolute determination of k2, where k2 is the overall rate constant for
reaction of CH3O• with NO2:
k2
R1.2 CH3O + NO2 (+M) Products
By combining the two measurements, I have been able to determine the absolute value of
k1 at temperatures down to 250 K.
8
In order to examine the extent of tunneling effect for k1, I also investigated the deuterium
kinetic isotope effect (KIE) for the methoxy + O2 reaction by substituting CD3O• for CH3O• over
the temperature range of 250-335 K. The experimental approach was the same as that described
above. The KIE is defined here as kH/kD—the rate constant ratio between the reaction involving
the normal reactant and that involving deuterated reactant. The KIEs for the methoxy + O2 and
methoxy + NO2 reactions are k1/k3 and k2/k4 respectively.
k3
R1.3 CD3O + O2 DCDO + DO2
k4
R1.4 CD3O + NO2 (+M) Products
In this thesis (Chpater 5), I also determined the pressure dependent rate constants under 30
– 700 Torr and in the temperature range of 250-333 K for methoxy + NO2 (R1.2 and R1.4),
which is a prototypical radical-radical recombination reaction. The context of this research will
be discussed in Chapter 2 (section 2.3).
9
References
1 Wayne, R. P. Chemistry of Atmospheres (3rd edition); Oxford University Press Inc., New York, 2000,
2 ff
2 Wayne, R. P. Chemistry of Atmospheres (3rd edition); Oxford University Press Inc., New York, 2000,
pp. 348-350, pp. 65-366.
3 Seinfeld, J. H.; Pandis, S. N. Atmospheric Chemistry and Physics (2nd edition); John Wiley & Sons,
Inc., 2006, 6ff, 1044ff
4 Haagen-Smit, A. Industrial & Engineering Chemistry 1952, 44, 1342-1346.
Inc., 2006, pp. 204-205.
6 Orlando, J. J.; Tyndall, G. S.; Wallington, T. J. The Atmospheric Chemistry of Alkoxy
Radicals. Chem. Rev., 2003, 103, 4657-4690.
7 Jenkin, M. E.; Hayman, G. D. Atmos. Environ. 1999, 33, 1275-1293.
8 Jenkin, M. E.; Clemitshaw, K. C. Atmos. Environ. 2000, 34, 2499-2527.
9 Lim, Y. B.; Ziemann, P. J. Environ. Sci. Technol. 2005, 39, 9229-9236.
10 Ziemann, P. J.; Atkinson, R. Chem. Soc. Rev., 2012, 41, 6582-6605.
11 Kroll, J. H.; Seinfeld, J. H. Atmos. Environ., 2008, 42, 3593-3624.
12 Finlayson-Pitts, B. J.; Pitts, Jr., J. N. Chemistry of the Upper and lower atmosphere: Theory,
Experiment, and Applications; Academic Press, 2000, 349 ff
13 Cooper, C. D.; Alley, F. C. Air Pollution Control: A Design Approach (3rd edition); Waveland Press,
Inc., 2002, 50 ff
14 Kanakidou, M.; Seinfeld, J.; Pandis, S.; Barnes, I.; Dentener, F.; Facchini, M.; Dingenen, R. V.;
Ervens, B.; Nenes, A.; Nielsen, C. Atmospheric Chemistry and Physics 2005, 5, 1053-1123.
15 Jimenez, J.; Canagaratna, M.; Donahue, N.; Prevot, A.; Zhang, Q.; Kroll, J. H.; DeCarlo, P. F.; Allan,
J. D.; Coe, H.; Ng, N. Science 2009, 326, 1525-1529.
10
16 Wantuck, P. J.; Oldenborg, R. C.; Baughcum, S. L.; Winn, K. R. J. Phys. Chem. 1987, 91, 4653-4655.
17 Gutman, D.; Sanders, N.; Butler, J. J. Phys. Chem. 1982, 86, 66-70.
18 Lorentz, K.; Rhasa, D.; Zellner, R.; Fritz, B. Ber. Bunsenges. Phys. Chem., 1985, 89, 341-342.
19 Balla, R. J.; Nelson, H.; McDonald, J. Chem. Phys. 1985, 99, 323-335.
20 Hartmann, D.; Karthäuser, J.; Sawerysyn, J.; Zellner, R. Ber. Bunsenges. Phys. Chem. 1990, 94, 639-
645.
21 Deng, W.; Wang, C.; Katz, D. R.; Gawinski, G. R.; Davis, A. J.; Dibble, T. S. Chem. Phys. Lett. 2000,
330, 541-546.
22 Deng, W.; Davis, A. J.; Zhang, L.; Katz, D. R.; Dibble, T. S. J. Phys. Chem. A 2001, 105, 8985-8990.
23 Fittschen, C.; Frenzel, A.; Imrik, K.; Devolder, P. Int. J. Chem. Kinet. 1999, 31, 860-866.
24 Mund, C.; Fockenberg, C.; Zellner, R. Ber. Bunsenges. Phys. Chem. 1998, 102, 709-715.
25 Hein, H.; Hoffmann, A.; Zellner, R. Phys. Chem. Chem. Phys. 1999, 1, 3743-3752.
26 Hein, H.; Hoffmann, A.; Zellner, R. Ber. Bunsenges. Phys. Chem. 1998, 102, 1840-1849.
27 Zhang, L.; Kitney, K. A.; Ferenac, M. A.; Deng, W.; Dibble, T. S. J. Phys. Chem. A 2004, 108, 447-
454
28 Zhang, L.; Callahan, K. M.; Derbyshire, D.; Dibble, T. S. J. Phys. Chem. A 2005, 109, 9232-9240.
29 Wu, F. and Carr, R. W. J. Phys. Chem. A 2001, 105, 1423-1432.
30 Wu, F. and Carr, R. W. Chem. Phys. Lett.,1999, 305, 44-50.
31 Wu, F. and Carr, R. W. J. Phys. Chem. 1996, 100, 9352-9359.
32 Atkinson, R. Atmos. Environ. 2007, 41, 8468-8485.
34 Bofill, J. M.; Olivella, S.; Solé, A.; Anglada, J. M. J. Am. Chem. Soc. 1999, 121, 1337-1347.
35 Setokuchi, O.; Sato, M. J. Phys. Chem. A 2002, 106, 8124-8132.
36 Wennberg, P.; Hanisco, T.; Jaegle, L.; Jacob, D.; Hintsa, E.; Lanzendorf, E.; Anderson, J.; Gao, R.;
Keim, E.; Donnelly, S. Science 1998, 279, 49-53.
11
37 McKeen, S.; Gierczak, T.; Burkholder, J.; Wennberg, P.; Hanisco, T.; Keim, E.; Gao, R.; Liu, S.;
Ravishankara, A.; Fahey, D. Geophys. Res. Lett. 1997, 24, 3177-3180.
38 Jaeglé, L.; Jacob, D. J.; Brune, W. H.; Wennberg, P. O. Atmos. Environ., 2001, 35, 469-489.
12
Chapter 2. Background
In this Chapter, I will first give a literature review on the atmospheric chemistry of alkoxy
radicals, and present the current challenges encountered in obtaining general rules (structure-
reactivity relationships (SAR)) to predict rate constants for alkoxy + O2 reactions, and why an
SAR for the O2 reaction would be valuable for understanding the unimolecular reactions. Next, I
will discuss the importance of the reaction of methoxy + NO2, and review kinetic studies on it.
The rate constant of this reaction, as is typical of radical-radical recombination reactions, is
pressure dependent; and, therefore, it is necessary to introduce the basis of pressure dependent
rate constants. In addition, I will clarify the basis of the kinetic isotope effect (KIE), and this will
help the reader to understand the meaning of KIE values I determine for both the methoxy + O2
and methoxy + NO2 reactions. Lastly, I will introduce the principles of my experimental
methods: laser flash photolysis (LFP) for generating alkoxy radicals and laser induced
fluorescence (LIF) for sensitive detection of alkoxy radicals.
2.1 Temperature profile of atmosphere
The Earth’s atmosphere is divided into different layers based on the variation of
temperature profile with altitude. Temperature structure below 50 km is shown in Figure 2.1.
The lowest region, called the troposphere, extends from the Earth’s surface up to around 8-17 km
depending on the latitude and time of the year. The temperature in this region declines with
increasing height due to the adiabatic cooling from the absorption of solar radiation. A strong
and rapid vertical mixing characterizes this layer so that the species emitted from the Earth’s
surface become vertically well-mixed throughout the troposphere in no more than a few days.
13
Conversely, transport from tropopause to the next highest layer, the stratosphere, occurs on a
time scale of years. Due to the slow air exchange between troposphere and stratosphere, air
pollution caused by ground level emissions, and therefore alkoxy radical chemistry, are mostly
relevant to troposphere. The exception is the methoxy radical, which is largely produced from
oxidation of long-lived (~10 years) methane.
Figure 2.1 Temperature structure of atmosphere below 50 km from U.S. Standard Atmosphere. 1
2.2 Atmospheric chemistry of alkoxy radicals
As is discussed in Chapter 1, there are three major fates for alkoxy radicals (RO•) in the
atmosphere, and these competing fates greatly influence formation of ozone as well as the gas-
particle-partitioning of eventual products of photooxidation of VOCs. This section starts with a
review of kinetic studies of the three main reactions of RO•: reaction with O2, unimolecular
14
isomerization, and unimolecular decomposition. I will then proceed to touch upon several other
unimolecular reactions that have been demonstrated for RO•.
Both experimental and theoretical methods have been employed to study the kinetics of
alkoxy radical reactions. Experimental methods include both relative rate methods2,3,4,5,6,7 and
absolute rate methods.8,9,10,11,12,13,14,15,16,17,18,19,20,21 In relative rate methods, a radical precursor
(alkylnitrites or alkylperoxides) mixed with reactants in a static cell are photolyzed to directly or
indirectly produce alkoxy radicals. Key stable products and reactants are detected by one or more
methods, such as Fourier transform infrared (FTIR) spectroscopy,22 gas chromatography with
various detectors (including mass spectroscopy),5 or chemical ionization mass spectrometry.23
By determining the fractional product yields one can derive the rate constant ratio between a
target reaction and a reference reaction. The relative rate approach is most commonly used for
measuring rate constant ratios for a pair of bimolecular reactions (usually with O2 and either NO
or NO2) or for a bimolecular reaction versus a unimolecular reaction. In other relative rate
experiments, alkoxy radicals are produced in the oxidation of volatile organic compounds
(VOCs) initiated by OH radical or atomic Cl or F. This indirect approach is usually done with
ample concentrations of O2 and only enough NO to drive conversion of ROO• to RO•, with a
goal of determining kO2/kunimolecular.24 This approach may also yield insight into the relative
importance of competing pairs of unimolecular reactions.
Absolute rate measurements of RO• kinetics most commonly use pulsed laser flash
photolysis (LFP) to generate RO•, combined with laser induced fluorescence (LIF) to monitor
the concentration of alkoxy radicals as a function of time.8,9,10,11,12,13,14,15,16,17,18,19,20,21 When
using this technique for unimolecular reactions, exponential decay rate constant of the LIF signal
15
of a RO• is the rate constant of disappearance of RO•, and is thus the rate constant of the
reaction. For second order elementary reactions, such as RO• + O2 and RO• + NO2, large excess
of the more stable reactant (O2 or NO2) has been applied, and the exponential decay rate of LIF
signal is expected as the pseudo first order decay rate of RO•. Several concentrations of O2 or
NO2 are normally used, and the slope of linear fit of the pseudo first order decay rate versus the
concentration of O2 or NO2, gives the second order rate constant. In both cases, absolute
concentration of RO• is not needed. LFP/LIF is the method I used, and I will discuss it in detail
in section 2.6.
Quantum chemical calculations are carried out to produce inputs to computations of rate
constants. Quantum chemical calculations can predict the structures, vibrational frequencies, and
energies of reactants, products, and transition states. A great advantage of using computational
chemistry to predict rate constants is that it allows one to calculate the rate constants of the same
type of reactions for a whole group of radicals via the same method, thus helping to develop the
structure-reactivity relationships. Also, computational chemistry is useful in providing data at
temperatures that are challenging to reach in experiments, or for radicals that are difficult to
selectively generate or monitor in experiments. 25,26,27
Kinetic results from all three methods can validate each other. A long-term goal of kinetic
studies of alkoxy radicals is to build up general rules to predict rate constants for reactions of
alkoxy radicals. These rules would enable modelers to quickly generate approximate rate
constants for use in kinetic modeling of VOC oxidation processes leading to ozone and
secondary organic aerosol.
16
2.2.1 RO• + O2
The reaction of RO• + O2 leads to the formation of a carbonyl compound and the HO2
radical. This is evidenced by early relative rate study on methoxy radical from Cox et al.2 Direct
product study for the methoxy + O2 reaction showed the product of formaldehyde with a yield of
85±15% at 298 K,28 and direct measurement for ethoxy + O2 obtained a 89+22
−12 % yield of HO2.
29
Three absolute measurements of 𝑘CH3O• + O2 over 298 ≤ T ≤ 973 K were conducted in 1980s
using the LFP-LIF technique, with the CH3O• produced by photolysis of either CH3ONO or
CH3OH.8,9,10,11 The results of these studies are in general agreement in their range of overlap
(298 ≤ T ≤ 610 K), and were fitted by an Arrhenius expression:24
1150±190
E2.1 𝑘CH3 O• + O2 = 7.82+4.68
−2.93 × 10
−14
exp[− ]
𝑇
At 298 K this expression obtains 𝑘CH3O• + O2 = 1.6 × 10-15 cm3 molecule-1 s-1. The rate
constants obtained at T > 610 K greatly exceed that obtained by an extrapolation E2.1.11 Notably,
no direct measurement of 𝑘CH3O• + O2 has been done below 298 K. This is because the rate
constant for the reaction CH3O• + O2 is small and become smaller as the temperature decreases
so that the use of high concentration of O2 (~50 Torr at room temperature) is required.
Unfortunately, O2 efficiently quenches fluorescence of electronically excited CH3O•,11 and the
high [O2] that would be needed at temperatures lower than ~293 K may reduce the fluorescence
signal below the level needed to achieve a good signal:noise.
Relative rates studies were carried out to determine the rate constant for CH3O• + O2 in the
temperature range 296 ≤ T ≤ 450 K by various groups in the 1970s.2,3,4,5,6,7 Rate constant ratios
17
of the reaction CH3O• + O2 to the reference reactions CH3O• + NO or NO2 were derived through
product analysis and kinetic modeling from a proposed mechanism. At the time these relative
studies were carried out, there was limited understanding of the mechanism and kinetics of the
reference reactions. Orlando et al.24 re-analyzed representative data using the updated
mechanisms and kinetics of the reference reactions as recommended by the IUPAC Gas Kinetic
Subcommittee.30 The rate constants ratios obtained in these experiments exhibited a lot of
scatter, and only the result of Cox et al.2 at room temperature agreed with absolute rate studies to
within 10%.
The three direct studies of ethoxy (C2H5O•) + O2 in the temperature range of 280-420 K
showed agreement on the small temperature dependence (activation energy ~1 kcal/mol) for the
reaction.9,12 However, the rate constants reported by Hartmann et al.29 are about 30-50% higher
than those from the other two studies.9,12 The single relative rate study31 investigated kO2 versus
kNO over a larger temperature range that extended down to 225 K. The scaled kO2 (according to
the updated kNO) from this relative study agreed well with those of Hartmann et al.29 at room
temperature and above, but exhibit slightly less temperature dependence below room temperature
than above room temperature.
Absolute rate studies on RO• + O2 were also carried out for 1-propoxy,12,13
isopropoxy,13,14,15 1-butoxy,16 2-butoxy,15, 1-pentoxy,16,17,18,19 3-pentoxy,17,32 cyclohexoxy,20
trans-4-methylcyclohexoxy.21 Results of these studies are shown in Table 2.1. The studies of
Hein et al.16,18,32 for 1-butoxy, 2-butoxy, 1-pentoxy and 3-pentoxy at 298 K are somewhat
indirect in that rate constants were obtained from fitting the profiles of OH radical and NO2
18
concentration versus time observed after pulsed laser formation of the parent alkyl radical. This
fitting requires assumptions about the reaction mechanism and kinetics.
Table 2.1 Absolute rate data in the form of Arrhenius expression for non-substituted alkoxy radicals from
C2 to C7. Both k and A are in units of cm3 molecule-1 s-1. Cited errors are 2σ.
Alkoxy radical A(×1014) Ea/R (K) T range (K) k×1015 (298 K) reference
Ethoxy 2.86 378 296-353 8.0 9

7.1 552±64 295-411 11.1 29
2.4 325±120 286-390 8.1 12
5.13 550 280-420 8.1 24 a
1-propoxy 1.4 108±60 223-303 9.1 13
2.5 241±60 289-381 11.1 12
2.6 253±60 223-381 11.1 12+13 b
2-propoxy 1.51 196±141 294-384 7.8 14
1.4 217±48 218-311 6.8 13
1.6 265±24 288-364 6.6 12
1.9 310±24 218-364 6.9 12+13 b
1-Butoxy 293±3 14±7 16
2-Butoxy 293 6.5±2 18

0.133 -(659±83) 223-311 12±4 15
0.12 -(553±192) 221-266 7.7 c 17
291-295 9±2 19
1-pentoxy 293±3 ≤100 16
3-pentoxy 293±3 7.2±3.5 32
0.41 -(319±76) 220-285 12±6 17
cyclohexoxy 580 1720±96 225-302 18 20
c
cyclohexoxy-d11 3.7 760±400 228-267 2.9 20
trans-4-methyl-
14 810±400 228-292 9.2 21
cyclohexoxy
a
Recommendation by Orlando et al.24
b
Linear fitting of combined results from references 12 and 13
c
Extrapolating the corresponding Arrhenius expression to room temperature
19
For the reactions of ethoxy and 1- and 2-propoxy radicals with O2, both A-factors and
activation energies are similar, especially those for the 1- and 2-propoxy radical reactions. Based
on the temperature-dependent data for ethoxy and 1- and 2-propoxy and the room temperature
rate constants for 1- and 2-butoxy and 3-propoxy, Atkinson33 recommended
E2.2 𝑘𝑂2 = 2.5 × 10−14 𝑒 −300/𝑇 cm3 molecule-1 s-1
for both primary (RCH2O•) and secondary (RCH(O•)R’) alkoxy radicals. This equation leads to
𝑘𝑂2 = 9 × 10−15 cm3 molecule-1 s-1 at 298 K. Aschmann and Atkinson34 made an estimate of the
rate constant of 1.4 × 10−14 cm3 molecule-1 s-1 (independent of temperature) for reaction of
ether-derived alkoxy radical ROC(O•)< (e.g., CH3OCH(O•)CH3) with O2, based on this reaction
being much more exothermic than the reactions of primary and secondary RO• with O2.
The results of Deng et al.17 for reaction of both 2-butoxy and 3-pentoxy with O2 exhibited
small negative temperature dependencies, in contrast with the positive temperature dependencies
of the reactions of smaller alkoxy radicals with O2. The reason for the surprising difference is
unknown. It is possible that either Deng et al.’s measurement had some systematic error, or the
structure of larger alkoxy radicals have some effect on the reactivity. Zhang et al. investigated
the absolute rate constants of cyclohexoxy + O2 and trans-4-methyl-cyclohexoxy + O2 in the
temperature range of ~220-300 K.20,21 The room temperature rate constant of trans-4-methyl-
cyclohexoxy + O2 is consistent with recommended value for smaller alkoxy radical according to
equation E2.2, but the observed activation energy is ~3 times larger than the recommended
value.33 For cyclohexoxy + O2, compared to the recommended values, the room temperature rate
constant is greater by a factor of ~10, and the pre-exponential factor is larger by two orders of
20
magnitude, and activation energy of 3.4 kcal/mol is also six times larger. The compelling
difference between this cyclic alkoxy radical and other acyclic alkoxy radicals might be
attributed to ring strain in the transition state.20 However, the enormously difference in A-factor
and activation energy between the cyclohexoxy and trans-4-methyl-cyclohexoxy radical
reactions with O2 tends to contradict the proposition that ring strain is affecting the rate constant.
20,21
The two values of kO2 obtained by Zhang et al. are the only absolute rate constant
determinations for six-member ring RO• reaction with O2, but these results do not provide
reliable guidance on the effect of a six-member ring on the rate constant for alkoxy + O2
reactions.
The only two substituted alkoxy radicals for which data are available are CH2ClO•35,36 and
CFCl2CH2O•.37 Wu and Carr used flash photolysis coupled with time-resolved mass
spectrometry to measure rate constants for these radicals reacting with O2 in the temperature
range of 265-306 and 251-341 K, respectively. Both reactions show a positive temperature
dependence, with the Arrhenius expression 𝑘𝐶𝐻2 𝐶𝑙𝑂+𝑂2 = 2.06+2.76

−1.18 × 10
−12 −(940±240)/𝑇
𝑒 and
𝑘𝐶𝐹𝐶𝑙2 𝐶𝐻2 𝑂+𝑂2 = 2.53+1.81

−1.06 × 10
−15 −(960±160)/𝑇
𝑒 , which give rate constants of 8.6 × 10−14 and
9.9 × 10−17 cm3 molecule-1 s-1 at room temperature, respectively. The rate constant for
CFCl2CH2O• is much lower than that for other alkoxy radicals, even though the halogen
substitutions are on the β-C rather than on the α-C atom (the one bonded to the radical center).
So the effects of halogenation on kO2 are not understood.
In summary, the limited and scattered contradictory data from absolute rate measurements
of kO2 provide very little insight into the effects of structure on kO2. The effect of halogenation on
kO2 is unclear, as is the difference in kO2 between the cyclohexoxy and trans-4-methyl-
21
cyclohexoxy radical. In the atmosphere, there is a much wider range of alkoxy radicals,
including those derived from alkenes (like isoprene) or oxygenated VOCs (OVOCs). The large
emissions of these VOCs and OVOCs make them potentially very important to the formation of
tropospheric ozone;38 their large emissions make them important for formation of secondary
organic aerosol (SOA).39,40 Absolute rate studies for functionalized alkoxy radicals from these
VOCs are quite challenging due to experimental constraints. These constraints include (a) the
RO• concentration is not measurable if it is produced from the parent VOC in the lab; such
reaction system has a number of reactants so that the alkoxy radicals would easily react with
them. (b) it is very difficult to synthesize the alkyl nitrite (RONO) precursors to the RO•
produced from alkenes or oxygenated VOCs; for instance, RONO with a carbonyl group is
soluble in water, and therefore it is quite difficult to purify the product. (c) even if we synthesize
the RONO, photolysis may not generate the RO•, or the RO• may not fluoresce.41,42 Product
analysis is limited in what can be measured. For instance, the problem of formation of multiple
isomers of RO (in unknown yields, together with the fact that two different isomeric ROs may
lead to similar (or the same) product, makes it hard to determine relative rates.
As mentioned in Chapter 1, the lack of absolute rate constants kRO•+O2 limits the
determination of kunimolecular. This prevents us from establishing structure-reactivity relations
(SARs) for the unimolecular reactions, which in turn limit our understanding of the tropospheric
chemistry.
Due to the experimental limitations, it would be very beneficial if theory could help derive
kO2 for a wide range of alkoxy radicals. Several theoretical studies have tried to elucidate the
mechanism of CH3O• + O2. Jungkamp and Seinfeld proposed the reaction occurs via formation
22
of a short-lived trioxy radical intermediate followed by HO2 elimination; this mechanism was
consistent with the unusually low Arrhenius pre-exponential factor (A-factor) for the reaction.25
However, Bofill et al.26 found an error in their analysis, and argued against this hypothesis due to
the enormous barrier (50 kcal/mol) they found for HO2 elimination from the trioxy radical.
Instead, Bofill et al.26 proposed that a direct H-atom abstraction occurs through a five-member
ring-like transition structure with an intramolecular non-covalent O…O bonding. This non-
covalent interaction lowers the energy of the transition state by ~8 kcal/mole as compared to an
acyclic transition state for the same reaction. Based on this mechanism, Bofill et al.26 and
Setokuchi et al.27 calculated the rate constant 𝑘CH3O• + O2 , and included tunneling effects using the
asymmetric Eckart potential, a one-dimensional tunneling approximation. This led to a tunneling
coefficient, , of 9 at 298 K, yielding a rate constant 𝑘CH3O• + O2 of 2.7 × 10-15 cm3 molecule-1 s-1,
which is comparable with the experimental value (1.6 × 10-15 cm3 molecule-1 s-1).24 In contrast,
Setokuchi et al.27 applied multidimensional tunneling methods to this problem. They found less
significant tunneling (≈2) than Bofill et al.26 at room temperature, although their calculation
showed that tunneling became more important as the temperature decreased (≈8 at 200 K). In
addition, their calculated 𝑘CH3O• + O2 was 9.8 × 10-16 cm3 molecule-1 s-1 at room temperature,
which is consistent with the experimental value (1.6 × 10-15 cm3 molecule-1 s-1).24 Note that Bofill
et al.26 and Setokuchi et al.27 calculated different barrier heights, thus obtaining different classical
rate constants.
Although both calculated values of 𝑘CH3O• + O2 at 298 K are similar, the computed extent of
tunneling differs significantly. Recently my group members carried out the quantum calculations
on methoxy+O2,43 which confirmed Bofill et al.’s reaction mechanism,26 and obtained analogous
23
tunneling effects to results of Setokuchi et al.27 The experimental determination of 𝑘CH3O• + O2 and
𝑘CD3 O• + O2 at our temperature range may be valuable for validating the calculation method. A
validated computational method for RO• + O2 would be valuable for calculating rate constants
for the O2 reactions of larger and functionalized alkoxy radicals derived from atmosphere VOCs.
2.2.2 Unimolecular isomerization
2.2.2.1 1,5-H shift isomerization
Isomerization reactions of alkoxy radicals most often occur as an intramolecular 1,5-H
shift via a six-member ring transition state. The simplest system for 1,5-H shift isomerization is
1-butoxy radical, as in reaction R2.1:
R2.1 CH3CH2CH2CH2O• → •CH2CH2CH2CH2OH
The preference of six-member ring can be rationalized in terms of both strain energy and
entropy. Baldwin et al.44 estimated that the strain energy of a five member ring (1,4-H shift) is ~5
times that of six member ring transition state. The entropy difference between reactant and
transition state for a 1,5-H shift is 8 cal/(mol K) smaller than for a 1,6-H shift.45 On account of
these factors, the 1,5 H-shift has a higher rate constant than the 1,4 and 1,6 H-shift reactions.
Experimental evidence comes from the work of Hornumg et al.,46 who used isotopic labeling to
show that isomerization of 1-pentoxy occurred exclusively from the 4-position.
The hydroxyl-alkyl radical formed from R2.1 is expected to react with O2 to form a peroxy
radical, which can then react with NO to form a hydroxyl-alkoxy radical, as shown in Scheme
24
2.1. This alkoxy radical could react with O2 or decompose, but is more likely to undergo a
second isomerization. This second isomerization would ultimately form a hydroxycarbonyl
compound,22 as is also shown in Scheme 2.1. Atkinson et al. identified hydroxycarbonyls formed
as isomerization products in the oxidation of alkanes in the series from n-butane through n-
octane.47,48 These studies showed that the yields of unsubstituted carbonyls, which are products
from reactions of alkoxy radicals with O2, decreased gradually from butane to octane. By
contrast, yields of hydroxycarbonyls (products of alkoxy radical isomerization) increased
substantially with the length of the carbon chain. Similarly, for β-hydroxyalkoxy radicals formed
in the oxidation of 1-alkenes, the yields of both formaldehyde (product of decomposition) and
the complementary carbonyl (product of reaction with O2) decreased with the length of the
alkene chain.49,50,51 This indicates the isomerization of β-hydroxyalkoxy radicals can compete
with other channels, despite the fact that the β-hydroxy group appears to reduce the rate constant
for alkoxy radical isomerization.52,53
O2 NO 1,5-H shift
CH2CH2CH2CH2OH OCH2CH2CH2CH2OH HOCH2CH2CH2CHOH
O2
HOCH2CH2CH2CHO + HO2
Scheme 2.1 Reactions following R2.1, including isomerization (1,5 H-shift) of the δ-hydroxy alkoxy
radical
Absolute experimental rate constants for isomerization have only been obtained for 1-
butoxy, 1-pentoxy and 2-pentoxy at 293±3 K and 37.5 Torr pressure by Hein and coworkers.16,18
Their approach employed time-resolved and simultaneous measurement of [NO2] and [OH]
following a laser pulse that initiated the oxidation. Rate constants for isomerization and reaction
25
with O2 were extracted from numerical kinetic simulations of the kinetics. However, the
theoretical work on the isomerization of 1 and 2-pentoxy by the same research group suggest that
isomerization rate constants for these species are far from the high-pressure limit at the low
pressure of these experiments.54 Therefore the rate constants obtained by Hein and coworkers are
in the fall-off region, and not directly comparable with relative rate constants obtained at the
atmospheric pressure of air.
Isomerization rate constants for 1-butoxy,22,55,56,57 1-pentoxy,58 2-pentoxy,58,59 and 5-
methyl-2-hexoxy58 radicals were measured at atmospheric pressure (i.e. ~1 atm) relative to the
corresponding O2 reaction rate constant, in the range of 243-319 K.58 The calculations of
Somnitz and Zellner54 indicated that at 298 K and 760 Torr of air the isomerization rate constants
for 1-butoxy and 1- and 2-pentoxy radicals are within 20% of the high-pressure limit. Atkinson33
transformed the above-mentioned relative rate constants results to absolute rate constants by
applying the recommended rate constant for O2 reactions, 𝑘𝑂2 = 2.5 × 10−14 𝑒 −300/𝑇 cm3
molecule-1 s-1.33 The results indicated that the isomerization rate constant increased as the site of
hydrogen abstraction changed from the primary site (CH3 group, kprim), to a secondary site (CH2
group, ksec) and to a tertiary site (CH group, ktert). Based on these values, Atkinson33 derived an
SAR. The A-factor in this SAR accounted for the number of abstractable hydrogen atoms in the
alkoxy radical:
E2.3 𝑘𝑝𝑟𝑖𝑚 = 1.2 × 1011 𝑒 −3905/𝑇 s-1
E2.4 𝑘𝑠𝑒𝑐 = 8.0 × 1010 𝑒 −3090/𝑇 s-1
E2.5 𝑘𝑡𝑒𝑟𝑡 = 4.0 × 1010 𝑒 −2520/𝑇 s-1
26
Different neighboring groups to the site of H-abstraction are expected to have different
effects on the isomerization rate constant. Atkinson33 treated this problem using the same method
used to estimate rate constants for hydrogen abstraction by OH radical from organic
compounds.60 In this approach, one assigns a coefficient to each neighboring group (substituent),
and multiplies the isomerization rate constant (E2.3 - E2.5) by the corresponding coefficient for
the neighboring substituent. Atkinson estimated the value of the coefficient to be 1.3 for
neighboring groups including -CH2-, >CH-, and >C<, unlike in OH reactions where there is
sufficient data to distinguish the effects of these different groups. He adopted coefficients from
OH reactions for other groups, including –OH, -CH2OH, >CHOH, and -OR. In addition, for an
isomerization occurring across an ether linkage, Aschmann and Atkinson estimated that the rate
constant of isomerization decreased by a factor of ~30 at 298 K due to additional ring strain,34
while the strain energy due to a >C=O group was estimated to be negligible.61
The estimations described above are based on only a limited number of relative rate
experimental results, and the accuracy of absolute isomerization rate is confined using estimated
kO2 discussed previously. Also, the coefficients that quantify the effect of neighboring groups are
mostly values adopted from analogous models for hydrogen atom abstraction by the OH radical.
For these reasons, the SAR described above is very rough, and its predictions may be seriously in
error. Therefore, more experimental and theoretical work are needed to establish the rate
constant of isomerization for larger alkane-derived alkoxy radicals, especially functionalized
alkoxy radicals. A reliable computational method for kRO+O2 validated by our experimental
kmethoxy+O2 results would enable one to obtain more accurate kiso from available relative rate data,
as well as kiso for functionalized alkoxy radicals. This would, in turn, validate computational
27
methods for computing isomerization rate constants, thus helping us to build SAR for
isomerization reactions.
A number of theoretical studies based on quantum chemical calculations and Rice-
Ramsperger-Kassel-Marcus (RRKM) theory have been carried out for isomerization reaction of
1-butoxy, 1-pentoxy and 2-pentoxy radical.52,54,62,63,64 Ferenac et al.63 also calculated the barrier
for isomerizations (1,5 H-shift) of several primary oxygenated alkoxy radicals, and found that
the barrier was more dependent on the placement of the functional group than on the nature of
that group, except for the ketone-derived alkoxy radicals, for which the placement of the
carbonyl group made little difference. Recently Zheng et al. used a more reliable approach to
calculate a barrier of ~12 kcal/mol for 1,5 H-shift isomerization of 1-butoxy.65 Their value is
50% larger than estimated by Atkinson from the relative rate study and his estimate of kO2.33 The
influence of tunneling on the rate constants for 1-butoxy isomerization have been investigated by
several studies using various tunneling treatments.65, 66, 67,68 Among these studies, the most recent
work from Xu et al.68 provided a tunneling coefficient of 11 at 300 K using a small curvature
tunneling (SCT) treatment, which they believed to be the most reliable method for tunneling.
This tunneling coefficient is only one-fourth that from the same group’s previous results65 using
the SCT. However, as is suggested by the authors, the latter value (11) is expected to be more
reliable than the former one.68
Isoprene (CH2=CH-C(CH3)=CH2) is an important biogenic VOC, because its
photooxidation greatly impacts ozone and SOA formation in the troposphere.69 The possibility
for the isomerization of alkoxy radicals from isoprene oxidation has been investigated by several
investigations.69,70,71 OH addition in either of the outer carbons followed by O2 and NO attack on
28
the δ-carbon leads to the formation δ-hydroxy alkoxy radicals, i.e. HOCH2C(CH3)=CHCH2O•
and HOCH2CH=C(CH3)CH2O•. These two radicals are likely to undergo 1,5-H shift
isomerization. According to Dibble’s calculation, the isomerization rate constants for both
radicals are over three orders of magnitude faster than the assumed rate constant for reaction
with O2, and about six orders of magnitude faster than the rate constant of decomposition.70
2.2.2.2 Other types of H-shift isomerization reactions
In the oxidation of isoprene, dihydroxy-alkoxy radicals are produced from the 1,5-H shift
isomerization of δ-oxy radicals previously mentioned and their subsequent oxidation. The
dihydroxy-alkoxy radicals have two hydrogen bonds that are donated in series: an enol group
donates a hydrogen bond to a -CH2OH group, which donates in turn to the oxygen radical center.
Dibble72,73 proposed a double H-atom transfer mechanism for this type of alkoxy radical as is
depicted in Scheme 2.2, and calculated a barrier of ~5 kcal/mol.
OH HO O HO O HO
O OH OH
OH HO O HO O HO
O OH OH
Scheme 2.2 Double H-atom transfer.72
29
As for other types of isomerization, 1,6-H-shift isomerization via seven-member ring
transition state is not favored on entropic grounds.45 However, in the study of oxidation of rigid
framework α-pinene, Orlando et al.74 noted a possibility of 1,6-H-shift isomerization, and Peeters
et al. noted a possibility of 1,7- H-shift in α-pinene chemistry.75
2.2.3 Unimolecular decomposition
2.2.3.1 Decomposition of unsubstituted alkoxy radical
The decomposition of alkoxy radicals usually takes place via the mechanism of β C-C
bond cleavage, producing a carbonyl compound and an alkyl radical. The decomposition of 2-
butoxy radical is shown below as an example:
R2.2 CH3CH(O•)CH2CH3 → CH3CHO + •CH2CH3
Much of the data concerning decomposition of alkoxy radicals comes from absolute
measurement on a handful of unsubstituted alkoxy radicals, namely, ethoxy76, 2-propoxy14,77, 2-
butoxy,18,19 tert-butoxy,78,79 3-pentoxy32 and cyclohexoxy80 radicals. All these studies except the
one for the cyclohexoxy radical demonstrate that the decomposition rate constant is pressure
dependent, and at the temperatures studied the decomposition reactions of all these radicals
(except for cyclohexoxy) are in the fall-off region at atmospheric pressure and below. The results
from these studies are listed in Table 2.2. In the studies of Hein et al.,18,32 rate constants were
obtained at 37.5 Torr total pressure and 293 K, and theoretical calculations indicated that the
measured rate constants were lower than those at 760 Torr of air by a factor of ~5.54 The study
for cyclohexoxy radical decomposition did not observe pressure dependence of the rate constant
30
over the range from 5-55 bar, indicating the high pressure limit had been reached, which is
confirmed by the calculation results from Zhang et al.20
Table 2.2 Absolute rate constant Arrhenius expression for decomposition at high pressure limit.
Radical k (298 K)a Aa Ea/R (K) T range (K) reference
Ethoxy 5 1.1×1013 8456 391-471 76

b
2-Propoxy 8.2×102 1.2×1014 7662 330-408 77
2-Butoxy 4.4×104 1.1×1014 6447 291-348 19

c
2.9×104 6.7×1012 5737 291-348 19
t-butoxy 1.4×103 1.4×1013 6856 303-393 78
2.5×103 1.0×1014 7277 323-383 79
cyclohexoxy 6.3×104 3.8×1013 6026 293-341 80
a
Units for k and A are s-1.
b
At 37.5 Torr, kdecomp for 2-butoxy and 3-pentoxy are 3.5×103 and 5.0×103 s-1 respectively. 18,32
c
Measured over 7.5 - 600 Torr of He.
The other major source of information concerning decomposition rate constants of alkoxy
radicals is from the rate constant ratio kdecomp/kO2 studies, which are available for wide range of
alkoxy radicals including 2-butoxy,55,56,81,82 3-pentoxy,59,82 cyclohexoxy,83 2,2-dimethyl-1-
propoxy,84 2-methyl-2-butoxy,85 2-methyl-3-pentoxy,85 and 1-tertbutoxy-1-ethoxy radical.34
Most of these studies were only carried out at room temperature and atmospheric pressure.
Values of kdecomp can be obtained by combining the rate constant ratios and the estimated values
of kO2 according to E2.2.
31
2.2.3.2 Structure activity relationship for decomposition of alkoxy radicals
The SAR for decomposition rate constants of alkoxy radicals has been made at their high
pressure limit. Based on the absolute rate constant studies (Table 2.2) and theoretical
calculations,19,54,91,86 Atkinson estimated an A-factor of 5 × 1013 s-1 per reaction path degeneracy,
and an Arrhenius expression 𝑘𝑑𝑒𝑐𝑜𝑚𝑝 = 𝐴𝑛𝑒 −𝐸𝑑 /(𝑅𝑇) (n is the reaction path degeneracy, and Ed
is the activation energy).33 Using this equation, Atkinson33 derived values for the activation
energy at the high pressure limit for the decomposition of selected alkoxy radicals for which
experimental rate constants (either absolute or relative) were available.
Atkinson87 proposed a relationship between activation energy (Ed) and the heat of reaction
ΔH for the decomposition reaction in the form
E2.6 Ed = a + b (ΔHr)
where ΔHr is the reaction enthalpy. Choo and Benson88 presented data indicating that the
parameter a in E2.6 depended on nature of the alkyl radical leaving-group. The parameter a
decreased monotonically along the series •CH3, •C2H5, •CH(CH3)2, and (CH3)3C• and roughly in
proportion to the ionization potential (IP) of the alkyl radical product. Atkinson33 obtained values
of ΔHr for use in E2.6 using the limited and uncertain thermodynamic data then available.30, 89, 90
Based on the experimentally derived Ed, Atkinson carried out least-squares analysis of Ed versus
ΔH for all cases with •CH3 as the leaving group; this yielded an Ed = 12.7 + 0.4(ΔHr), with Ed
and ΔHr in the unit of kcal/mol. By assuming b in E2.6 was constant, values of a for each leaving
radical (methyl, ethyl, 1-propyl, 2-propyl and tert-butyl) corresponding to different
decomposition reactions were calculated. Since a is proportional to the ionization potential (IP)
32
energy, a least-squares analysis of these values of a against the corresponding values of IP led to
a (in kcal/mol) of 1.81(IP) +/- 4.92, where IP is in eV.33 Using this equation, values of a were
estimated for a series of oxygenated leaving groups, such as •CH2OH, CH3O• and CH3C•(=O).
The limitation of this approach lies in the large uncertainties in the estimation of ΔHr.33 An
uncertainty of 1 kcal/mol in Ed will lead to roughly a factor of 5 in the predicted rate constant
near room temperature. Furthermore, this estimation is based on the assumption that b is
constant. Without validation of this assumption for larger leaving groups other than the methyl
group using adequate experimental data, the predictions of Atkinson may be seriously in error.
Lastly, this estimation method does not perform well in predicting the decomposition activation
barrier for oxygenated alkoxy radicals.33
Another SAR for the activation energy of decomposition of alkoxy radicals has been
reported by Peeters et al.,91 and it is based on available activation barriers derived from theory
and validated barriers derived from experimental data. In this approach, decomposition of the
ethoxy radical is taken as the reference reaction, with an adopted activation barrier of 17.5
kcal/mol. A set of substituent factors are derived for α- and β-alkyl-, -OH, and carbonyl-
substituted alkoxy radicals, and expressed by E2.7:
E2. 7
𝐸𝑑 = 17.5 − 2.1 × 𝑁𝛼 (𝑎𝑙𝑘) − 3.1 × 𝑁𝛽 (𝑎𝑙𝑘) − 8.0 × 𝑁𝛼,𝛽 (𝑂𝐻) − 8.0 × 𝑁𝛽 (𝑂 =) −
12 × 𝑁𝛼 (𝑂 =)
where Ed is in units of kcal/mol, Nα(alk) and Nβ(alk) are the numbers of alkyl substituents on the
α- and the β-carbon respectively, Nα,β(OH) the total number of OH substituents on the α- and β-
33
carbons together, and Nβ(O=) and Nα(O=) are the numbers of oxo functional groups on the β-
and the α-carbon, respectively. The SAR reproduces the available experimental and theoretical
activation barriers/energies within 0.5 to 1 kcal/mol. An advantage of this SAR over Atkinson’s
SAR is that this one eliminates the need for use of ΔH, which may not be known or, if known,
reliable. Compared to Atkinson’s SAR, this SAR does better job in predicting the decomposition
rate constants of oxygenated alkoxy radicals.91 One limitation of this SAR is that it is only valid
for activation energy above 7 kcal/mol.91 Another limitation in this SAR approach lies in the
increasing or decreasing ring strain upon decomposition, or by resonance stabilization of the
leaving radical, which are not accounted for in this approach.91
Due to the uncertainties of above two estimation methods, as well as the limited data
available for decomposition reactions, predictions by current SARs can be problematic. There is
a great need for experimentally and/or theoretically-determined rate constants for wider range of
alkoxy radicals, especially those with functional groups.
2.2.3.3 Decomposition of functionalized alkoxy radical
β-hydroxyalkoxy radicals (e.g., HOCH2CH2O•) are generated in the OH-initiated oxidation
of alkenes, and these radicals represent a significant portion of the alkoxy radical pool in the
troposphere. Alkenes are largely emitted from natural sources such as vegetation and soils and
from anthropogenic sources such as fossil fuel burning. Reaction of the β-hydroxyalkoxy radical
with O2 leads to the formation of β-hydroxy-carbonyl (e.g., HOCH2CHO), while the
decomposition of the β-hydroxyalkoxy radical leads to formation of two carbonyl compounds
ultimately. The decomposition products are more volatile than the products from O2 reactions
34
and unimolecular isomerization, thus contributing less to SOA formation. Under atmospheric
conditions, a number of product studies on ethene oxidation have found that decomposition of
HOCH2CH2O• competes with the O2 reaction,92,93,94,95,96 while several studies23,97,98 have
indicated that unimolecular decomposition is the dominant atmospheric fate of β-hydroxyalkoxy
radicals derived from C > 2 alkenes. Product studies69,71 and theoretical studies71,72,73 on isoprene
oxidation also suggest that decomposition is the dominant pathway for β-hydroxyalkoxy radicals
formation. Among these studies, only Orlando et al.92 measured the rate constant ratio between
decomposition and reaction with O2 for β-hydroxy-alkoxy radical. By using the estimated kO2
they computed a barrier to decomposition of 10-11 kcal/mol. The estimated kO2 results in
significant uncertainty in the determination of the decomposition rate constant for β-
hydroxyalkoxy radicals. A reliable kO2 for β-hydroxyalkoxy radicals is needed in order to build
an SAR for decomposition reactions of this type of radical.
The major pathway for the formation of alkoxy radicals is the RO2• + NO reaction. This
reaction is exothermic by ~10-11 kcal/mol,30 so it is likely that these alkoxy radicals are formed
with significant amount of internal energy, and are chemically activated. Chemically activated
radicals can decompose promptly prior to being quenched to a thermal distribution of energy.
The fraction of radicals undergoing prompt decomposition is nearly independent of temperature,
but depends on the barrier height for decomposition and the size of the radical. Theoretical
studies have reported that 38% of the HOCH2CH2O• radicals formed in the HOCH2CH2OO• +
NO reaction will decompose promptly;93 while the proportion of chemical activation rises to
80% for HOCH2CH(O•)CH3 formed from OH-initiated oxidation of propene in the presence of
NO due to the lower decomposition barrier.97 Chemical activation appears to be important for
35
alkoxy radicals with a barrier to decomposition of ≤ 9 kcal/mol,83,99 whereas it is unimportant for
alkoxy radicals with higher barriers to decomposition, e.g. cyclohexoxy radical with Ed=11.5
kcal/mol,83 and 2-butoxy with Ed=11.3 kcal/mol.81
Ethers have been widely used as solvents and additives to diesel fuels and gasoline. The
chemistry of alkoxy radicals formed from the photooxidation of ether also affect the formation of
ozone.100 The first type of radical of this class is the alkoxy methoxy radical (R-O-C(O•)H2),
whose dominant fate is through reaction with O2,101,102,103,104 with the exception of long chain
radicals (>C4) that can undergo isomerization.100 Reaction of alkoxy methoxy radicals with O2
produce formate esters, which can be further oxidized. The second type of radical of this class is
R2CHOC(O•)HR. Decomposition for this type of radical is dominanted by β C-C bond
cleavage,105 leading to the formation of an alkyl radical (R) and aldehyde. A third type of radical
of this class is the β-oxy-substituted alkoxy radical (e.g., CH3OCH2CH2O•).101,102 The β-alkoxy-
substituents are good leaving groups, so C-O bond-cleavage decomposition is the dominant fate
of these radicals.34 This is consistent with a relatively low barrier of ~6-9 kcal/mol for
decomposition of CH3OCH2CH2O• determined by theoretical calculation.63 There are very few
experimental or theoretical studies on the rate constants of the reactions of ether-derived alkoxy
radicals. A reliable structure-activity relationship for this class of radicals will be needed to
predict the potential for ethers to form ozone and SOA in polluted air.
β-oxo-substituted alkoxy radicals (RC(=O)CH(O•)R’) are formed from the atmospheric
degradation of carbonyl compounds. Decomposition of these radicals leads to the formation of
acyl radicals (RC•(=O)) and carbonyl compounds (R’CH(=O)) that are smaller than the parent
36
carbonyl compound (R2.3), and are hence less likely to form SOA than products from its
competing pathways.
R2.3 RC(=O)CH(O•)R’ → RC•(=O) + R’CH(=O)
Acyl radicals are good leaving groups, making the C-C bond cleavage rapid. Investigations
of CH3C(=O)CH(O•)CH3 and CH3C(=O)CH2O• have shown that decomposition predominates
under atmospheric conditions,55, 99 consistent with a calculated low barrier of ~6-7 kcal/mol. This
low barrier suggests that chemical activation is important for decomposition of these two alkoxy
radicals.99 The theoretical study of Ferenac et al., found the position of carbonyl functional group
relative to the radical site has a big effect on decomposition barriers.63 For larger ketones, OH
oxidation occurs largely at sites that are distant from the carbonyl group, and hence the influence
of the carbonyl group on the chemistry of the resulting alkoxy radical is negligible.106 Again,
very little kinetic data (if any) are available for reactions of this radical class.
Ester-derived alkoxy radicals are also an important class of oxygenated alkoxy radicals.
Atmospheric esters primarily come from their use as solvents in paints, adhesives, and cleaning
agents, and from the food industry.107 Moreover, ethyl acetate is emitted during the combustion
of esterified rapeseed oil used recently as a diesel fuel.108 As previously pointed out, esters may
also be produced from the decomposition of ether-derived alkoxy radicals. The chemistry of
ester-derived alkoxy radicals varies with the position of alkoxy radical site. The radical type
RC(=O)OCH(O•)R’ (named as acyloxy alkoxy radical) undergoes a rearrangement through a
five member ring transition state leading to RC(=O)OH and R’C•(=O). RC(=O)OH is water
soluble and is an important component of aqueous aerosols, while RC•(=O) directly leads to the
37
formation of peroxyacetyl nitrate (PAN). PANs are powerful respiratory and eye irritants present
in photochemical smog. This process is called α-ester rearrangement, and it was first observed by
Tuazon et al. for ethyl acetate, 109 and the mechanism was proposed as Scheme 2.3:
O H
O H
C C
C C
CH3 O CH3
O CH3 O CH3
O
O O
C + C
CH3 OH CH3
Scheme 2.3 An example of α-ester rearrangement.
The α-ester rearrangement has been confirmed by several product studies for methyl-,110
ethyl-,108 n-propyl-,108, 109

and isobutylacetate,111 as well as methylpropionate,112
methylpivalate,113 and methylformate.114 Studies on ethyl acetate concluded that the α-ester
rearrangement is the exclusive fate for the corresponding acyloxy alkoxy radical
(CH3C(O)OCH(O•)CH3). However, for α-acyloxy alkoxy radicals derived from methyl esters,
reaction with O2 is in competition with the α-ester rearrangement.110, 112, 114 Therefore, it appears
that the more heavily substituted the carbon atom of the alkoxy radical center, the more readily
the α-ester rearrangement occurs. For ester-derived alkoxy radicals where an α-hydrogen atom is
not available, such as CH3C(O)OC(CH3)2O• (formed from isopropyl acetate),109,111
decomposition through C-C bond cleavage to form an anhydride and an alkyl radical becomes
the dominant degradation pathway. For long chain esters, the alkoxy radical site may occur in
38
other positions than the α-carbon. In this situation, the fate of these alkoxy radicals is C-C bond
cleavage and/or reaction with O2, depending on the structure.108,112
Temperature-dependent relative rate studies for the α-ester rearrangement have been
carried out in the temperature range 253-324 K for the acyloxy methoxy radicals HC(O)OCH2O•
and CH3C(O)OCH2O•.107 The estimated activation energy for the α-ester rearrangement in both
radicals is ~10 kcal/mol relative to an estimated activation energy of 0.5 - 1 kcal/mol for the O2
reaction. Several theoretical investigations have been made on α-ester rearrangements.63,115,116
Good and Francisco115 calculated a barrier of about 13 kcal/mol in the case of the methyl formate
radical. Rayez et al.116 found a relatively low barrier for ethyl acetate of 6-7 kcal/mol. Ferenac et
al.63 found different values for the barriers depending on the level of theory. In addition, Rayez et
al.116 also calculated the rate constant of α-ester rearrangement for the case of ethyl acetate using
RRKM theory, and found it three orders of magnitude larger than the experimentally determined
estimate.107 The differences in both the barrier and rate constant of α-ester rearrangement
between the results of experiment107 and theory63,117,118 leave a great deal of uncertainty about the
kinetics of such reaction. If one could obtain a reliable kO2 for this type of radical from a
validated SAR (via both experimental and theoretical calculation), one would be able to extract
reliable rate constant data from relative rate studies.
As discussed above, the decomposition of oxygenated alkoxy radicals generate either
highly water-soluble compounds such as acid, aldehyde, or PAN. However, we can only obtain
the relative importance of the alkoxy radical reactions from estimation methods.33,91 Considering
the variety of functional groups and positional isomers that are possible for the alkoxy radical,
there are far too few absolute or relative rate constant measurements for these radicals. To study
39
the kinetics of these reactions experimentally, an environmental chamber is usually considered,
although such studies will only provide rate constant ratios, such as kO2/kunimo. A good prediction
of kO2 via a reliable computational method would be extremely valuable for the determination of
kunimo for these oxygenated alkoxy radicals. This determination, in turn, would help us build the
SAR for the unimolecular reactions of these radicals. My research will contribute to this
important goal.
2.3 Methoxy + NO2
The major fate of the methoxy radical in the atmosphere is the reaction with O2. However,
reaction with NO2 can also be important in some severely polluted regions such as those near a
power plant plume.119 In the laboratory, the kinetics of methoxy + NO2
kNO2
R2.4 CH3O• + NO2 (+M) → Products
is important to interpret smog chamber experiment results, where NOx (NO + NO2)
concentrations are often much higher than in the atmosphere. Early relative rate studies
employed R2.4 as the reference reaction for studies of the methoxy + O2 reaction.14,120,121 One
goal of my work is to combine the relative rate studies carried out in our group and my absolute
rate measurement of the rate constant of R2.4 (and the corresponding CD3O• + NO2 reaction) to
determine absolute rate constants for methoxy + O2 near ambient pressure from 335 K down to
250 K.122 The methoxy + NO2 reaction (R2.4) is the prototype for kinetic and mechanistic
studies of other RO• + NO2 reactions.
40
Direct kinetic investigations of R2.4 (using LIF detection of CH3O•) have been carried out
at pressures varying from 0.6 to 600 Torr over the temperatures ranging from 220-473 K, with
Ar, CF4 or He as bath gases, as shown in Table 2.3. The discharge-flow method was employed to
produce CH3O• for low-pressure studies by McCaulley et al. (0.6-5 Torr) 123 and Biggs et al. (1-
10 Torr) 124
, while pulsed laser photolysis was used to generate CH3O• for higher pressure
studies by Frost and Smith (6-125 Torr) 125, Wollenhaupt et al. (10-200 Torr), 126 and Martínez et
al. (50-600 Torr) 127. All of these studies showed similar pressure-dependent behavior of kNO2,
except that of Martínez et al.127, whose values are ~30% larger than the rest over the 50-600
pressure range at room temperature.
Table 2.3 The conditions of previous direct rate measurement on CH3O• + NO2. Note: DF-LIF is
discharge-flow/laser induced fluorescence; LFP-LIF is laser flash photolysis/laser induced
fluorescence.
Reference Pressure range Temperature range Bath Gas Method
(Torr) (K)
9 0.6-4.0 220-470 He DF-LIF
10 1-10 298 He DF-LIF
11 30-125 295, 390 Ar, CF4, He LFP-LIF
13 10-200 233-356 Ar LFP-LIF
14 50-600 250-390 He LFP-LIF
R2.4a proceeds via formation of an activated methylnitrate (CH3ONO2) intermediate,
which is quenched by bath gas in competition with dissociation to reactants. Martínez et al.127
used He as a bath gas, and, given that He is expected to be less effective in deactivating the
energized complex CH3ONO2* than Ar and CF4,128,129,130 the rate constants determined by
Martínez et al.127 would be expected to smaller than those determined in the other studies at the
41
same pressure and temperature. Thus, the results of Martínez et al. are an anomaly. Also, there is
some discrepancy in the temperature dependence between the study of Wollenhaupt et al.126 and
that from Martínez et al.127 showed a much larger temperature dependence of k0 and a slightly
126
smaller temperature dependence of k∞ than Wollenhaupt et al. It is noteworthy that R2.4
appears not to have reached the high-pressure limit at 600 Torr of Ar or He.
R2.4a CH3O• + NO2 (+M) → CH3ONO2
R2.4b → HCHO + HONO (minor)
There have been no direct kinetic measurements on R2.4 in the presence of N2 bath gas
except at 50 Torr, so the experiments carried out previously do not mimic tropospheric
conditions. For these reasons, it is valuable to reinvestigate kNO2 at pressures closer to 1
atmosphere while using N2. In my study, I investigated this reaction over the pressure range 30 -
700 Torr in N2 and the temperature range 250-333 K.
The use of LIF detection of the loss of CH3O• in the studies described above means that
the experiment does not directly inform us about the nature of the products. It is widely agreed
that reaction between CH3O• and NO2 can undergo two channels—recombination (R2.4a)
producing methyl nitrate and disproportionation (R2.4b) yielding formaldehyde and nitrous acid.
The recombination channel is believed to be dominant at pressures greater than 1 Torr, and is the
source of the observed significant pressure dependence in the overall rate constant.123,124,125,126,127
1150+550
McCaulley et al. reported a rate constant of 9.6+17.3
−2.7 × 10
−12
exp (− −170
) cm3 molecule-1 s-1
𝑇
for the disproportionation channel, corresponding to 2.0×10-13 cm3 molecule-1 s-1 at 298 K.123
42
It is still unknown whether the two reaction channels proceed via the same energized
complex CH3ONO2* formed directly from association of CH3O• + NO2. McCaulley et al.123
proposed that channel R2.4b undergoes direct intermolecular H-abstraction independent of
channel R2.4a due to the strongly favored redissociation of CH3ONO2* over the internal H-shift.
Several theoretical studies have investigated the energetics as well as the mechanisms of
R2.4.131,132,133 The energetics of the reactions related to CH3O• + NO2 from these theoretical
studies were in general agreement, and are shown in Figure 2.2. One exception is that Lesar et
al.132 who found the disproportionation barrier (labeled TS in Figure 2.2) is several kcal/mol
higher than the reactant energy. If the energy barrier for the disproportionation of CH 3ONO2 is
no higher than, or no more than a few kcal/mol lower than, the energy of CH3O• + NO2, the
dominance of the disproportionation channel at pressures less than 1 Torr is reasonable. At these
low pressures, the rates of collision with the bath gas is too low to quench CH3ONO2* to
energies below the barrier to disproportionation.133 At higher pressures, quenching would
outcompete disproportionation. On the other hand, the rate constants from McCaulley et al.123 is
consistent with a disproportionation barrier that is higher in energy than the energy for CH3O• +
NO2. Such a barrier height is still consistent with the observation of a pressure-dependence of the
association channel (R2.4a).
While all the above-referenced theoretical studies were able to find the transition state for
the disproportionation of methyl nitrate, only Pan et al.131 reported trying (but failing) to find a
transition state for a direct hydrogen atom-abstraction mechanism for R2.4b. In addition to the
formation of methyl nitrate (exothermic by ~40 kcal/mol), the association of CH3O• and NO2
could also form weakly bonded (exothermic by ~ 13 kcal/mol) CH3OONO, which can easily
43
dissociate back to CH3O• + NO2.131,132,133 In addition, Pan et al. 131 found transition states for the
dissociation of CH3OONO to HCHO and HONO (or HNO2). However, these channels are not
kinetically favorable. Likewise, the dissociation of CH3OONO to CH3OO• and NO is not
favored thermodynamically.
Figure 2.2 Simplified energy diagram for reaction CH3O• + NO2. TS is transition state.
Another importance of the reaction of methoxy radical with NO2 is that it is an example of
radical-radical recombination reaction. The experimental pressure-dependent rate constant of this
type of reaction is useful to constrain their RRKM/Master Equation (ME) simulations of the
reaction.134,135 Barker et al.135 did the simulation by using previous experimental results at 297 K
for the CH3O• + NO2 reaction.126 By fitting to the fall-off curve they estimated the collisional
energy transfer parameter (α):
1 −(𝐸−𝐸 ′ )
E2. 8 𝑃(𝐸 ′ , 𝐸) = 𝑁(𝐸) 𝑒𝑥𝑝 [ ] 𝑓𝑜𝑟 (𝐸 − 𝐸 ′ ) ≥ 0
𝛼(𝐸)
where P(E’,E) is the probability density for energy transfer from a high vibrational energy, E, to
low energy, E’, in a deactivation step, N(E) is a normalization factor, and the energy transfer
44
α(E) is approximately a linear function of internal energy and is almost identical to the average
energy transferred in deactivating collisions.135
Although the fitted rate constants agreed well with the experimental values, it was pointed
out that the accuracy of fitting the collisional energy transfer parameter is subject to large
uncertainties in assumptions for treating the transition state, with no well-defined saddle point on
the potential energy surface.135 Furthermore, by using a similar model, they found a quite
different α value for 2-C5H11ONO2 that was 1/40 of that for CH3ONO2. Since the two systems
are the same type of reaction, this finding raises a question that the large discrepancy might be
due to erroneous assumption made in constructing models or erroneous experimental data.135
Therefore, they suggested that experimental re-investigations of the kinetics of both reactions in
an extended pressure range would enable validation or refinement of their fitted values of α.135
For this to happen, it would also require another study with RRKM/ME modeling and fitting.
In Chapter 5, I report on the pressure dependent behavior of the kinetics for reaction R2.4
in the temperature range of 250-335 K and the pressure range of 30-700 Torr with bath gas N2.
This provides the kinetic data of this reaction at pressures up to 700 Torr in the N2 bath gas for
the first time, which is more relevant to the atmosphere than the conditions of previous studies. I
also investigated the kinetics of the isotopologue of CH3O•, i.e. perdeuterated methoxy (CD3O•)
+ NO2, and this offers a check on the reproducibility of my rate constant determinations for R2.4,
and may further facilitate modeling of collisional energy transfer for other barrierless radical-
radical recombination reactions.
45
2.4 Rate theory for recombination reaction (P-dependence)
The main focus of my research is reactions R2.4a. This is a termolecular recombination
reaction. CH3O• and NO2 associate to form an energized complex, which can be stabilized to
CH3ONO2 through deactivating collisions with abundant third-body molecules, which are
usually inert gas molecules such as He, Ar or N2. This process competes with the dissociation of
the energized complex. Such a reaction system corresponds to a reversed Lindemann
mechanism,136 and can be described as follows:
k1
k1
A + B AB*
R2.5 A + B k-1 AB*
k-1
k2
AB* + M k2 AB + M
R2.6 AB* + M AB + M
where A and B are reactants, AB* is the energized product, and M is any third body (mostly bath
gas). Application of the steady-state assumption for [AB*] yields the apparent recombination
rate constant, krec, for stabilization, defined as
E2.9 d[AB]⁄d𝑡 = 𝑘rec [A][B]
where
𝑘1 𝑘2 [M]
E2.10 𝑘rec =
𝑘2 [M]+𝑘−1
In the limit of high or low pressure, E 2.10 can be simplified considerably:
0 𝑘1 𝑘2
E2.11 In the limit as [M] ⟶ 0, 𝑘𝑟𝑒𝑐 = 𝑘−1
[M]
46
∞
E2.12 In the limit as [M] ⟶ ∞, 𝑘𝑟𝑒𝑐 = 𝑘1
where k0(T) is the temperature dependent termolecular rate constant at the low pressure limit and
k∞(T) is the temperature dependent bimolecular rate constant at the high pressure limit.
Figure 2.3 Schematic fall-off curve for the apparent recombination rate constant krec implicitly defined by
E2.8.
Based on Equation E2.10, krec depends on the total pressure, i.e., the concentration of third-
body molecules. By plotting krec vs. pressure, a fall-off curve is usually observed, as illustrated in
Figure 2.3. At very low pressures, krec is linearly dependent on pressure and equals the product of
0
[M] and the low-pressure-limiting rate constant (𝑘𝑟𝑒𝑐 ). At very high pressures, krec is independent
∞
of pressure and is called high-pressure-limiting rate constant (𝑘𝑟𝑒𝑐 ). The region in between the
two limits is described as the fall-off region, where krec increases sub-linearly with pressure. The
breadth of the fall-off region depends on the third body molecules in two ways. First, the size of
47
third-body molecules influences the observed rate constant. The larger and heavier molecule has
a higher collisional cross section, and more readily converts the rotational energy of energized
species to its own translational energy137 Second, polyatomic molecules (such as N2, CF4, SF6)
are more efficient than monoatomic molecules, because the former have vibrational and
rotational modes into which energy can be transferred from the energized intermediate. Negative
temperature dependences are typical for recombination reactions, since the energized complex
would be expected to have more total energy (on average) as the temperature is increased. The
larger the energy with which the energized complex is formed, the higher the rate constant for
dissociation back to reactant, and the lower the probability for collisional deactivation.138
The pressure dependent rate constants can be fit by the Troe expression
𝑘 0 (𝑇)[𝑀]
E2.13 𝑘([𝑀], 𝑇) = (1+𝑘 0 (𝑇)[𝑀]/𝑘 ∞(𝑇)) 𝐹𝑐𝑒𝑛𝑡 𝑝
2 −1
𝑘 0 (𝑇)[𝑀]⁄
𝑝 = (1 + (𝑙𝑜𝑔10 ( 𝑘 ∞ (𝑇))) )
Fcent is a parameter that describes broadening of the fall-off curve, which results from the energy
and angular momentum dependence of k.139,140 A fixed value of Fcent of 0.6 is recommended by
Troe in the range of 100-400 K.141 This recommendation is adopted by the NASA Panel for Data
Evaluation, although not by the IUPAC Subcommittee for Gas Kinetic Data Evaluation, 30 which
allows Fcent to vary when fitting experimental data. Golden142 pointed out that both the NASA
and IUPAC formulations are adequate to represent pressure dependent rate constant as long as
one does not extrapolate too far out of the data range.
48
2.5 Kinetics of CD3O• reactions
To date, there has only been one early kinetic study on CD3O• + O2 at 298 K.143 In this
study, only the rate constant ratio between the reaction of CD3O• + O2 and the reaction of CD3O•
+ CD3O• was determined. Since the absolute rate constant for the reference reaction was (and is)
unknown, the absolute rate constant for the reaction of CD3O• + O2 could not be obtained. There
has been no previous study on the reaction CD3O• + NO2.
Let us consider general information about isotope effects on rate constants. The effect due
to isotopic substitution is quantified in the kinetic isotope effect (KIE), expressed by the ratio
kH/kD in the case of substituting D for H (kH is rate constant for the normal reactant and kD is the
rate constant for the deuterated reactant). Kinetic isotope effects can be primary and secondary.
Kinetic isotope effects are termed primary when the substituted isotope is transferred in the
reaction, and primary KIEs are usually significant for deuteration (factors of three even in the
absence of large tunneling for H-atom transfer). KIEs are termed secondary when the isotopic
replacement is transferred in the reaction. In such cases, deuteration only slightly affects the rate
constant.144 KIEs in the reaction of methoxy + O2 are considered primary KIEs.
The reaction of methoxy + O2 occurs via a transition state that features an energy barrier,
as is shown in Figure 2.4. Compared to CH3O•, deuterium substitution reduces the C-D stretch
frequencies below the values for the C-H stretch frequencies due to the heavier mass of D than
H. This lowers the zero-point energy (ZPE) of the reactant CD3O• as compared to CH3O•. By
contrast, the ZPE of the transition state, i.e., [CD3O…O2]‡ and [CH3O…O2]‡, is expected to
change somewhat less upon deuteration. The result is a larger activation energy of the CD3O• +
49
O2 than CH3O• + O2. Meanwhile the vibrational frequency has only a modest influence on the
ratio of partition functions.138 Therefore, we expect a slower rate constant for CD3O• + O2 than
for CH3O• + O2.
Figure 2.4 Schematic energy profile for two isotopic variants of the methoxy + O2 reaction. (TS denotes
transition state; Horizontal line represent zero-point energy level of corresponding species)
The recombination reaction CD3O• + NO2 + M does not involve any deuterium
participation in bond breaking or formation, and therefore the KIE is categorized as secondary.
The KIE for this reaction would be most strongly influenced by the fact that, at the energy of the
separated reactants, CD3ONO2 has a larger density of states than CH3ONO2, making the
dissociation of CD3ONO2 back to the reactants slower than for CH3ONO2. Hence a slightly
inverse KIE would be expected to be observed, i.e., the rate constant become larger upon
deuteration. McCaulley et al.145 studied the similar reaction CD3O• + NO over pressure range
50
0.75-5.0 Torr at 294 K, and found that kCD3O•+NO increased with pressure and was slightly larger
than kCH3O•+NO (~1.2 times at 5 Torr). Based on the fact that the contribution of pressure-
independent disproportionation channel becomes less as pressure increases, it was inferred that
the rate of recombination was increased by the deuterium substitution.
2.6 Principle of LFP/LIF method
The experimental method used in this study is the pulsed laser flash photolysis/pulsed laser
induced fluorescence (LFP/LIF) method. The principle of LFP/LIF includes two successive
steps: (1) generating methoxy radical in a sharp busrt (~20 ns) through the photolysis of a
precursor by the photolysis laser and (2) exciting methoxy radical by the probe laser and
collecting the fluorescence. The fluorescence signal is proportional to methoxy radical
concentration.
2.6.1 CH3ONO photolysis
The methoxy radical will be produced by the photodissociation of methyl nitrite, and the
reaction proceeds as R2.7:
R2.7 CH3ONO (S0) + hν → CH3ONO (S1) → NO + CH3O•
The ultraviolet absorption spectrum of methylnitrite is characterized by two regions: (1) a
weak and relatively broad absorption band in the region of 300-400 nm, with vibrational
structure, which is assigned to the S1←S0 (π*← n) transition and (2) a strong but structureless
51
absorption band ranging from 160-280 nm (peaking around 220 nm), which is assigned to
S2←S0 (π*← π) transition.146 The S1 state is only bound by about 40 kcal/mol. When CH3ONO
is excited to the S1 state, the CH3O—N=O bond breaks in a time so short (25-320 fs), that it only
allows for one or two stretches of the N=O.147, 148

It is these stretches that give rise to the
vibrational structure in the absorption spectrum. The potential energy surface of the S2 state is
directly dissociative along the CH3O-NO coordinate, which causes this bond to break within ~20
fs. 149 In my experiments, I use 351 nm as the photolysis wavelength. I could use 248 or 193 nm,
but this risks that too much of the excess energy of the photolysis over the CH3O-N=O bond
energy would be imparted to CH3O•.149 The less internal energy, the less the potential that
CH3O• decomposes and less time is needed for the bath gas to remove excess energy initially
imparted to CH3O• by photolysis.
2.6.2 Principle of LIF spectroscopy
The working principle of LIF spectroscopy can be depicted by simplified Jablonski
diagram shown in Figure 2.5.
52
Figure 2.5 Principle of laser induced fluorescence (LIF) by simplified Jablonski diagram
In this process, the alkoxy radicals produced from photolysis are first excited from the
̃ 2 E) to the fourth vibrational level of the first excited state with the
electronic ground state (X
̃2 A1 ). Through collisional relaxation, the excited alkoxy radicals are partially

same multiplicity (A
̃2 A1 state. The emission from a given vibrational

quenched to lower vibrational levels of the A
̃2 A1 state can occur to several vibrational levels at A

state of the A ̃2 A1 state, so the observed
fluorescence spans a wide range of emission wavelengths. By tuning the probe laser across a
range of wavelengths while collecting the total fluorescence intensity, one obtains the LIF
excitation spectrum, which provides information on the vibrational energy levels of the excited
state.144 In the kinetic study, the probe laser is tuned to a specific vibrational transition of the
methoxy radical chosen to maximize S/N.
53
Absolute rate constant measurements of the alkoxy radical were first carried out by Inoue
et al.150 for CH3O• and CD3O•. They reported the transition origin (energy/
frequency/wavelength corresponding to the transition between the zero-point levels of the two
excited states) to be at 31540 cm-1 (317.05 nm) for CH3O• and 31546 cm-1 (316.99 nm)
respectively; extensive progressions in the C-O stretching mode (3' = 678 cm-1 for CH3O• and
655 cm-1 for CD3O•) dominate the spectrum. No predissociation occurs below 35900 cm-1 in
CH3O•. Miller and coworkers151 investigated the LIF spectrum of methoxy in a supersonic free
jet expansion environment, where rotational temperatures were sufficiently low (3-25 K) to
obtain a rotationally resolved spectrum. The results of this study yielded a more reliable value of
the transition origin (31650 cm-1), with a 3' spacing of 660 cm-1.
In another study of jet-cooled methoxy radical from Miller’s group,152 a threshold for
̃ 2 E,
photodissociation of methoxy radical was found to be the energy for the transition (X
̃2 A1 , ν’3=6). This threshold is consistent with the theoretical prediction that153,154 the
ν”3=0) (A
2
A1 state could cross with three repulsive states 4A2, 2A2 and 4E (likely in that energy order).
Surface hopping from the X2E state to one of these quartet states leads to the dissociation of
CH3O to CH3 (2A1) + O (3P) (the ground states of both products). Below this threshold, the
fluorescence quantum yield is near unity,152 while above this threshold, photodissociation occurs
and the fluorescence quantum yield starts dropping rapidly.
54
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62
Chapter 3. Experimental Design
This chapter outlines the methods used for my research: (1) the synthesis of photolytic
precursor (methyl nitrite) to methoxy radicals, (2) the experimental apparatus (lasers, optics,
electronic, and sample handling), and (3) obtaining laser induced fluorescence spectra, together
with carrying out the absolute rate constant measurement.
3.1 Preparation of gaseous reactant
The photolytic precursor of methoxy radical, methyl nitrite (CH3ONO or CD3ONO), was
synthesized from the corresponding methyl alcohol (CH3OH, Sigma-Aldrich, 99.8%; CD3OH
Sigma-Aldrich, 99.8%). Synthesis of alkyl nitrites can be considered an O-nitrosation reaction,
as in shown in reaction R3.1.

H2SO4
R3.1 ROH + NaNO2 RONO + H2O
The mechanism for reaction R3.1 is illustrated in Scheme 3.1. In the mechanism, NO2- is
first protonated to HNO2, which is subsequently protonated to H2ONO+. The OH group of the
alcohol attacks N of the H2ONO+ to form RO(H)NO+ and H2O. Deprotonation of RO(H)NO+
yields RONO. Based on this mechanism, the -NO group replaces the hydrogen on the hydroxyl
group, and therefore the structure and isotopic content of the alkyl group in the alcohol is
retained in the alkyl nitrite.1
63
NaNO2 + H+ HNO2 + Na+
H O N O + H+ H O N O
H
R OH + H O N O R O N O + H2O
H H
R O N O R O N O + H+
H
Scheme 3.1. Mechanism of alkyl nitrite synthesis. R represents the alkyl group.1
The synthesis was carried out by dropwise addition of 57.6% (by mass) sulfuric acid
solution into an aqueous solution of NaNO2 and methanol.2,3 The reaction was carried out at 0
°C. The resultant gaseous products were transferred by N2 gas over a NaOH solution (absorb
NO2) and then over anhydrous CaCl2 (dessicant), and finally collected in a dry-ice trap at -50 °C.
The isolated product was purified by freeze-pump-thaw distillation, resulting in a pale yellow,
glassy solid, which was stored at -196 °C until it was used. Gaseous methyl nitrite was
characterized by Fourier transform infrared spectroscopy (FTIR)4,5 (Bruker Tensor 27) and UV-
Vis spectroscopy (Agilent 8453).3,6 The UV-Vis absorption spectrum of methylnitrite is
characterized by two featured regions: a weak absorption band in the region of 300-400 nm with
broad peaks at 310, 318, 328, 339, 351, 364 nm above a featureless baseline, and strong but
structureless absorption band extending from 160-280 nm and peaking around 220 nm. 3,6 The
identity of methylnitrite can also be confirmed by FTIR at peaks 1676 cm-1, 800 cm-1, 600
cm-1.4,5
NO2 was produced by mixing NO (American Gas group, >99.5%) with a large excess of
O2 (MG welding products 99.999%), which was allowed to react overnight. The resulting
64
brownish gas was purified by free-pump-thaw distillation at -196 °C until a pure white solid was
obtained. The resultant NO2 was characterized by FTIR7 and UV-Vis spectroscopy.8 The UV-Vis
absorption spectrum of NO2 between 200 – 800 nm can be separated into two regions; a strong
absorption below 250 nm and the broad absorption between 300 and 800 nm peaking around 400
nm. The presence of nitrogen dioxide is indicated in the FTIR spectrum by the peaks centered at
2220 and 1621 cm-1.7
For the kinetic experiments, purified CH3ONO (CD3ONO) was purged into a 10-Liter
blackened glass bulb, which was then diluted with N2 (Haun Welding Supply, 99.999%) to a
total pressure of 1000 Torr. This resulted in ~3% methyl nitrite in N2. NO2 was diluted with N2
in another blackened glass bulb, with a molar percentage of ~2%. Concentrations of NO2 in the
gas bulb were determined from their UV-vis absorptions by applying Beer’s law: A = σ × c × l,
where A is the absorbance of gaseous sample at wavelength λ, σ is the corresponding absorption
cross section, c is the concentration, and l is the path length of UV cell. Using known absorption
cross sections of six wavelengths in the range of 380-440 nm for NO2,8 the concentration at each
wavelength was determined. The average value of the six concentrations was used as the
concentration of NO2 in the bulb. The relative standard error for 2σ was ~5%, and it was
considered as the uncertainty for the NO2 concentration. Likewise, the concentration of methyl
nitrite was calculated using known absorption cross sections at peaking wavelengths in the range
of 310-370 nm. 3
65
3.2 Absolute rate constant—LFP/LIF method
3.2.1 Experimental apparatus
Pulsed laser flash photolysis coupled with pulsed laser induced fluorescence (LFP/LIF) is
the method used for measuring the rate of decay of methoxy radical concentration for kinetics
experiments. The LIF apparatus is shown in Figure 3.1. A pulsed XeF excimer laser (GAM Laser
Inc., EX100H) with energy of 10 mJ/pulse and repetition rate of 2 Hz was used to generate
CH3O• by photolyzing CH3ONO at 351 nm. The resulting methoxy radicals were probed at ~293
nm (0.1 mJ/pulse) by the frequency-doubled (Inrad Autotracker III) narrow band emission from
a dye laser (Lambda Physik FL3002,), which was pumped at 308 nm using the XeCl excimer
laser (Lextra 100) operating at 2 Hz. The dye solution was made by dissolving rodhmin 6G in
methanol solution in recommended ratios.9 The energy of the dye laser beam was optimized by
adjusting its alignment, and its output was cleaned by a cutting-edge iris before entering the
frequency doubler. The frequency doubler was mounted by a mounting with a platform to make
sure that the dye laser beam passed through the crystal inside the doubler. A harmonic separator
(Inrad 752-104) was placed after the doubler to separate the frequency-doubled beam from dye
laser beam. The harmonic separator was also carefully mounted to the same level as frequency
doubler. A level was used to ensure the surfaces holding the frequency doubler and harmonic
separator were horizontal. The final frequency-doubled beam was again cleaned by an iris. The
purity and wavelength of the dye laser beam and the output of the frequency doubler were
checked using a monochromator (SpectraPro®-300i), which was calibrated with mercury lamp
(Oriel 6035).
66
The LIF cell was made of a jacketed 1.9 liter PyrexTM tube with an inner diameter of 57
mm. Two glass side arms with an inner diameter of 19.7 mm were installed at each side of the
cell. Quartz windows were attached at Brewster’s angle to the end of each arm. In order to
minimize scattered light inside the cell, two cone apertures were installed in the left arm (shown
in Figure 3.1) to reduce the scattered light. For the same purpose, both arms were painted black
with Krylon® black, and the Wood’s horn on the bottom was coated with black AquadagTM. The
temperature in the LIF-cell was controlled between 250 and 335 K by flowing cooled ethanol or
heated ethylene glycol through the jacket and measured with a calibrated thermocouple
thermometer (Digi-Sense® Dual Input J-T-E-K®). The temperature of the cooling and heating
liquid was regulated by a temperature bath (Neslab ULT-80).
The photolysis laser beam was introduced into the cell from the right side arm (80 mm) by
a 2 m focusing lens and a laser mirror. The probe laser beam was directed into the cell from the
left side arm (120 mm) through optics combination. As a result, the two laser beams counter-
propagated nearly collinearly through the LIF cell. Adequate overlapping of the two beams was
ensured by tracking the beam path using a white card held inside the LIF cell through the port
where the Wood’s horn was usually installed. The white card fluoresces upon being struck by the
either the photolysis or probe laser. The diameter of the photolysis laser beam was adjusted to be
three times that of the probe laser beam so that the interference of diffusion was minimized. Red-
shifted emission from the excited radicals was collimated by two convex lenses (f=10 cm),
passed through a long-pass filter that can pass light with a wavelength greater than 345 nm, and
entered the photomultiplier tube (PMT, R212UH, Hamamatsu Photonics) mounted on top of the
cell window orthogonal to both laser beams. The filter acts effectively to shield the PMT from
67
scattered light from both laser beams. The signal from the PMT was amplified 10 times by a pre-
amplifier (Model: Ortec 9305), and transmitted to a boxcar averager No.1 (Model SR250,
Stanford Research Systems, Inc.), as indicated in Figure 3.1, for signal integration. The
sensitivity of the boxcar averager was set to 0.1V/1V, meaning that the output integrated signal
was magnified by 10 times. Subsequently the output signal was transmitted to an analog-digital
converter computer interface (Model SR245, Stanford Research Systems, Inc.), and finally to a
computer installed with data acquisition software (SR272, Stanford Research Systems, Inc.).
Simultaneously, the LIF signal was transmitted to an oscilloscope for real-time monitoring.
Figure 3.1 The experimental apparatus used in this study.
68
The time delay between two laser pulses was controlled by a Digital Delay Generator
(DDG, Model DG535, Stanford Research Systems, Inc.), which was triggered by the photolysis
laser. The timing sequence of a 2 Hz operation with a preset of delay time of 5 μs between
photolysis laser and pump laser is depicted in Figure 3.2. The photolysis laser triggers itself and
sends a triggering signal to the DDG at time T0. ~1 μs (internal delay of photolysis laser) later,
the photolysis is fired at time TP. The internal delay of the DDG is 85 ns. Depending on the
preset delay time for each experiment, the pump laser was triggered at T1 = T0 + 85 ns + delay
time. After ~1 μs (the internal delay of pump laser), the pump laser pumps dye laser and the
fluorescence signal of the excited methoxy radical is formed at time Tf = T1 + 1 μs. In order to
avoid the weak scattered light signal from probe laser, the gate of boxcar averager was opened
50 ns after the fluorescence pulse, and the gate width for integration of the fluorescence signal
was 25 ns.
69
Figure 3.2 Timing sequence of an LIF experiment with a specific delay time of 5 μs between two laser
pulses preset in DDG.
To obtain a background LIF signal, I also conducted the LIF experiment by blocking the
photolysis laser while having only probe laser beam pass through the LIF cell. No fluorescence
signal was found for either CH3ONO or CD3ONO, however, there was electronic noise and very
weak scattered light (~10 ns) that creates a background signal that was subtracted from LIF
signal obtained with both lasers passing through the cell. To correct for fluctuations in the
intensity of the probe laser pulses, scattered light from the interaction of the probe beam with a
prism was detected by photodiode 2. The resulting signal was transmitted to boxcar 2 and finally
to the computer.
70
3.2.2 Gas handling system
Buffer gas N2, ~3% methyl nitrite in N2, and ~2% NO2 in N2 gas flowed through stainless
steel tubing into the reservoir of the heater/chiller before passing into the LIF cell. The flow rate
of each gas was regulated by a calibrated pressure-based Mass-FLO® controller (MKS type
1640). The reaction pressure in the cell was measured, close to LIF zone, by absolute pressure
transducers (MKS Baratron® types). The partial pressures of NO2 and methyl nitrite in the cell
were determined using
E3.1 Pi = (Ri / Rt) × Si% × Pt
where Pi is the partial pressure of species i (i.e., NO2 or methyl nitrite), Pt is the total pressure in
the cell, Ri is the flow rate of a mixed gas (in a bulb) containing species i balanced with N2, Rt is
total flow rate (i.e., sum of the flow rates of all mixed gas and bath gas), and Si% is the molar
percentage of species i in corresponding mixed gas.
The gas mixture flowed into the LIF cell from one end, passed through the LIF detection
zone in the middle, and was pumped out of the cell from other end. Continuous gas flow through
the cell helped ensure that a fresh sample of methylnitrite was present in the monitored volume
for each successive laser shot. The flow was slow enough for the gas to be essentially unmoving
over the course of the delay times used to study the reaction kinetics.
3.2.3 Estimation of radical concentration
The concentration of methoxy radical is a crucial factor as it determines what
concentration of the other excess reactant, e.g. NO2, we need to use in order to achieve pseudo-
71
first order conditions. Pseudo-first order conditions allowed me to carry out the experiment
without needing to accurately determine absolute radical concentrations, a task which is
challenging for radicals, generally, and especially when using fluorescence. The determination of
the methoxy radical concentration is calculated based on the characteristics of the photolysis
laser beam and the concentration and absorption cross-section of methylnitrite.
We took a single UV spectrum of methyl nitrite in a 10 cm pyrex cell covered with two
quartz windows on both sides. The gas cell was held by a special long-path length cell holder
(Agilent 89076C).
The concentration of methoxy radical was calculated using the following equation (E3.2):
E3.2 [𝑚𝑒𝑡ℎ𝑜𝑥𝑦] = [𝑚𝑒𝑡ℎ𝑦𝑙 𝑛𝑖𝑡𝑟𝑖𝑡𝑒]𝑝ℎ𝑜𝑡𝑜𝑙𝑦𝑧𝑒𝑑 =
[N × ϕ × (1 − e−A351nm )]/V
where φ is the quantum yield (~1) of methoxy radical formation from the photolysis of
methylnitrite,11,10 V is the volume of the photolysis region, given by the product of photolysis
laser beam cross section area and laser path length (L) inside the LIF cell, equal to 0.6 cm ×1 cm
× L cm. The absorbance at 351 nm, A351nm, equals 351 × [methylnitrite] × L. 351 is the
absorption cross section of methyl nitrite at 351nm. 351 for CH3ONO is ~3×10-19 cm2 11,12 and
the same value was applied to CD3ONO, since I observed no obvious difference between the UV
spectra of CH3ONO and CD3ONO, shown in Figure 3.3 This agrees with our observation at the
National Center for Atmospheric Research (NCAR).13 The term (1 − 𝑒 −𝐴351𝑛𝑚 ) represents the
fraction of photons that are absorbed by methyl nitrite, as it is derived from another form of
Beer’s law 𝐴 = ln(𝐼0 /𝐼), where I0 and I are the intensity of the incident light and transmitted
72
light respectively. N is the number of photons generated by each pulse of the excimer
(photolysis) laser, given by equation E3.3:
E3.3 𝑁 = 𝐸𝜆/(ℎ𝑐)
where λ is the wavelength of photolysis laser (351 nm), E is the energy per pluse, in my case,
E≈10 mJ, c is the speed of light (3×108 m s-1), and h is Plank’s constant (6.63×10-34 J s). With
these values, there were ~1.8×1016 photons per pulse.
With the methyl nitrite concentration of ~3×1015 molecule cm-3 in our experiment, I
calculated a methoxy radical concentration of ~2×1013 molecule cm-3 at 295 K. Also from the
calculation, the radical concentration was found to not be affected by the laser pathlength, and
therefore we can assume that the radical concentration was uniform along the path of the
photolysis laser beam.
73
Absorbance (arbitrary unit) 0.20 CH3ONO
CD3ONO
0.16
3
0.12
2
0.08 1
0.04 0
200 300
0.00
300 400 500

Wavelength (nm)
Figure 3.3 UV spectra of CH3ONO and CD3ONO. Black line (lower) —CH3ONO; red line (upper)—
CD3ONO.
3.2.4 Kinetic experiment
̃2 A1 ← X
I obtained the fluorescence excitation (A ̃ 2 E) spectra (LIF) of CH3O• and CD3O• in
the range of 285.7-302.7 nm at 295 K, as shown in Figure 3.4 and 3.5 respectively. Both Figures
also include the spectra from literature for comparison. This work well reproduced the LIF
spectra of Inoue et al.14 My work show clear progressions in the C—O stretching mode (ν3),
which are ~ 610 cm-1 and ~ 630 for CH3O• and CD3O• respectively. In addition, my spectrum of
CH3O• also shows general agreement with the absorption spectrum of CH3O•.15
74
Figure 3.4 LIF spectrum of CH3O• in the range of 285.7-302.7 nm at 295 K. Details of excitation
transitions (𝐴̃2 𝐴1 ← 𝑋̃ 2 𝐸) in ν3 mode (C—O stretching) are denoted on top of each peak. For example, in
the case of 5’←0’’, 0’’ represents the first vibrational level of ground electronic state, 5’ represents the
sixth vibrational level of first excited electronic state. Blue line—this work; black dot line—the work of
Inoue et al.14
75
Figure 3.5 LIF spectrum of CD3O• in the range of 285.7-302.7 nm at 295 K. Details of excitation
transitions (𝐴̃2 𝐴1 ← 𝑋̃ 2 𝐸) in ν3 mode (C—O stretching) are denoted on top of each peak. The meanings
of the denotations are the same as Figure 3.4. Red line—this work; black dot line—the work of Inoue et
al.14
We chose the peaks at 293.06 nm and 293.35 nm for monitoring CH3O• and CD3O•,
̃2 A1 , ν’3= 4) ← (X
respectively. In both isotopologues, these are the (A ̃ 2 E, ν”3=0) transitions,
where ν3 is the C-O stretching mode.16 The effective excitation fluorescence regions for NO2,
NO and HCHO are >400 nm,17 <240 nm,18,19 and >320 nm20,21,22 respectively. Therefore, major
species besides methoxy radicals do not seem able to produce interfering fluorescence signals.
76
All the LIF signals were corrected for their corresponding background signals, as described
above. The rate of the disappearance of methoxy radicals was detected through the change of
fluorescence intensity as a function of delay time between the photolysis and probe beams. The
shortest delay time used was 5 µs, with the maximum ranging from 30 µs to 125 µs depending
on the concentrations of NO2. At each delay time, 100 to 120 laser shots were taken and
averaged.
In the kinetic experiments, the time-resolved CH3O• temporal profiles were recorded as a
function of the delay time between the photolysis laser and the probe laser pulse. These profiles
were generally well represented by a single-exponential decay, which indicated that pseudo-first
order conditions held for loss of the methoxy radical. The lifetime, , of the decay at any set of
conditions (temperature, [NO2], and total pressure) was the inverse of the pseudo-first order rate
constant for loss of methoxy radical under those conditions. The second-order removal rate
constants of methoxy radicals were determined from changes in the exponential decay rate, with
systematic changes in the NO2 partial pressure at fixed temperature and total pressure.
Obtaining the apparent second order rate constant for methoxy + NO2 at any one total
pressure involved two steps for data analysis:
a. Obtain the pseudo-first order rate constant (k1st) by linear fitting of ln
(fluorescence intensity) versus delay time between the two lasers at each fixed
temperature and pressure.
b. Obtain the bimolecular rate constant by linear fitting of Ln [k1st] from a versus
NO2 concentrations at each fixed temperature and pressure.
77
Two additional steps were carried out to characterize the pressure dependence of the apparent
second-order rate constant.
c. Determine k0 and k∞ at each temperature through non-linear least-squares fitting
of the fall-off curves with Troe expression8,23 (equation E2.13 from Chapter 2) by
plotting k2nd vs. total pressure.
d. Get the temperature dependence of k0 and k∞ by plotting log [k0] and log [k∞]
against log [temperature], together with a weighted-linear least squares fitting to the T-m
behavior. This is described in more detail in Chapter 5.
78
References
1 D. L. H. Williams, Nitrosation; Cambridge University Press, 1988, pp 153

2 A. H. Blatt, Organic Syntheses; Wiley: New York, 1966, pp 108-109.
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12, 231-240.
4 Rook, F. L. J. Chem. Eng. Data 1982, 27, 72-73.
5 Rook, F. L.; Jacox, M. E. J. Mol. Spectrosc. 1982, 93, 101-116.
6 J. G. Calvert and J. N. Pitts, Photochemistry; Wiley: New York, 1966, 455 ff
7 Mabbott, G. A. J. Chem. Educ. 1995, 72, 471.
8 S. P. Sander, J. Abbatt, J. R. Barker, J. B. Burkholder, R. R. Friedl, D. M. Golden, R. E. Huie, C. E.
Kolb, M. J. Kurylo, G. K. Moortgat, V. L. Orkin and P. H. Wine "Chemical Kinetics and Photochemical
Data for Use in Atmospheric Studies, Evaluation No. 17," JPL Publication 10-6, Jet Propulsion
Laboratory, Pasadena, 2011 http://jpldataeval.jpl.nasa.gov.
9 Instruction Manual for Dye Laser FL 3001/3002, Lambda Physik Inc., Gӧttingen, Germany, 1987.
13 Hu, H.; Dibble, T. S.; Tyndall, G. S.; Orlando, J. J. J. Phys. Chem. A 2012, 116, 6295-6302.
15 Wendt, H.; Hunziker, H. J. Chem. Phys. 1979, 71, 5202.
92, 5914-5921.
17 Anastasi, C.; Hancock, D. U. J.Chem.Soc., Faraday Trans.2 1988, 84, 1697-1706.
79
18 Verbiezen, K.; van Vliet, A.; Meerts, W.; Dam, N.; ter Meulen, J. Combust. Flame 2006, 144, 638-
641.
19 Schulz, C.; Yip, B.; Sick, V.; Wolfrum, J. Chem. Phys. Lett.1995, 242, 259-264.
20 Smith, G. D.; Molina, L. T.; Molina, M. J. J. Phys. Chem. A 2002, 106, 1233-1240.
21 König, R.; Lademann, J. Chem. Phys. Lett.1983, 94, 152-155.
22 Schneider, A.; Mantzaras, J.; Bombach, R.; Schenker, S.; Tylli, N.; Jansohn, P. Proceedings of the
Combustion Institute 2007, 31, 1973-1981.
23 Troe, J. J. Phys. Chem. 1979, 83, 114-126.
80
Chapter 4. Rate constants and kinetic isotope effects for methoxy radical
reacting with NO2 and O2
This chapter is part of a manuscript of the same title above. Experimental setup and result
analysis regarding relative rate study are not shown here, and they belong to Hongyi Hu’s
thesis.1
4.1 Introduction
Alkoxy radicals (RO•) are important intermediates in the photooxidation of volatile
organic compounds (VOCs) in the troposphere. The fates of alkoxy radicals (unimolecular
decomposition and isomerization, and reaction with O2) greatly impact ozone formation in the
troposphere as well as gas-particle partitioning of the eventual stable products. 2,3 To date, direct
kinetic studies of RO• + O2 have been limited to alkoxy radicals derived from C1-C7
alkanes4,5,6,7,8,9,10,11,12,13,14,15,16 and two halogenated alkanes.17,18,19 The rate constants for O2
reactions are unknown for alkoxy radicals derived from oxygenated VOCs or non-alkane
hydrocarbons. Many previous studies of alkoxy radical kinetics have only determined the rate
constant ratio kunimolecular/kO2, and relied upon an estimate of kO2 to determine kunimolecular. The lack
of absolute rate constants, kO2, obstructs the determination of kunimolecular, thus preventing us from
establishing accurate structure-activity relations (SARs) for the unimolecular reactions. SARs are
needed to enable prediction of the tropospheric fate of larger and functionalized alkoxy radicals,
for which experimental data is largely absent and difficult to obtain from experiments.20, 21
81
Methoxy radical (CH3O•) is the prototype for all alkoxy radicals. Similarly, the kinetics
and mechanism of the methoxy + O2 reaction (R4.1):
k1 HCHO + HO •
R4.1 CH3O• + O2 → 2
is the prototype for other RO• + O2 reactions. Both absolute and relative rate studies have been
carried out to determine k1. Rate constant ratios for R4.1 versus CH3O• + NO or NO2 have been
reported for 296 ≤ T ≤ 450 K by various groups in the 1970s.22,23,24,25,26,27 These efforts used
product analysis following the photolysis of methyl nitrite (CH3ONO) or the pyrolysis of
dimethyl peroxide (CH3OOCH3) or CH3ONO. The rate constants ratios obtained in these
experiments exhibited a lot of scatter, and only the result of Cox et al.24 at room temperature
agree well with absolute rate studies. Three absolute measurements of k1 over 298 ≤ T ≤ 973 K
have been conducted using laser flash photolysis-laser induced fluorescence (LFP-LIF), with
either CH3ONO or CH3OH used as the CH3O• precursor.4,5,6 These results are in general
agreement in their range of overlap (298 ≤ T ≤ 610 K), and were fitted by an Arrhenius
1150±190
expression: 𝑘1 = 7.82+4.68
−2.93 × 10
−14
exp[− ] cm3 molecule-1 sec-1, with quoted
𝑇
uncertainties of two standard deviations.20 At 298 K this expression yields k1=1.6 × 10-15 cm3
molecule-1 s-1. Rate constants obtained at T > 610 K greatly exceed those obtained by an
extrapolation of an Arrhenius plot of the data at lower temperatures.4 Notably, no measurement
of k1 has been done below room temperature. This is because k1 is small and becomes smaller as
temperature decreases, so that the use of high concentration of O2 (>50 Torr) is required.
Unfortunately, O2 efficiently quenches fluorescence of excited CH3O• (Ã 2A1).28
82
To overcome the difficulty of directly measuring k1 below room temperature, we combined
measurements of the ratio k1/k2 with absolute determination of k2, where k2 is the overall rate
constant for reaction of CH3O• with NO2:
k2
R4.2b → HCHO + HNO2 (minor)
Measurements were carried out over the temperature range 250 – 335 K. The rate constant ratio
(k1/k2) was measured at the National Center for Atmospheric Research (NCAR) in a smog
chamber based on product analysis by Fourier Transform InfraRed (FTIR) spectroscopy. The
absolute rate constants k2 were measured at SUNY-ESF using LFP-LIF. By combining the two
measurements, the absolute rate constant k1 was determined as a function of temperature.
The kinetics of methoxy + NO2 (R4.2) can be important to interpret smog chamber
experiments, where NOx concentrations are often much higher than in the atmosphere. It is
widely agreed that reaction between CH3O• and NO2 can proceed via two channels—
recombination (R4.2a) producing methyl nitrate (CH3ONO2), and disproportionation (R4.2b)
yielding formaldehyde and nitrous acid. McCaulley et al. reported a rate constant of 9.6+17.3
−2.7 ×
1150+550
10−12 exp(− −170
)cm3 molecule-1 s-1 for the disproportionation channel, which is only
𝑇
significant at rather low pressures, e.g., k2b/k2a≈ 0.1 at 1 Torr and 298 K.29,30 At higher pressures,
the recombination channel becomes even more dominant.30,31,32,33,34 Direct kinetic investigations
of R4.2 (using LIF detection of CH3O•) have been carried out at pressures up to 600 Torr over
the temperature range 220-473 K with Ar, CF4 or He as bath gases. All of these studies showed
broadly similar pressure dependent behavior of k2, however, the values from the two studies with
83
the largest pressure range differ by 30%. Due to the inconsistency of previous results on k 2, it is
valuable to re-examine this rate constant, and especially to use a buffer gas more representative
of air than Ar, CF4, or He.
We also investigated the deuterium kinetic isotope effect (KIE) of the methoxy + O2
reaction by substituting CD3O• (R4.3 and R4.4) for CH3O•:
R4.3 k3 DCDO + DO2

CD3O• + O2 →
k
R4.4a CD3O• + NO2 (+M) →4 CD3ONO2
R4.4b → DCDO + DNO2 (minor)
over the temperature range of 250-333 K, using the same method of combining measurement of
the ratio k3/k4 with the absolute measurement of k4. KIE is defined as kH/kD—the rate constant
ratio between the reaction involving non-deuterated reactant and that involving deuterated
reactant; here KIEs for methoxy + O2 and methoxy + NO2 are k1/k3 and k2/k4 respectively.
No kinetic studies have been reported for either R4.3 or R4.4, except for one relative rate
study that estimated k3 as (8.0-21) × 10-18 cm3 molecule-1 s-1 at 298 K.35 This value would imply
a KIE of about 100 at 298 K. Even if this KIE estimate is high, the value of k 3 is expected to be
significantly smaller than k1, thus making the measurement of k3 by LFP-LIF method extremely
difficult even at room temperature.
A long-term goal of ours is to better understand the kinetics of alkoxy + O2 reactions,
including the role of tunneling. Several theoretical studies have tried to elucidate the mechanism
of CH3O• + O2. Jungkamp and Seinfeld proposed that the reaction occurs via formation of a
short-lived trioxy radical intermediate followed by HO2 elimination;36 this mechanism is
consistent with the unusually low Arrhenius pre-exponential factor (A-factor) for the reaction.
84
However, Bofill et al.37 found an error in their analysis, and reported an enormous barrier (50
kcal/mol) to HO2 elimination from the trioxy radical intermediate. Instead, Bofill et al. found that
H-abstraction, while direct, occurs through a five-member ringlike transition state structure that
accounts for the low A-factor.37 Based on this mechanism, both Bofill et al.37 and Setokuchi and
Sato38 calculated k1 in good agreement with experiment. Curiously, these two groups found very
different tunneling corrections () to the rate constant at 298 K. Bofill et al. found  =9 using the
asymmetric Eckart model, while Setokuchi and Sato found  of about 2 using a
multidimensional tunneling approach. Recently two of us carried out calculations on three
isotopologues of the methoxy+O2 reaction.39 These calculations further confirmed the reaction
mechanism of Bofill et al.,37 and obtained analogous tunneling corrections to those of Setokuchi
and Sato.38 These calculations also agreed remarkably well with our previously reported
experimental branching ratio for hydrogen- versus deuterium-abstraction in the reaction CH2DO
+ O2.40
Immediately below I present a description of the absolute rate constant measurement of
R4.2 and R4.4 at SUNY-ESF. Next we present absolute rate constant measurements of R4.2a
and R4.4a and combine the absolute and relative rate data to yield rate constants for R4.1 and
R4.3. This is followed by a comparison with previous rate constant determinations and a
discussion of KIEs.
85
4.2 Experiment
4.2.1 Preparation of gaseous reactants
The photolytic precursor of the methoxy radical, methyl nitrite (CH3ONO or CD3ONO)
was synthesized from the corresponding methyl alcohol (CH3OH, Sigma-Aldrich, 99.8%;
CD3OH Sigma-Aldrich, 99.8%).41,42 The reaction was initiated by dropwise addition of
concentrated (58 % by mass) sulfuric acid solution into an aqueous solution of NaNO2 and
methanol at 0 °C. The resultant gaseous products were transferred by N2 gas over a NaOH
solution and then over anhydrous CaCl2, and finally collected in a dry-ice trap at -78 °C. The
isolated product was purified by freeze-pump-thaw distillation, resulting in a pale yellow, glassy
solid, which was stored at -196°C until it was used. Gaseous methyl nitrite was characterized by
Fourier Transform InfraRed (FTIR)43 and ultraviolet (UV)-visible spectroscopy.42,44 NO2 was
produced by mixing NO (American Gas group, >99.5%) with a large excess of O2 (MG Welding
Products 99.999%), and purified by free-pump-thaw distillation at -196 °C until a pure white
solid was obtained. The resultant NO2 was checked for purity via FTIR45 and UV-visible
spectroscopy.46
In both experiments, the purified CH3ONO (CD3ONO) was first transfered into a
blackened glass bulb, which was then diluted by bath gas N2 (in the LFP-LIF experiment: Haun
Welding Supply, 99.999%; in the chamber experiment: General Air, liquid nitrogen boil-off) to a
total pressure of 1000 Torr. This resulted in ~3% methyl nitrite in N2. NO2 was diluted with N2
in another blackened glass bulb, with molar percentage of ~2%. Concentrations of CH3ONO and
NO2 in the gas bulbs were determined using UV-visible absorption cross sections42,46 of multiple
86
peaks in the range of 310-370 nm and 380-440 nm respectively. Concentrations determined with
different peaks spanned a range of 5%.
4.2.2 Experimental setup for absolute rate constant measurement—LP/LIF method
The LIF apparatus has been shown in Figure 3.1 of Chapter 3. A pulsed XeF excimer laser
(GAM Laser Inc., EX100H) with energy of 10 mJ/pulse and repetition rate of 2 Hz was used to
generate CH3O• or CD3O• by photolyzing CH3ONO or CD3ONO at 351 nm. The resulting
methoxy radicals were probed at ~293 nm by the frequency-doubled (Inrad Autotracker III)
narrow band emission from a dye laser (Lambda Physik FL3002, 0.1 mJ/pulse), which was
pumped at 308 nm using the XeCl excimer laser (Lextra 100) operating at 2 Hz.
The two laser beams counter-propagated collinearly through the LIF cell. The photolysis
laser beam diameter was adjusted to be three times that of the probe laser beam. Red-shifted
emission from the radicals was collimated by two convex lenses (f=10 cm), passed through a
long-pass filter (>345 nm), and entered the photomultiplier tube (PMT, R212UH, Hamamatsu
Photonics) mounted on top of the cell window orthogonal to the laser beams. The signal from the
PMT was amplified (Ortec 9305) before being transmitted to a boxcar averager (SR250, Stanford
Research Systems, Inc.) and then to the computer data acquisition system (SR245 and SR272).
Simultaneously, the LIF signal was transmitted to an oscilloscope for real-time monitoring.
The time delay between the two laser pulses was controlled by a Digital Delay Generator
(DG 535). In order to avoid the scattered light signal from the probe laser, the gate of the boxcar
averager was opened 50 ns after the initial rise of the fluorescence signal. The gate width was 25
ns. Scattered light from the interaction of the probe beam with a prism was detected by
87
photodiode 2. The resulting signal was transmitted to Boxcar 2 to enable normalization of LIF
signal to the energy of the probe laser.
The LIF cell consisted of a jacketed 1.9 liter PyrexTM tube with an inner diameter of 57
mm. Two glass side arms with an inner diameter of 19.7 mm were installed at each side of the
cell. Quartz windows were attached at Brewster’s angle to the end of each arm. In order to
minimize scattered light inside the cell, two conical apertures were installed in the left arm to
reduce the scattered light. Both arms were painted black with Krylon ® black, and the Wood’s
horn on the bottom was coated with black AquadagTM. The temperature in the LIF-cell was
controlled between 250 and 335 K by flowing cooled ethanol or heated ethylene glycol through
the jacket and measured with a calibrated thermocouple thermometer (Digi-Sense® Dual Input J-
T-E-K®). The temperatures in the LIF region inside the reaction cell were measured for each
pressure with a thermocouple prior to a LIF experiment. This ensured accurate temperature
measurement. During the LIF experiment, the thermocouple was removed from the LIF region to
avoid it causing scattering from the probe laser. The temperature of the cooling and heating
liquid was regulated by a temperature bath (Neslab ULT-80).
̃2 A1 ← ̃
The fluorescence excitation (A X 2 E) spectra of CH3O• and CD3O• in the range of
285.7-302.7 nm at 296 K show clear progressions in the C—O stretching mode (ν3), which agree
well with the work of Inoue et al.47 We chose the peaks at 293.06 nm and 293.35 nm for
̃2 A1 , ν’3= 4)
monitoring CH3O• and CD3O•, respectively. In both isotopologues, these are the (A
̃ 2 E, ν”3=0) transitions, where ν3 is the C-O stretching mode.48 All the LIF signals were
← (X
corrected for their corresponding background signals, which were obtained by blocking the
photolysis beam while only passing the probe beam. The disappearance of methoxy radicals was
88
detected through the change of fluorescence intensity as a function of delay time between the
photolysis and probe beams. The shortest delay time used was 5 µs, with the maximum ranging
from 30 µs to 125 µs depending on the concentration of NO2. For each delay time, 100 to 120
laser shots were taken and averaged.
The initial methoxy radical concentration can be estimated from the absorption cross
section of methyl nitrite at 351 nm (~ 3×10-19 cm2 molecule-1),33 the quantum yield for the
formation of CH3O• (~1) at 351 nm,24,25 the photolysis laser fluence (~17 mJ pulse-1 cm-2) and
the methyl nitrite concentration (~3×1015 molecule cm-3). The resulting initial concentration of
CH3O• and CD3O• is ~2×1013 molecule cm-3 at 295 K. In order to keep the pseudo-first order
condition and avoid side reactions, we used a large excess of NO2 (0.9-6×1015 molecule cm-3) for
the kinetic experiments. The temperature range for both absolute and relative rate constant
measurements was 250-335 K, and the total pressure was 700 Torr. We also tested the pressure
dependence of the rate constant for methoxy + NO2 over the range of 30-700 Torr at 295 K. The
gas flow rate in the LIF cell is 3360 sccm, and its residence time 10 s at 295 K and 700 Torr.
4.3 Results and discussion
4.3.1 Absolute rate constants for CH3O• + NO2 and CD3O• + NO2
We first determined the absolute rate constant for CH3O• + NO2 (R4.2) at 295 K as a
function of pressure (30 Torr-700 Torr). Figure 4.1 shows a typical plot of the logarithm of LIF
intensity versus delay time between the photolysis laser pulse and the probe laser pulse for
several NO2 concentrations at 295 K and 700 Torr. Pseudo-first order reaction rate constants, k’,
89
were obtained from the slopes of linear-least squares fits to data at each NO2 concentration. The
slope of plots of k’ against NO2 concentration, determined by linear least-squares fitting,
provided the bimolecular reaction rate constants k2 and k4 at each temperature and pressure.
Figure 4.2 shows the plot of bimolecular reaction rate constant for each [NO2] as a function of
pressure at 295 K. The high linearity of the data shown in Figure 4.1 and the small and consistent
intercept in Figure 4.2 confirm the applicability of the pseudo-first order approximation to our
experiment and that complications from secondary chemistry are minimal.
The non-zero y-intercepts in Figure 4.2 can be interpreted as the sum of the loss rates of
CH3O• for all loss processes other than reaction with NO2. These processes include (1) diffusion,
and (2) reaction primarily with CH3ONO (k298=4.45×10-13 cm3 molecule-1 s-1),49 NO (k298 =
3.6×10-11 cm3 molecule-1 s-1),46 CH3O• (k298=1-4×10-11 cm3 molecule-1 s-1),50,51,52 and CH3ONO2
(k298=2.84×10-14 cm3 molecule-1 s-1).46
90
2.0
15 -3
[NO2] (×10 molecule cm )
1.5 0.84
1.67
2.79
1.0 4.18
5.57
Ln[Intensity]
0.5
0.0
-0.5
-1.0
-1.5
0 20 40 60 80 100
Time delay (s)
Figure 4.1 Typical linear decay of ln(LIF intensity) as a function of the delay time for CH3O•+NO2 at
total pressure 700 Torr and 295 K. NO2 concentrations in molecule cm-3 are: 9.2×1014, 1.84×1015,
3.06×1015, 4.58×1015, and 6.10×1015. Error bars are 2σ.
120000
700 Torr
500 Torr
100000
215 Torr
100 Torr
80000 50 Torr
30 Torr
k' (s )
60000
-1
40000
20000
0
15 15 15 15 15 15 15
0 1x10 2x10 3x10 4x10 5x10 6x10 7x10
-3
[NO2] (molecule cm )
Figure 4.2 Plot of k’ versus [NO2] at 295 K under different pressures. Error bars are 2σ precision of the
fitted slope of ln(intensity) versus time.
91
By constructing a plot of k2 and k4 versus concentration of bath gas (N2), the pressure
dependence of k2 and k4 was obtained. Figure 4.3 shows results for k2 along with literature
results at room temperature, which reveals a characteristic fall-off curve. Our results coincide
well with the results of Wollenhaupt et al. (10-200 Torr Ar)33 and Frost and Smith (6-125 Torr
Ar),31,32 which are the basis for the JPL Data Evaluation.46 Our results also agree well with the
extrapolated values at 500 Torr and 700 Torr from the above two studies (hollow symbols in
Figure 4.3). This is reasonable because the efficiencies of N2 and Ar as third bodies are usually
close.53,54,55 However, the rate constants of Martínez et al.34 using He bath gas are higher than
those of the other two studies and of the present work. Helium is a less efficient collider than Ar
or N2, consequently the rate constant in helium would be expected to be smaller than that in N2
and Ar.53,54,55 The agreement of our results with those of Wollenhaupt et al.33 and Frost and
Smith31,32 suggest some error in the results of Martínez et al.34
We also studied the pressure dependence of k4 (CD3O• + NO2), and results for both k2 and
k4 are plotted in Figure 4.4. One can see that the rate constant for the CD3O• isotopologue are
higher than those for CH3O•, but that rate constants appear to be converging at higher pressures.
Similar behavior has been observed for CH3O• and CD3O• in their reactions with NO.56 A more
complete survey of the joint pressure and temperature dependence of k2 and k4 will be presented
in Chapter 5.
92
-11
2.0x10
-11
1.8x10
kCH3O+NO2 (cm3 molecule -1 s-1)
-11
1.6x10 D
-11
M
1.4x10 E
-11 W
1.2x10 y
This work, M=N2
-11
1.0x10 Wollenhaupt et al. M=Ar C
-12
8.0x10 Martinez et al. M=He R
Wollenhaupt extrapolation
-12 P
6.0x10 Martinez extrapolation
Frost et al, M=Ar P
-12
4.0x10 Frost el al, M=He
-12 Frost et al., M=CF4
2.0x10
0.0
18 19 19 19 19
0.0 5.0x10 1.0x10 1.5x10 2.0x10 2.5x10
[M] (molecule cm-3)
Figure 4.3 Pressure dependence of k2 for CH3O• + NO2 at room temperature. Cited errors are 2σ of
precision in the fitted slopes of plots of k’ versus [NO2]. The black solid line is the fit of the Troe
expression to our results (see details below). All these data are listed in Appendix I.
93
-11
1.8x10
-11
1.6x10
kCH3O+NO2 (cm3 molecule -1 s-1)
-11
1.4x10
-11
1.2x10
-11
1.0x10 CH3O+NO2
CD3O+NO2
-12
8.0x10
-12
6.0x10
-12
4.0x10
-12
2.0x10
0.0
18 19 19 19 19
0.0 5.0x10 1.0x10 1.5x10 2.0x10 2.5x10
[M] (molecule cm-3)
Figure 4.4 Comparison of the pressure dependent behavior for rate constants for CH3O + NO2 (squares)
and CD3O + NO2 (triangles) at room temperature. Cited errors are 2σ statistical errors in the fitted slopes
of plots of k’ versus [NO2].
In order to obtain absolute rate constants k1 and k3 for methoxy + O2 as a function of
temperature, we measured k2 and k4 for methoxy + NO2 at each temperature that was used in the
relative rate measurements at NCAR. Table 4.1 lists all the rate constants over 250-333 K at 700
Torr. Rate constants k2 and k4 are the overall values directly measured by the LFP/LIF method.
The rate constant, k2b, for the disproportionation reaction29 is thought to be less than 2% of the
value of k2 measured in these experiments. There has been no study of the disproportionation
reaction for CD3O + NO2; however, it is reasonable to assume that deuterium substitution will
lower the disproportionation rate constant. In the analysis that follows, we assume that the rate
constant observed for the methoxy + NO2 reaction is equal to the rate of the association reaction
to form methyl nitrate.
94
Table 4.1 Rate constants for CH3O + NO2 at 700 Torr. Cited errors are statistical 2σ, and the 5%
uncertainty for [NO2] measurement is not included.
CH3O• + NO2 CD3O• + NO2
k2 (700 Torr) T(K),700Torr k4 (700 Torr)

T(K),700Torr cm3 molecule-1 s-1 cm3 molecule-1 s-1
×10-11 ×10-11
250 2.16±0.04 250 2.18±0.07
265 2.01±0.06
278 1.89±0.03 277 1.93±0.03
295 1.70±0.03 294 1.75±0.04
316 1.62±0.02 319 1.65±0.03
333 1.47±0.03 335 1.53±0.03

340 320 300 280 260 
-11
2.3x10
-11
2.2x10
kNO2(cm molecule s )
-11
-1 -1
2.1x10 CH3O+NO2
-11 CD3O+NO2
2.0x10
-11
1.9x10
3
-11
1.8x10

-11
1.7x10
-11
1.6x10
-11
1.5x10
0.0030 0.0032 0.0034 0.0036 0.0038 0.0040

-1
1/T (K )
Figure 4.5 Temperature dependence of methoxy + NO2 at 700 Torr. Error bars are statistical 2σ, and the
5% uncertainty for [NO2] measurement is not included.
95
Figure 4.5 shows the measured rate constant for both k2 and k4 as a function of temperature
at 700 Torr. By plotting ln(k) against 1/T, the Arrhenius expressions for the temperature
dependence of the rate constants for CH3O and CD3O are obtained from linear least squares
fitting as:
E4.1 𝑘2 = 4.86+0.40
−0.37 × 10
−12
𝑒𝑥𝑝[(374 ± 24)⁄𝑇 ]
E4.2 𝑘4 = 5.59+0.59
−0.53 × 10
−12
𝑒𝑥𝑝[(348 ± 29)/𝑇]
Both plots show slightly negative temperature dependencies of the rate constants. Values of k 2
are consistently slightly lower than k4, but this difference is not statistically significant. This can
be rationalized in that CD3ONO2 has a slightly larger density of states than the normal
isotopologue, making the decomposition of CD3ONO2 back to reactants slower.57,58
4.3.2 Rate constant for CH3O• + O2 and CD3O• + O2 and tunneling effect
By combining the rate constant ratios k1/k2 and k3/k4 determined in our group (shown in
Table 4.2) with the absolute rate constants k2 and k4 determined from this work, we can calculate
the absolute rate constants k1 and k3 at 700 Torr over the whole temperature range of our
experiment. Results of these calculations are shown in Table 4.3. The temperature dependence of
k1 and k3 at 700 Torr is plotted in Figure 4.6. The uncertainty in k1 and k3 arise from
uncertainties in both the relative rate study and the measurement of methoxy + NO2. The
following Arrhenius expressions are derived for k1 and k3:
E4.3 𝑘1 = 1.3+0.9
−0.5 × 10
−14
exp[−(663 ± 144)/𝑇]
E4.4 𝑘3 = 8.2+7.7
−4.0 × 10
−15
exp[−(974 ± 210)/𝑇]
96
The uncertainties above are 2 and include uncertainty in the methyl nitrate concentration
(10%), plus uncertainties of concentration measurement of NO2.
Table 4.2 The ratios of the rate constants for the CH3ONO+O2/NO2 (i.e.: k1/k2) and CD3ONO + O2/NO2
(i.e., k3/k4) experiments at all measured temperatures, 700 Torr. The error bars for all numbers
are 2σ.
T(K) 250 265 278 295 316 333
k1/k3(×105) 4.32±0.18 5.26±0.09 7.08±0.15 7.83±0.13 10.49±0.16 12.14±0.21
T(K) 277 294 319 335
k2/k4(×105) 1.26±0.05 1.70±0.08 2.32±0.06 2.92±0.10
Table 4.3 Absolute rate constant for k1 and k3 (unit: cm3 molecule-1 s-1) and kinetic isotope effect (KIE).
The quoted errors (2σ) include statistical uncertainties from linear fitting of both relative and
absolute rate methods, uncertainty in the methyl nitrate concentration (10%), plus uncertainties
of concentration measurement of NO2 from absolute rate method.
T(K) 250 265 278 295 316 333
k1(×1015) 0.940.11 1.060.12 1.340.15 1.330.15 1.700.19 1.790.20
T(K) 277 294 319 335
k3(×1016) 2.430.29 2.980.37 3.830.44 4.460.53
KIE (this
work) 5.500.48 4.460.40 4.440.35 4.000.34
KIE (Ref.
4.6 4.2 4.0 3.7 3.5 3.3
39)
97
340 320 300 280 260 T (K)
-15
2x10
-15
1.6x10
kO2 (cm molecule s )
-1
-15
1.2x10
-1
CH3O+O2
-16 CD3O+O2
3
4x10
3.0 3.2 3.4 3.6 3.8 4.0

-1
1000/T (K )
Figure 4.6 Temperature dependence of methoxy + O2 at 700 Torr. The solid lines represent linear least
squares fits to the data.
98
600 500 400 300 T (K)
Wantuck et al.
Lorentz et al.
-14
10 Gutman et al.
This work
k1 (cm molecule s )
-1
Orlando et al.
-1
Wiebe et al.*
Cox et al.*
3
-15
10
1.6 2.0 2.4 2.8 3.2 3.6 4.0

-1
1000/T (K )
Figure 4.7 Temperature dependent rate constant for CH3O+O2 in the range of 250 – 610 K. The solid line
represents the Arrhenius fit suggested in reference 20, Among the previous experimental data, the ones
from Wiebe et al.25 and Cox et al.24 at 298 K (denoted by*) were derived by combining originally
determined relative rate constant with the absolute rate constant for the reference reaction measured in
current work (R4.2) or elsewhere (CH3O•+NO).46
Both k1 and k3 show positive temperature dependencies. Note that the reaction methoxy +
O2 is bimolecular,37,38,39 and therefore the rate constant is expected to be independent of pressure.
Consequently, it is valid to compare our results with previous absolute rate measurements carried
out at lower pressures (<100Torr).4,5,6,20 Compared with the Arrhenius fitting of previous
experimental data on R4.1 over 298-610 K by Orlando et al.,20 the pre-exponential factor for k1
in this study (7.86+5.66

−3.29 ×10
-15
cm3 molecule-1 s-1) is one order of magnitude smaller, and the
activation energy (4.2 ± 1.3 kJ/mol) is approximately 40% smaller. By putting our results
together with previous experimental data, the temperature dependent k 1 over 250-610 K is
99
displayed in Figure 4.7. There is reasonable agreement between our results and the absolute
results from Lorentz et al.6 in the overlapping temperature range (298-333 K); however, our data
(below room temperature) exhibit less temperature dependence for k1, which is the source of our
lower activation energy and pre-exponential factor.
Kinetic isotope effects (KIEs) are reported in Table 4.3 for each temperature for which k1
and k3 were obtained. The KIE values obtained here (4.0-5.5) are far smaller than the value of
~100 determined in the one previous study.35 Note that the theoretical study of Reference 39,
also listed in Table 4.3, yielded slightly smaller KIEs and a more modest temperature
dependence of the KIEs than observed in the present experiments. Another way to express the
KIE is in Arrhenius form:
312±255
E4.5 𝑘1 /𝑘3 = (1.6+2.1
−0.9 ) ∙ exp( )
T
59,60,61
Several studies proposed that AH/AD <1.0 or <0.562 signal that tunneling is important.
Similarly, the difference in activation energy ED-EH >1.2 - 1.4 kcal/mol 59,60,61 is also proposed as
a criterion for tunneling. If both criteria are met, the tunneling coefficient, , for the normal
hydrogen is typically greater than ~5. 61

In our results, neither AH/AD (1.6+2.1
−0.9 ) nor ED-EH
(0.6±0.5 kcal/mol) suggest that  is nearly as high as 5. This is qualitatively consistent with the
conclusion from our measurements of the branching ratios for the reaction of CH2DO• + O2.40
An oddity of the theoretical findings in Reference 39 was that tunneling was computed to be of
similar importance for both the CH3O• + O2 and CD3O• + O2 reactions. For the normal
isotopologue,  rose from 1.9 to 3.4 as the temperature fell from 330 K to 250 K. Over the same
range,  in the deuterated isotopologue rose from 1.8 to 2.5.
100
Along with previous studies of the methoxy + NO2 reaction, our work suffers from
ignorance of the extent of formation of methyl peroxy nitrite (CH3OONO):
R4.2c CH3O• + NO2 (+M)  CH3OONO
This weakly bound species63,64,65 may dissociate on a timescale slower than that of the
LFP-LIF experiments, but certainly dissociates faster than the timescale of the chamber
experiments. While there is experimental evidence enabling quantification of the HOONO
formation channel in the HO + NO2 reaction,66 there is not even any direct experimental
evidence for the occurrence of R4.2c. Obviously, if formation of CH3OONO (or CD3OONO)
occurs to a significant extent, the Troe fitting to k2 and k4 determined in our LFP-LIF experiment
is being applied to the sum of two reactions, and extrapolation of that fitting beyond the range of
the experimental conditions could be misleading. In the analogous OH + NO2 reaction the
branching fraction for peroxy nitrate (HOONO) formation decreases with increasing temperature
(at 700 Torr).46 If this trend holds for R4.2, and if methyl peroxy nitrite is stable on the timescale
of our LFP-LIF experiments, then R4.2c would contribute more error in our determination of k1
at the lower temperatures used in our experiments than at the higher temperatures. As our results
for k1 disagree with previous work more at high temperatures than at low temperatures, R4.2c
appears unlikely to be a major factor confounding our results.
4.4 Conclusion
Our relative rate measurement for CH3O• + O2/NO2 and absolute rate measurement for
CH3O• + NO2 combined together have enabled us to determine the absolute rate constant (k1) for
101
CH3O• + O2 in the temperature range of 250 – 333 K and at 700 Torr. This enabled us to carry
out the first determination of k1 below room temperature, and the first pressure or temperature
dependent study of k1 in N2. These data are thus of greater relevance to the atmosphere than
previously reported values of k1. Our results show reasonable agreement with previous absolute
rate studies in the overlapping temperature range (298-333 K). However, they exhibit less
temperature dependence for k1 than previous data obtained at higher temperatures.
By carrying out the same experiments for the isotopologue CD3O• in the temperature range
of 277 – 335 K, we have been able to determine the kinetic isotope effect (kH/kD = k1/k3) for
methoxy + O2. The measured KIEs do not seem greatly affected by tunneling. The KIEs
reported here are similar to, if slightly higher than, those computed from theory.39
We hope that the experimentally determined k1(T) for methoxy + O2 will be helpful to
further validate a computational method for this reaction. A validated method would, hopefully,
enable reliable and affordable computational studies for RO• + O2 reactions for larger and
functionalized alkoxy radicals derived from atmospherically important compounds such as
isoprene or oxygenated VOCs. This would, in turn, enable one to extract absolute rate constants
for decomposition and isomerization of these alkoxy radicals from relative rate experiments.
These absolute rate constants are necessary to build structure-reactivity relations for
unimolecular reactions of the broad array of functionalized alkoxy radicals that are difficult to
study experimentally.
102
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Molecule. Ph.D dissertation, State University of New York College of Environmental Science and
Forestry, April 2013
6 Lorentz, K.; Rhasa, D.; Zellner, R.; Fritz, B. Ber. Bunsenges. Phys. Chem. 1985, 89, 341.
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17 Wu, F. and Carr, R. W. J. Phys. Chem. A 2001, 105, 1423-1432.
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19 Wu, F. and Carr, R. W. J. Phys. Chem. 1996, 100, 9352-9359.
23 Batt, L.; Robinson, G. Int J Chem Kinet 1979, 11, 1045-1053.
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27 Mendenhall, G. D.; Golden, D. M.; Benson, S. W. Int. J. Chem. Kinet. 1975, 7, 725-737.
29 McCaulley, J.; Anderson, S.; Jeffries, J.; Kaufman, F. Chem. Phys. Lett. 1985, 115, 180-186.
30 Biggs, P.; Canosa-Mas, C.E.; Fracheboud, J.M.; Parr, A.D.; Shallcross, D.E.; Wayne, R. P.; Caralp, F.
A. J. Chem. Soc., Faraday Trans. 1993, 89, 4163-4169.
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40 Hu, H.; Dibble, T. S.; Tyndall, G. S.; Orlando, J. J. J. Phys. Chem. A, 2012, 116, 6295-6302.
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42 Taylor, W.; Allston, T.; Moscato, M.; Fazekas, G.; Kozlowski, R.; Takacs, G. Int. J. Chem. Kinet.
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44 Calvert, J. G.; Pitts, J. N. Photochemistry; Wiley: New York, 1966, 455ff.
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46 Sander, S. P.; Abbatt, J.; Barker, J. R.; Burkholder, J. B.; Friedl, R. R.; Golden, D. M.; Huie, R. E.;
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50 Biggs, P.; Canosa-Mas, C. E.; Fracheboud, J.; Shallcross, D. E.; Wayne, R. P. J. Chem. Soc., Faraday
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51 Meier, U.; Grotheer, H.; Riekert, G.; Just, T. Ber. Bunsenges. Phys. Chem. 1985, 89, 325-327
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106
Chapter 5. Pressure dependence and kinetic isotope effects in the absolute
rate constant for methoxy radical reacting with NO2
5.1 Introduction
Methoxy radical (CH3O•) is a very important intermediate in the oxidation of methane (and
other alkane VOCs) for both combustion chemistry1 and atmospheric chemistry.2 In the
atmosphere, it is primarily produced from the reaction transferring an oxygen atom from methyl
peroxy radical (CH3O2•) to NO, forming NO2 as a co-product. The major atmospheric fate of
methoxy radical is reaction with O2; however, reaction with NO2 may be important in a power
plant plume.3 Methoxy radical (CH3O•) is the prototype for all alkoxy radicals, whose fates
greatly impact ozone formation and gas-particle partitioning of the eventual stable products of
oxidation of volatile organic compounds.4,5 Similarly, the methoxy + NO2 reaction (R5.1):
k1a
k1b
R5.1b → HCHO + HONO (minor)
is the prototype for kinetic and mechanistic studies of other RO• + NO2 reactions. Early relative
rate studies of the methoxy+O2 reaction had employed R5.1 as the reference reaction.6,7,8 Chapter
4 report similar relative rate studies together with the absolute rate measurement of R5.1 (and
CD3O• + NO2) to determine (for the first time at near-ambient pressure) the temperature
dependence of the absolute rate constant of methoxy+O2.
107
Direct kinetic investigations of R5.1 (using LIF detection of CH3O•) have been carried out
previously at pressures ranging from 0.6 to 600 Torr over the temperatures range 220-473 K,
with Ar, CF4 or He as bath gases. The discharge-flow method was employed to produce CH3O•
for low-pressure studies by McCaulley et al. (0.6-5 Torr)9 and Biggs et al. (1-10 Torr),10 while
pulsed laser photolysis was used to generate CH3O• for higher pressure studies by Frost and
Smith (6-125 Torr),11,12 Wollenhaupt et al. (10-200 Torr),13 and Martínez et al. (50-600 Torr). 14
All of these studies showed similar pressure-dependent behavior of k1, except Martínez et al.,14
whose values are ~30% larger than the rest over the whole pressure range at room temperature.
R5.1a proceeds via formation of an activated methylnitrate (CH3ONO2) intermediate, which is
quenched by bath gas in competition with dissociation to reactants. Martínez et al.14 used He as a
bath gas, and, given that He is expected to be less effective in deactivating the energized
complex CH3ONO2* than Ar and CF4,15,16,17 the rate constants determined by Martínez et al.14
would be expected to smaller than those determined in the other studies at the same pressure and
temperature. Thus, the results of Martínez et al.14 are an anomaly. Also, among the two studies
with largest pressure range, Martínez et al.14 showed a much larger temperature dependence of k0
and a slightly smaller temperature dependence of k∞ than Wollenhaupt et al.13 It is noteworthy
that R5.1a appears not to have reached the high-pressure limit at 600 Torr of Ar or He. There
have been no direct kinetic measurements on R5.1 in the presence of N2 bath gas except at 50
Torr, so previous experiments did not mimic atmospheric conditions. For these reasons, it is
valuable to reinvestigate k1 at pressures closer to one atmosphere and especially to use N2 as a
bath gas, since it is more representative of air than Ar, CF4, or He.
Reaction between CH3O• and NO2 is known to proceed via two channels—recombination
(R5.1a) producing the methyl nitrate (CH3ONO2) and disproportionation (R5.1b) yielding
108
formaldehyde and nitrous acid. The disproportionation is found to be a minor process, and only
play an important role at rather low-pressures (< 1 Torr).9,10 McCaulley et al. reported a rate
constant (independent of pressure) of 9.6+17.3

−2.7 × 10
−12
exp[−(1150+550 3 -1 -1
−170 )] cm molecule s for
the disproportionation channel.9
There have been several theoretical investigations of the energetics and mechanisms of
R5.1a and R5.1b.18,19,20,21 A composite of these energetics is shown in Figure 5.1. Of these
studies, only Lesar et al.19 found the disproportionation barrier to be higher (by 2-6 kcal/mol
depending on levels of theory) than the reactant energy, yet this seems more consistent than other
reports with the pressure-independent rate constant (k1b) reported by McCaulley et al.9 While all
the three theoretical studies were able to find the transition state for the disproportionation of
methyl nitrate, only Pan et al.18 reported trying (they failed) to find a transition state for a direct
hydrogen atom-abstraction mechanism for R5.1b.
Figure 5.1 Simplified energy diagram for reaction CH3O• + NO2.
109
The reaction of methoxy radical with NO2 is a radical-radical recombination reaction, and
is analogous to the very important reaction OH + NO2.22 Barker and coworkers took advantage
of known pressure-dependent rate constants at 297 K for various NO2 addition reactions to
constrain RRKM/Master Equation simulations of these reactions.23,24 This enabled them to
estimate the collisional energy transfer parameters (α) by fitting to the fall-off curves. Although
the fitted rate constants agreed well with the experimental values, Barker and coworkers
concluded that the accuracy of fitting the collisional energy transfer parameter is subject to large
uncertainties in assumptions for treating the transition state, which is not a well-defined saddle
point on the potential energy surface. They suggested that an experimental reinvestigation of the
kinetics of these reactions in extended pressure ranges would enable validation or refinement of
their fitted values of α.
In this work, we present the experimental kinetics of R5.1 in the temperature range of 250-
335 K and pressure range of 30-700 Torr of N2. We also investigated the kinetics of the
perdeuterated isotopologue of CH3O•, i.e. CD3O• + NO2 (R5.2). This offers a check on the
kinetic results of R5.1 and may facilitate further investigations of collisional energy transfer in
the methoxy + NO2 reaction.
k2a
R5.2a CD3O• + NO2 (+M) → CD3ONO2
k2b
R5.2b → DCDO + DONO (minor)
110
5.2 Experimental
Many of the experimental details appear in Chapter 4. Methyl nitrites (both CH3ONO and
CD3ONO) were synthesized from the reaction of methyl alcohol with saturated aqueous NaNO2
solution upon addition of concentrated sulfuric acid at 0°C.25 The gaseous products were passed
over NaOH solution and anhydrous CaCl2 before being collected in a dry-ice trap. The isolated
product was purified by freeze-pump-thaw distillation, resulting in a pale yellow solid, which
was stored at ~196°C until it was used. Vapor of the purified methyl nitrite was characterized by
FTIR26 to verify the absence of residual alcohol and ultraviolet (UV) spectroscopy to verify its
identity.27 A 10-L blackened glass bulb was filled with approximately 3 Torr of methyl nitrite,
which was then diluted by N2 (Haun Welding Supply, 99.999%) to a total pressure of 1000 Torr.
NO2 was produced by mixing NO (American Gas group, >99.5%) with excess of O2 (MG
Welding Products 99.999%), followed by purification by free-pump-thaw distillation at -196°C
until a pure white solid was obtained. FTIR and UV-visible spectroscopy were used to verify the
identity and purity of the NO2. NO2 was diluted with N2 to a mole fraction of about 0.02 in a
second glass bulb. Concentrations of methyl nitrite and NO2 in the bulbs were determined using
UV-visible absorption cross sections28,29 at each of six wavelengths in the range of 310-370 nm
and 380-440 nm respectively, with an uncertainty (2σ) of 5%.
The LIF apparatus is shown in Figure 3.1 of Chapter 3. Methoxy radicals were generated
by photolyzing methylnitrite at 351 nm with a XeF excimer laser (GAM Laser Inc., EX100H)
possessing a pulse energy of 10 mJ and repetition rate of 2 Hz. The resulting methoxy radicals
111
were probed at ~293 nm (0.1 mJ/pulse) by frequency-doubled (Inrad Autotracker III) narrow
band emission from dye laser (Lambda Physik FL3002). The laser repetition rate was 2 Hz.
The probe and photolysis laser beams counter-propagate collinearly through the reactor.
The size of the photolysis laser beam was adjusted to be three times that of the probe laser beam
̃2 A1 , ν’3= 4) ← (X
(3mm×6mm). The (A ̃ 2 E, ν”3=0) transition at 293.06 nm (CH3O•) and 293.35
nm (CD3O•) were used for probing the methoxy radicals.30,31 Emission from radicals excited by
the probe beam was detected orthogonal to the laser beams by a photomultiplier tube mounted
over a window of the cell. Prior to reaching the PMT, the emission was collimated using two
f=10 cm convex lenses and passed through a long-pass filter (>345 nm). The PMT signal was
transmitted to a boxcar averager and then to computer data acquisition system. The gate (25 ns)
of the boxcar averager was opened 50 ns after the initial rise of the fluorescence pulse. The
fluorescence intensity was normalized to the energy of the probe laser for each laser shot. 100 to
120 laser shots were averaged for each delay time. A digital delay generator was used to control
the delay time between the two laser pulses. The shortest delay time was 5 µs, and delay times
extended to 30 µs to 260 µs depending on the experimental conditions. The change of
fluorescence intensity as a function of delay time was used to determine the rate of the
disappearance of methoxy radicals.
The LIF cell is made of jacketed and insulated 1.9 liter Pyrex TM tube with two glass side
arms. The end of each arm has a quartz window attached at Brewster’s angle. The temperature
inside the LIF cell could be controlled between 250 and 335 K. The temperature in the cell was
measured in the LIF region at each pressure and temperature with a calibrated thermocouple
thermometer. The flow rate of each gas was regulated by pressure-based Mass-FLO® controllers.
112
The reaction pressure in the cell was measured, close to where fluorescence was being imaged on
the PMT, by absolute pressure transducers.
For a typical concentration of methyl nitrite (~3×1015 molecule cm-3) and photolysis laser
fluence (~17 mJ pulse-1 cm-2), one pulse of the photolysis laser can produce approximately
2×1013 molecule cm-3 methoxy radicals (CH3O• or CD3O•) at 295 K. This estimate is based on
the cross section of CH3ONO (~ 3×10-19 cm-2),13 the quantum yield for the formation of CH3O•
(~1) at the photolysis wavelength 351 nm.8,32 NO2 concentrations were 0.9-6×1015 molecule
cm-3, sufficient to ensure create pseudo-first order conditions. Pseudo-first order rate constants
were measured over the range 250-335 K and 30-700 Torr.
5.3 Results and discussion
Under pseudo-first order condition, the temporal profile for the disappearance of methoxy
radicals should be well expressed by a single-exponential decay, e.g., for CH3O•:
[𝐶𝐻 𝑂]
E5.1 𝑙𝑛 ([𝐶𝐻3𝑂] 𝑡 ) = −𝑘′𝑡
3 0
Where k’ is the pseudo-first order rate constant for loss of methoxy radical. Under pseudo-first
order condition, methoxy radical predominantly reacts with NO2; other loss processes include
side reactions and diffusion. Hence, k’ can be described by:
E5.2 𝑘 ′ = 𝑘1 [𝑁𝑂2 ] + 𝑘𝑜𝑡ℎ𝑒𝑟 𝑙𝑜𝑠𝑠
The concentration of NO2 was treated carefully due to the dimerization of NO2 into N2O4,
especially at low temperatures. The recommended equilibrium constant29 as a function of
113
temperature (Equation E5.3) was used to correct the concentration of NO2, and all the data shown
below reflect this correction:
[𝑁 𝑂4 ] 6643
E5.3 2
𝐾2𝑁𝑂2→𝑁2 𝑂4 = [𝑁𝑂 2 = 5.9 × 10−29 𝑒𝑥𝑝 ( ) 𝑐𝑚−3
2] 𝑇
As indicated from Equation E5.3, the extent of the conversion of NO2 to N2O4 increases as the
temperature gets lower and the initial NO2 concentration increases. The highest dimerization
fraction is ~14 % (for the highest NO2 concentration at 250 K). The modest size of this
correction ensures that dimerization causes minimal error in the determination of k1 and k2.
Figure 5.2 shows typical plots of the logarithm of LIF intensity versus delay time between
the photolysis laser pulse and probe laser pulse at several NO2 concentrations. The pseudo-first
order rate constants, k’, for loss of CH3O• were obtained from the absolute value of the slope of
the linear least squares fit to the data at each [NO2]. Values of k’ for each experimental condition
are listed in Appendix I. Figure 5.3 plots of k’ versus [NO2] at six different total pressures at 250
K, along with linear-least squares fits to the data. The slopes of these fits correspond to the
effective bimolecular reaction rate constants (k1) for the specified temperatures and total
pressures.
114
2.0
15 -3
1.5 [NO2] (10 molecule cm )
0.63
1.19
1.0 1.89
2.73
Ln[Intensity]
3.52
0.5
0.0
-0.5
-1.0
0 20 40 60 80 100
Delay Time (s)
Figure 5.2 Typical linear decays of ln(LIF intensity) as a function of the delay time for CH 3O•+NO2 at
total pressure 700 Torr and 250 K. NO2 concentrations in molecule cm-3 are: 6.3×1014, 1.19×1015,
1.89×1015, 2.73×1015, and 3.52×1015.
115
30 Torr
8.0x10
4 50 Torr
100 Torr
500 Torr
4
700 Torr
6.0x10
k' (s )
-1
4
4.0x10
4
2.0x10
0.0
15 15 15 15
0 1x10 2x10 3x10 4x10
-3
[NO2] (molecule cm )
Figure 5.3 Plot of k’ for CH3O• + NO2 versus [NO2] at 250 K under different pressures. Error bars are 2σ
of the fitted slope of ln(intensity) versus time.
The data shown in Figure 5.2 and Figure 5.3 exhibit the high linearity expected from
pseudo-first order conditions. The non-zero y-intercepts in Figure 5.3 can be interpreted as the
sum of the loss rates of CH3O• for all loss processes other than reaction with NO2. These loss
processes include diffusion and side reactions with CH3ONO (k=4.45×10-13 cm3 molecule-1 s-
1 33
), CH3ONO2 (k298=2.84×10-14 cm3 molecule-1 s-1),29 NO (k=3.6×10-11 cm3 molecule-1 s-1)29 and
CH3O• itself (k=1 - 4×10-11 cm3 molecule-1 s-1) 34,35,36 However, the consistently small intercept
compared with the values of k’ exclude a dominant loss by these side process. In addition, the y-
intercepts in Figure 5.3 range tend to increase with increasing pressure (from 1100 s-1 at 50 Torr
to 6000 s-1 at 700 Torr). This might be explained by the longer residence time of gas in the cell
under higher pressure. Pressures were increased by a combination of increasing the flow rate of
116
N2 buffer gas and increasing the residence time of gas in the cell. Residence time ranges from
10-25 s. The longer residence time allows more reaction product (CH3ONO2) to accumulate, so
there is more loss of CH3O• by reaction with CH3ONO2. We also tested the validity of the
measured absolute rate constants by reducing the fluence of photolysis laser (by 40% at 295 K
and 30 Torr) as well as varying the total gas flow rate while holding total pressure constant.
Neither of these two factors was found to affect the measured rate constants by more than 2%.
Table 5.1 Pressure dependent rate constants k1 and at k2 different temperatures. Error bars represent 2σ
k1 (×10-11 cm3 molecule s-1)
Pressure 250 K 265 K 278 K 295 K 316 K 333 K
700 2.16±0.12 2.01±0.12 1.89±0.10 1.70±0.09 1.62±0.08 1.47±0.08
500 2.00±0.11 1.89±0.10 1.77±0.09 1.63±0.09 1.54±0.08 1.44±0.08
215 -- 1.62±0.09 -- 1.52±0.09 1.32±0.07 1.15±0.06
100 1.54±0.08 -- 1.36±0.07 1.23±0.07 1.17±0.06 1.04±0.06
50 1.38±0.08 1.23±0.07 1.15±0.06 1.08±0.06 1.00±0.05 0.89±0.05
30 1.28±0.07 1.11±0.06 0.97±0.05 0.90±0.05 0.81±0.04 0.48±0.03
117
k2 (×10-11 cm3 molecule s-1)
Pressure 250 K 277 K 294 K 319 K 335 K
700 2.18±0.13 1.93±0.10 1.75±0.10 1.65±0.09 1.53±0.08
500 2.14±0.12 1.84±0.10 1.71±0.09 1.55±0.08 1.43±0.08
215 2.06±0.11 -- 1.53±0.08 -- 1.29±0.07
100 1.91±0.10 1.60±0.08 1.45±0.08 1.26±0.07 1.19±0.06
50 1.83±0.10 1.52±0.08 1.24±0.07 1.15±0.06 1.07±0.06
30 1.72±0.09 1.40±0.07 1.15±0.06 0.98±0.05 0.86±0.05
Measured absolute rate constants k1 and k2 at different temperatures and pressures over the
range 250 to 335 K and 30-700 Torr are listed in Table 5.1. As discussed previously, two
channels are known to exist for the reaction CH3O• + NO2: recombination and
disproportionation. Therefore, the measured absolute rate constants k1 and k2 are the sum of both
channels. By comparing the value of k1b reported by the Arrhenius expression of McCaulley et
al.9 with the overall rate constant k1 measured from our experiments, we find the contribution of
k1b to k1 is quite modest. The ratio k1b/k1 increases with increasing temperature and decreasing
pressure, and hence k1b/k1 ranges from ~0.4% at 250 K and 700 Torr to ~6% at 333 K and 30
Torr. Therefore, we approximate k1a as equal to the measured k1, so that all data is included
when calculating the fall-off curve. There is no available study of k2b for CD3O• + NO2 (R5.2b),
however, the reaction mechanism is expected to be similar to CH3O• + NO2 reaction. Thus we
assumed that the disproportionation channel for R5.2 could also be neglected, i.e. k2a=k2. The
effects of neglecting channel b will be discussed in more detail below.
We plot the pressure-dependent rate constants k1a and k2a at all temperatures in Figure 5.4
and Figure 5.5 respectively. The solid lines result from non-linear fitting to falloff curves,
described by Troe expression:16
118
𝑘 0 (𝑇)[𝑀]
E5.4 𝑘([𝑀], 𝑇) = (1+𝑘 0 (𝑇)[𝑀]/𝑘 ∞(𝑇)) 𝐹𝑐𝑒𝑛𝑡 𝑝
2 −1
𝑘 0 (𝑇)[𝑀]⁄
𝑝 = (1 + (𝑙𝑜𝑔10 ( 𝑘 ∞ (𝑇))) )
where k0 is the termolecular rate constant in the low-pressure limit and k∞ is the bimolecular rate
constant in the high-pressure limit. Values of k∞ and k0 are presented in
Table 5.2. Fcent is a parameter that describes broadening of the fall-off curve, which results from
the energy and angular momentum dependencies of k.16,37 A fixed value of Fcent equal to 0.6 is
recommended by Troe in the range of 100-400 K,17 and this recommendation is adopted by the
NASA Panel for Data Evaluation (although not by the IUPAC Subcommittee for Gas Kinetic
Data Evaluation,38 which allows Fcent to vary when fitting experimental data). Golden39 pointed
out that both the NASA and IUPAC formulations are adequate to represent pressure dependent
rate constant as long as the range of temperature and/or pressure is not too wide. Wollenhaupt et
al.13 attempted fitting their data using E5.4 by taking Fcent as a variable, and it was found that Fcent
was ~0.6 and insensitive to temperature in their temperature range (233 – 356 K).
119
-11
2.4x10
-11
2.0x10
kCH3O+NO2 (cm molecule s )
-1
-11
1.6x10
-11
1.2x10
3
-11
-12 1.2x10
8.0x10
250 K
265 K -12
8.0x10 250 K
-12 265 K
4.0x10 278 K 278 K
295 K
295 K -12
4.0x10
316 K
333 K
316 K
0.0
333 K 0.0 8.0x10
17 18
1.6x10
19 19 19
0 1x10 2x10 3x10
-3
[N2] (molecule cm )
Figure 5.4 Fall-off curve non-linear fitting of absolute rate constant k1a measured for CH3O• + NO2 in this
study to equation E5.4. T=250-333 K, P=30-700 Torr. Error bars represent 2σ statistical error propagated
with 5% uncertainty in [NO2] concentration determination. The insertion demonstrates the magnification
of the low pressure data.
.
120
-11
2.4x10
-11
2.0x10
kCD3O+NO2 (cm molecule s )
-1
-11
-1
1.6x10
-11
1.2x10
-3
-11
1.6x10
-12
8.0x10
250 K
B
277 K E
H
-12
4.0x10 294 K -12
8.0x10 K
N
319 K
335 K 0.0 17
8.0x10
18
1.6x10
0.0
19 19 19
0 1x10 2x10 3x10
-3
[N2] (molecule cm )
Figure 5.5 Fall-off curve non-linear fitting of absolute rate constant k2a measured for CD3O• + NO2 in this
study to equation E5.4. T=250-335 K, P=30-700 Torr. Error bars represent 2σ statistical error propagated
with 5% uncertainty in [NO2] concentration determination. The insertion demonstrates the magnification
of the low pressure data.
121
Table 5.2 High pressure and low pressure limit rate constants resulted from fits of our data to Troe
expression (E5.4). Quoted errors are 2σ. Unit: k0—cm-6 molecule-2 s-1; k∞—cm-3 molecule-1 s-1
CH3O• + NO2 CD3O• + NO2
T(K) k0 (×10-29 ) k∞(×10-11) T(K) k0 (×10-29) k∞(×10-11)
250 (6.32±1.66) (2.35±0.11) 250 (24.2±4.02) (2.32±0.03)
265 (5.45±1.24) (2.21±0.09)
278 (4.59±0.52) (2.14±0.05) 277 (15.4±3.54) (2.05±0.05)
295 (4.42±0.66) (1.96±0.05) 294 (9.09±1.46) (1.91±0.04)
316 (4.15±0.82) (1.84±0.06) 319 (7.30±1.19) (1.81±0.05)
333 (3.46±0.94) (1.70±0.09) 335 (6.99±1.92) (1.66±0.05)
333a (2.47±0.68) (1.75±0.12)

a
fall-off curve fitting of pure recombination rate constant k1a, each data point being obtained by
subtracting k1b (calculated by the Arrhenius expression of McCaulley et al.9) from measured k1 at 333 K
in this study.
Disregarding the disproportionation channel is expected to negligibly affect the fitting of
the fall-off curve. The largest effect of k1b on k1 occurs at 333 K, where k1b/k1 is largest (~2-6%).
For this purpose, we computed k1b from the Arrhenius expression of McCaulley et al. 9 and
subtracted that value from k1 to get k1a. This value of k1a was fitted to the Troe equation (E5.4).
The resulting parameters are listed in
Table 5.2 along with the values computed neglecting reaction R5.1b. Accounting for k1b
reduces k10 (333 K) by 29%, but increases k1∞ (333 K) by only 3%. The large uncertainty in k1b 9
122
and its limited effect of k1 makes us unwilling to apply this correction to calculating the Troe
parameters.
A negative temperature dependence is expected for recombination reactions, and such a
temperature dependence can be observed from Figure 5.4 and Figure 5.5, as well as Table 5.1
0
and Table 5.2 above. k0 decreases by a factor of 2.6 (k1a ) or 3.5 (k 02a ) as the temperature is raised
from 250 K to 333 or 335 K. The temperature dependence of k∞ is much smaller over the same
temperature range, and only decreases by ~20%. As is commonly done,13,14,29 we fit the
T
temperature dependencies of k0 and k∞ to power law expressions: k 0 = k 0298K × (298)−n and
T
k∞ = k∞
298K × (298)
−m
, respectively. Figure 5.6 and Figure 5.7 show log-log plots of k versus
temperature for k0 and k∞, respectively; each figure shows data for both CH3O• and CD3O•.
0
k1a
0
k2a
k (cm molecule s )
-1
-28
1.0x10
-2
6
240 260 280 300 320 340

Temperature (K)
123
0
Figure 5.6 Temperature dependence of low pressure limit rate constant k1a and k 02a for CH3O• + NO2 and
T
CH3O• + NO2 respectively. The fitting is based on the equation k 0 = k 0298K × (298)−n . Error bars are 2σ
statistical error from non-linear fitting.
-11
2.4x10 k1a
k2a
-11
2.2x10
k (cm molecule s )
-1
-1
-11
2x10
3
-11
1.8x10
-11
1.6x10
240 260 280 300 320 340
Temperature (K)
∞
Figure 5.7 Temperature dependence of high pressure limit rate constant k1a and k ∞
2a for CH3O• + NO2 and
T
CD3O• + NO2 respectively. The fitting is based on the equation k ∞ = k ∞
298K × (298)
−m
. Error bars are 2σ
statistical error from non-linear fitting.
Linear fitting of log(k) versus log(T) leads to expressions for k0 and k∞, which are shown in
E5.5 – E5.8. The intercept of each fit gives the corresponding low- and high-pressure limiting
rate constants at 298 K, while slopes of each fit indicate the temperature dependent factors n and
m.
0
E5.5 k1a = 4.29+0.40
−0.37 × 10
−29
(T/298)−(1.65±1.11) cm6 molecule-2 s-1
∞
E5.6 k1a = (1.95 ± 0.03) × 10−11 (T/298)−(1.13±0.18) cm3 molecule-1 s-1
124
E5.7 k 02a = 9.97+1.00
−0.91 × 10
−29
(T/298)−(4.79±0.92) cm6 molecule-2 s-1
E5.8 k∞
2a = (1.91 ± 0.02) × 10
−11
(T/298)−(1.11±0.09) cm3 molecule-1 s-1
Figure 5.6 illustrates that the extent of the temperature dependence of k 02a (n = 4.79 ± 0.92)
0
is greater than for k1a (n = 1.65 ± 1.11), while, as shown in Figure 5.7, the temperature
∞
dependencies of k1a and k ∞
2a (m = 1.13±0.18 and 1.11±0.09, respectively) are similar and both
are much smaller than that of the corresponding low-pressure limiting rate constants. k ∞
2a is
∞
slightly larger than k1a ; however, a t-test of the data in Figure 5.7 indicates no significant
0 0
difference exist between k ∞ ∞
2a and k1a . While k 2a is consistently larger than k1a in the
temperature range in our study, the difference between k 02a and k1a
0
decreases with increasing
temperature. The ratio of rate constants for two isotopologues, kH/kD, defines the kinetic isotope
effect (KIE). Values of KIE at each temperature are presented in Table 5.3 and Figure 5.8 for
both k0 and k∞. As implied above, the KIE for k0 is quite temperature dependent, ranging from
0.26 to 0.57. However, as indicated in Figure 5.8, uncertainties for these values (20 - 40%) are
large. Linear regression of log(KIE0) versus log(T) leads to a slope of 3.2±1.7, and the
uncertainty is too large to conclude that a significant temperature dependence exists in the KIE
for k0. By contrast, the KIE for k∞ (1.01 - 1.04) is equal to 1.00 within the noise and shows no
hint of any temperature dependence. Similar temperature dependencies of k0 and k∞ have been
observed for the reactions CH3O•/CD3O• + NO.40
125
Table 5.3 KIE values (kH/kD) at different temperatures for k0, and k∞. Cited errors are 2σ.
T (K) 250 278 295 316 333
KIE for k0 0.26±0.08 0.30±0.08 0.49±0.11 0.57±0.15 0.50±0.19
KIE for k∞ 1.01±0.05 1.04±0.03 1.02±0.03 1.02±0.04 1.03±0.06
1.2
1.0
KIE at low pressure limit

0.8
KIE at high pressure limit
KIE value
0.6
0.4
0.2
0.0
240 260 280 300 320 340
Temperature (K)
Figure 5.8 Temperature dependent KIE for k0 and k∞. Cited errors are 2σ.
The KIE for methoxy + NO2 can be interpreted using RRKM theory. The energy
dependent rate constant for a unimolecular reaction can be expressed as
1 𝐺 ≠ (𝐸−𝐸0 )
E5.9 𝑘(𝐸) = ℎ 𝜌(𝐸)
126
where h is Planck’s constant, G≠ (E − E0 ) is the sum of states of the transition state, and ρ(E) is
the density of states (number of states per energy interval) of the reactant. Here, the reactant is
the energized methylnitrate molecule (CH3ONO2* or CD3ONO2*) formed from methoxy + NO2,
and the reaction is methylnitrate dissociation back to methoxy + NO2. Since the dissociation of
methylnitrate has no intrinsic barrier other than the endothermicity of the dissociation reaction,
there is no well-defined saddle point. Instead, there exist transition states (maxima in Gibbs Free
Energy) whose positions along the reaction coordinate (CH3O-NO2 distance) vary with energy.
As the energized methylnitrate molecules are formed in the thermal reaction of methoxy + NO2,
the energized molecules probably mostly possess energies that are no more than a few kBT above
the energy of methoxy + NO2. The higher mass of D than H lowers the frequencies of several
vibrational modes of CD3ONO2, and so CD3ONO2 possesses a larger density of states than does
CH3ONO2 at the same energy. At the same time, the sum of states for the transition state does
not change significantly, because the most of the vibrational modes affected by isotopic
substitution have frequencies whose energies correspond to more than a few k BT. Following
E5.9, this larger density of states for CD3ONO2 than CH3ONO2 leads to a lower rate constant for
dissociation back to methoxy + NO2 for CD3ONO2* than CH3ONO2*. This higher density of
states may also increase the rate constant for collisional deactivation of CD 3ONO2* over that for
CH3ONO2*. These two effects result in a larger probability of collisional deactivation for
CD3ONO2* than for CH3ONO2*. This explains the difference of rate constants especially in the
condition of low-pressure limit (k 02a > k1a

0
), shown in Figure 5.6.
At the high-pressure limit, all the energized methylnitrate molecules will be quenched
rather than dissociating to reactants, and thus the first step, association of radicals to energized
complex, becomes the rate limiting step in R5.1a and R5.2a. The rates of the association step are
127
expected to be very nearly the same for both isotopologues, and this explains the similarity of
∞
k1a and k ∞
2a at all temperatures.
Table 5.4. Troe parameters resulting from different studies for the methoxy + NO2 reaction. Cited errors
are 2σ. Units for k0 and k∞ are cm6 molecule-2 s-1 and cm3 molecule-1 s-1 respectively.
Troe Parameters Experimental Condition
k0 (298 K) k∞(298K) P range T range

n m M Reference
(×10-29) (×10-11) (Torr) (K)
4.29±0.4 1.65±1.11 1.95±0.03 1.13±0.08 N2 30-700 250-333 this work
this work
9.97±1.0 4.79±0.92 1.91±0.02 1.11±0.09 N2 30-700 250-335
(CD3O)
5.22±0.3 4.36±0.36 1.84±0.05 1.87±0.15 Ar 10-200 233-356 Wollenhaupt13
3.98±0.05 1.74±0.4 2.41±0.02 0.88±0.32 He 50-600 250-390 Martínez14
5.3 4.0±2.0 1.9 1.8 JPL29
2.68±0.8 4.5±1.3 He 0.6-4.0 220-473 McCaulley9
5.30±0.2 1.4±0.1 He 1-10 298 Biggs10
9.00±1.9 1.9±0.3 Ar 6-100 295 Frost11,12
5.20±1.9 2.1±0.6 He 30-100 295 Frost11,12
11.00±3.0 2.0±0.5 CF4 30-125 295 Frost11,12
Displayed in Table 5.4 are the parameters (k0, n, k∞, m) of the temperature dependent fall-
off curves derived in this study together with those from previous investigations. I previously
showed in Chapter 4 that our rate constants at 295 K coincide well with the results of the two
studies with Ar and N2 being the third body gas, 11,12,13 upon which the JPL evaluation is based.
128
As k0 is expected to differ with the types of third-body molecule, we first make
0
comparisons between k1a (298 K) in the same bath gas. With He as the third body, the value of
0
k1a (298 K) from Martínez et al.14 is ~30% greater than that from McCaulley et al.,9 but ~30%
smaller than that given by Biggs et al.10 Note that Martínez et al.’s14 result was based on their
own data along with the values from the two low-pressure studies.9,10 With Ar as the third body
0
gas, k1a (298 K) from Wollenhaupt et al.13 is ~70% smaller than that was reported by Frost and
11,12
Smith, which can be primarily attributed to different value of Fcent (0.44) that Frost and
Smith11,12 used in fall-off curve fitting. Applying Fcent = 0.6 to the fitting of Frost and Smith’s
0
data11,11 to Troe expression (E5.4) yields a similar value of k1a (298 K) as that of Wollenhaupt et
0
al. 13 Our value of k1a (298 K) is in reasonable agreement with that from Martínez et al. (in He)14
and Wollenhaupt et al. (in Ar).13
The high-pressure limiting rate constant at any one temperature should be independent of
∞
the type of third body molecule. We obtained a very similar value of k1a (298 K) to that from the
∞
two measurements in N2 and Ar bath gas. 11,12,13 Our value of k1a (298 K) is 42% larger than that
of Biggs et al., in He.10 The pressure (≤10 Torr) used in Biggs et al.’s study is rather low for an
∞
accurate k1a determination, because the rate constant does not approach the high-pressure limit
∞
until pressures of a few hundred Torr. Our value of k1a (298 K) is 24% smaller than that of
Martínez et al.14 As mentioned previously, this is inconsistent with the prediction that He is less
efficient third body collider than N2 and Ar.
By comparing the value of the power law exponent, n, in the low-pressure limit from
different studies shown in Table 5.4, we find that the values of n obtained fall into two groups:
the present results and those of Martínez et al.14 agree within 5%, while those from Wollenhaupt
129
et al.13 and McCaulley et al.9 agree within 3%. However, the latter pair of studies yield n 2.6
times larger than the former pair of studies. The value we derive for the power law exponent, m,
in the high-pressure limit (1.13) is 35% larger than that of Martínez et al.;14 however, the two
results agree within their errors. Our m value is 65% smaller than that derived by Wollenhaupt et
al.13
We present a plot comparing rate constants from our study and previous work in order to
facilitate a global comparison. In view of the fact that the temperatures used in our experiments
are different from other studies, we plot the calculated fall-off curves using Troe expression
(E5.4) using our fitted parameters at temperatures that match the available literature data.
-11
2.4x10
-11
2.0x10
k1a (cm molecule s )
-11
-1
1.6x10
-1
-11
1.2x10
233K
-3
-12
8.0x10 262K
297K
297K - N2
-12
4.0x10 356K
390K (Frost) - Ar
0.0
18 19 19 19
0.0 5.0x10 1.0x10 1.5x10 2.0x10
-3
[M] (molecule cm )
Figure 5.9 Comparison of fall-off curves (lines) for k1a calculated using parameters resulted from current
study together with the absolute T,P-dependent k1a values (symbols) determined by Wollenhaupt et al. in
Ar,13 except for results of Frost and Smith in Ar (noted in the legend).11,12
130
Figure 5.9 illustrates that, despite that our results near room temperature agree well with
those of Wollenhaupt et al.,13 our results shows less temperature dependence both above and
below room temperature. On the other hand, our fall-off curve does a good job of predicting
Frost and Smith’s results11,12 at 390 K in N2 (the only other study carried out in N2).
Neither our work nor previous studies of the methoxy + NO2 reaction take into account the
extent of formation of methyl peroxy nitrite (CH3OONO).18,19,20 The absence of direct
experimental evidence for formation of CH3OONO makes such quantification impossible.
However, there is experimental evidence enabling quantification of the HOONO formation
channel in the HO + NO2 reaction.22,41 If formation of CH3OONO (or CD3OONO) occurs to a
similar (large) extent as HOONO formation, use of the Troe formula to fit the data would be
invalid.
5.4 Conclusion
We determined rate constants for the reaction of CH3O• (CD3O•) + NO2 in N2 bath gas
over a range for pressures and temperatures. From these results, we derived the parameters of the
Troe expression with Fcent= 0.6: k0(298 K), n, k∞(298 K), and m. Our fitting results is consistent
with all the literature data in Ar bath gas at room temperature, and the limited data in N2 at 390
K. Wollenhaupt et al.13 reported stronger temperature dependences of k0 and k∞ than found here,
but we have no explanation for this discrepancy.
The highest pressure investigated in this study (700 Torr) is close to the high-pressure limit
for k1a and k2a, and the third body gas N2 should reliably mimic third body effects in air.
Therefore, our results can be used to reliably interpret chamber studies carried out in air near 1
131
atm total pressure. Energy transfer is increasingly a source of error in Master Equation (ME)
studies of the fate of chemically activated species and calculations of pressure-dependent rate
constants. The rate constants obtained from this study may be valuable to constrain Rice-
Ramsperger-Kassel-Marcus (RRKM)/ME studies of collisional energy transfer in these
reactions.
132
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29 Sander, S. P.; Abbatt, J.; Barker, J. R.; Burkholder, J. B.; Friedl, R. R.; Golden, D. M.; Huie, R. E.;
Kolb, C. E.; Kurylo, M. J.; Moortgat, G. K.; Orkin, V. L.; Wine, P. H. Chemical Kinetics and
Photochemical Data for Use in Atmospheric Studies, Evaluation No. 17; JPL Publication 10-6, Jet
Propulsion Laboratory, Pasadena, 2011 http://jpldataeval.jpl.nasa.gov.
92, 5914-5921
33 Albaladejo, J.; Jiménez, E.; Notario, A.; Cabañas, B.; Martínez, E. J. Phys. Chem. A 2002, 106, 2512-
2519.
34 Biggs, P.; Canosa-Mas, C. E.; Fracheboud, J.; Shallcross, D. E.; Wayne, R. P. J. Chem. Soc., Faraday
Trans. 1997, 93, 2481-2486.
35 Meier, U.; Grotheer, H.; Riekert, G.; Just, T. Ber. Bunsenges. Phys. Chem. 1985, 89, 325-327
36 Hassinen, E.; Koskikallio, J. Acta Chem. Scand. A 1979, 33, 625-630.
134
37 Cobos, C. J.; Troe, J. Z. Phys. Chem. 2003, 217, 1031-1044.
38 Atkinson, R.; Baulch, D.; Cox, R.; Crowley, J.; Hampson, R.; Hynes, R.; Jenkin, M.; Kerr, J.; Rossi,
M.; Troe, J. Evaluated kinetic data: http://www.iupac-kinetic.ch.cam.ac.uk 2007, .
39 Golden, D. M. Int. J. Chem. Kinet. 2005, 37, 625-632.
40 McCaulley, J. A.; Moyle, A. M.; Golde, M. F.; Anderson, S. M.; Kaufman, F. J. Chem. Soc., Faraday
Trans. 1990, 86, 4001-4009.
41 Hippler, H.; Nasterlack, S.; Striebel, F. Phys. Chem. Chem. Phys. 2002, 4, 2959-2964.
135
Chapter 6. Conclusions
Absolute rate constants for the reaction of CH3O• + NO2 in 700 Torr N2 and in the
temperature range of 250 – 333 K have been determined in my research. By combining these
results with our group’s relative rate measurement for CH3O• + O2/NO2, rate constants for the O2
reaction have been determined for the first time below room temperature. The absolute rate
constants were measured using laser flash photolysis to generate radicals and laser-induced
fluorescence for time-resolved detection. I obtained the LIF spectra of CH3O• and CD3O• from
286 to 303 nm, which were in good accord with those of Inoue et al.1 The kinetic results on
CH3O• + O2 show reasonable agreement with previous absolute rate studies in the overlapped
temperature range (298-333 K), however, exhibit less temperature dependence for k1 than
previous data obtained from 293-610 K. By carrying out the same experiments for the
isotopologue CD3O• in the temperature range of 250 – 335 K, I have been able to determine the
kinetic isotope effect (kH/kD = k1/k3) for methoxy + O2. Comparison of the rate constants from
these two isotopologues of methoxy radical does not reveal evidence for an influence of
tunneling on the CH3O + O2 rate constant. For future work, it would be interesting to extend the
study to temperatures lower than that used here. This would shed more light on KIE and the
influence of tunneling.
The experimentally determined k1 (T) for methoxy + O2 will be invaluable to validate a
computational method for the same reaction carried out in our group. The low temperature rate
constants will enable the validation of tunneling treatment used in the computations. A validated
method would, hopefully, enable reliable (and affordable) computational studies for RO• + O2
reactions for those RO• that do not fluoresce or which we cannot cleanly produce. These include
136
1-butoxy radical (which is the prototype for alkoxy radical isomerization) and larger and
functionalized alkoxy radicals derived from atmospherically important species such as alkenes or
oxygenated VOCs. The availability of absolute rate constants for RO• + O2 reactions would, in
turn, enable one to extract absolute rate constants for decomposition and isomerization of these
alkoxy radicals from relative rate experiments. Absolute rate constants for these unimolecular
reactions are necessary for alkoxy radicals from certain compounds of special importance (such
as isoprene) and to build structure-reactivity relations for unimolecular reactions of broad classes
of functionalized alkoxy radicals that are difficult to study experimentally. In addition, the
validated computational method would also enable us to extend the temperature range of the
branching ratio of the two product channels for the reaction of singly deuterated methoxy radical
(CH2DO•) with O2, which is an important step for deuterium enrichment in molecular
hydrogen.2,3
I have also investigated the kinetics for the reaction of CH3O• + NO2 in N2 bath gas. The
pressure dependent recombination channel is dominant throughout the temperature (250-333 K)
and pressure range (30-700 Torr) range studied, and the rate constants are well-fit by the Troe
expression. My work is the first to obtain rate constants for this reaction at pressures up to 700
Torr in the bath gas N2. In addition I investigated the kinetics of the perdeuterated isotopologue
of CH3O•, i.e. CD3O• + NO2, and this offers a reasonable reproducibility of the kinetic results of
reaction CH3O• + NO2. Fitting results from this study can reproduce all the literature data in N2
or the similar third body gas Ar at room temperature. However, significant discrepancies exist at
higher and lower temperatures among our result and those from two other studies.4,5 In one case,
we strongly suspect that some systematic error exists in the study of Martínez et al.,5 due to the
higher rate constants they obtained in He than those obtained under similar conditions by others
137
in Ar and in the present work in N2. There is no definitive basis for preferring our rate constants
versus those from Wollenhaupt et al.;4 however, Wollenhaupt et al.’s results4 may be less reliable
because of their indirect method of generating methoxy radical. Their method is more subject
than my approach to error due secondary chemistry.
The highest pressure 700 Torr investigated in this study is close to the high pressure limit
for both CH3O• + NO2 and CD3O• + NO2, and the third body gas N2 is more realistic to mimic
atmospheric environment, and therefore our results can be used to reliably interpret the chamber
studies from our group as well as from previous relative rate studies on methoxy chemistry.6, 7
Furthermore, the pressure dependent rate constants for the association reaction methoxy +
NO2 measured in this study will be valuable to validate the current Rice-Ramsperger-Kassel-
Marcus (RRKM)/Master Equation (ME) study on the same reaction.8 As has been done
previously using literature results obtained at lower pressure than my work, one can estimate the
value of the collisional energy transfer parameter (α) by fitting rate constants from an RRKM
ME simulation to experimental rate constants. RRKM theory is required to compute the rate
constant for association/dissociation, k(E), as a function of energy, E. As the formation of
methylnitrate from methoxy + NO2 does not have a maximum in potential energy along the
reaction path, standard RRKM theory does not apply. What is required, instead, is a variational
treatment to find transition states at each energy that are minima in k(E) along the reaction path.
The quantum calculations required to do this rigorously are computationally demanding and
repetitive.8 As Barker et al.8 pointed out, the accuracy of collisional energy transfer parameters
derived from RRKM/ME simulations is subject to large uncertainty when using approximate
treatments of variational transition states. Compared to the experimental study from Wollenhaupt
et al.,4 the highest pressure applied in this study is closer to high pressure limit and the chemistry
138
of the reaction system is much cleaner. Therefore, my study offers more reliable high pressure
limiting rate constants as a function of temperature. A better understanding of energy transfer in
the methoxy + NO2 system may provide insights to help constrain models of energy transfer for
other barrierless radical-radical recombination reactions. Such reactions include R• + O2, RO2• +
NO, and OH + NO2, which are of atmospheric and combustion importance.
139
References
2 Hu, H.; Dibble, T. S.; Tyndall, G. S.; Orlando, J. J. J. Phys. Chem. A 2012, 116, 6295-6302.
3 Nilsson, E. J. K.; Johnson, M. S.; Taketani, F.; Matsumi, Y.; Hurley, M. D.; Wallington, T. J.
Atmos. Chem. Phys. 2007, 7, 5873-5881.
5 Martıń ez, E.; Albaladejo, J.; Jiménez, E.; Notario, A.; Dıaz de Mera, Y. Chem. Phys. Lett. 2000, 329,
191-199.
6 Barker, J. R.; Benson, S. W.; Golden, D. M. Int. J. Chem. Kinet. 1977, 9, 31-53.
8 Barker, J. R.; Lohr, L. L.; Shroll, R. M.; Reading, S. J. Phys. Chem. A 2003, 107, 7434-7444.
140
Appendix I.
Experimental conditions and pseudo first order rate constant of methoxy + NO2 under each
condition. Error bars are 2σ of the fitted slope of ln(intensity) versus time. Unit for NO2
concentration is molecule cm-3.
Table I-1-1 CH3O + NO2 at 250 K
700 Torr 500 Torr

-1
[NO2] k' (s ) k’error (2σ) [NO2] k' (s-1) k’error (2σ)
5.98E+14 1.91E+04 4.60E+02 6.62E+14 1.98E+04 3.29E+02
1.17E+15 3.31E+04 3.71E+02 1.29E+15 3.36E+04 4.47E+02
1.88E+15 4.78E+04 7.48E+02 2.10E+15 4.76E+04 8.10E+02
2.73E+15 6.55E+04 1.15E+03 3.04E+15 6.83E+04 1.50E+03
3.53E+15 8.13E+04 1.99E+03 3.94E+15 8.52E+04 1.54E+03
100 Torr 50 Torr

[NO2] k' (s-1) k’error (2σ) [NO2] k' (s-1) k’error (2σ)
7.82E+14 1.32E+04 3.47E+02 7.94E+14 1.46E+04 4.76E+02
1.46E+15 2.38E+04 3.63E+02 1.47E+15 2.38E+04 6.46E+02
2.36E+15 3.80E+04 9.84E+02 2.36E+15 3.56E+04 1.13E+03
3.41E+15 5.37E+04 7.22E+02 3.41E+15 4.91E+04 1.57E+03
4.41E+15 6.88E+04 1.54E+03 4.41E+15 6.56E+04 1.44E+03
30 Torr
[NO2] k' (s-1) k’error (2σ)
7.85E+14 1.22E+04 2.92E+02
1.46E+15 2.17E+04 2.91E+02
2.36E+15 3.32E+04 5.84E+02
3.41E+15 4.56E+04 1.10E+03
4.41E+15 5.67E+04 1.27E+03
141
700 Torr 500 Torr

-1
6.01E+14 1.90E+04 4.27E+02 6.98E+14 2.43E+04 2.99E+02
1.20E+15 3.04E+04 5.35E+02 1.33E+15 3.29E+04 3.66E+02
1.98E+15 4.69E+04 1.37E+03 2.19E+15 5.07E+04 7.28E+02
2.94E+15 6.86E+04 2.64E+03 3.26E+15 7.45E+04 8.07E+02
3.89E+15 8.40E+04 2.63E+03 4.30E+15 9.08E+04 2.07E+03
50 Torr 30 Torr
7.52E+14 1.36E+04 2.65E+02 7.52E+14 1.36E+04 2.65E+02
1.44E+15 2.36E+04 6.01E+02 1.44E+15 2.36E+04 6.01E+02
2.37E+15 3.37E+04 7.45E+02 2.37E+15 3.37E+04 7.45E+02
3.52E+15 4.75E+04 1.34E+03 3.52E+15 4.75E+04 1.34E+03
4.64E+15 6.05E+04 1.62E+03 4.64E+15 6.05E+04 1.62E+03
215 Torr
7.57E+14 1.49E+04 3.96E+02
1.44E+15 2.69E+04 4.26E+02
2.37E+15 4.28E+04 9.97E+02
3.52E+15 5.96E+04 1.49E+03
4.65E+15 7.59E+04 1.83E+03
142
700 Torr 500 Torr

-1
8.13E+14 1.69E+04 3.41E+02 7.06E+14 2.00E+04 2.83E+02
1.41E+15 3.00E+04 1.00E+03 1.40E+15 3.25E+04 4.44E+02
2.13E+15 4.07E+04 6.95E+02 2.32E+15 4.92E+04 7.01E+02
3.18E+15 6.24E+04 9.76E+02 3.46E+15 6.88E+04 1.22E+03
4.23E+15 8.24E+04 2.39E+03 4.59E+15 8.59E+04 2.43E+03
50 Torr 30 Torr
6.95E+14 1.13E+04 2.06E+02 6.95E+14 1.04E+04 2.38E+02
1.38E+15 1.91E+04 3.68E+02 1.38E+15 1.72E+04 2.53E+02
2.29E+15 2.83E+04 8.16E+02 2.29E+15 2.78E+04 5.43E+02
3.42E+15 4.38E+04 9.11E+02 3.42E+15 3.66E+04 8.23E+02
4.54E+15 5.49E+04 1.28E+03 4.54E+15 4.68E+04 7.20E+02
100 Torr
6.97E+14 1.14E+04 2.39E+02
1.38E+15 2.04E+04 3.83E+02
2.29E+15 3.38E+04 5.28E+02
3.42E+15 4.80E+04 9.88E+02
4.54E+15 6.26E+04 1.36E+03
143
700 Torr 500 Torr

-1
9.20E+14 2.09E+04 3.20E+02 9.43E+14 2.01E+04 3.83E+02
1.84E+15 3.67E+04 7.90E+02 1.88E+15 3.62E+04 3.84E+02
3.06E+15 5.59E+04 1.42E+03 3.13E+15 5.70E+04 1.07E+03
4.58E+15 8.46E+04 1.38E+03 4.69E+15 8.01E+04 1.73E+03
6.10E+15 1.08E+05 1.83E+03 6.24E+15 1.03E+05 2.81E+03
100 Torr 215 Torr

-1
8.35E+14 1.36E+04 2.01E+02 8.35E+14 1.58E+04 5.72E+02
1.67E+15 2.33E+04 3.31E+02 1.67E+15 2.83E+04 1.37E+03
2.78E+15 3.81E+04 5.52E+02 2.78E+15 4.57E+04 1.19E+03
4.18E+15 5.44E+04 1.60E+03 4.18E+15 6.78E+04 2.30E+03
5.57E+15 7.25E+04 1.96E+03 5.57E+15 8.48E+04 3.92E+03
50 Torr 30 Torr
-1
8.35E+14 1.21E+04 3.71E+02 8.35E+14 1.18E+04 3.13E+02
1.67E+15 2.01E+04 4.28E+02 1.67E+15 1.79E+04 4.19E+02
2.78E+15 3.28E+04 8.02E+02 2.78E+15 2.98E+04 6.17E+02
4.18E+15 4.90E+04 9.86E+02 4.18E+15 4.17E+04 9.60E+02
5.57E+15 6.25E+04 2.59E+03 5.57E+15 5.51E+04 2.41E+03
144
700 Torr 500 Torr

-1
8.49E+14 1.54E+04 2.58E+02 9.22E+14 1.82E+04 2.04E+02
1.49E+15 2.66E+04 5.28E+02 1.61E+15 2.83E+04 5.33E+02
2.33E+15 3.99E+04 8.28E+02 2.52E+15 4.30E+04 6.08E+02
3.28E+15 5.55E+04 6.02E+02 3.55E+15 5.84E+04 1.08E+03
4.22E+15 6.83E+04 1.18E+03 4.56E+15 7.43E+04 1.18E+03
100 Torr 215 Torr

3.44E+14 5.37E+03 1.31E+02 3.51E+14 6.52E+03 1.71E+02
6.89E+14 9.44E+03 1.25E+02 6.92E+14 1.14E+04 1.27E+02
1.15E+15 1.48E+04 2.35E+02 1.15E+15 1.73E+04 4.40E+02
1.72E+15 2.15E+04 5.14E+02 1.72E+15 2.45E+04 7.56E+02
2.29E+15 2.82E+04 5.63E+02 2.29E+15 3.19E+04 7.29E+02
50 Torr 30 Torr
-1
3.5E+14 4.68E+03 1.01E+02 3.51E+14 5.02E+03 1.90E+02
6.94E+14 8.01E+03 2.00E+02 6.87E+14 7.70E+03 1.86E+02
1.15E+15 1.27E+04 2.46E+02 1.14E+15 1.15E+04 1.56E+02
1.71E+15 1.84E+04 5.05E+02 1.72E+15 1.60E+04 2.58E+02
2.29E+15 2.40E+04 6.07E+02 2.29E+15 2.05E+04 5.16E+02
145
700 Torr 500 Torr

-1
8.27E+14 1.53E+04 2.61E+02 9.18E+14 1.67E+04 3.80E+02
1.45E+15 2.53E+04 6.13E+02 1.61E+15 2.65E+04 8.72E+02
2.27E+15 3.52E+04 6.89E+02 2.52E+15 3.98E+04 5.30E+02
3.20E+15 5.13E+04 8.39E+02 3.55E+15 5.56E+04 1.75E+03
4.12E+15 6.35E+04 1.39E+03 4.56E+15 6.75E+04 2.35E+03
100 Torr 215 Torr

3.31E+14 5.74E+03 1.59E+02 3.31E+14 6.47E+03 1.99E+02
6.39E+14 8.31E+03 2.19E+02 6.45E+14 9.59E+03 2.36E+02
1.06E+15 1.32E+04 2.51E+02 1.06E+15 1.47E+04 3.23E+02
1.58E+15 1.90E+04 3.69E+02 1.59E+15 2.10E+04 4.98E+02
2.10E+15 2.36E+04 7.85E+02 2.12E+15 2.68E+04 6.26E+02
50 Torr 30 Torr
-1
3.18E+14 4.25E+03 1.18E+02 3.19E+14 3.75E+03 1.00E+02
6.39E+14 7.16E+03 1.59E+02 6.33E+14 4.85E+03 1.47E+02
1.06E+15 1.12E+04 4.41E+02 1.06E+15 6.78E+03 1.66E+02
2.10E+15 2.17E+04 8.80E+02 1.58E+15 9.74E+03 2.20E+02
2.10E+15 1.25E+04 2.10E+02
146
Table I-2-1 CD3O + NO2 at 250 K
700 Torr 500 Torr

-1
6.92E+14 2.47E+04 5.81E+02 9.30E+14 2.79E+04 4.55E+02
1.13E+15 3.47E+04 1.09E+03 1.52E+15 4.19E+04 5.11E+02
1.87E+15 5.31E+04 1.34E+03 2.49E+15 6.21E+04 1.08E+03
2.78E+15 6.85E+04 1.82E+03 3.65E+15 8.41E+04 2.54E+03
3.72E+15 8.91E+04 3.60E+03 4.86E+15 1.10E+05 3.03E+03
100 Torr 215 Torr

6.93E+14 1.65E+04 1.81E+02 6.93E+14 1.58E+04 3.56E+02
1.35E+15 2.97E+04 5.06E+02 1.35E+15 2.83E+04 9.03E+02
2.18E+15 4.47E+04 6.27E+02 2.18E+15 4.73E+04 7.60E+02
3.15E+15 6.28E+04 1.20E+03 3.15E+15 6.65E+04 1.56E+03
4.07E+15 8.04E+04 1.80E+03 4.07E+15 8.29E+04 1.97E+03
50 Torr 30 Torr
-1
6.93E+14 1.83E+04 3.26E+02 6.93E+14 1.61E+04 1.98E+02
1.35E+15 3.10E+04 6.18E+02 1.35E+15 2.83E+04 2.54E+02
2.18E+15 4.58E+04 1.02E+03 2.18E+15 4.17E+04 7.17E+02
3.15E+15 6.25E+04 1.52E+03 3.15E+15 5.80E+04 9.99E+02
4.07E+15 7.95E+04 1.65E+03 4.07E+15 7.30E+04 8.83E+02
147
700 Torr 500 Torr

-1
7.18E+14 1.95E+04 2.94E+02 9.13E+14 2.08E+04 3.18E+02
1.20E+15 3.11E+04 5.03E+02 1.52E+15 3.42E+04 6.42E+02
2.03E+15 4.93E+04 8.67E+02 2.58E+15 5.22E+04 6.23E+02
3.09E+15 7.30E+04 1.18E+03 3.91E+15 7.78E+04 1.44E+03
4.25E+15 9.37E+04 1.56E+03 5.40E+15 9.96E+04 1.84E+03
50 Torr 30 Torr
-1
6.97E+14 1.34E+04 1.56E+02 6.95E+14 1.27E+04 2.35E+02
1.38E+15 2.41E+04 2.85E+02 1.38E+15 2.43E+04 4.62E+02
2.29E+15 3.77E+04 4.90E+02 2.29E+15 3.44E+04 6.46E+02
3.43E+15 5.44E+04 8.47E+02 3.43E+15 5.22E+04 9.01E+02
4.55E+15 7.28E+04 1.41E+03 4.55E+15 6.55E+04 1.33E+03
100 Torr
6.95E+14 1.39E+04 1.06E+02
1.38E+15 2.45E+04 3.20E+02
2.29E+15 4.00E+04 7.07E+02
3.43E+15 5.84E+04 1.03E+03
4.55E+15 7.50E+04 1.56E+03
148
700 Torr 500 Torr

-1
7.24E+14 1.66E+04 2.91E+02 7.32E+14 1.87E+04 2.93E+02
1.45E+15 2.94E+04 2.90E+02 1.45E+15 3.20E+04 3.54E+02
2.41E+15 4.48E+04 1.14E+03 2.42E+15 4.76E+04 5.00E+02
3.59E+15 6.81E+04 8.16E+02 3.62E+15 6.73E+04 1.28E+03
4.75E+15 8.23E+04 2.38E+03 4.80E+15 8.83E+04 1.97E+03
100 Torr 215 Torr

6.04E+14 1.05E+04 1.98E+02 6.13E+14 1.05E+04 1.88E+02
1.20E+15 1.95E+04 2.39E+02 1.21E+15 1.93E+04 3.63E+02
2.00E+15 3.14E+04 5.59E+02 2.00E+15 3.19E+04 2.12E+02
3.00E+15 4.49E+04 6.83E+02 3.00E+15 4.72E+04 4.93E+02
3.99E+15 5.90E+04 1.03E+03 3.99E+15 5.86E+04 1.42E+03
50 Torr 30 Torr
-1
6.07E+14 9.42E+03 1.36E+02 6.07E+14 8.82E+03 1.42E+02
1.20E+15 1.69E+04 2.06E+02 1.20E+15 1.59E+04 2.63E+02
2.01E+15 2.69E+04 4.60E+02 2.01E+15 2.47E+04 3.99E+02
3.01E+15 4.04E+04 1.22E+03 3.01E+15 3.70E+04 7.98E+02
4.00E+15 5.09E+04 1.13E+03 4.00E+15 4.72E+04 9.33E+02
149
700 Torr 500 Torr

-1
8.36E+14 1.65E+04 2.05E+02 9.26E+14 1.78E+04 2.68E+02
1.39E+15 2.63E+04 3.67E+02 1.54E+15 2.78E+04 3.24E+02
2.36E+15 4.22E+04 5.71E+02 2.61E+15 4.39E+04 6.34E+02
3.60E+15 6.28E+04 1.10E+03 3.99E+15 6.54E+04 1.07E+03
4.99E+15 8.37E+04 2.19E+03 5.52E+15 8.88E+04 1.23E+03
50 Torr 30 Torr
-1
5.63E+14 8.94E+03 2.30E+02 5.61E+14 7.49E+03 1.90E+02
1.12E+15 1.51E+04 3.06E+02 1.11E+15 1.37E+04 3.01E+02
1.85E+15 2.35E+04 4.97E+02 2.77E+15 2.85E+04 6.30E+02
2.77E+15 3.44E+04 8.64E+02 3.69E+15 3.86E+04 6.78E+02
3.69E+15 4.57E+04 1.16E+03
100 Torr
5.74E+14 8.84E+03 2.59E+02
1.11E+15 1.53E+04 4.07E+02
1.85E+15 2.52E+04 5.43E+02
2.77E+15 3.69E+04 8.38E+02
3.69E+15 4.77E+04 1.10E+03
150
700 Torr 500 Torr

-1
7.79E+14 1.64E+04 3.23E+02 9.14E+14 1.63E+04 4.15E+02
1.30E+15 2.43E+04 6.24E+02 1.52E+15 2.63E+04 4.29E+02
2.20E+15 3.83E+04 7.69E+02 2.58E+15 4.18E+04 1.09E+03
3.36E+15 5.93E+04 2.10E+03 3.95E+15 5.94E+04 1.43E+03
4.68E+15 7.46E+04 1.78E+03 5.48E+15 8.03E+04 1.78E+03
100 Torr 215 Torr

-1
5.36E+14 7.84E+03 2.29E+02 5.37E+14 8.45E+03 1.96E+02
1.07E+15 1.42E+04 2.63E+02 1.07E+15 1.51E+04 4.15E+02
1.76E+15 2.26E+04 5.59E+02 1.78E+15 2.35E+04 6.23E+02
2.64E+15 3.25E+04 5.29E+02 2.70E+15 3.66E+04 4.77E+02
3.51E+15 4.41E+04 1.06E+03
50 Torr 30 Torr
-1
5.32E+14 7.09E+03 1.73E+02 5.32E+14 6.84E+03 1.73E+02
1.06E+15 1.23E+04 1.56E+02 1.06E+15 1.07E+04 2.47E+02
1.76E+15 2.09E+04 3.75E+02 1.76E+15 1.68E+04 3.05E+02
2.64E+15 2.97E+04 7.93E+02 2.64E+15 2.54E+04 4.21E+02
3.51E+15 3.79E+04 8.44E+02 3.51E+15 3.26E+04 7.32E+02
151
Curriculum Vitae
Jiajue Chai
EDUCATION
Ph.D candidate in Physical and Atmospheric Chemistry, SUNY
College of Environmental Science and Forestry, Syracuse, NY 08/2008-12/2013
Visiting M.S. student in Chemical Engineering

National Taiwan University, Taipei 02/2008-06/2008
M.S. program in Applied Chemistry

Renmin University of China, Beijing, China 09/2006-08/2008
B.S. in Environmental Sciences

Nanjing Forestry University, Nanjing, China 09/2002-06/2006
RESEARCH EXPERIENCE
 Radical kinetics and Laser spectroscopy 09/2008-present
1) Constructed a system for Pulsed Laser Flash Photolysis/ Pulsed Laser-induced Fluorescence
(PLP/PLIF) spectroscopy study for atmospheric-related alkoxy and hydroxyl radicals:
- repaired, aligned and optimized two excimer lasers (GAM Laser EX100, and Lambda Physik
Lextra 100) and a dye laser (Lambda Physik FL3002)
- established data acquisition system
- debugged and aligned frequency doubler (Inrad Autotracker III)
2) Synthesized and purified organic nitrites as precursors to alkoxy radicals and characterized them by
1
H NMR, FTIR, UV-vis, and GC-MS
3) Obtained LIF spectra of alkoxy radicals without and with functional group (e.g. methoxy, vinoxy,
cyclohexoxy, benzyloxy)
4) Obtained rate constants as a function of temperature and pressure for the reaction methoxy+NO2 and
its deuterated isotopologues; carried out quantum calculation for the same reaction
- Used results of relative rate study to determine rate constants of methoxy+O2 as a function of
temperature, which is of great atmospheric significance; kinetic modeling
 Supercritical CO2-based cleaning of micro-electronics 11/2006-05/2008

Established a physical model for the cleaning of micro-porous surfaces in CO2-based solvent using
Mathematica programming
 Adsorption behavior of heavy metal 12/2005-05/2006

Determined adsorption behavior of heavy metal on hydrous manganese dioxide and humic substance
152
PUBLICATIONS
J. Chai, H. Hu, T. S. Dibble, G. S. Tyndall, and J. J. Orlando, Rate constants and kinetic isotope effects
for methoxy radical reacting with NO2 and O2. (in preparation for J. Phys. Chem. A)
J. Chai and T. S. Dibble, Pressure dependence and kinetic isotope effects in the absolute rate constant for
methoxy radical reacting with NO2. (in preparation for Int. J. Chem. Kinet.)
J. Chai, X. Zhang et al., Microscopic model of nano-scale particles removal in high pressure CO2-based
solvent, Journal of Supercritical Fluids 49 (2009) 182-188
X. Tan, J. Chai, The model of nano-scale copper particles removal from silicon surface in high pressure
CO2+H2O and CO2+H2O+IPA cleaning solutions, Journal of Nanoscience and Nanotechnology 11
(2011) 10782-10786
CONFERENCE PRESENTATIONS (co-author)

22nd International Symposium on Gas Kinetics in Boulder, CO (Poster) 6/2012
244th ACS National Meeting in Philadelphia, PA (Poster) 8/2012
38th ACS Northeastern Regional Meeting in Rochester, NY (Oral) 10/2012
68th OSU International Symposium on Molecular Spectroscopy in Columbus, OH (Oral) 6/2013
HONORS
China Aerospace Science and Technology Corporation Scholarship 10/2007
NJFU Excellent Undergraduate Student Award 06/2006
NJFU Scholarship for Excellent student leadership 11/2005
Governmental Scholarship of Jiangsu Province 11/2004
National Fellowship of China 10/2003
TEACHING AND MENTORING EXPERIENCE

Guest lecturer in Physical Chemistry and Atmospheric Chemistry
Mentored one graduate and one undergrad student 2010-present
Teaching Assistant for general chemistry lab at SUNY-ESF 08/2008-12/2009
Teaching Assistant for general chemistry lecture at SUNY-ESF 08/2013-12/2013
EXTRACURRICULAR
President of Chinese Student and Scholar Association at SUNY-ESF 07/2009-07/2010
Volunteer judge for 2013 7th and 8th Grade Science Fair 3/13/2013
Volunteer at ESF booth at New York State Fair 2009-2013
153

CH3ONO and CD3ONO Synthesis Method

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DIRECT KINETIC STUDY OF METHOXY RADICAL REACTING WITH NO2 AND O2

AND DEUTERIUM KINETIC ISOTOPE EFFECTS

Approved: Department of Chemistry

Theodore S. Dibble, Major Professor S. Scott Shannon, Chair

Greg Boyer, Department Chair S. Scott Shannon, Dean

List of Figures ....................................................................................................... viii

Curriculum Vitae .................................................................................................152

Table I-1-2 CH3O + NO2 at 265 K ............................................................................................ 142

Table I-1-3 CH3O + NO2 at 278 K ............................................................................................ 143

Table I-1-4 CH3O + NO2 at 295 K ............................................................................................ 144

Table I-1-5 CH3O + NO2 at 316 K ............................................................................................ 145

Table I-1-6 CH3O + NO2 at 333 K ............................................................................................ 146

Table I-2-1 CD3O + NO2 at 250 K ............................................................................................ 147

Table I-2-2 CD3O + NO2 at 277 K ............................................................................................ 148

Table I-2-3 CD3O + NO2 at 294 K ............................................................................................ 149

Table I-2-4 CD3O + NO2 at 319 K ............................................................................................ 150

Table I-2-5 CD3O + NO2 at 335 K ............................................................................................ 151

Figure 3.1 The experimental apparatus used in this study. ........................................................... 68

stratosphere (formed following photolysis of O2 by short-wavelength ultraviolet light) filters out

properties of the atmosphere. The Earth’s atmosphere is a tremendously active photochemical

(VOCs) in the presence of NOx.

and deterioration of vegetation, in contrast to the beneficial UV-blocking effect of stratospheric

rural regions and the upper troposphere.

Scheme 1.2 VOC oxidation mechanism. (VOC is represented here by RH)

produced, the O3 level is increased.6

unimolecular decomposition, unimolecular isomerization, or reaction with O2.6 Reaction with O2

subsequently transformed to second generation alkylperoxy radicals (R’O2•).6 As shown in

1.2 Secondary organic aerosols (SOAs)

tropical forested area.14

1.3 Research objectives

oxidation.6,15 Therefore, elucidation of the chemistry of alkoxy radicals is essential. To date,

alkanes16,17,18,19,20,21,22,23,24,25,26,27,28 and two halogenated alkanes.29,30,31 With these exceptions, the

obtained from experiment.6, 32

reaction methoxy + O2 (R1.1):

significant non-Arrhenius behavior.33 Theoretical calculations by Bofill et al.34 suggested that

rate measurement does not apply.

of upper tropospheric chemistry, photolysis of various oxygenated organic compounds such as

source of unexpectedly high concentrations of OH and HO2 radicals. This higher-than-expected

temperature range extended below room temperature, which is presented in Chapter 4. It is

concentration of O2 (>50 Torr) is required. Unfortunately, O2 efficiently quenches fluorescence

reaction of CH3O• with NO2:

k1 at temperatures down to 250 K.

methoxy + NO2 reactions are k1/k3 and k2/k4 respectively.

be discussed in Chapter 2 (section 2.3).

pp. 348-350, pp. 65-366.

Inc., 2006, 6ff, 1044ff

4 Haagen-Smit, A. Industrial & Engineering Chemistry 1952, 44, 1342-1346.

Inc., 2006, pp. 204-205.

6 Orlando, J. J.; Tyndall, G. S.; Wallington, T. J. The Atmospheric Chemistry of Alkoxy

Radicals. Chem. Rev., 2003, 103, 4657-4690.

7 Jenkin, M. E.; Hayman, G. D. Atmos. Environ. 1999, 33, 1275-1293.

8 Jenkin, M. E.; Clemitshaw, K. C. Atmos. Environ. 2000, 34, 2499-2527.

9 Lim, Y. B.; Ziemann, P. J. Environ. Sci. Technol. 2005, 39, 9229-9236.

10 Ziemann, P. J.; Atkinson, R. Chem. Soc. Rev., 2012, 41, 6582-6605.

11 Kroll, J. H.; Seinfeld, J. H. Atmos. Environ., 2008, 42, 3593-3624.

Experiment, and Applications; Academic Press, 2000, 349 ff

J. D.; Coe, H.; Ng, N. Science 2009, 326, 1525-1529.

29 Wu, F. and Carr, R. W. J. Phys. Chem. A 2001, 105, 1423-1432.

30 Wu, F. and Carr, R. W. Chem. Phys. Lett.,1999, 305, 44-50.