Gas-Phase Ion Chemistry: Kinetics and Thermodynamics

byCharles M. Nichols
B. S., Chemistry – ACS Certified
University of Central Arkansas, 2009

A thesis submitted to the
Faculty of the Graduate School of the
University of Colorado in partial fulfillmentof the requirements for the degree ofDoctor of Philosophy
Department of Chemistry and Biochemistry2016This thesis entitled:
Gas-Phase Ion Chemistry: Kinetics and Thermodynamics
Written by Charles M. Nichols
has been approved for the Department of Chemistry and Biochemistry by:_______________________________________
Veronica M. Bierbaum_______________________________________
W. Carl Lineberger

Date: December 08, 2015

A final copy of this thesis has been examined by all signatories, and we find that both the content
and the form meet acceptable presentation standards of scholarly work in the above mentioned
discipline.

Nichols, Charles M. (Ph.D., Physical Chemistry)
Gas Phase Ion Chemistry: Kinetics and Thermodynamics
Thesis directed by Professors Veronica M. Bierbaum and W. Carl Lineberger
Abstract:

This thesis employs gas-phase ion chemistry to study the kinetics and thermodynamics of chemical
reactions and molecular properties. Gas-phase ion chemistry is important in diverse regions of theuniverse. It is directly relevant to the chemistry occurring in the atmospheres of planets and moons as wellas the molecular clouds of the interstellar medium. Gas-phase ion chemistry is also employed todetermine fundamental properties, such as the proton and electron affinities of molecules. Furthermore,gas-phase ion chemistry can be used to study chemical events that typically occur in the condensed-phase,such as prototypical organic reactions, in an effort to reveal the intrinsic properties and mechanisms ofchemical reactions.
Mass spectrometry is used to investigate the reactivity and fundamental properties of ions. Kinetic
data for reactions between cations or anions with neutral atoms or molecules is acquired using theselected ion flow tube (SIFT). These data are interpreted mechanistically and thermodynamically using
literature data and theoretical computations. Additionally, the proton affinities of a series of anions are
determined using a triple quadrupole mass spectrometer by investigating the collision induced
dissociation of proton-bound, anionic heterodimers.
The SIFT is coupled with a variety of ion-generation techniques. This thesis reports the design andoperation of a new electrospray ion source for the SIFT in detail. The electrospray ionization techniqueallows ions that are pre-formed in the condensed-phase to be transmitted into the gas-phase for studies of
their reactivity.

To my friends, family, teachers, and Alex. Thanks for all of your support.

Acknowledgements
The completion of this thesis would not be possible without the assistance of so many people in mylife. First and foremost, I’d like to thank my scientific mentor and advisor, Veronica ‘Ronnie’ Bierbaum. I
couldn’t have picked a better research advisor for my time at the University of Colorado. Thank you foryour support and guidance over the last six years. I wouldn’t be the speaker, writer, or scientist that I am
today without your help. I feel lucky to have Carl Lineberger as a co-advisor. His comments and insight
into public speaking, experimental phenomena, and data analysis have been invaluable. Mathias Webersupported me through three oral examinations, a process that proved to be one of the hardest obstacles in
my life. I must also thank William Old for helping me initiate the kinetic method experiments on East
campus, and for helping me find the postdoc position at Vanderbilt University. To Barney Ellison, thanks
for your patience when the lab sink would overflow or the water line of a diffusion pump would pop off
and flood your lab. I owe you a copy of Pedley. To my undergraduate advisor, Bill Taylor, thanks for
introducing me to gas-phase ion chemistry, and becoming a friend over the years.
I don’t know what I would have done without my coworkers in the Bierbaum lab. To John, Callie,
Nick, Ditte, Jennifer, Zhibo, Zhe-Chen, Nadine, and Scott, thanks for making the group office anawesome workplace! I appreciate all of the scientific and instrumental discussions, as well as the snackand coffee breaks. Many thanks to the Lineberger, Ellison, and Weber research labs for being part of theion chemistry “supergroup.” The supergroup has been a great way to practice public speaking, receive
feedback on my work, and to keep up with some of the research that is taking place on campus. The workthat follows would not have been possible without funding, and therefore I must thank the National
Science Foundation and the Air Force Office of Scientific Research (Grants CHE-1012321, CHE1300886, ACI-1053575, and FA-9550-12-1-0125).
I must thank the support staff for their assistance. In the CIRES shop, Don has always been fun to talkwith about instrumentation and electronics in general. The staff in the chemistry department also provided
much assistance from purchasing orders to academic advising. Furthermore, I’m lucky to have access tov

the JILA staff; the electronics, machine shop, and computer staff have helped me resolve many problemson the fly.
I must give a special thanks to Casey Christopher. Casey helped me with homework in the first threesemesters, teaching me linear algebra and other math tricks. Without his generous assistance, I’m not surethat I would be here today. Casey and Jon were my roommates for four unforgettable years; thanks for
putting up with me, and reintroducing me to the joy of board-games. The rest of the Boulder crew,
Michael, Brett, Joe, Phil, Kristin, Laura, Daniel, Mindy, Jessie, Jessica, Joel, Kelsey, Rajine, Silvio, andeveryone else… I had such a good group of friends during my time in Boulder; I’ll always remember oursnowboarding, hiking, backpacking, and camping adventures. Alex has been both supporting and loving
since we began dating approximately three years ago. I don’t know what I would do without you. I can’t
wait to start a new life down in Nashville together. Finally, I need to thank my family. To my mom anddad, Roy and Melissa, thanks for all of your unconditional support throughout my life. My two sisters,
Jessica and Rebekah, and brother-in-laws, thanks for putting up with and supporting me. To my two
awesome nephews and my adorable niece, hopefully I’ll get to see more of you while living in Nashville.
Thanks to all of my cousins, cousin-in-laws, and cousins’ babies, I’m lucky to have such a large
supportive family. My cousin Jim in particular has joined me on so many adventures that I think of him
like a brother. Finally, thanks to my grandma, Tater Bug. You have been one of the biggest supporters in
my life. Thanks for everything, including the strong family values you have instilled in me.

viPublications
John M. Garver, Zhibo Yang, Charles M. Nichols, Benjamin B. Worker, Scott Gronert, Veronica M. Bierbaum;“Resolving the α-Effect in Gas-Phase SN2 Reactions: A Marcus Theory Approach,” Int. J. Mass Spectrom., 2012,
316, 244
John M. Garver, Zhibo Yang, Nadine Wehres, Charles M. Nichols, Benjamin B. Worker, Veronica M. Bierbaum;“The α-Effect in Elimination Reactions and Competing Mechanisms in the Gas-Phase,” Int. J. Mass Spectrom.,
2012, 330, 182
Charles M. Nichols, Zhibo Yang, Benjamin B. Worker, Denver R. Hager, Nico M. M. Nibbering, Veronica M.
Bierbaum; “Gas-Phase Reactions of the Atomic Oxygen Radical Cation with Halogenated Compounds,” Phys.
Chem. Chem. Phys., 2013, 15, 561
Charles M. Nichols, Zhibo Yang, Veronica M. Bierbaum; “Gas-phase Organic Reactions of the Atomic Oxygen
Radical Cation,” Int. J. Mass Spectrom., 2013, 353, 1
Ditte L. Thomsen, Jennifer N. Reece, Charles M. Nichols, Steen Hammerum, Veronica M. Bierbaum;“Investigating the α-Effect in Gas-Phase SN2 Reactions of Microsolvated Anions,” J. Am. Chem. Soc., 2013, 135,
15508
Ditte L. Thomsen, Charles M. Nichols, Jennifer N. Reece, Steen Hammerum, Veronica M. Bierbaum. “The αEffect and Competing Mechanisms: The Gas-Phase Reactions of Microsolvated Anions with Methyl Formate,” J.
Am. Soc. Mass Spectrom., 2014, 25, 159.
Ditte L. Thomsen, Jennifer N. Reece, Charles M. Nichols, Steen Hammerum, Veronica M. Bierbaum; “The αEffect in Gas-Phase SN2 Reactions of Microsolvated Anions: Methanol as a Solvent,” J. Phys. Chem. A, 2014, 118,
8060.
Scott Gronert, John M. Garver, Charles M. Nichols, Benjamin B. Worker, Veronica M. Bierbaum;“Dehalogenation of Arenes via SN2 Reactions at Bromine: Competition with Nucleophilic Aromatic Substitution,” J.
Org. Chem., 2014, 79, 11020
Charles M. Nichols, William M. Old, W. Carl Lineberger, Veronica M. Bierbaum, “Gas-Phase Acidities of Nitrated
Azoles as Determined by the Extended Kinetic Method and Computations,” J. Phys. Chem. A, 2015, 119, 395
Daniel J. Nelson, Wilson K. Gichuhi, Charles M. Nichols, Veronica M. Bierbaum, and W. Carl Lineberger;“Structure and Thermochemistry of o-, m-, and p-Dehydromethylphenol” – In preparation.
Charles M. Nichols, Zhe-Chen Wang, Zhibo Yang, W. Carl Lineberger, Veronica M. Bierbaum, “Experimental andtheoretical studies of the reactivity and thermochemistry of dicyanamide: N(CN) 2─” – In preparation.
Charles M. Nichols, Zhe-Chen Wang, W. Carl Lineberger, Veronica M. Bierbaum, “Gas-Phase Reactions of the
deprotonated nucleobases with H, N, and O atom” – In preparation.

vii

Table of Contents

1

Introduction: Gas-Phase Ion Chemistry ........................................................................................... 1
1.1

Overview ............................................................................................................................... 1

1.2

Potential Energy Surfaces...................................................................................................... 5

1.3

Thermochemistry................................................................................................................... 7

1.4

2

1.3.1

Gas-Phase Acidity ............................................................................................................. 7

1.3.2

Electron Affinity/Electron Binding Energy ...................................................................... 9

1.3.3

Thermochemical Cycles .................................................................................................. 10

1.3.4

Temperature Corrections................................................................................................. 12

1.3.5

Thermochemistry and Mass Spectrometry for Ion-Neutral Reactions ............................ 12

Kinetics ................................................................................................................................ 13
1.4.1

Collision Theory ............................................................................................................. 13

1.4.2

Reaction Efficiencies ...................................................................................................... 15

1.5

Electron Spin ....................................................................................................................... 15

1.6

Summary of This Work ....................................................................................................... 17

1.7

References ........................................................................................................................... 18

Experimental and Computational Methods ................................................................................... 20
2.1

2.2

Selected Ion Flow Tube ....................................................................................................... 20
2.1.1

Ion Production Region .................................................................................................... 21

2.1.2

Ion Selection Region ....................................................................................................... 29

2.1.3

Reaction Flow Tube ........................................................................................................ 30

2.1.4

Ion Detection Region ...................................................................................................... 31

2.1.5

Ion-Neutral Reaction Rate Constant ............................................................................... 31

2.1.6

Product Fraction .............................................................................................................. 38

Triple Quadrupole Mass Spectrometer ................................................................................ 40
viii

3

4

5

2.3

Computational Methods ...................................................................................................... 41

2.4

References ........................................................................................................................... 42

Gas-phase Reactions of the Oxygen Radical Cation, O•+............................................................... 44
3.1

Introduction ......................................................................................................................... 44

3.2

Experimental Methods......................................................................................................... 45

3.3

Computational Methods ...................................................................................................... 46

3.4

Results and Discussion ........................................................................................................ 46
3.4.1

O•+ + Perfluorinated Compounds .................................................................................... 46

3.4.2

O•+ + Methyl Halides ...................................................................................................... 49

3.4.3

O•+ + Organic Molecules................................................................................................. 53

3.5

Conclusion ........................................................................................................................... 61

3.6

References ........................................................................................................................... 62

Reactivity and Thermochemistry of Dicyanamide: N(CN)2─ ........................................................ 66
4.1

Introduction ......................................................................................................................... 66

4.2

Experimental ....................................................................................................................... 67

4.3

Calculations ......................................................................................................................... 69

4.4

Results and Discussion ........................................................................................................ 69
4.4.1

Formation of 1,5-dinitrobiuret ........................................................................................ 70

4.4.2

Dicyanamide Proton Affinity and Solvent Molecule Affinity ........................................ 72

4.5

Conclusion ........................................................................................................................... 76

4.6

References ........................................................................................................................... 77

Gas-Phase Acidities of the Nitrated Azoles ..................................................................................... 81
5.1

Introduction ......................................................................................................................... 82

5.2

Experimental ....................................................................................................................... 86

5.3

Calculations ......................................................................................................................... 88

5.4

Results ................................................................................................................................. 88
ix

6

5.5

Discussion ........................................................................................................................... 95

5.6

Conclusion ........................................................................................................................... 98

5.7

References ........................................................................................................................... 99

Gas-phase Investigations of the α-Effect ....................................................................................... 102
6.1

Introduction ....................................................................................................................... 102

6.2

Marcus Theory and the α-Effect in Gas-Phase SN2 Reactions .......................................... 104

6.3

Investigating the α-Effect in SN2 and E2 Reactions .......................................................... 108

6.4

The α-Effect within Micro-Solvated Systems ................................................................... 111

6.5

The α-Effect and Competing Mechanisms with Methyl Formate ..................................... 117

6.6

Conclusion ......................................................................................................................... 123

6.7

References ......................................................................................................................... 124

7

Appendix A ...................................................................................................................................... 128

8

Appendix B ...................................................................................................................................... 131

9Bibliography .................................................................................................................................... 139x

List of Figures

Figure 1.1 Artist rendition of the Cassini spacecraft’s flyby of Titan .......................................................... 2
Figure 1.2 The pillars of creation ................................................................................................................. 3
Figure 1.3 Potential energy surface for an SN2 reaction .............................................................................. 5
Figure 1.4 Thermochemical Cycles ........................................................................................................... 11
Figure 2.1 Schematic of Selected Ion Flow Tube (SIFT) .......................................................................... 20
Figure 2.2 Electron Ionization (EI) source diagram (a) and photograph (b).............................................. 22
Figure 2.3 Electrospray ionization (ESI) source diagram. ......................................................................... 24
Figure 2.4 Ion funnel photograph............................................................................................................... 25
Figure 2.5 Electrospray ionization source wiring schematic ..................................................................... 26
Figure 2.6 Electrospray ionization source spectra ..................................................................................... 28
Figure 2.7 EI source positive ion spectra of CO: (a) no mass selection (b) m/z 16 mass selected ............ 30
Figure 2.8 [O•+] vs. reaction distance for the O•+ + CH3F reaction ............................................................ 34
Figure 2.9 Microwave discharge apparatus ............................................................................................... 35
Figure 2.10 Microwave discharge titration plot ......................................................................................... 36
Figure 2.11 Product Fraction plot for O•+ + CH3F ..................................................................................... 39
Figure 2.12 4000 QTrap triple quadrupole mass spectrometer .................................................................. 40
Figure 3.1 Stationary points for the reaction of O•+ + CH3Cl; S=3/2 ......................................................... 51
Figure 3.2 Stationary points for the reaction of O•+ + CH3Cl; S=1/2 ......................................................... 52
Figure 3.3 Stationary points for the reaction of O•+ with acetaldehyde ..................................................... 56
Figure 3.4 Lowest energy ionic products for reaction 3.20 and 3.24 ......................................................... 58
Figure 4.1 Stationary points for the reaction of dicyanamide with two nitric acid molecules ................... 71
Figure 4.2 Lowest energy structures for dicyanamide and protonated dicyanamide ................................. 73
Figure 4.3 Lowest energy structures of dicyanamide clustered with nitric acid ........................................ 75
Figure 5.1 Selected azoles and their nitrated derivatives. .......................................................................... 83
xiFigure 5.2 CID spectra for the proton-bound heterodimer of 2-nitroimidazole and 3-nitrobenzoic acid atfour collision energies ......................................................................................................................... 87
Figure 5.3 Plots of ln(ki/k0) vs. ΔH°acid (kcal/mol) for 4-nitropyrazole. ................................................. 91
Figure 5.4 Plot of GAapp vs. Teff values for 4-nitropyrazole. ...................................................................... 92
Figure 6.1 α-nucleophiles......................................................................................................................... 102
Figure 6.2 Marcus Theory for a SN2 potential energy surface ................................................................. 105
Figure 6.3 Marcus Theory relationship of ΔH‡activation and ΔH‡driving force to ΔHrxn .................................... 107
Figure 6.4 Brønsted correlation of E2 reaction efficiencies of oxyanions with (CH3)3CCl..................... 110
Figure 6.5 Brønsted correlation of water-, methanol-, and un-solvated nucelophiles reacting with CH3Cl.......................................................................................................................................................... 113
Figure 6.6 Brønsted correlation of methanol-solvated nucleophiles with CH3Br and CH3Cl ................. 116
Figure 6.7 Brønsted plot: Micro-solvated and un-solvated ions with methyl formate ............................ 120
Figure 6.8 Lowest energy transition states for micro-solvated ions with methyl formate ....................... 122

xii

List of Tables

Table 1.1 Bracketing results for didehydro–1–, 2–, and 3–cresol ................................................................ 8
Table 3.1 Reaction rate constants and product fractions for O•+ with perfluorinated compounds ............. 47
Table 3.2 Reaction rate constants and product fractions for O•+ with methyl halides ................................ 49
Table 3.3 Bond dissociation energies of methyl halides ............................................................................ 50
Table 3.4 Reaction rate constants and product fractions for O•+ with organic molecules .......................... 54
Table 5.1 Reference acids .......................................................................................................................... 89
Table 5.2 Results of kinetic method and computational studies ................................................................ 93
Table 5.3 Experimental and theoretical thermochemical properties of nitrated azoles .............................. 94
Table 6.1 α-effect and Marcus barriers for SN2 reactions ........................................................................ 107
Table 6.2 α-effect for SN2 and E2 reactivity ............................................................................................ 108
Table 6.3 Water- methanol- and un-solvated nucleophiles in reactions with methyl chloride ................ 112
Table 6.4 Methanol solvated nucleophiles reactivity with methyl bromide............................................. 115
Table 6.5 The α-effect for reactions of un-solvated anions with methyl formate .................................... 118
Table 6.6 The α-effect for reactions of water-solvated anions with methyl formate ............................... 121

xiii

1 Introduction: Gas-Phase Ion Chemistry
1.1

Overview
Gas-phase ion chemistry is the study of the interactions of ions with molecules or atoms in the gas-

phase, including kinetics, mechanisms, and thermochemistry. Gas-phase ion chemistry is naturally
occurring throughout the universe in diverse regions such as the atmospheres of planets and moons, andin the interstellar medium. In addition, gas-phase ion chemistry can be employed to study the intrinsic
properties of molecules and reactions commonly found in the condensed-phase. Therefore, gas-phase ion
chemistry can directly or indirectly tell us about the behavior of chemical species in virtually all aspectsof the physical world!
Ions exist in abundance along the edge of planetary atmospheres where the flux of high-energy
photons is large. In Earth’s atmosphere, cations reach concentrations of 10 5 cm–3 at altitudes of ~100 km.1
Anions are also found in the Earth’s atmosphere, and they reach their peak concentration at lower
elevations due in part to the pressure-dependent three body stabilization process that forms the primary2anion O2–, equation 1.1.
e– + O2 + M → O2– + M

1.1

The [O] increases at higher altitudes, therefore reducing the concentration of all anions by removing O 2–
through an associative detachment reaction, equation 1.2.
O2– + O → O3 + e–

1.2

Equations 1.1 and 1.2 provide examples of reactions that are of interest to gas-phase ion chemists. In
particular, our laboratory measures rate constants for these types of chemical reactions.
Moving across the solar system, Saturn’s moon, Titan, is rich in gas-phase ion chemistry. Thanks tothe success of the Cassini spacecraft, the atmosphere of Titan has been directly sampled and analyzed,
1

Figure 1.1. Many believe that Titan’s atmosphere resembles that of early Earth. Both atmospheres
currently have large amounts of molecular nitrogen, but unlike modern Earth, Titan’s atmosphere is alsorich in methane (CH4), shrouding the moon in an organic haze. Because of its unique chemical makeup,
the ion-chemistry in Titan’s atmosphere differs from that on Earth. Ion laboratories around the world are
studying properties and reactions of ions found in Titan’s atmosphere in an attempt to better understandits composition and chemistry.3

Figure 1.1 Artist rendition of the Cassini spacecraft’s flyby of Titan(Image credit: NASA/JPL)The ‘space between the stars,’ or interstellar medium, was originally thought to be empty and devoidof any particles. However, researchers now know that this is not the case. Currently, over 200 chemical
species have been detected in space, and this list continues to grow each year.4 More than 30 of theseobserved species are ions! Gas-phase ion chemistry is very important within these regions as gas-phaseion neutral reactions will often proceed without a barrier. As the temperature in interstellar densemolecular clouds, Figure 1.2, can be quite low—on the order of ~10 K—the barrier-less reactions
between ions and molecules or atoms become very important because endothermic reactions andexothermic reactions with high activation barriers cannot proceed. Similar to the chemistry of theatmosphere, experimentally determined reaction rate constants are important for understanding theinterplay within the complex chemical environments of the interstellar medium. Experimentally2determined ion-neutral rate constants are invaluable to astrochemical modelers, who attempt tounderstand the intrinsic origins and relative abundances of the molecular species observed in space.

Figure 1.2 The pillars of creation
Image captured by the Hubble telescope in 2014. Ion-chemistry is rich in interstellar clouds, or
nebulae. (Image credit: NASA, ESA, and the Hubble Heritage Team (STScI/AURA))
Gas-phase ion chemistry can also provide information about the fundamental, intrinsic properties ofmolecules and reactions that occur in solvents because condensed-phase environments are difficult tounderstand on a fundamental level. In the condensed-phase, solvent molecules interact with the chemical
species under study. These solvent interactions can vastly complicate quantitative physical descriptions of
the chemistry taking place. For these reasons, molecular properties such as acidity, electron-affinity, andbond dissociation energy, are commonly determined in the gas-phase to ensure that the values obtained
are intrinsic in origin. One example of how condensed-phase interactions can mask intrinsic
thermodynamic properties involves the acidity ranking of simple organic alcohols. In the condensedphase, the acidity of the alcohol decreases as the size of the alcohol increases, i.e., H 2O > CH3OH >C2H5OH > i-C3H7OH > t-C4H9OH; the anionic conjugate bases are stabilized by solvation, and the3smaller species are more effectively solvated. However, the gas-phase acidity order is reversed: tC4H9OH > i-C3H7OH > C2H5OH > CH3OH > H2O.5 In the gas phase in the absence of solvent, the larger
conjugate bases can more effectively stabilize/delocalize the negative charge.
Similar to molecular properties, reaction kinetics and mechanisms may be investigated in the gasphase in efforts to understand the intrinsic nature of the processes observed. In addition, solvent
molecules can be added incrementally to gas-phase reactions in a process known as micro-solvation,
therefore linking observations in the gas-phase to those occurring in solution. One of the first microsolvation studies investigated the reaction of hydroxide and methyl bromide. 6 The rate constant for thehydroxide-methyl bromide reaction decreases by a factor of ten upon the addition of a single watermolecule to hydroxide, and the addition of three water molecules reduces the measured rate constant byover three orders of magnitude!
While the studies in this thesis may seem disparate, they are all united under the theme of gas-phaseion chemistry. I have studied highly reactive cations, such as the atomic oxygen radical cation, which
reacts rapidly with every molecule that has been investigated. I have characterized anions, such asdicyanamide, that are so stable that they do not even react with atoms and electronically excitedmolecules. I have investigated fundamental physical properties of ions, such as the proton affinities of thenitrated azolates, and I have studied phenomena pertaining to prototypical organic chemical reactions,such as the intrinsic origin of the α-effect. Furthermore, the ion chemistry reported in this thesis isrelevant to Earth’s atmosphere; Saturn’s moon, Titan; the interstellar medium; comets; ionic liquids; andphysical-organic chemistry. This thesis, therefore, will offer a flavor of many of the aspects of researchthat are typical of gas-phase ion chemistry.

4

1.2

Potential Energy Surfaces
The changes in potential energy occurring during a chemical reaction are often summarized in a

potential energy surface. For the generic SN2 reaction portrayed in equation 1.3, Figure 1.3 illustrates thepotential energy as a function of reaction coordinate.
X– + CH3Y → CH3X + Y–

1.3

Figure 1.3 Potential energy surface for an SN2 reaction
The double well potential for the SN2 reaction of X– + CH3Y. The transition state barrier-height isindicated as well as the overall exothermicity of the reaction.

The energy of the reactants at infinite separation is usually assigned a value of zero. As the reactants
approach, the attractive interaction causes a decrease in potential energy until the reactant ion-neutral5complex is formed with a binding energy of ΔHcomp. The reaction then proceeds through a transition statein which there is partial bond formation between the nucleophile and the substrate and partial bondcleavage of the leaving group. In the SN2 process, the transition state is characterized by a Walden
inversion. The reaction then proceeds to form the product ion-neutral complex, which separates into theionic and neutral product. Since the products lie at lower energy than the reactants, the process is
exothermic and the overall enthalpy change, ΔHrxn, is negative in value.
ΔH‡barrier represents the overall barrier height of the transition state. For this example, the transition
state barrier falls below the energy of the reactants, and the reaction can proceed without an externalsource of energy. However, if ΔH‡barrier is positive in value, the reaction will not occur; this situation
explains the lack of reactivity of methyl bromide with hydroxide that is hydrated by three watermolecules. Solvation lowers the energy of the reactants (with localized charge) more dramatically thansolvation lowers the energy of the transition state (with diffuse charge); therefore, the transition state lies
above the energy of the reactants.
As ΔH‡barrier becomes smaller (or more negative), the rate of the chemical reaction usually increases.
For example, catalysts increase the rate of a chemical reaction by providing an alternate pathway andlowering ΔH‡barrier.
Throughout this thesis, calculations are carried out to characterize the structures and energies of thereactants, ion-neutral complexes, transition states, and products. These results are presented in the formof potential energy surfaces that provide insight into the experimental results.

6

1.3

Thermochemistry

1.3.1

Gas-Phase Acidity

A thermodynamic value that can be experimentally determined with the selected-ion flow tube (SIFT;
to be described in Chapter 2) is the gas-phase acidity of a molecule. The gas-phase acidity, ∆acidG(AH), is
defined as the change in free energy associated with equation 1.4.
AH →A─ + H+

1.4

Alternatively, the proton affinity of the anion, PA(A‒) or ΔaH298(AH), is defined as the enthalpy change
for equation 1.4.7 Historically, the SIFT has been used to determine both ∆aG298(AH) and ΔaH298(AH).8, 9These quantities are directly related by the equation:
ΔaH298(AH) = ΔaG298(AH) + TΔaS298(AH)

1.5

In general, ∆aG298(AH) is lower than ΔaH298(AH) by ~7 kcal/mol, which is the typical value of
TΔaS298(AH).
In favorable cases, we can experimentally determine ∆aG298(AH) by measuring the forward andreverse rate constants for proton transfer between an ion and a neutral acidic molecule. That is to say, we
would measure the two rate constants of proton transfer, k1 and k2, for the following reactions:A─ + HB → AH + B─

1.6

B─ + HA → BH + A─

1.7This allows us to determine the equilibrium constant, Keq = k1/k2, for reaction 1.8.
A─ + HB ⇄ AH + B─

1.8

Subsequently, we can determine the free energy change using equation 1.9.
7∆(∆aG298) = -RTlnKeq = ∆aG298(AH) - ∆aG298(BH)

1.9

In this example, ∆aG298(BH) is the known acidity of the ‘reference acid,’ and the previously unknown
value, ∆aG298(AH), can be reported. However, this approach requires that the two molecules are close in
intrinsic gas-phase acidity, within approximately 4 kcal/mol.9 The experimentally determined ∆aG298(AH)
can be converted to ΔaH298(AH) by using TΔaS298(AH) values that are determined by electronic structure
calculations.
The SIFT is also used for directly determining ΔaH298(AH) values, which are often referred to as the
proton affinity of the conjugate base, PA(A‒). For these studies, the ‘bracketing’ technique is often
applied.8 In this method, a series of reference acids with known ∆aH298 are used, and their reactivity withan anion of unknown PA is monitored. If proton transfer from the reference acid to the anion is observed,
the proton transfer process is assumed to be exothermic. Experimentally, this corresponds to a decrease in
the parent anion signal and a proportional increase in the anion signal of the deprotonated reference acid.
By reporting the PA at which the proton transfer switches from exothermic to endothermic, a ‘bracket’ isplaced on the unknown PA.
Table 1.1 shows the results of a bracketing experiment used to determine the acidities of didehydro‒
1–, 2–, and 3–cresol. These studies were carried out in collaboration with Dan Nelson, and will be thesubject of a forthcoming publication.
Table 1.1 Bracketing results for didehydro–1–, 2–, and 3–cresol
Reference Acid
Experimental
Proton Transfer Observed?
∆aH298 (kcal/mol) 1-cresol 2-cresol 3-cresol
3-trifluoromethylphenol
3-mercaptopropionic acid1-chlorophenol
3-fluorophenol
propionic acidacetic acidphenolpentafluoro-1-propanol

337.0±2.2339.4±2.2343.4±2.4346.8±2.2347.4±2.2
348.1±2.2350.0±2.0355.4±6.1

Yes
Yes
No
No
No
No
No
No

Yes
Yes
Yes
Yes
Yes
Yes
No
No
8

Yes
Yes
No
No
No
No
No
No

To determine the ‘bracketed’ PA of the didehydro–cresols, the average acidity of the two ‘bracketing’acids is reported. For didehydro–1– and didehydro–3–cresol the two bracketing acids are 3–
mercaptopropionic acid and 1–chlorophenol, and the experimentally determined acidity is ∆aH298 = 341.4± 4.3 kcal/mol. For didehydro‒2–cresol, the bracketing acids are acetic acid and phenol, and ∆aH298 =
349.1 ± 3.1 kcal/mol. Note that the reported error for these experiments is quite large because theuncertainty in the PA from the two reference acids is considered.
Cook’s kinetic method can also be used to measure the PA of an anion. In this thesis, Cook’s kinetic
method was used to determine the PA of six nitrated azolates (Chapter 5), using a commercial triple
quadrupole instrument. Experimental and theoretical formulation of the kinetic method will be discussed
in Chapter 5.
For all of the methods discussed above, the selection of appropriate reference acids is very important.
1.3.2

Electron Affinity/Electron Binding Energy

Using the negative ion photoelectron spectrometer (NIPES) the electron binding energy of an anion
can be measured, i.e., the energy required to remove an electron from the anion. To perform thisexperiment, anions are generated and subsequently irradiated with photons of well-defined energy. If theenergy of the photon is sufficiently large it can detach an electron from the anion. The kinetic energy ofthe electron can then be measured with a hemispherical analyzer. This process follows equation 1.10.
A– + hν → A + e–

1.10

The intensities of the photoelectrons are monitored versus their kinetic energies.

The resulting

photoelectron spectrum provides electron binding energies, the splitting of electronic energy levels, andvibrational frequencies. At the origin of the spectrum, the difference between the photon energy (E hν) andthe kinetic energy of the ejected electron (KEe–) gives the electron affinity of A, equation 1.11.
EA(A) = Ehν – KEe‒

1.11
9

1.3.3

Thermochemical Cycles

Hess’s law states that, regardless of the multiple stages or steps of a reaction, the total enthalpy
change for the reaction is the sum of all changes. When both ∆ aH0(AH) and EA(A) are experimentallymeasured, a third thermochemical value, the bond dissociation energy, D0(AH), can be inferred. The
routine of determining D0(AH) from ∆aH0(AH) and EA(A) is known as the negative ion cycle; to perform
this operation the ionization energy of H, IE(H), must also be considered, equation 1.12.
H → H+ + e–

1.12This value is well known, IE(H)=13.59844 eV. The Hess’s law application of the negative ion cycle isprovided in Figure 1.4a, and a schematic of the thermochemical cycle is provided in Figure 1.4b.

10

Figure 1.4 Thermochemical Cycles(a) Using Hess’s law to determine D0(AH) from EA(A) and ΔaH0(AH). (b) A schematic overview of thenegative ion cycle as well as the proton affinity and appearance energy cycle.

111.3.4Temperature Corrections

There is a small discrepancy between the PA values determined from our experiment, ΔaH298(AH)
and those used to determine D0(AH) from the negative ion cycle, i.e. ΔaH0(AH). The difference in thesevalues is due to the temperature at which they are measured. The acidity experiments are performed atroom temperature when using the SIFT, but the EA(A) values determined using NIPES are intrinsicallymeasured at 0 K. In order to determine the bond dissociation energy, D0(AH), the thermal contribution for
ΔaH298(AH) must be corrected to 0 K. To perform this conversion, the heat capacities and their integral
from 0–298 K must be considered, equation 1.13.
298∆𝑎 𝐻0 (𝐴𝐻) = ∆𝑎 𝐻298 (𝐴𝐻) − ∫0

[𝐶𝑝 (𝐴− ) + 𝐶𝑝 (𝐻 + ) − 𝐶𝑝 (𝐴𝐻)]𝑑𝑇

1.13

When the heat capacities for a specific molecule or ion are unknown these values are theoreticallydetermined with electronic structure calculations.
1.3.5

Thermochemistry and Mass Spectrometry for Ion-Neutral Reactions

This thesis will explore a wide variety of ion-neutral reactions illustrated generically in equation 1.14.
A– + B → C– + D

1.14

Equation 1.14 represents a reaction between an anion, A–, and a neutral molecule, B, which advance totheir anion-neutral products, C– and D. This reaction also applies to cations where the charge is positive
instead of negative. In the Bierbaum lab, mass spectrometry is used to monitor the reaction progress; thedetails of the experiment will be described in the next chapter. Mass spectrometry only provides detectionof the ionic reactants and products, which can present some ambiguity for equation 1.14. As a reaction is
studied using mass spectrometry, the reactants, A– and B in equation 1.14, are usually well defined, but
the exact structure of C– and D must be inferred. The reactant ion, A–, is usually formed under conditions
that allow confident assignment of identity using the mass-to-charge ratio, m/z. The neutral reactant B is

12generally supplied by commercial vendors. As products are formed, only the ionic product, i.e., C – can be
identified by its m/z. By understanding the nature of the reactants, the m/z of C – will usually provide
enough information to confidently determine its identity, but this task can sometimes be difficult wheninvestigating the reactivity of large anions or molecules, especially when redundant atoms are present.
Furthermore, the structure of the neutral product, D, must be inferred; this process can be complex, but
thermodynamic arguments are used to suggest a reasonable structure for D.
Thermodynamics and a balanced chemical equation are used to report products for these ion-neutral
reactions investigated through mass spectrometry. In the case of common ions and neutrals, the heats offormation of the inferred products will be readily available. When investigating systems where the heatsof formation are unknown, theoretical heats of formation can be determined using computational
software. In some cases, potential energy surfaces, such as Figure 1.3, are computationally generated torationalize the structure and formula of the products. The experiments of this thesis explore only roomtemperature exothermic reactions; therefore, if the proposed products are higher in energy than thereactants, they can be discarded as physically unreasonable.
1.4

Kinetics
Kinetics is the study of the rates and mechanisms of chemical reactions. For many of the chemicalsystems studied within this thesis, we report experimental reaction rate constants and compare them totheoretical collision rate constants.
1.4.1

Collision Theory

In order for an ion-neutral reaction to occur in the gas-phase, the ion and neutral must meet, or‘collide.’ Therefore, a theoretical collision rate constant is useful to the experimentalist measuring
reaction rate constants. Fortunately, theoretical ion-neutral collision rate constants have been studied
somewhat comprehensively.

13

The development of capture collision theory began over a century ago when Langevin discussed theattractive potential between an ion, described as a point charge, and a polarizable neutral molecule. 10 With
contributions from Gioumousis and Stevenson,11 Langevin’s theory resulted in equation 1.15 that can beused to predict the collision rate constant, kL, between an ion and a polarizable neutral atom or molecule.
kL = 2πe(αn/µ)1/2

1.15

In equation 1.15, e, αn, and µ correspond to the elementary charge, the polarizability of the neutral atom
or molecule, and the reduced mass of the ion-neutral pair, respectively. Despite the simplicity of theformula, the Langevin collision rate provides a very good first-order approximation for the reaction rate
constant of an efficient reaction.
The Langevin collision rate constant is derived by considering only the polarizability of the neutral,
αn. In the last 40 years, several other ion-neutral capture collision theories have emerged by a morethorough treatment of the ion-neutral potential. Average dipole orientation (ADO) capture collision
theory includes the neutral permanent dipole moment, µD, in the treatment of the potential.12 The result isa theoretical collision rate constant that better matches experimentally determined reaction rate constants.
ADO theory was eventually simplified to parameterized trajectory theory (PTT) by applying an empirical
fit to the ADO model in an attempt to make ADO theory more accessible to experimentalists. 13 Later, in
2003, Eichelberger et al. formulated point-polarizable ion (PPI) collision theory which incorporated thepolarizability of the ion, αi, into the ion-neutral potential.14 This approach was developed because theexperimentally measured rate constants for linear carbon anion chains, Cn– and CnH–, with atomic oxygen
were higher than the theoretical collision rate constants. The theory was successful, and PPI provided abetter match to the experimental results. However, PPI only predicts substantial deviations from PTT
when the polarizability of the ion is very large (anions), and the polarizability of the neutral is very low(atoms.) Recently, Kent Ervin published a paper summarizing the effect of several variables on thetheoretical capture collision rate constant.15 His new treatment uses numerical methods to calculate the14theoretical capture collision rate by considering the polarizability of the neutral, αn; polarizability of theion, αi; dipole moment of the ion, µi; and quadrupole moment of the ion, Qi.
In summary, the treatment of the ion-neutral effective potential affects the resulting theoreticalcollision rate constant. Theoretical collision rate constants provide, at best, an upper limit to the reaction
rate constant for an ion-neutral pair that reacts upon every collision. When selecting a capture collision
theory, the experimentalist must determine which variables are necessary to accurately describe theattractive potential between the ion-neutral pair.
1.4.2

Reaction Efficiencies

Throughout the thesis reaction efficiencies will be reported. Reaction efficiencies compare
experimentally determined reaction rate constants to theoretical collision rate constants by taking the ratioof the rate constant (kexp) to the collision rate constant (kcol). Reaction efficiencies effectively normalize
experimentally determined rate constants, allowing for comparison of multiple ion-neutral reactions thathave different collision rate constants.
Equation 1.16 shows an example calculation to determine the reaction efficiency of O•+ with CF4. The
theoretical collision rate constant for O•+ and CF4 is 1.30 × 10-9 cm3/s, and the measured reaction rate
constant is 1.01 × 10-9 cm3/s.

𝑅𝑒𝑎𝑐𝑡𝑖𝑜𝑛 𝐸𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑐𝑦 =

1.5

𝑘𝑒𝑥𝑝
1.01 × 10−9=
= 0.78𝑘𝑐𝑜𝑙
1.30 × 10−9

1.16Electron Spin
The total sum of electron spin is highly relevant to some of the reactions in this thesis, particularly

reactions involving the atomic cation O●+ (4S) as well as the atomic neutrals H (2P), N (4S), and O (3P).
The ground-states of these atomic species are open-shell, i.e., they have one or more unpaired electrons

15

resulting in a total electron spin that is non-zero. The ground-state term symbols are presented in theparentheses to the right of the species; the superscripted number located to the left of the upper case letter
designates the multiplicity. The multiplicity indicates the total electron spin of the species. If all electrons
are paired, the atom or molecule is described as closed-shell with a multiplicity equal to one. If the atom
or molecule has a single unpaired electron, such as H atom or hydroxyl radical, the multiplicity equals
two. If the atom has two unpaired electron, such as O, the atom can have a multiplicity of 1 or 3
depending on the alignment (spin-up or spin-down) of the two unpaired electrons. Hund’s rule indicatesthat the lowest energy configuration for O atom will have its two unpaired electrons spin-aligned, giving
the ground state configuration of O atom a multiplicity of 3. Indeed, Hund’s rule holds for O atom, andthe electron configuration for O atom with a multiplicity of one is higher in energy than the configuration
with a multiplicity of three. Multiplicities of 1, 2, 3, 4, etc. are designated singlets, doublets, triplets, andquartets, respectively. While a multiplicity larger than 4 is possible, it is less common and does notprevail in the studies presented herein.
As a general rule, the total electron-spin is conserved for a reaction, i.e., ∆S = 0. However, small
species with high spin-states such as O●+ (4S), N (4S), and O (3P) are known to sometimes undergo spin
conversion reactions where ∆S = –1 in order to access more exothermic products. Spin-conversion
reactions occur on two or more potential energy surfaces. For example, the atomic oxygen radical cation,
O●+, has five valence electrons, and the ground-state electronic configuration is a quartet. On the contrary,a majority of molecules that contain an odd number of electrons will have a doublet ground state.
Therefore, when O●+ reacts with a molecule, the reaction will often move from a quartet to a doublet
potential energy surface.
An early account of spin-conversion for O●+ (4S) was reported in 1966 by Fehsenfeld, Fergeson, andSchmeltekopf.16 They observed spin-conversion for the reaction of O+(4S) + CO2(1Σ) → O2+(2Π) +CO(1Σ). In moving from reactants (S=3/2) to products (S=1/2), the total electron spin of the system has

16changed by –1, i.e., ∆S = ‒1. Therefore, this reaction violates the principle of spin-conservation. More
examples of spin-conversion reactions with O●+ (4S) will be discussed in Chapter 3.
1.6

Summary of This WorkA variety of gas-phase ion chemistry experiments relevant to environmental, materials, and physical-

organic chemistry are reported within this thesis. Kinetics, mechanisms, and thermodynamics are themajor foci of the ion chemistry that follows. This thesis will report a new electrospray ion source for theselected ion flow tube (SIFT), reactions of the atomic oxygen radical cation, the reactivity andthermochemistry of dicyanamide, the gas-phase acidities of the nitrated azoles, and experimental
investigations into the intrinsic origin of the α-effect.
Kinetic data for ion-neutral reactions are obtained using the SIFT apparatus, which will be describedin detail in the following chapter. In the SIFT, the rate constants are measured in the field-free reaction
flow tube. Therefore, the SIFT provides a benchmark for studying the kinetics of ion-neutral reactions.
Furthermore, a variety of ion generation techniques may be coupled to the SIFT, permitting the study of awide variety of ions in the gas phase. For the experiments in this thesis, the ions are generated withelectron ionization and electrospray ionization. Development of the electrospray ion source is a major
component of this thesis, and its implementation will be discussed in detail in the next chapter.
The reactions of the atomic radical cation O●+ (4S) will be discussed in Chapter 3. O●+ is an importantion in the atmosphere and the interstellar medium. With a recombination energy of 13.62 eV, O●+ ishighly reactive. Furthermore, O●+ is characterized by a large multiplicity, and it will often undergointeresting spin-conversion reactions as outlined in section 1.5.
Chapter 4 will discuss the reactivity, or lack thereof, of the dicyanamide anion, N(CN)2–. This ion is acommon anionic component of ionic liquids. Ionic liquids containing N(CN) 2– have been of great interest
to the Air Force Office of Scientific Research due to their hypergolic behavior when mixed with nitricacid. Despite the hypergolic behavior, N(CN)2– is an intrinsically stable anion in the gas phase. The studyof dicyanamide contrasts that of O●+ in many ways because dicyanamide is extremely unreactive. In
17addition to experimental studies of this anion, the hypergolic ignition mechanism of dicyanamide andnitric acid has been explored through electronic-structure calculations.
In Chapter 5, a study of the gas-phase acidities of the nitrated azoles will be reported. The acidities of
the nitrated azoles were determined using Cook’s kinetic method and a commercial triple quadrupole CID
experiment. The details of the triple quadrupole instrument and of Cook’s kinetic method will be
discussed in Chapters 2 and 5, respectively. The gas-phase acidities of six nitrated azoles (2- and 3nitropyrrole, 3- and 4-nitropyrazole, and 2- and 4-nitroimidazole) were experimentally measured withhigh precision. Because of agreement between the experimental and computational gas-phase acidities,
the gas phase acidities and electron affinities of 24 nitrated azole compounds were investigatedcomputationally. From these investigations, trends in the nitro-group placement and aza substitution’s
effect on the gas-phase acidity of the aromatic heterocycles are reported.
The final chapter of this thesis will summarize several gas-phase investigations of the α-effect. The αeffect describes an increased reactivity of α-nucleophiles, or nucleophiles in which the atom adjacent tothe attacking center contains a lone pair of electrons. A multifaceted phenomenon, the origin of the αeffect has been the topic of many discussions over the last 50 years. The effect was first observed insolution, and this has prompted researchers to examine whether the α-effect is intrinsic in origin, or if it isa result of solvation effects. In Chapter 6, the reactivity of both un-solvated and micro-solvated αnucleophiles will be investigated through gas-phase SN2, E2, BAC2, and PT organic reaction mechanisms.
The results suggest that the α-effect does have a small intrinsic contribution, but the micro-solvation
studies confirm that solvation plays an additional role in the α-effect.
1.7

References[1] M.R. Torr, D.G. Torr, The role of metastable species in the thermosphere, Rev. Geophys., 20 (1982)
91-144.
[2] R.P. Wayne, Chemistry of Atmospheres, 2, Oxford University Press, New York, 1991.

18

[3] A. Somogyi, M.A. Smith, V. Vuitton, R. Thissen, I. Komaromi, Chemical ionization in theatmosphere? A model study on negatively charged "exotic" ions generated from Titan's tholins by
ultrahigh resolution MS and MS/MS, Int. J. Mass Spectrom., 316 (2012) 157-163.
[4] D.E. Woon, Interstellar and circumstellar molecules, http://www.astrochymist.org, 2015.
[5] J.I. Brauman, L.K. Blair, Gas-phase acidities of alcohols, J. Am. Chem. Soc., 92 (1970) 5986-5992.
[6] D.K. Bohme, G.I. Mackay, Bridging the gap between the gas-phase and solution - transition in the
kinetics of nucleophilic displacement reactions, J. Am. Chem. Soc., 103 (1981) 978-979.
[7] S.G. Lias, J.E. Bartmess, Gas-phase ion thermochemistry, http://webbook.nist.gov/chemistry/ion/,
2015.
[8] N. Eyet, V.M. Bierbaum, Gas-phase acidities of thiocarboxylic acids, Int. J. Mass Spectrom., 265(2007) 267-270.
[9] A.J. Gianola, T. Ichino, R.L. Hoenigman, S. Kato, V.M. Bierbaum, W.C. Lineberger,
Thermochemistry and electronic structure of the pyrrolyl radical, Journal of Physical Chemistry A,
108 (2004) 10326-10335.
[10] P. Langevin, A fundamental formula of kinetic theory, Ann. Chim. Phys., 5 (1905) 245-288.
[11] G. Gioumousis, D.P. Stevenson, Reactions of gaseous molecule ions with gaseous molecules. V.
Theory, Journal of Chemical Physics, 29 (1957) 294-299.
[12] T. Su, M.T. Bowers, Ion-polar molecular collisions. Effect of ion size on ion-polar molecular rate
constants. Parameterization of the Average-Dipole-Orientation Theory, Int. J. Mass Spectrom. Ion
Process., 12 (1973) 347-356.
[13] T. Su, A.A. Viggiano, J.F. Paulson, The effect of the dipole-induced dipole potential on ion polar
molecule collision rate constants, Journal of Chemical Physics, 96 (1992) 5550-5551.
[14] B.R. Eichelberger, T.P. Snow, V.M. Bierbaum, Collision rate constants for polarizable ions, J. Am.
Soc. Mass Spectrom., 14 (2003) 501-505.
[15] K.M. Ervin, Capture Collisions of Polyynide Anions with Hydrogen Atoms: Effect of the Ion Dipole,
Quadrupole, and Anisotropic Polarizability, Int. J. Mass Spectrom., 378 (2014) 48-53.
[16] F.C. Fehsenfeld, E.E. Ferguson, A.L. Schmeltekopf, Thermal-Energy Ion-Neutral Reaction Rates.
III. The Measured Rate Constant for the Reaction O+(4S)+CO2(1Σ)→O2+(2Π)+CO(1Σ), Journal ofChemical Physics, 44 (1966) 3022-3024.

19

2 Experimental and Computational Methods
2.1

Selected Ion Flow Tube
The flowing-afterglow (FA) technique was developed in Boulder in the 1960s. 1 This techniquerevolutionized the field of gas-phase ion chemistry by providing a means to measure gas-phase rate
constants of ion-neutral reactions. Versatility of the FA technique was limited by the co-existence ofmany ions and precursor gases in the reaction region.2 This situation was remedied by the advent of theselected ion flow tube (SIFT) technique.3 The SIFT technique allowed chemists to study the reactivity ofions of a single mass-to-charge ratio with a neutral species in the absence of precursor gases.
The selected ion flow tube (SIFT) instrument was used for most of the studies described in thisthesis.4 A schematic of the instrument is shown in Figure 2.1.

Figure 2.1 Schematic of Selected Ion Flow Tube (SIFT)

20

The instrument consists of four regions; ion-production, ion-selection, reaction flow tube, and iondetection.
2.1.1

Ion Production Region

The SIFT technique employs a variety of ion sources. The Bierbaum lab has implemented electron
ionization, electrospray ionization, dc glow discharge, Penning ionization, and laser induced acoustic
desorption techniques to produce ions for the SIFT. Ions pertaining to the studies herein were generatedusing electron ionization and electrospray ionization. This thesis reports the design and implementation ofa new electrospray ion source; this source will be discussed in detail.
2.1.1.1

Electron Ionization

The housing for the electron ionization (EI) source (Figure 2.2) consists of a stainless steel flow tube
with an inner diameter of 4.75 cm. The EI source is mounted on a hollow stainless steel rod that can beaxially moved within the flow tube. Helium is introduced upstream of the EI filament through a metering
valve. The helium gas serves as both a discharge medium and a carrier gas. Precursor gases are
introduced through either stationary inlets located axially along the flow tube or one of two axiallyadjustable ring inlets. The ring inlets contain 24 equally spaced holes that promote even mixing of the
precursor gas. The absolute pressure of the helium and precursor gases is maintained at ~0.3 Torr by a
Stokes roots pump (190 L s-1). The gases are continuously pumped away generating a flow of helium overthe EI filament. About ~2-5V with 2.5A of current is applied across the yttrium oxide coated iridium
filament. The electrons are extracted from the filament using a repeller plate and an extraction grid. The
grid is biased with an attractive potential ~70V more positive than the filament voltage. The current on
the extraction grid is maintained at ~25 µA by a feedback loop. The EI source is floated at the potential
applied to the stainless steel flow tube.

21

Figure 2.2 Electron Ionization (EI) source diagram (a) and photograph (b).

The EI source may be used to generate a wide variety of ions; it contains several precursor inlets
permitting ion synthesis. One example is the production of hydrated methoxide, CH 3O–[H2O]n. This isachieved through a multistep ion synthesis. First, O– is generated from electron ionization of N2O,
equation 2.1.
N2O + e– → O– + N2

2.1

Next, O– extracts a hydrogen from methane to make hydroxide, equation 2.2.
O– + CH4 → OH– + CH3

2.2

OH– subsequently abstracts a proton from methanol to generate methoxide, equation 2.3.
OH– + CH3OH → H2O + CH3O–

2.3

Further downstream, water may be introduced to hydrate methoxide, equation 2.4. The methoxide-water
cluster is stabilized by collisions with He.
CH3O– + H2O + He → CH3O–[H2O] + He

2.4

22

Hydroxide, OH–, is used to make a wide variety of deprotonated species. Amide, NH2–, may also be used
for deprotonation synthesis; it is formed directly through electron ionization of ammonia, NH 3. However,the EI source is not limited to anions; it can generate a rich variety of cations as well. In particular, the EI
source was used to generate the oxygen radical cation, O•+, which will be discussed later in this thesis.
2.1.1.2

Electrospray Ionization

The electrospray ionization (ESI) source for the SIFT apparatus (Figure 2.3) is based on Moision andArmentrout’s design for their guided ion beam apparatus.5 Ions are generated at ambient pressure bypushing a solution through a conductive union biased at 2-3 kV. The flow rate of the solution is
maintained by a syringe pump; typical flow rates are 0.2-50 µL min-1. The charged solution exits througha silica carbide capillary, which has an inner diameter of ≤ 50 μm. This ‘emitter’ is mounted to an X-Y-Z
stage (not shown in Figure 2.3) and positioned in front of a stainless-steel capillary inlet that leads into thevacuum chamber. The capillary inlet is interchangeable, but for the experiments reported in this thesis, acapillary with an inner diameter of 0.76 mm and a length of 98 mm was used. While these exact
dimensions are somewhat arbitrary, using a shorter capillary increased the ion signal intensity. The
capillary inlet is seated in an aluminum block that can be heated from 50 °C to 200 °C using heater
cartridges and a temperature controller. The entrance capillary is typically biased at 100–300 V.

23

Figure 2.3 Electrospray ionization (ESI) source diagram.

The ions enter the vacuum chamber maintained at ~0.25 Torr by a high-capacity roots-blower pumpwhere they are then concentrated by an ion funnel. The ion funnel is based on the design of Smith andcoworkers.6 The ion funnel consists of 101 electrodes, but does not contain a ‘jet disruptor.’ The first 100
lenses have both RF and DC potentials applied to them. The RF amplitude is approximately 25 V peakto-peak and oscillates at a frequency that can be adjusted; for these experiments, a frequency of 1.5–2.8
MHz was used. The RF potential applied to each adjacent electrode is 180 degrees out-of-phase. The RF
potentials focus the ions near the center axis of the funnel. The applied DC voltages create a potential
gradient along the ion funnel. Both the entrance and exit DC potentials are adjustable, and the DC
potential applied across the ion funnel is typically 10–150 V. The potential gradient transmits the ions
towards the differentially pumped mass-selection region. The 101st electrostatic lens of the funnel is

24

supplied with an independently adjustable DC potential. A photograph of the ion funnel is provided in
Figure 2.4. Note that each of the potentials listed are referenced to a common ground (Figure 2.5).

Figure 2.4 Ion funnel photograph

Figure 2.5 shows the wiring layout for the ESI source; the common ground is the source housing. The
ion funnel requires five electrical inputs. Two of the inputs provide RF potentials to the lenses of the ion
funnel; the remaining electrical inputs provide DC potentials to the entrance, exit, and final lens of the ion
funnel. The final, or 101st, lens of the ion funnel is capacitively coupled to ground to remove RF noise. A
100 nF capacitor is used for this purpose.

25

Figure 2.5 Electrospray ionization source wiring schematic

A signal generator (BK Precision, model 4040A) is used to provide a sine wave of tunable frequency.
The sine wave is then sent through an RF amplifier (Electronics & Innovation, model 240L), which
increases the amplitude from ~1 to ~25 Vp2p. The amplified wave then enters an RF balun (indicated by asquare of dashed lines in Figure 2.5) that sends out two RF signals, which are 180° out of phase. For
instructions on building the RF balun, as well as dimensions on the ion funnel electrodes and spacers,
please see Appendix A.
The aluminum heating block is electrically isolated from the chamber housing. The heating block isbiased at the same potential as the ‘entrance capillary’ that sits within it. It uses three heater cartridges(Omega, model CCS-01115/120) that are connected externally, and the temperature is regulated by athermal controller (Love Controls, series 16A). The temperature heating range is 30-175 °C. The entrance
capillary (1.6 mm O.D.) is seated within the heating block using an NPT to Swagelok union. The entrance
capillary is removable, which facilitates cleaning and replacement.

26

The emitter assembly includes a syringe pump (Harvard Apparatus, model Pump11Elite), which
depresses a syringe (Hamilton, Model 1750 RN SYR) connected to fused silica capillary tubing(Polymicro technologies; 75 μm I.D., 360 μm O.D.) using a microunion assay (Idex, model P-720) and atubing sleeve (Idex, model F-185X). Downstream, the fused silica carbide tubing is connected to a
conductive union (Idex, model M-572) that is connected to a thinner silica carbide tube (Polymicrotechnologies; 50 μm I.D., 180 μm O.D.). The emitter tip I.D. is drawn thinner by heating and stretching.
The emitter tips are cut with a fused silica carbide tubing cutter (Idex, model FS-311). The newly formed
tips are inspected with a microscope to ensure that they are not malformed. Using pre-manufactured
emitter tips is also an excellent option.
The ESI source can transmit ions from solution into the gas-phase. Therefore any ion that can begenerated in solution can, in principle, be transferred into the gas-phase. Typical solution concentrations
range from 1 × 10-5 to 4 × 10-4 M. Several solvents, i.e., water, methanol, acetonitrile, etc., can be usedwith the electrospray. Most ESI experiments performed with the SIFT employ acetonitrile as the solvent.
Furthermore, the pH of the solution can be adjusted to form protonated or deprotonated species.
The ESI source performs extremely well. Stable ions can be generated with high intensities
comparable to the EI source. Figure 2.6a shows a spectrum of 1-butyl-2,3-dimethylimidazolium cationthat was generated from a 2.4 × 10-4 M solution of 1-butyl-2,3-dimethylimidazolium tetrafluoroborate inmethanol (99.8%, EMD Millipore). This ion was mass-selected, injected into the flow tube, and detected.
Figure 2.6b shows a spectrum of dicyanamide anion that was generated from a 6 × 10 -5 M solution of
sodium dicyanamide in acetonitrile (99.8%, Sigma-Aldrich). In this case, all negative ions were injected
and detected, but the desired dicyanamide anion dominates the mass spectrum.

27

Figure 2.6 Electrospray ionization source spectra
The cation (a) 1-butyl-2,3-dimethylimidazolium and anion (b) dicyanamide are generated in highabundance with the ESI source.

28

2.1.2Ion Selection Region

The ion selection region designates the experiment a selected ion flow tube; it differentiates FA from
SIFT techniques. The ions exit the source region through either an electrostatic nose cone (for EI) or anelectrodynamic ion funnel (for ESI) and enter a differentially pumped region maintained at ~10–5 Torr by
a 10-inch diffusion pump (Edwards Diffstak, 2300 L s–1) backed by a mechanical pump (Welch 1374, 650
L min–1). The ions pass through six focusing lenses directing them into a quadrupole mass filter (Extrel #4-270-9, four rods 16 mm O.D. × 220 mm). The quadrupole is housed in a smaller chamber differentiallypumped to ~10–6 Torr by a 6-inch diffusion pump (Edwards Diffstak, 805 L s–1) and backed by amechanical pump (Welch 1397, 500 L min–1). RF and DC potentials are applied to the quadrupole by an
Extrel power supply (# 011-1, #011-14) with a resonant frequency of ~1.5 MHz. The quadrupole is used
to mass select (or SIFT) ions of a single m/z with unit resolution. The quadrupole may transmit ions up to~400 m/z. The mass selected ion current is focused by three electrostatic lenses into an ion injection
orifice. The injection orifice contains a Venturi gas-inlet consisting of two concentric annuli. Helium isforced through the concentric circular inlets, and the ratio of flow between the two annuli is adjusted tominimize back streaming of helium into the ion selection region. The Venturi inlet is discussed in detailin the original report of the SIFT apparatus.4
Carbon monoxide (CO) is the precursor gas used to generate the positive ion spectra in Figure 2.7.
Figure 2.7a illustrates that the EI source produces many ions even when a single precursor gas (CO) isused. Figure 2.7b shows the EI source of mass selected and optimized m/z 16 (O•+).

29

Figure 2.7 EI source positive ion spectra of CO: (a) no mass selection (b) m/z 16 mass selected

2.1.3

Reaction Flow Tube

After mass selection, the ions are injected into the reaction flow tube. The ions are entrained inhelium as they pass through the Venturi inlet; the ions are quickly thermalized through multiple collisions
with the helium carrier gas (~106 collisions s–1). The flow tube is a long (117 cm), field-free stainless steel
cylinder with an inner diameter of 7.3 cm. The ions in the flow tube are not transferred by electrostatic
lensing, but instead are carried by a ~170 std cm3 s–1 flow of helium maintained at a pressure of ~0.45
Torr by a high-capacity Stokes roots pump (755 L s–1). The ions take ~10 ms to travel the flow tube (vion ~104 cm/s). The first ~45 cm of the flow tube is used to establish laminar flow. The remaining ~70 cm of30

the flow tube is used to study ion-neutral reactions. Neutral reactants are introduced through sevenevenly-spaced ring inlets (section 2.1.5.1), a microwave discharge apparatus (section 2.1.5.2), or a lowvolatility neutral reagent source (Chapter 5).
2.1.4

Ion Detection Region

When the product ions and remaining reactant ions reach the end of the reaction flow tube, they are
guided into the ion detection region through an electrostatic nose cone with an orifice of 0.5 mm. The
nose cone is mounted onto a carrier plate biased with an independent electrostatic potential. Upon exiting
the nose cone lens, the ions enter a region containing five electrostatic focusing lenses that isdifferentially pumped by a 6-inch diffusion pump (Edwards Diffstak, 805 L s–1) backed by a mechanicalpump (Welch 1397, 500 L min–1). Next, the ions enter a triple quadrupole mass spectrometer (TQMS;
Extrel, Model 5221 C50 TQMS: 1.2 MHz, 300W for Q1 and Q3, 200W for Q2, 19 mm O.D. rods). The
triple quadrupole assembly lays across two chambers each differentially pumped by 10-inch diffusion
pumps (Edwards Diffstak, 2300 L s–1) and backed by mechanical pumps (Welch 1374, 650 L min–1); the
TQMS chambers have pressures of 10–6 and 10–7 Torr, respectively. The TQMS is coupled to an electron
multiplier (DeTech, 402A-H) with a gain of ~108. The ion pulses are processed by Extrel Software(Merlin Automation System, version 3.0). The reported SIFT experiments use the TQMS for single MSexperiments only. To perform collision induced dissociation experiments reported in this thesis, acommercial instrument is employed (section 2.2).
2.1.5

Ion-Neutral Reaction Rate Constant

Both the FA and SIFT are used to measure ion-neutral reaction rate constants. To understand thisprocess, consider the generalized ion-neutral reaction:
I± + B → Products

2.5

31

I± represents an ion (cation or anion), and B represents a neutral reactant. B may be a molecule or an
atom. Bimolecular reactions are defined as second-order because their rate depends on the concentrationof two reacting species. Consequently, ion-neutral reactions follow the kinetic rate equation:𝑑[𝐼 ± ]= −𝑘𝑒𝑥𝑝 [𝐼 ± ][𝐵]𝑑𝑡

2.6

The reaction rate depends on both [I±] and [B], but the experiments within this thesis follow pseudo–firstorder reaction kinetics because [I±] <<< [B] by at least four orders of magnitude. [I±] is on the order of~105 ions cm–3; [B] ~ 1011 particles cm–3. Rearranging equation 2.6:
𝑡

𝑡
1
±
𝑑[𝐼
]
=
−𝑘
[𝐵]
∫
𝑑𝑡
𝑒𝑥𝑝
±
0 [𝐼 ]
02.7∫

Equation 2.7 is integrated over the reaction time, t; kexp and [B] are removed from the integral because
they are constant under pseudo–first-order reaction kinetics. Evaluating the integral yields the pseudo–first-order reaction rate law:𝑙𝑛

[𝐼 ± ]𝑡= −𝑘𝑒𝑥𝑝 [𝐵]𝑡
[𝐼 ± ]02.8

Equation 2.8 reveals kexp can be experimentally determined by monitoring the change in [I±] with respect
to varying the reaction time, t, or varying the neutral concentration, [B].
2.1.5.1

Reaction Rate Constant: Seven-Inlet Method

Volatile neutral reactants are introduced into the flow tube through seven evenly-spaced ring inlets.
The first ring inlet lies 13.5 cm upstream from the ion-detection region, and the remaining ring inlets arespaced by 9.7 cm. Each ring inlet contains seventeen radially spaced holes promoting even mixing. The
inlets are coupled to a glass-rack by solenoid valves. The neutral-reagent flow rate is established with a32metering valve and measured by monitoring the change in pressure in the glass rack (of known volume)with respect to time (ΔP/Δt.)
Due to their axial distribution along the reaction flow tube, each neutral ring-inlet corresponds to adifferent reaction distance. Furthermore, the SIFT is a constant flow system where the velocity of helium
carrier gas, and therefore the velocity of the ion (vion), is constant. Thus, the distance of the ring-inlet fromthe ion-detection chamber renders a temporal relationship to the reaction kinetics, and reaction rate
constants are measured by monitoring the parent ion signal, [I±], as a function of reaction distance.
To formulate the experimental rate constant, equation 2.8 is rearranged to yield the relationship:ln[𝐼 ± ]𝑡 = −𝑘𝑒𝑥𝑝 [𝐵]𝑡 + ln[𝐼 ± ]0

2.9

Incorporating the relationship between velocity and time into equation 2.9 yields the expression:ln[𝐼 ± ]𝑡 =−𝑘𝑒𝑥𝑝 [𝐵]𝑧
+ ln[𝐼 ± ]0
𝑣𝑖𝑜𝑛

2.10

Where z is the reaction distance, or axial–flow tube distance between the ion-detection chamber and thering-inlet. This parallels the equation y = mx + b where m = −kexp×[B]/vion. The ion-neutral reaction rate
constant, kexp, for reaction 2.5 is determined by plotting ln[I±] vs. reaction distance. Figure 2.8 shows anexample for the reaction of O•+ with CH3F.

33

Figure 2.8 [O•+] vs. reaction distance for the O•+ + CH3F reaction

For Figure 2.8, each of the data points represents one of the seven neutral reagent ring-inlets. Because[CH3F] and vion are known, the reaction rate constant, kexp[O●+ + CH3F], can be determined from the slope
in Figure 2.8.
2.1.5.2

Reaction Rate Constant: Microwave Discharge

N(4S) and O(3P) are introduced as neutral reagents by using a microwave discharge connected to thereaction flow tube.7 The experimental setup is shown in Figure 2.9. A microwave power supply (Opthos34

MPG-4) provides ~40-50W of microwave radiation to the Evenson cavity. Molecular nitrogen is passed
through the Evenson cavity generating N atoms, equation 2.11.
N2 + hν → 2N

2.11

The dominant active species released into the flow tube from the microwave discharge of N 2 is N(4S).
While concentrations of metastable excited species, N(2D), N(2P), and N2(3Σu+) are considerable near thedischarge plasma, these excited species are destroyed within a few milliseconds and will not survive thepath to the flow tube.8-11

Figure 2.9 Microwave discharge apparatus

O atoms are generated by titrating [N] with NO, equation 2.12.
N + NO → N2 + O

2.12

35

To determine [N] and [O] for kinetic formulation, a titration plot is generated by monitoring the log of theintensity of the parent ion as a function of NO flow, Figure 2.10.

Figure 2.10 Microwave discharge titration plot

The black triangle represents the intensity of the parent ion signal before the microwave discharge is
ignited. The depletion in ion signal between the black triangle and the red circle occurs due to the reactionof the ion with N atom. The first red circle represents [N]0. The flow of NO is established using ametering valve, and monitored using a Tylan flow meter. As the flow of NO is increased, [N] is converted
to [O]. The lowest point in the titration plot occurs where [O] is at a maximum. At this point, [NO] added =[N]0 =[O]titration point. After the titration plot is obtained, a reaction rate constant for equations 2.13 and 2.14
can be determined.
I± + N → Products

2.13

I± + O → Products

2.14

While the seven inlet method monitors reactant ion falloff vs. reaction distance and determines the rate
constant by evaluating the slope, the microwave discharge method determines the reaction rate constantfrom a single experimental reaction distance (distance ∝ t), equation 2.15.

36

−𝑘𝑒𝑥𝑝 =𝑙𝑛[𝐼 ± ]𝑡[𝐼 ± ]0⁄
[𝐵]𝑡

2.15

Equation 2.15 arrives from the rearrangement of equation 2.8, where [B] represents either [N] or [O], bothof which are determined from the titration plot.
2.1.5.3

Determination of Variables

The velocity of the helium carrier gas (vHe) is determined from the helium flow rate (FHe), pressure(PHe), and the cross sectional area of the reaction flow tube (πr2), equation 2.16.
𝑐𝑚3
𝐹𝐻𝑒 (𝑆𝑇𝑃 𝑠 )𝑐𝑚
760(𝑇𝑜𝑟𝑟)𝑇(𝐾)𝑣𝐻𝑒 ( ) = 2××2𝑠𝜋𝑟 (𝑐𝑚 )𝑃𝐻𝑒 (𝑇𝑜𝑟𝑟)
𝑎𝑡𝑚
273.16(𝐾)2.16

The velocity of the ions relative to the velocity of helium was determined experimentally. In fact, the ionsmove through the flow tube slightly faster than the carrier gas. A correction constant (α), determined bothcomputationally and empirically, is multiplied by vHe to determine vion, equation 2.17.
𝑐𝑚
𝑐𝑚
𝑐𝑚
𝑣𝑖𝑜𝑛 ( ) = 𝛼 × 𝑣𝐻𝑒 ( ) = 1.6 × 𝑣𝐻𝑒 ( )
𝑠
𝑠
𝑠

2.17

The flow of the neutral reagent, FB, may be directly measured with a flow meter. Otherwise, FB isdetermined experimentally by monitoring the change in pressure within a known volume with respect totime, equation 2.18.

𝐹𝐵 (𝑇𝑜𝑟𝑟

𝑐𝑚3∆𝑃(𝑇𝑜𝑟𝑟)) = 𝑉𝑘𝑛𝑜𝑤𝑛 (𝑐𝑚3 ) ×
𝑠
∆𝑡(𝑠)

2.18

The concentration of the neutral reagent, [B], in the flow tube is determined using the flow of the neutral
reagent (FB), helium velocity (vHe), and the cross sectional area of the reaction flow tube (πr2), equation
2.19.
37

𝑐𝑚3

𝐹𝐵 (𝑇𝑜𝑟𝑟)
𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠𝑠
16 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠[𝐵] ( 𝑐𝑚3 ) = 2 2(𝑐𝑚3 ∙𝑇𝑜𝑟𝑟)
𝑐𝑚 × 3.535 ∙ 10
𝜋𝑟 (𝑐𝑚 )𝑣𝐻𝑒 ( )

2.19

𝑠

‘Particles’ is a unitless quantity, and may represent atoms or molecules. The reaction rate constant will
now be considered. Rearrangement of equation 2.10 yields equation 2.20.

𝑘[𝐼± +𝐵] (𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠 ∙ 𝑐𝑚3−𝑑𝑙𝑛[𝐼 ± ] 𝑣𝑖𝑜𝑛
×
)=
𝑠
𝑑𝑧
[𝐵]

2.20

Substituting for [B] and vion yields equation 2.21.
2−∆𝑙𝑛[𝐼 ± ]𝛼𝐹𝐻𝑒𝑘[𝐼± +𝐵] =×2∆𝑧
𝐹𝐵 𝑃𝐻𝑒𝜋𝑟 2

2.21

Equation 2.21 formulates the rate constant in terms of experimentally measurable variables.
2.1.6

Product Fraction

Rate constants in the SIFT are measured by monitoring Δln[I±], but how are products
measured/reported? Indeed, ion-neutral reactions will often generate various products through competingproduct generation channels. When this occurs the product fraction must also be reported. Consider theion-neutral reaction shown in Scheme 2.1. While this scheme considers a real example using cations theprinciples illustrated also apply to anions.

Scheme 2.1

38

As indicated in Scheme 2.1, the total ion spectra would show peaks corresponding to O+, CH2F+, CH3F+,
and CH4F+. The primary and secondary reactions need to be disentangled. This can be accomplished bymonitoring the product ion intensities, [CH2F+], [CH3F+], and [CH4F+], as a function of reaction (or inlet)
distance, and extrapolating to ‘zero’ reaction distance. At ‘zero’ reaction distance, the concentration of thesecondary product ions will approach zero. The ratio of the primary product ions at ‘zero’ reaction
distance is reported as the product fraction (PF). Figure 2.11 shows a product fraction plot for the reactionof O•+ + CH3F. The PF is reported as the y-intercept for the linear curves. CH4F+ is a secondary productsince its intensity extrapolates to zero.

Figure 2.11 Product Fraction plot for O•+ + CH3F

39

The PF may also be determined as a function of reactant ion depletion. Furthermore, the PF may be used
for determining the relative reaction rate constants because ktotal = k1+ k2. The relative reaction rate
constants are directly proportional to the relative abundance of the primary products, equations 2.22 and
2.23.

2.2

𝑃𝐹C𝐻2 𝐹+ =[C𝐻2 𝐹 + ]𝑘1=++[C𝐻2 𝐹 ] + [𝐶𝐻3 𝐹 ]𝑘1 + 𝑘2

2.22

𝑃𝐹C𝐻3 𝐹+ =[C𝐻3 𝐹 + ]𝑘2=[C𝐻2 𝐹 + ] + [𝐶𝐻3 𝐹 + ]𝑘1 + 𝑘2

2.23

Triple Quadrupole Mass Spectrometer
In addition to the SIFT technique, collision induced dissociation experiments were performed on acommercial triple quadrupole mass spectrometer (AB Sciex, 4000 QTrap). 12 This apparatus is located at
the University of Colorado’s core mass spectrometry facility on east campus. While the detection regionof the SIFT contains a triple quadrupole, the collision energies are not well defined, and ion intensities are
more difficult to optimize. Figure 2.12 provides a schematic for the QTrap.

Figure 2.12 4000 QTrap triple quadrupole mass spectrometer

Ions are generated using a ‘turbo spray’ electrospray ionization source; the spray trajectory is off axis
relative to the path of the ions. Due to the high sensitivity of the QTrap, dilute solutions with analyte40

concentrations of 7.5 µM or smaller in methanol were used to generate ions. Infusion rates were
controlled by a syringe pump, and typical flow rates varied from 5 to 40 µL/min. After generation, theions are extracted from the spray by attractive potentials. The ions pass through a ‘curtain’ of N 2 gas(represented by blue arrows). The curtain gas flow, which is tunable in the software, desolvates the ions.
Beyond the curtain gas is a differentially pumped region maintained at approximately 2 Torr. The
entrance region is separated from the next differentially pumped region by a skimmer. The skimmer
chamber contains a radio frequency (RF)-only quadrupole ion guide, Q0, and is maintained at ~3 × 10-3
Torr. An aperture separates the next differentially pumped chamber which contains a Brubaker lens andQ1; this chamber is maintained ~1 × 10–6 Torr. Q0, Q1, and the Brubaker lens are capacitively coupled to
one another by an RF drive voltage. Q1 is used to mass-select the precursor ion of interest by applying a
DC potential offset. The collision cell, q2, is placed in an isolated enclosure containing 3 mm entrance
and exit apertures. N2 gas (blue arrow) can be introduced into the q2 enclosure to perform collisioninduced dissociation. For the experiments reported in this thesis, the collision gas pressure is maintained
as low as possible (~6 × 10-3 Torr) in order to approach single collision conditions. The collision energy
for the QTrap is defined by the DC potential differences between Q0 and q2. The third quadrupole, Q3, isused to analyze the collision induced dissociation fragment ions. Q3 is coupled to an electron multiplier.
The QTrap was utilized to study the gas-phase acidities of the nitrated-azoles using Cook’s kinetic
method. This technique will be discussed in detail in Chapter 5.
2.3

Computational Methods
Throughout this thesis, computational methods are used to theoretically examine the experimental

systems under investigation. Many ab initio methods have been employed. The methods employed are
Becke-Lee-Yang-Parr exchange-correlation functional (B3LYP), coupled cluster (CCSD), and MøllerPlesset perturbation theory (MP2) in conjunction with either Pople or Dunning basis sets. Most of the41

theoretical computations reported within this work were carried out using the Gaussian 09 software
suite.13
Computational methods are employed to optimize geometries for global minima of reactants,
products, and intermediates. The energies of the optimized geometries are then obtained. Frequency
calculations are then carried out to establish the optimized structure as either a transition state(characterized by one imaginary frequency) or intermediate (no imaginary frequencies).
Structures of transition states are important for explaining the kinetics and dynamics of a reaction
system. Transition state optimization can sometimes be challenging. However, using the relaxed potential
energy surface scans along a fixed reaction coordinate (i.e., the elongation of a bond or the rotation about
an angle) can help to find the transition state by revealing a first-order saddle point. The geometry of thesaddle point can be used for further optimization, and the frequency calculation will determine if thestructure is a true transition state with only one imaginary frequency. As a proper transition state, theimaginary vibrational frequency should be consistent with the reaction coordinate of interest. Transition
states are further confirmed by employing the intrinsic reaction coordinate (IRC) analysis.
For each of the studies presented in this thesis, appropriate computational details will be discussed in
the relevant chapters.
2.4

References[1] E.E. Ferguson, F.C. Fehsenfeld, A.L. Schmeltekopf, Flowing afterglow measurements of ion-neutral
reactions, Adv. in Atom. Mol. Phys., 5 (1969) 1-56.
[2] V.M. Bierbaum, Go with the flow: Fifty years of innovation and ion chemistry using the flowing
afterglow, Int. J. Mass Spectrom., 377 (2015) 456-466.
[3] N.G. Adams, D. Smith, The selected ion flow tube: A technique for studying ion-neutral reactions, Int.
J. Mass Spectrom. Ion Phys., 21 (1976) 349-359.
[4] J.M. Van Doren, S.E. Barlow, C.H. DePuy, V.M. Bierbaum, The tandem Flowing Afterglow-SIFTDRIFT, Int. J. Mass Spectrom. Ion Process., 81 (1987) 85-100.

42

[5] R.M. Moision, P.B. Armentrout, An electrospray ionization source for thermochemical investigation
with the guided ion beam mass spectrometer, J. Am. Soc. Mass Spectrom., 18 (2007) 1124-1134.
[6] R.T. Kelly, A.V. Tolmachev, J.S. Page, K.Q. Tang, R.D. Smith, The ion funnel: Theory,
implementations, and applications, Mass Spectrom. Rev., 29 (2010) 294-312.
[7] Z.C. Wang, C.A. Cole, N.J. Demarais, T.P. Snow, V.M. Bierbaum, Reactions of azine anions withnitrogen and oxygen atoms: Implications for Titan's upper atmosphere and interstellar chemistry, J.
Am. Chem. Soc., 137 (2015) 10700-10709.
[8] S.N. Foner, R.L. Hudson, Mass spectrometric studies of metastable nitrogen atoms and molecules in
active nitrogen, J. Chem. Phys., 37 (1962) 1662-1667.
[9] C.L. Lin, F. Kaufman, Reactions of metastable nitrogen atoms, J. Chem. Phys., 55 (1971) 3760-3770.
[10] J.A. Meyer, D.W. Setser, D.H. Stedman, Energy transfer reactions of N2(A3Σu+). II. Quenching andemission by oxygen and nitrogen atoms, J. Phys. Chem., 74 (1970) 2238-2240.
[11] R.A. Young, G.A. St. John, Experiments on N2(A3Σu+). I. Reaction with N, J. Chem. Phys., 48 (1968)
895-897.
[12] J.W. Hager, A new linear ion trap mass spectrometer, Rapid Commun. Mass Spectrom., 16 (2002)
512-526.
[13] M.J. Frisch, G.W. Trucks, H.B. Schlegel, G.E. Scuseria, M.A. Robb, J.R. Cheeseman, G. Scalmani,
V. Barone, B. Mennucci, G.A. Peterson, H. Nakatsuji, M. Caricato, X. Li, H.P. Hratchian, A.F.
Izmaylov, J. Bloino, G. Zheng, J.L. Sonnenberg, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J.
Hasegawa, T. Nakajima, Y. Honda, O. Kitao, H. Nakai, T. Vreven, J.A. Montgomery Jr., J.E. Peralta,
F. Ogliaro, M. Bearpark, J.J. Heyd, E. Brothers, K.N. Kudin, V.N. Staroverov, R. Kobayashi, J.
Normand, K. Raghavachari, A. Rendell, J.C. Burant, S.S. Iyengar, J. Tomasi, M. Cossi, N. Rega, J.M.
Millam, M. Klene, J.E. Knox, J.B. Cross, V. Bakken, C. Adamo, J. Jaramillo, R. Gomperts, R.E.
Stratmann, O. Yazyev, A.J. Austin, R. Cammi, C. Pomelli, J.W. Ochterski, R.L. Martin, K.
Morokuma, V.G. Zakrzewski, G.A. Voth, P. Salvador, J.J. Dannenberg, S. Dapprich, A.D. Daniels,
Ö. Farkas, J.B. Foresman, J.V. Ortiz, J. Cioslowski, D.J. Fox, Gaussian 09, Gaussian, Inc.,
Wallingford CT, 2009.

43

3 Gas-Phase Reactions of the Oxygen Radical Cation, O•+

3.1

Introduction
The oxygen radical cation, O•+, is abundant in many regions of the Universe. In the Earth’s

atmosphere, this ion dominates the ionic constituents in the F-region of the ionosphere, reaching
concentrations of 4 × 105 particles cm–3. O•+ is commonly generated by photoionization from UVradiation, but it may also be generated by charge exchange with H+.1, 2 These phenomena also account for
the existence of O•+ in the diffuse clouds of the interstellar medium (ISM) and other non-terrestrial
atmospheres.3 Furthermore, O•+ has been detected in flames4 and in industrial plasma-etching processes
using O2.5, 6This chapter investigates the reactivity of O•+ with a variety of neutral molecules. It summarizes two
research publications. The first publication explored the reactions of O•+ with the perfluorinatedcompounds CF4, SF6, and SF5CF3, as well as the methyl halides CH3F, CH3Cl, CH3Br, and CH3I to
examine trends in reactivity.7 The second O•+ publication explored reactivity of organic compounds withvarious functional groups including methanol, acetaldehyde, acetic acid, methyl formate, dimethyl ether,and acetone.8 The majority of these reactions are reported here for the first time. While these studies
contribute to the general knowledge of gas-phase reactions, they also taught me how to investigate ionneutral reactions both experimentally and theoretically. Furthermore, the wide range of neutral molecules
investigated (from perfluorinated compounds to organic molecules) permits some profound insight into
the overall reactivity of O•+.
All in all, the chemical reactivity of O•+ is quite interesting! With its large recombination energy (RE)of 13.62 eV, O•+ can ionize every carbon-containing molecule except carbon monoxide/dioxide.
Furthermore, O•+ will sometimes undergo electron-spin conversion processes, an occurrence that is often

44

described as forbidden. Even in situations where spin-conserved reaction channels generate highly
exothermic products, the spin-conversion reactions channels are likely in competition.
3.2

Experimental Methods
For the following experiments, electron ionization (EI) was used to produce the atomic oxygen

radical cation, O•+. The selected ion flow tube (SIFT) apparatus and the EI ion source are described indetail in Chapter 2. Chapter 2 also presents a mass-selected spectrum of O•+.
When generating O•+ to study ion-neutral reactions, it is important to ensure that only a single
electronic-state is present. Using the methods of Smith et al.,9 a high percentage of ground-state, O•+(4S),
ions are produced. In the EI source, a mixture of 0.10 Torr CO (Airgas, CP grade; ionization energy =14.01 eV) and 0.30 Torr He (Airgas, 99.999%) undergoes electron ionization to form both ground- andexcited-state O•+. Since the ionization energy of CO is between the recombination energies of the groundand excited-state O•+ ions, the excited-state ions, O•+(2D), O•+(2P), etc., are rapidly destroyed in the ion
source by charge-transfer reactions with CO (kexp = 1.3 × 10-9 cm3 s-1) 10, while the ground state, O•+(4S),
does not react. Therefore, by introducing a large concentration of CO in the source, the excited-state O•+
ions are quenched. To confirm the absence of excited states in the experiments, O+· is injected into thereaction flow tube, and CO was introduced downstream. The observed rate constant, < 3.3 × 10-12 cm3 s-1,
indicates that less than 0.25% of O•+ introduced into the reaction flow tube is electronically excited.
The chemical compounds used in this study were obtained from commercial vendors: helium, Airgas,
99.999%; CF4, Airgas, 99.99%; SF6, Airgas, 99.99%; SF5CF3, SynQuest, 99%; CH3F, Matheson, 99%;
CH3Cl, Matheson, 99.9%; CH3Br, Matheson, 99%; CH3I, Sigma, 99%; acetaldehyde, Sigma, ≥99.5%;
acetic acid, Mallinckrodt, 99.9%; acetone, Fischer, 99.5%; dimethyl ether, Sigma, ≥99%; methanol,
Sigma, ≥99.9%; methyl formate, Sigma, 99%. Several compounds employed as neutral reactants exist as
liquids. For these compounds, the vapor pressures of the liquids are high enough to establish a stable

45

gaseous flow rate; each liquid sample is submerged in a room-temperature water bath to ensure flow
stability.
3.3

Computational Methods
Energy diagrams for the reaction of O•+ with methyl chloride and acetaldehyde were calculated at the

CCSD(T)/aug-cc-pVDZ//B3LYP/aug-cc-pVDZ level of theory using Gaussian 09.11 The productfractions for the unimolecular dissociation of acetaldehyde ionized by charge transfer with O •+ were
determined using RRKM theory implemented by the CRUNCH program.12-14
3.4

Results and Discussion
The gas-phase ion-neutral reactions of O•+ (4S) with CF4, SF6, SF5CF3, CH3F, CH3Cl, CH3Br, CH3I,

CH3COH (acetaldehyde), CH3COOH (acetic acid), CH3COCH3 (acetone), CH3OH (methanol), andCH3OCOH (methyl formate) have been studied; rate constants and product fractions are reported. The
results are discussed in three sections: perfluorinated compounds, methyl halides, and organic molecules.
Using this diverse set of neutral reagents, some interesting trends in the reactivity of O•+ are reported.
3.4.1

O•+ + Perfluorinated Compounds

The perfluorinated compounds are characterized by large ionization energies (IE) comparable to therecombination energy of O•+. Rate constants and product fractions for the reaction of O•+ with CF4, SF6,
and SF5CF3, which have been determined in this study, are reported in Table 3.1.

46

Table 3.1 Reaction rate constants and product fractions for O•+ with perfluorinated compounds
Reactant[IE (eV)]aCF4[≤14.7]SF6 [15.3]

Ionic
Product
Fractionb

1.00

Rxn. #

ΔHrxnd(kcal mol-1)

CF3+[S] + OF[Q]

3.1a

9.4

CF3+[T] + OF[D]

3.1b

132.3

CF3 [S] + O[T] + F[D]

3.1c

123.4

CF3+ [S] + OF[D]+1.00
0.07(±0.05)

SF5CF3[12.9]

Productsc

0.88(±0.05)

3.1d

–10.6

+

3.2

–30.6

+

3.3

–30.1

+

3.4a

–30.0

+

3.4b

–91.1

SF5 [S] + OF[D]CF3 [S] + SF5[D]+O[T]
SF3 [S] + CF4[S] + F[D] + O[T]SF3 [S] + CF4[S] + FO[D]kexp.e(×10-9 cm3 s-1)

Efficiencyf(kexp./kcol.)1.0(±0.3)g0.78

1.3(±0.4)h0.82

1.4(±0.4)i

0.60

0.05(±0.01) SF5+[S] + CF3[D] + O[T]
3.5
–16.5
Ionization energies from Lias et al. (Ref. 22). bProduct Fraction error reported as 1σ of the data sets. cS,D,T, and Q
denote the multiplicity. dThe enthalpies of reaction from Lias et al. (Ref. 22) and Afeefy et al. (Ref. 23). eRate
constants determined in this study; Error bars reported as total error (±30%). fSu and Chesnavich (Ref. 24).
g
Previous values (×10-9 cm3 s-1): k=1, Richter and Lindinger (Ref. 6); k=1.3(±0.4), Fisher and Armentrout (Ref. 5);
k=1.4(±0.4), Morris et al. (Ref. 27). hPrevious values (×10-9 cm3 s-1): k=1.8(±0.5), Morris et al. (Ref. 27);k=1.5(±0.5), Fehsenfeld (Ref. 25). iPrevious value k=1.9(±0.4) ×10-9 cm3 s-1, Atterbury et al. (Ref. 26).
a

15-19

The reactions of O•+ with the three fully fluorinated compounds, CF4, SF6, and SF5CF3 have been studied
previously.5, 6, 18-20 Our rate constants agree with previous studies within reported error.
It is well documented that reactions involving atomic species with high spin states often undergospin-conversion processes,21, 22 i.e., they break the spin selection rule, ΔS=0. Previous studies of the O•+ +CF4 reaction report CF3+ and OF as the products at thermal energies.5, 6, 20 The ground state multiplicities
of OF and CF3+ are a doublet and a singlet, respectively. To gain insight into the electron-spin dynamics,
this reaction is examined computationally. In Table 3.1, the multiplicity of each product is indicated inbrackets (neutral reagents are singlets). The total spin of the reactants, O•+ and CF4, is S=3/2, andreactions 3.1a-3.1c in Table 3.1 present the three spin conserved reaction channels for generation of CF3+.
Interestingly, these channels are all endothermic by at least 9 kcal/mol. The only exothermic productchannel, reaction 3.1d, follows a spin-forbidden process. Therefore, the total spin of the system haschanged from 3/2 to 1/2. Despite the necessity for spin-conversion, the reaction of O•+ and CF4 occurs
with an efficiency that approaches the collision limit.

47

Reaction 3.2 between O•+ (4S) and SF6 likely occurs by a similar mechanism to the CF4 reaction, i.e.,
it likely undergoes spin-conversion. The ground state products of SF5+ and OF are a singlet and doublet,
respectively; their enthalpies of formation were taken from the literature. The energies of the excitedstates of SF5+ and OF were not reported, nor were they computed due to limited computational resources.
However, noting that the computational difference in energy between CF 3+[S] and CF3+[T] is ~140kcal/mol, the spin-conserved products for the reaction between O•+ (4S) and SF6 would likely be highly
endothermic. Therefore, it is expected that the reaction with SF6 occurs through a spin-conversion
process.
In addition, the reaction of O•+ with SF5CF3 has been investigated. This process differs from thereactions with CF4 and SF6 because the ionization energy of SF5CF3 (12.9 eV) is low enough to permit
charge transfer with O•+. Because the ground state of SF5CF3+ is repulsive in the Franck-Condon region,
the radical cation rapidly dissociates.23 Three major ionic products, CF3+, SF3+, and SF5+, were observed,
and trace amounts of the ion SF4+ were also observed. Previous results suggest that this reaction occurs
only by charge transfer.19 While the reaction of O•+ with SF5CF3 is rapid, it does not occur at the collisionrate, suggesting that some of the ion-neutral complexes dissociate before a reaction takes place. Sincecharge transfer does not occur upon every collision, fluoride abstraction may occur in addition to chargetransfer as it yields more exothermic products. If the product SF3+ is formed by fluoride abstraction asshown in reaction (3.4b), the generation of this ion is three times more exothermic. Reactions (3.3) and(3.5) are likely the result of charge transfer processes. Because the HOMO of SF 5CF3 most closely
resembles the single bond between the carbon and sulfur atom,24 removing an electron from this orbital
will cause the molecule to fragment. Products of reaction 3.3 and 3.5 differ only in the location of thecharge.
Our product fraction for the reaction with SF5CF3 disagrees sharply with the previous study.19 Earlier
experiments report that the product fraction is dominated by the formation of CF3+ at 0.83, while SF3+
constituted only 0.16. However, the dominant product in our studies is SF3+ with a product fraction of48

0.88; both spin-allowed and spin-forbidden channels are likely in competition to generate SF3+. The originof the differences in product branching fractions is unknown; our experiments carefully excluded excitedstates of O•+, determined initial product fractions, and minimized mass discrimination. All of these factors
can affect product ratios. There is no reason to doubt the present result; this is the recommended productfraction.
3.4.2

O•+ + Methyl Halides

Rate constants and product fractions for the reaction of O•+ with CH3F, CH3Cl, CH3Br, and CH3I arereported in Table 3.2.
Table 3.2 Reaction rate constants and product fractions for O•+ with methyl halides
Reactant[IE (eV)]a

Product
Fractionb

CH3F[12.50]0.57(±0.06)
0.43(±0.06)

Products c

Rxn. #

ΔHrxn(kcal mol-1)dCH2F+[S] + H[D] + O[T]

3.6a

–4.7

CH2F+[S] + HO[D]3.6b–111.6

3.7

–29.6

3.8a

–3.8

3.8b

–68.1

+

CH3F [D] + O[T]+0.88(±0.03)
CH3Cl[11.26]

CH3Br[10.54]
CH3I[9.61]

CH3 [S] + Cl[D] + O[T]+

CH3 [S] + ClO[D]+0.08(±0.02)

HCO [T] + H[D] + HCl[S]3.9–126.30.04(±0.02)

H2CO+[D] + HCl[S]1.00
0.94(±0.04)

3.10–149.9

+

3.11a

–17.8

+

CH3 [S] + BrO[D]

3.11b

–74.1

CH3+[S] + I[D] + O[T]

3.12a

–30.7

3.12b

–170.1CH3 [S] + Br[D] + O[T]+

CH3 [S] + IO[D]kexp.
(×10-9 cm3 s-1)e

Efficiency(kexp./kcol.)f

2.7(±0.8)0.922.2(±0.7)

0.72

2.2(±0.7)

0.77

1.4(±0.3)

0.51+

0.06(±0.04)
I [T] + CH3[D] + O[T]
3.13
–16.0
Ionization energies from Lias et al. (Ref. 22). bProduct Fraction error reported as 1σ of the data sets. cS,D,T,
and Q denote the multiplicity. dThe enthalpies of reaction from Lias et al. (Ref. 22) and Afeefy et al. (Ref.
23). eError bars reported as total error (±30%). fSu and Chesnavich (Ref. 24).
aThis study reports for the first time the reactivity of O•+ with CH3F, CH3Cl, CH3Br, and CH3I. The
adiabatic ionization energies (AIEs) of the methyl halides are all below the RE of O•+; the exothermicity
for charge transfer increases as the halides become larger. Each reaction proceeds at 50% or more of thecollision rate with the CH3F reaction being the most efficient.

49

For CH3Cl, CH3Br, and CH3I, the dominant product ion is CH3+. In contrast, the dominant product of
the CH3F reaction is CH2F+. The relevant bond energies for the methyl halides are provided in Table
3.3.25
Table 3.3 Bond dissociation energies of methyl halidesa-1
BDE298 (kcal mol )
Bond
CH3–F
CH3–Cl
CH3–Br

110.0
83.7
70.3

CH3–I
H–CH2F
H–CH2Cl
H–CH2Br

54.0
101.3
100.1
102.1

H–CH2I

103.2

aRef. 32

Since the carbon-fluorine bond is the only carbon-halogen bond that is stronger than the carbon-hydrogenbond, cleavage of the carbon-hydrogen bond is more favorable for reaction (3.6). The major product ionsof these reactions can be formed by the spin-allowed process (charge transfer followed by bond cleavage[(3.6a), (3.8a), (3.11a), (3.12a)]) or the spin-forbidden process (atom abstraction [(3.6b), (3.8b), (3.11b),(3.12b)].) It is possible that these mechanisms are competing, but the spin-forbidden pathway may
dominate since it is more exothermic in every case.
Although charge transfer is energetically feasible for all methyl halide reactions, the charge transfer
product is only observed for CH3F (reaction 3.7). Charge transfer is facilitated by good energy resonancebetween the RE of the cation and the AIE of the neutral molecule, and the existence of nonzero FranckCondon factors.26 For reactions of O•+ with the methyl halides, the best energy resonance occurs for
CH3F; therefore charge transfer deposits the smallest amount of energy into the internal modes of CH3F+.
The lack of energy resonance during charge transfer of O•+ with the heavier methyl halides, and theweaker carbon-halogen bonds leads to their fragmentation upon production. The richness of product50

channels in the reaction of O•+ with CH3Cl warrants a deeper investigation of the energy landscape in thefollowing section.
3.4.2.1

Stationary Points for the O•+ + CH3Cl Reaction

CCSD(T)/aug-cc-pVDZ level calculations were used to develop two reaction coordinate plots for thereaction of O•+ with CH3Cl. For Figure 3.1 and Figure 3.2, the total spin for each point on the potential
energy surface is 3/2 and 1/2, respectively.

Figure 3.1 Stationary points for the reaction of O•+ + CH3Cl; S=3/2
Values on the potential energy surface are in kcal mol-1. The energies were determined at the
CCSD(T)/aug-cc-pVDZ//B3LYP/aug-cc-pVDZ level of theory, they are corrected for zero point energy
using the harmonic-oscillator approximation, and reported at 0 K. At each stationary point, S=3/2. The
symbols S, D, T, and Q denote the multiplicity. The caption on the left side of the plot represents thecurve crossing geometry.
51

Figure 3.2 Stationary points for the reaction of O•+ + CH3Cl; S=1/2
Values on the potential energy surface are in kcal mol-1. The energies were determined at the
CCSD(T)/aug-cc-pVDZ//B3LYP/aug-cc-pVDZ level of theory, they are corrected for zero point energy
using the harmonic-oscillator approximation, and reported at 0 K. At each stationary point, S=1/2. The
symbols S and D denote the multiplicity.

In each case, all possible spin allowed product combinations, even those which are endothermic and thus
not energetically feasible, are included. The major product of the CH3Cl reaction is the ion CH3+.
Formation of this ion occurs by either dissociative charge transfer, or chloride abstraction as illustrated on
the two reaction coordinate plots. It is the only possible experimentally observed product from the chargetransfer channel of reactivity. However, we have shown with the O•+ + CF4 reaction that spin-forbidden
52

atom abstraction processes are feasible. It is possible that this product is generated through the spinforbidden chloride abstraction process, as this process is more exothermic by an order of magnitude. This
would proceed via curve crossing between the two spin states; this crossing would likely occur during theformation of a primary complex as the energies overlap for the reactants and primary complexes in bothplots. An ion complex geometry for spin conversion, or curve crossing, was calculated and is included on
the left side of Figure 3.1. This geometry embodies a minor perturbation between the primary complexesof each reaction coordinate plot. It represents a point in which the two spin states are energeticallyequivalent.
Formation of the CH3+ ion likely dominates the product distribution of the CH3Cl reaction because theminor product ions, H2CO+ and HCO+, are formed through complex rearrangement processes. If thecomplex is short-lived, these products are unlikely to be generated. The product ions H2CO+ and HCO+
were not observed in reactions involving CH3Br and CH3I because more energy is transferred during thereactions, and the complexes dissociate even more rapidly.
3.4.3

O•+ + Small Organic Molecules

Reaction rate constants and product fractions were determined for each reaction. The results of theseexperiments as well as pertinent thermochemical data (ionization energies, appearance energies, reaction
enthalpies, etc.) are presented in Table 3.4.9, 17, 27-38

53

Table 3.4 Reaction rate constants and product fractions for O•+ with small organic molecules
Reactant[I.E. (eV)]a
Acetaldehyde
CH3COH[10.23]

Acetic Acid
CH3COOH[10.65]Acetone
CH3COCH3[9.70]

Dimethyl
Ether
CH3OCH3[10.02]

Methanol
CH3OH[10.84]

MethylFormateCH3OCOH[10.84]A.E.
(eV)

Products

11.78

ΔHrxn
(kcal mol-1)

Rxn.
#

Product
Fraction

CHO+[S] + CH3[D] + O[T]3.14

0.99(±0.01)

–40.0

–41.1

10.67

CH3CO+[S] + H[D] + O[T]3.15

0.01(±0.01)

–60.4

–68.1

12.05

CH2OH+[S] + CHO[D] + O[T]3.16

0.03(±0.01)

–28.8

–32.2

11.54

CH3CO+[S] + OH[D] + O[T]3.17

0.06(±0.02)

–43.4

–48.2

3.18

0.05(±0.02)

–28.6

–34.7

3.19

0.24(±0.09)

–63.1

–68.5

12.3
10.65

+

COOH [S] + CH3[D] + O[T]+

CH3COOH [D] + O[T]

Calc.

b

Lit.

a

–

CO3H [S] + CH3[D]3.20

0.62(±0.10)

–34.0

–

–

CHO+[S] + 2H[D] + CH3CO[D]3.21

0.06(±0.02)

–17.3

–22.8

10.38

CH3CO+[S] + CH3[D] + O[T]3.22

0.18(±0.14)

–65.1

–74.0

9.70

CH3COCH3+[D] + O[T]3.23

0.09(±0.05)

–89.4

–84.4

–

CH3CO2+[S] + CH3[D]3.24

0.67(±0.20)

–55.6

–

≤12.85

CHO+[S] + H2[S] + CH3[D] + O[T]3.25

0.31(±0.12)

–37.4

–37.9

≤11.85

CH2OH+[S] + CH3[D] + O[T]3.26

0.34(±0.04)

–61.8

–67.2

+

11.12

CH3OCH2 [S] + H[D] + O[T]3.27

0.35(±0.11)

–53.6

–61.0

13.06

CHO+[S] + H2[S] + H[D] + O[T]3.28

0.06(±0.02)

–16.0

–15.6

3.29

0.13(±0.04)

–37.6

–40.2

12.05

+

+

CH2O [D] + H2[S] + O[T]11.55

CH2OH [S] + H[D] +O[T]3.30

0.73(±0.02)

–40.4

–44.9

10.84

CH3OH+[D] + O[T]3.31

0.08(±0.01)

–57.1

–64.0

12.91

CHO+[S] + CH3O[D] +O[T]3.32

0.04(±0.01)

–25.9

–32.1

3.33

0.02(±0.01)

–31.3

–35.1

3.34

0.94(±0.01)

–34.0

–39.8

13.6
12.4

+

+

CH2O [D] + H2[S] + CO[S] + O[T]+CH2OH [S] + H[D] + CO[S] + O[T]kexp. c(×10-9 cm3 s-1)

Efficiency(kexp./kcol.)2.9(±0.9)

0.82

3.0(±0.9)

1.09

2.8(±0.8) d0.67

1.7(±0.5)0.69

2.2(±0.6) e

0.76

2.3(±0.7)

0.80

aRef. 22. bReaction enthalpies are calculated using CCSD(T)/aug-cc-pVDZ//B3LYP/aug-cc-pVDZ level of theory and reported at 0 K. cError
reported as total error, ±30%. dPreviously measured rate constant is 7.1 × 10-9 cm3 s-1 (Ref. 44). ePreviously measured rate constant is 1.9 × 10-9cm3 s-1 (Ref. 45) and 2.7 × 10-9 cm3 s-1 (Ref. 20).

The RE of O•+ is well above the ionization energy of each of the organic neutrals. As a result, eachreaction reported proceeds rapidly, and charge transfer is a common process within these reactions. The
efficiency of each reaction is over 50% of the collision rate. The most efficient reaction is that of O•+ withacetic acid, which occurs at unit efficiency.
When dealing with ion-molecule reactions involving high energy ions, such as O•+, it can be useful tocompare the experimental results to experimental appearance energy data. The high RE of O•+ allows it toremove an electron from the neutral molecules, and therefore its reactivity can be compared to different
ionization experiments. Although investigating the difference in appearance energy of the product ions54

and the recombination energy of O•+ provides similar information as investigating Hrxn values, theresults of ionization experiments can be used to guide the analysis of exothermic ion-neutral reactions.
When certain products are not observed in ionization experiments, it suggests that these products are theresult of some ion-molecule interaction and rearrangement. Therefore, we have included appearance
energy data in Table 3.4.
3.4.3.1O•+ + Acetaldehyde

The reaction of O•+ with acetaldehyde generates two products ions of m/z 29 and 43, the formyl andacetyl cations, respectively. The appearance energies of both ions were determined previously byphotoionization experiments, and are 11.78 and 10.67 eV, respectively.28, 29 Therefore both reactions 3.14
and 3.15 likely occur through dissociative charge transfer processes. Reaction 3.14 and 3.15 show
cleavage of the C-C and C-H bonds, respectively. The distribution favors reaction pathway 3.14 (99%)
although it is less exothermic than reaction pathway 3.15 (1%) by 27 kcal mol-1. For this reaction, RRKM
theory provides some insight into this preference in reactivity.
Calculated stationary points along a multi-dimensional potential energy surface for the reaction of O•+
with acetaldehyde are shown in Figure 3.3. The relaxed potential energy scan along the C-C bond in thereaction intermediate (CH3COH+) indicates that reaction 3.14 occurs through a loose transition state. If
these reactions proceed by charge transfer, fragmentation of the acetaldehyde cation can be treated by
RRKM unimolecular dissociation theory where the energy available to the acetaldehyde ion is thedifference between the ionization energy of acetaldehyde and atomic oxygen (72.3 kcal mol-1, Figure 3.3).
The energy then available to each fragmentation channel, Eaval, is the difference between the chargetransfer energy and the zero point energy of the products, or transition state for reaction 3.14 and 3.15,
respectively; these values are 40.0 and 55.0 kcal mol-1 as shown in Figure 3.3. Based on these energies,
RRKM predicts a branching ratio of 99.6% and 0.4% for the two reactions, respectively.12-14 These values
are in excellent agreement with the experimental results.

55

Figure 3.3 Stationary points for the reaction of O•+ with acetaldehyde
Units are in kcal/mol. The unimolecular product distribution (RRKM theory) for the acetaldehyde
cation is determined with an excitation energy of 72.3 kcal/mol.

3.4.3.2O•+ + Acetic Acid

The reaction of O•+ with acetic acid generates five ionic products with m/z 31 (CH2OH+), 43(CH3CO+), 45 (COOH+), 60 (CH3COOH+), and 61 (CO3H+). All but one of these products, m/z 61, are theresult of charge transfer reactions. The appearance energies of the first four molecular ions are below therecombination energy of O•+; they are 12.05, 11.54, 12.3, and 10.65 eV, respectively.27, 29-31 Selim andHelal have determined that the m/z 31 ion forms as the hydroxylmethyl cation, CH2OH+, and not themethoxy cation, CH3O+, based on heats of formation.30 The heats of formation of the two ions are 168 and
201 kcal mol-1, respectively.36 The hydroxylmethyl cation is therefore formed by fragmentation followedby hydrogen atom rearrangement, a common process in ion fragmentation.34, 39 Reaction 3.17 produces
m/z 43, the acetyl cation. This ion is formed by fragmentation of the C-OH bond upon ionization.
56

Reaction 3.18 forms the m/z 45, protonated carbon dioxide. This occurs through dissociative chargetransfer where the C-C bond is broken upon ionization. Shuman et al. observed m/z 43 and 45 ions intheir threshold photoelectron-photoion coincidence (TPEPICO) spectroscopy experiment,40 which
examined the unimolecular dissociation of energy selected acetic acid cations. Their results confirm thatthese products are the result of charge transfer, and that the neutral fragments for reaction 3.17 and 3.18
are hydroxyl and methyl radical, respectively. Reaction 3.19 shows charge transfer where there is nofragmentation after ionization. These four charge transfer reactions together account for about one third of
the product fraction, where reaction 3.19 alone represents 24%.
The final product channel, reaction 3.20, produces the ion m/z 61, and represents 62% of the productfraction. The m/z 61 ion is especially intriguing, and no information appears to be available in theliterature. Our product fraction studies indicate that m/z 61 is not a secondary product. This productinstead is formed from an interesting reaction in which a methyl radical is replaced by O•+. There are noliterature values for the enthalpy of formation of the CO3H+ ion, but we have determined this reaction
channel to be exothermic based on theoretical calculations. The lowest energy geometry of this ion isshown in Figure 3.4; it is a singlet with the three oxygen atoms bound directly to the carbon. The
molecule is completely planar. A similar mechanism and product is observed in the O•+ + acetone
reaction.

57

Figure 3.4 Lowest energy ionic products for reaction 3.20 and 3.24
Optimized at B3LYP/aug-cc-pVDZ level. For each reaction, O•+ has replaced •CH3.

3.4.3.3

O•+ + Acetone

The reaction of O•+ with acetone generates four products of m/z 29, 43, 58, and 59. The formyl cation(m/z 29, reaction 3.21) is not observed in any appearance energy experiments for acetone. We suggest an
addition elimination reaction in which the oxygen cation attacks the methyl group, causing fragmentationof the C-C bond, and a release of two hydrogen atoms. Because the m/z 29 product is the result of achemical reaction, this cation could be generated as COH+, protonated carbon monoxide. However,literature values for the enthalpy of formation of COH+36 indicate that this reaction would be endothermic
by about 10 kcal mol-1.
Reaction 3.22 generates the acetyl cation, CH3CO+, through a dissociative charge transfer reaction inwhich one of the two equivalent C-C bonds has fragmented. The appearance energy of the acetyl cation is10.38 eV.29 Reaction 3.23 is a charge transfer reaction in which the acetone molecule remains intact after
ionization.

58

Reaction 3.24 generates m/z 59, one amu heavier than the acetone molecule. The structure of thisproduct is similar to that observed with the O•+ + acetic acid reaction, where O•+ replaces a methyl radical.
There is no information in the literature about the CH3CO2+ ion, but we have provided an optimized
geometry for this ion in Figure 3.4. It is a singlet, similar in geometry to the product of reaction 3.20. Theion is nearly planar with respect to the two carbon and two oxygen atoms.
The rate constant for the reaction of O•+ with acetone is much lower than the previously measured rate
constant which was obtained using a flowing afterglow.37 The SIFT often provides more reliable rate
constants than the flowing afterglow due to the separation of ionization and reaction regions. Moreover,
the rate constant previously obtained for the O•+ + acetone reaction (k = 7.1×10-9 cm3 s-1) is almost double
the collision rate constant, which is not physically reasonable.26
3.4.3.4

O•+ + Dimethyl Ether

The reaction of O•+ with dimethyl ether generates three products with m/z 29 (CHO+), 31 (CH2OH+),and 45 (CH3OCH2+). The appearance energies of these ions were determined by threshold photoelectronphotoion coincidence (TPEPICO) spectroscopy to be ≤12.85, ≤11.85, and 11.12 eV respectively.32
Therefore, the observed products are likely the results of dissociative charge transfer reactions.
Many studies have focused on the unimolecular dissociation of ionized dimethyl ether.32, 41, 42 The m/z
31 ion observed during these studies is often the focal point of discussion, for this ion may form as either
the methoxy, CH3O+, or hydroxyl methyl, CH2OH+, cation. Nishimura et al. investigated the dissociation
dynamics of energy selected dimethyl ether ions.41 The identities of the products observed in their
TPEPICO breakdown curves at 13.6 eV (the recombination energy of O•+) are in agreement with ourobservations. However, at 13.6 eV they observe a greater fraction of the m/z 45 ion than m/z 29 or 31,
while our results generate all products at roughly equal distributions. The authors41 utilized the nonstatistical spectator mechanism of direct dissociation theory (SSD) to determine that, at high energies, theformation of the m/z 31 ion is accompanied by a large kinetic energy release, indicating that at thesehigher energies, the m/z 31 product is the methoxy cation. However, based on their results, the kinetic
59

energy release is negligible around the recombination energy of O•+, and the m/z 31 ion formed in ourstudy is likely the hydroxyl methyl cation.
3.4.3.5
TheO•+ + Methanol
reaction

of

O•+

with

methanol

forms

four

products

of

m/z

29

(CHO+), 30 (CH2O+), 31 (CH2OH+), and 32 (CH3OH+). We propose these ions are formed through chargetransfer reactions. The appearance energies of these ions are 13.06, 12.05, 11.55 and 10.84 eV,
respectively.27, 33 The O•+ + methanol reaction has been studied previously.9, 38 The most recently
measured reaction rate constant (k = 2.7×10-9 cm3 s-1) and product branching ratio (m/z 29, 5%; m/z 30,
10%; m/z 31, 80%; m/z 32, 5%)9 are in good agreement with our rate constant and product fraction
results.
3.4.3.6

O•+ + Methyl Formate

The reaction of O•+ with methyl formate generates three product ions with m/z 29, 30, and 31. We
propose these products are formed by dissociative charge transfer reactions; the appearance energy ofeach product ion is less than the recombination energy of O•+. Fragmentation of the methyl formate cation
is well studied.31, 34, 35, 43, 44 Reaction 3.32 produces m/z 29, the formyl cation, CHO+, by fragmentation of
the C-O bond after ionization. Reaction 3.33 forms the formaldehyde cation of m/z 30, likely through asequential rearrangement/dissociation mechanism.
O+· + CH3OCHO → [CH3OH+]* + CO + O → CH2O+ + H2 + CO + O

3.33

Experiments show that at the onset of the unimolecular dissociation of the methyl formate cation at about
11.5 eV, the formyl hydrogen will migrate forming a methanol cation, CH 3OH+.43, 44 The first step ofreaction 3.33 involves ionization and re-arrangement of the methyl formate molecule, forming thismethanol cation. The methanol cation will then fragment to form the hydrogen molecule and theformaldehyde cation.
The dominant product ion is m/z 31. This product is formed by the sequential reaction:60

O+· + CH3OCHO → [CH3OH+]* + CO + O → CH2OH+ + H + CO + O

3.34

Again, the first step of this process involves ionization and formation of the methanol cation throughhydrogen migration. The methanol cation then fragments to lose a hydrogen atom, forming the hydroxyl
methyl cation, CH2OH+.
Because the initial step of the O•+ + methyl formate reaction involves generation of the methanol
cation, this reaction can be compared to O•+ + methanol. The results agree nicely considering that thedominant product of both of these reactions is the hydroxyl methyl cation, with minor contributions of theformyl and formaldehyde cations.
3.5

Conclusion
This study explored reactions of O•+ with a wide range of neutral molecules. The perfluorinatedcompounds CF4 and SF6 have large ionization energies of ~14.7 and 15.3 eV, much larger than the RE ofO•+. Both of these reactions occur through fluoride abstraction, a process that is spin-forbidden. Despite
the requirement of spin-conversion, these reactions occur rapidly, near the collision limit.
Each of the subsequent reactions reported in this chapter involve neutral molecules with ionizationenergies below the recombination energy of O•+. Most of these reactions are assumed to occur throughdissociative charge transfer product generation channels. In fact, the reaction of O •+ with acetaldehyde
follows this assumption so well that the fragmentation pattern for the acetaldehyde cation is successfullymodeled using RRKM theory where the internal excitation energy is supplied by the excess chargetransfer energy (~3.4 eV).
In many cases, the product channels are complex, and dissociative charge transfer appears to competewith atom-abstraction reactions. Although spin-conversion is often required for the atom-abstraction
process, this does not appear to limit the reaction efficiency. As a general comment, products accessed61through spin-forbidden reaction channels are typically more exothermic. Whether the reactions undergospin-conversion or not, the reactions of O•+ occur rapidly.
The high reactivity of O•+ indicates that ion-molecule chemistry is an important potential loss
mechanism for this radical cation in many environments, including the Earth’s atmosphere, the ISM, andcombustion processes. Inclusion of the rate constants and products observed in this study into models ofthese environments will provide more accurate characterization of the relevant chemistry.
3.6

References[1] J.H. Hoffman, A mass spectrometric determination of the composition of the nighttime topside
ionosphere, J. Geophys. Res., 72 (1967) 1883-1888.
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[19] C. Atterbury, R.A. Kennedy, C.A. Mayhew, R.P. Tuckett, A study of the reactions of trifluoromethyl
sulfur pentafluoride, SF5CF3, with several positive ions of atmospheric interest, Phys. Chem. Chem.
Phys., 3 (2001) 1949-1953.
[20] R.A. Morris, A.A. Viggiano, S.T. Arnold, J.F. Paulson, Chemistry of atmospheric ions reacting withfully fluorinated compounds, Int. J. Mass Spectrom. Ion Process., 149 (1995) 287-298.
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Fragmentation of energy-selected SF5CF3+ probed by threshold photoelectron photoion coincidence
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Elsevier, 2006, pp. 89-129.
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65

4 Reactivity and Thermochemistry of Dicyanamide: N(CN)2─
4.1

Introduction
Although they have been known for more than 100 years,1 ionic liquids, or salts with a melting point≤ 100 °C, have recently received renewed interest in the chemical literature. Ionic liquids exist with awide variety of cation and anion components; in fact, their properties can be customized by switching or
modifying the cations or anions. As a liquid salt, they have many applications. Some synthetic chemists
employ ionic liquids as reusable solvents because of their low vapor pressure,2 and ionic liquids haveeven been shown to participate in synthesis reactions.3 Because of their inherent stability, they are being
used as the electrolytic medium in batteries, capacitors, and solar cells.4, 5 In addition, a number ofmaterials chemists are investigating ionic liquids as next-generation bipropellant hypergolic fuels inefforts to replace hydrazine based fuels; storage of hydrazine based fuels is challenging because of their
toxic, corrosive, and volatile nature. Several ionic liquids exhibit hypergolic behavior, 6-12 providing a
possible replacement.
Dicyanamide, N(CN)2─, is a common anion of ionic liquids; as an anion, it exhibits hypergolicbehavior when paired with certain cations. Furthermore, dicyanamide generates ionic liquids withrelatively low viscosities, allowing for more efficient delivery in propulsion applications. When pairedwith 1-propargyl-3-methylimidazolium, the resulting ionic liquid melts at 17 °C and is stable up to 144°C. Upon mixing with white-fuming nitric acid, 1-propargyl-2-methylimidazolium dicyanamide will
undergo hypergolic ignition with a 15 ms delay time.6 An ignition delay time of ≤ 5 ms is preferred forreal-world applications; this can be achieved by altering the cation or creating ionic liquid mixtures.7
The hypergolic interaction between dicyanamide and nitric acid has been a hot topic of study in thelast seven years.6, 9, 13, 14 In 2008, Chambreau and coworkers proposed that 1,5-dinitrobiuret is formed
during the hypergolic reaction between dicyanamide containing ionic liquids and nitric acid.9 Since this66

report, several studies have theoretically investigated the decomposition dynamics of 1,5-dinitrobiuret.1517

While dissociation dynamics is a popular topic, fewer studies have focused on the formation pathway of

1,5-dinitrobiuret. This study will computationally investigate energetics of the previously proposed
formation pathway9 of 1,5-dinitrobiuret through stationary points, including transitions states, along a
multidimensional reaction coordinate.
In addition to its unique properties in the liquid phase, dicyanamide is one of the very few anions thathas been found not to exhibit reactivity with H atom,18 the most abundant element in the universe.
Dicyanamide is believed to be present in the atmosphere of Saturn’s moon Titan. 19 It is also likely acomponent of the dense interstellar clouds where CN−, C3N−, and C5N− have been observed.20-22 Becauseof its large electron binding energy, 4.135 eV,23 dicyanamide may survive in regions of the interstellar
medium where most anions would undergo photodetachment. This study investigates the reactivity of
dicyanamide with many astronomically abundant species, including N and O atoms.

4.2

Experimental
The experiments were carried out at 298 K with a Selected Ion Flow Tube (SIFT) using the home-built electrospray-ionization (ESI) source. To generate the dicyanamide signal, a 6 × 10-5 M solution of
sodium dicyanamide (≥96%, Fluka) in acetonitrile (99.8%, Sigma-Aldrich) is used with the ESI source.
Typical syringe pump flow rates are 0.2-50 μL/min.
The reactivity of dicyanamide with H2, O2, H2O2, DBr, HCl, NH3, N2O, SO2, COS, CO2, CH3OH,
H2O, CH4, N2, CF4, and SF6 was also investigated, but the reaction rate constants are below the detection
limit. The lack of reactivity for these neutral gases was confirmed by introducing a high flow of theneutral gas into the flow tube at the longest reaction distance. Neither dicyanamide depletion nor new
ionic products were observed.

67

We have also investigated the reactivity of dicyanamide with ground state N and O atoms, whichwere introduced into the reaction flow tube by use of a microwave discharge. Microwave generation of N
and O atoms is a well-established technique and has been described in Chapter 2. The N atoms are
generated by flowing (~0.6 STP cm3/s) N2 gas (99.999%, Airgas) through an Evenson cavity suppliedwith 50 W of microwave radiation. The overall efficiency of generating N atoms from this technique isapproximately 1–3%. The O atoms are then formed by titrating the N atoms with NO (4% in He, Airgas),
which forms ground state O atom by the reaction N + NO → N2 + O. The titration endpoint for thisprocess is found by plotting the natural log of the ion signal as a function of NO flow. The titration
endpoint reveals the concentration of O atoms introduced in the flow tube and subsequently determines
the initial concentration of N atoms in the absence of NO. The microwave discharge is also used to obtain
excited state O2 molecules, O2 (1∆). This experiment is performed by flowing O2 through the microwave
discharge. Although no reactivity of dicyanamide anion with N(4S), O(3P), and O2(1∆) was observed, thepresence of these neutral reactants was confirmed by studies with other ions that are known to be reactive.
To investigate the reactivity of dicyanamide with sulfuric acid, we coupled a low-volatility neutral
reagent inlet to our reaction flow tube by adapting the methods used by Viggiano et al. 24 A Pyrex bulb
was filled with glass wool, and then fitted on both ends with 6 mm Pyrex tubing. Approximately 1 mL of
sulfuric acid was added to the glass wool. A flow of N2 (~0.2 STP cm3/s, UHP, Airgas) enters the glass
bulb through a metering valve and flowmeter, and carries sulfuric acid into the flow tube. The glass bulb
and reagent inlet were heated to approximately 100 °C. Although no reaction was observed between
sulfuric acid and dicyanamide, the presence of sulfuric acid was verified by investigating its protontransfer reaction with NO3‒, which occurs at the collision rate (k~2.6 × 10-9 cm3/s).25 Proton transfer was
confirmed by observing the loss of NO3‒ and generation of HSO4‒.

68

4.3

Calculations
Intermediates and transition states along the potential energy surface for the interaction of

dicyanamide with two nitric acid molecules were computed at the B3LYP/Aug-cc-pVDZ level of theory.
The acidities of dicyanamide, and the optimized geometries of dicyanamide clusters with nitric acid arereported using the MP2/6-311++G(d,p) level of theory; MP2/6-311++G(d,p) provides good agreementwith experimental results in our previous studies.26-28 All theoretical thermochemical values are reported
in terms of enthalpy. The values have been corrected for zero-point energy using the rigid-rotor‒
harmonic-oscillator approximation at 298 K. Transition states were verified by the existence of one
imaginary vibrational frequency and confirmed with IRC calculations. Optimized intermediates containno imaginary frequencies. Cartesian coordinates for the optimized geometries are reported in Appendix B.

4.4

Results and Discussion
The dicyanamide anion is a very stable species with an experimental electron-binding energy of 4.135± 0.010 eV.23 Our work explored its reactivity with several neutrals. In efforts to characterize thereactivity of dicyanamide in the gas-phase, we explored its reactivity with HNO3, H2SO4, H, N, O, O2,
H2O2, DBr, HCl, NH3, N2O, SO2, COS, CO2, CH3OH, H2O, CH4, N2, CF4, and SF6.
However, only one neutral reagent, nitric acid, reacted rapidly enough with dicyanamide for productdetection by the SIFT technique. The reaction of dicyanamide with nitric acid occurs with a pseudo-firstorder reaction rate constant of 2.7 × 10-10 cm3/s. Only one direct product, the cluster of nitric acid withdicyanamide, was observed. However, subsequent addition of another nitric acid was observed as asecondary product.
For the unreactive neutral reagents, we can place an upper limit on their reaction rate constants. These
values are estimated based on discernible ion-signal depletion and the number density of the neutral
reagent that can be introduced into the reaction flow tube. For the neutral reagents generated from69

microwave discharge plasmas, i.e., N, O, and O2(1Δ), as well as for sulfuric acid which has low volatility,
we designate an upper limit to the reaction rate constant of 5 × 10 -12 cm3/s. For the neutral reagents withhigh vapor pressures, i.e. O2, H2O2, DBr, HCl, NH3, N2O, SO2, COS, CO2, CH3OH, H2O, CH4, N2, CF4,and SF6, we assign an upper limit to their reaction rate constant of 1 × 10 -12 cm3/s. The lack of reactivity
with O(3P) and O2(1∆) is surprising, as these reagents are highly reactive with a wide variety of reagents.
The absence of reactivity with H2SO4 allows us to put an upper limit on the acidity of the parent
molecule, HNCNCN, of ∆Hacid < 309.6 ± 5.4 kcal/mol.
The lack of reactivity of dicyanamide has some important interstellar implications, i.e., dicyanamide
could survive the harsh chemical conditions of the interstellar medium as it will not react with theabundant atomic species, H, O(3P), and N(4S). Furthermore, the neutral molecules NH3, H2O, CH3OH, H2,
COS, SO2, CH4, N2O, N2, and O2 have been detected in the ISM,29-38 but our experimental results show no
detectable reactivity with dicyanamide. Additionally, a recent rendezvous of Saturn’s moon Titan by the
Cassini spacecraft provided a negative-ion mass spectrum of Titan’s nitrogen- and methane-richatmosphere. Dicyanamide is inferred to be present, as there is a strong presence of m/z 66 and several
other peaks corresponding to adducts of m/z 66 with neutral species, including HCN, HCCH, and H2.19
The presence of dicyanamide is not surprising, since the components of dicyanamide, i.e., N and C atoms,
are abundant in Titan’s atmosphere. The intense ionization environment intrinsic to planetaryatmospheres could provide a synthesis route for dicyanamide, and upon its formation dicyanamide will belong lived due to its chemical stability. A laboratory study of the microwave spectrum of dicyanamide—
which is a bent molecule with a calculated dipole of approximately 1 Debye—would allow observational
astronomers to search for its presence in the interstellar medium.
4.4.1

Formation of 1,5-dinitrobiuret

The mixing of nitric acid with ionic liquids containing the dicyanamide anion results in hypergolicbehavior.6 For this reason, the fundamental interaction between nitric acid and dicyanamide is of great70

interest. An insightful study by Chambreau and coworkers acquired infrared spectra of the evolved vaporof 1-butyl-3-methylimidizole dicyanamide with nitric acid.9 Their work suggests a mechanism where
dicyanamide interacts with two nitric acid molecules to form deprotonated 1,5-dinitrobiuret; their studies
computationally defined the thermochemical energetics of several reactions. However, they did notcharacterize the entire mechanism; therefore, we have calculated stationary points along a
multidimensional surface for the gas-phase interaction of dicyamide with one and two nitric acidmolecules to form dicyanamide (Figure 4.1).

Figure 4.1. Stationary points for the reaction of dicyanamide with two nitric acid molecules
The multi-dimensional potential energy surface provided above shows the interaction of dicyanamide
with two nitric acid molecules to form deprotonated 1,5-dinitrobiuret [7B]. The energies listed are
calculated with the B3LYP/Aug-cc-pVDZ method, and presented in units of kcal/mol. Geometries areposted in Appendix B.

71

The multidimensional potential energy diagram of dicyanamide with nitric acid to form 1,5dinitrobiuret illuminates several interesting points. First, the formation of dicyanamide-nitric acid clusters
are highly exothermic by 25 kcal/mol [2A] and 45 kcal/mol [2B] for the addition of one or two nitric acidmolecules, respectively. Second, for our experiment, the interaction of dicyanamide with one nitric acid islimited to forming the hydrogen-bonded cluster [2A] because transferring the nitrate to the nitrile carbon[3A] must overcome an endothermic barrier of 3 kcal/mol [2A/3A]. Conversely, the addition of a second
nitric acid lowers the transition states below the energy of the reactants, therefore supporting
Chambreau’s proposed formation of 1,5-dinitrobiuret.9 Although it is possible that we are generating 1,5dinitrobiuret as a secondary product in our flow tube, we cannot confirm the structure; in our experiment,
products such as [2B], [3B], [4B], etc. cannot be differentiated because they possess the same mass-tocharge ratio. Finally, the terminal product on the potential energy surface, deprotonated 1,5-dinitrobiuret[7B], is 70 kcal/mol lower in energy than dicyanamide with two nitric acid molecules, suggesting a strong
thermodynamic driving force for the formation of 1,5-dinitrobiuret.
4.4.2

Dicyanamide Proton Affinity and Solvent Molecule Affinity

An important mechanistic step in the formation of 1,5-dinitrobiuret is proton transfer between nitricacid and dicyanamide. Chambreau’s method of forming 1,5-dinitrobiuret requires proton transfer from
nitric acid to dicyanamide, followed by subsequent nucleophilic attack on the carbon atom of protonateddicyanamide by nitrate. While our experimental findings provide us with an upper limit for the proton
affinity of dicyanamide (< 310 ± 5 kcal/mol), we have further investigated the acidity of dicyanamide
computationally with a higher level of theory than was previously reported in the literature, and thelowest energy structures for the two protonated forms of dicyanamide are shown in Figure 4.2.

72

Figure 4.2 Lowest energy structures for dicyanamide and protonated dicyanamide(a) dicyanamide, (b) terminally protonated dicyanamide, and (c) centrally protonated dicyanamide,
optimized at the MP2/6-311++G(d,p) level of theory. Geometries posted in Appendix B.

As illustrated, dicyanamide has two protonation sites: the terminal nitrile and central nitrogen. Our
computational studies, as well as computational and experimental reports from the literature, conclude
that it is more favorable to protonate dicyanamide on the terminal nitrile nitrogen than on the central
nitrogen.9, 39 The B3LYP/6-31+G(d,p) calculations reported by Chambreau et al. and the MP2/6311++G(d,p) calculations from this work indicate that protonation on the terminal nitrile nitrogen is morefavorable than on the central nitrogen by 7.6 and 1.2 kcal/mol, respectively.9 This discrepancy might arise
from differences in the optimized geometry of the centrally protonated tautomer. The optimization of73

centrally protonated dicyanamide (Figure 4.2c) at the MP2/6-311++G(d,p) level of theory positions thehydrogen atom out of the molecular plane. This is distinctive between the two levels of theory used, forthe same tautomer optimized at the B3LYP level, positioned the hydrogen atom within the molecularplane, thereby preserving the intrinsic C2V symmetry of dicyanamide. In moving from B3LYP/631+G(d,p) to MP2/6-311++G(d,p), the acidity of the terminally and centrally protonated tautomers
changed by -1.9 and 4.5 kcal/mol, respectively, further suggesting that the shift in acidity value for thecentrally protonated tautomer occurred in large part due to the differences in the optimized geometry.
Nevertheless, even when structures are nearly identical, different theoretical methods can provide
different energetic values. Further experimental and computational investigation would be valuable toprovide a better understanding of the proton affinity of dicyanamide.
The question remains: is it possible for dicyanamide to accept a proton from nitric acid? Fromthermodynamic considerations, when dicyanamide encounters a single nitric acid molecule at room
temperature, proton transfer is not observed. Theoretical results support this observation; despite the levelof theory, proton transfer from nitric acid to dicyanamide is endothermic by at least 15 kcal/mol.
We therefore have conducted calculations to examine higher order clusters of dicyanamide with nitricacid to explore how proton transfer energetics change as a function of the number of nitric acid moleculesin the cluster. Figure 4.3 shows the optimized geometry of dicyanamide with one, two, and three nitricacid molecules.

74

Figure 4.3 Lowest energy structures of dicyanamide clustered with nitric acid
Dicyanamide clustered with one (a), two (b), and three (c) nitric acid molecules optimized at the MP2/g311++G(d,p) level of theory. Geometries posted in Appendix B.

When a bimolecular cluster between dicyanamide and one nitric acid forms, the hydrogen-bonded
proton resides predominantly on the oxygen atom of nitrate (Figure 4.3a). However, when an additional
nitric acid is added to the cluster, the lowest-energy structure locates the shared proton predominantly on
dicyanamide (Figure 4.3b). Furthermore, adding a third nitric acid molecule onto the final-unoccupied
nitrate oxygen (Figure 4.3c) results in a structure where the dicyanamide hydrogen bond length is ~0.5 Å
shorter than the oxygen-hydrogen bond length.
To further investigate the effects of clustering, the thermochemistry of proton transfer between
dicyanamide and nitric acid was investigated for the following reaction.
75

NCNCN─(HNO3)n + HNO3 → NCNCNH + NO3─(HNO3)nThe thermodynamic values below are reported for proton transfer to the terminal nitrile nitrogen. For
this reaction, when n = 0, the zero-point-thermally-corrected enthalpy for reaction at the MP2/6311++G(d,p) level of theory is 16.6 kcal/mol endothermic. Therefore, in the bimolecular interaction of
dicyanamide with nitric acid, proton transfer would not occur. Increasing the solvation of dicyanamide
with nitric acid such that n = 1 decreases the endothermicity from 16.6 to 9.8 kcal/mol. Subsequently
increasing the solvation to n = 2 and n = 3 further reduces the endothermicity of proton transfer to 4.7 and3.0 kcal/mol, respectively. This trend and the results from Figure 4.3 suggest that proton transfer canoccur when dicyanamide encounters a cluster of nitric acid molecules such as is possible on the surface ofa nitric acid droplet.
The interaction of dicyanamide-containing ionic liquids with nitric acid has been of increasing
interest because these liquids hold promise as a next-generation hypergolic fuel. The results of this study
suggest that interactions at the gas-liquid or liquid-liquid interface are very important for hypergolicignition. Indeed, hypergolic reactivity occurs when a macroscopic droplet of ionic liquid is added to a
cuvette containing a macroscopic quantity of nitric acid.6 For our gas-phase studies, we have determinedthat solvation of dicyanamide by nitric acid is highly exothermic, and this excess energy can be used topromote energetic chemistry such as, but not limited to, proton transfer; even though the self-sustaining
hypergolic reactions are only observed when dicyanamide is paired with certain cations such as 1propargyl-3-methylimidazolium, ‘reactivity’ is observed even when dicyanamide is paired with inert
cations such as Na+ or 1-butyl-3-methylimidazolium.9, 13

4.5

Conclusion
The dicyanamide anion is extremely stable, and exhibits no reactivity with a variety of neutral

reagents, including H2SO4, H, N, O, and O2(1Δ). However, dicyanamide readily clusters with nitric acid,and the ionic adducts are detected in our studies. Our computational results indicate that the association76

of dicyanamide and nitric acid is exothermic by about 25 kcal/mol. This energy may be responsible for
initiating the hypergolic phenomenon that is observed when the two liquids are mixed.
The intrinsic stability of dicyanamide suggests that it may exist in many astrochemical environments;
measurement of the microwave spectrum of this anion would allow observational astronomers to search
for this species in the interstellar medium.

4.6

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80

81

5 Gas-Phase Acidities of the Nitrated Azoles

5.1

Introduction
Azoles, five-membered aromatic rings in which one or more CH is replaced by N, are importantorganic compounds relevant to biological and material sciences. Pyrrole and imidazole, two of thesimplest azoles, are found naturally as components of the amino acids tryptophan and histidine, thenucleobases adenine and guanine, and heme B. Because of the strong electron donor interaction of thenitrogen atoms, pyrazole and imidazole are known to coordinate with metal ions as stable ligands. 1
Furthermore, azole compounds are ubiquitous components of ionic liquids. Although ionic liquids havebeen known for over 100 years,2

there has been a resurgence of interest in their application andcharacterization. These compounds are unusual because their melting point is ≤ 100 °C. This property has
led to the use of ionic liquids as reusable green solvents because of their low vapor pressure. 3 Because of
their stability, these compounds are finding use as novel electrolytic compounds in batteries, capacitors,
and solar cells.4, 5 Additionally, ionic liquids are under study for use as a next-generation hypergolic fuel6
and as the ionic medium in ion thrusters.7 In some cases, modified azole compounds have found use as
both the cationic and anionic species in ionic liquids.8, 9 For example, 4-nitroimidazolate, an anion studied
herein, has been combined with 1-butyl-3-methylimidazolium to produce a compound that melts at –63°C and is thermally stable to 185 °C.9 In this work, we examine the gas-phase acidity of several nitrated
azoles in an effort to characterize the stability of their conjugate bases, i.e., their corresponding anions.
The nitrated molecules of this study, illustrated in Figure 5.1 below the parent azoles, vary by theplacement of the ring-nitrogen atom, as well as the placement of the nitro group.

82

Figure 5.1 Selected azoles and their nitrated derivatives.
The acidic hydrogen is indicated on the pertinent nitrogen atom.

The gas-phase acidity of molecules is of fundamental interest to many chemists. This property reflects
the intrinsic stabilities of molecules and their deprotonated anions. It is common practice to determine theacidity of molecules by relative methods that compare a molecule of unknown acidity to one for which
the acidity has been previously determined. Several mass-spectrometry approaches are used to study the
relative gas-phase acidities of molecules including proton-transfer equilibrium,10-16 bracketing,17, 18 andthe kinetic method.19-28 Historically, our lab has used the Flowing Afterglow-Selected Ion Flow Tube
(FA-SIFT) to determine the gas-phase acidity of molecules.14-17 However, we were not able to use the FASIFT for this study because both the nitroazoles and the reference organic acids are not sufficiently

83volatile. Instead, we employed Cook’s kinetic method20 and a triple quadrupole instrument to determine
the gas-phase acidities of the nitrated azoles.
The kinetic method is a competitive dissociation technique that examines the relative proton affinityof two anions (gas-phase acidity) or two neutrals (proton affinity). The technique monitors the collisioninduced fragmentation of anionic–proton-bound–hetero dimers. (Scheme 5.1)

Scheme 5.1 Competitive dissociation of anionic proton-bound heterodimers

In Scheme 5.1, HAi represents a series of reference acids with known gas-phase acidities, and HA0
represents a compound of unknown acidity. The ratio of the intensities of the two product ions (A i– andA0–) is directly related to the ratio of the two rate constants (ki and k0). N2 is the collision gas used for thisexperiment.
The kinetic method has been used extensively over the last 20 years. The extended kinetic method, an
improvement over Cook’s standard method,20 was introduced by Fensleau and coworkers.23 Thisapproach requires that the reference acids are functionally similar among themselves, but different fromthe unknown acid under investigation. This allows the assumption that the entropic contributions for
proton transfer between the unknown and reference acid anions are constant [ΔΔS°(HA0,HAi) =ΔS°acid(HAi) – ΔS°acid(HA0) ≈ constant]. This assumption becomes a good approximation when thereference acids are thoughtfully chosen to be similar in structure as well as functionality. The extended
kinetic method also requires that the tandem mass spectrometry experiments are performed over multiple
collision activation conditions. To perform the analysis, the logarithms of the ratio of the two rate84

constants (ln[ki/k0]) from Scheme 5.1 are plotted against proton affinities of the reference anions,
ΔH°acid(HAi). This provides the relationship:[𝐴𝑖 − ]𝑘𝑖
∆𝐻 ° 𝑎𝑐𝑖𝑑 (𝐻𝐴𝑖 ) 𝐺𝐴𝑎𝑝𝑝 (𝐻𝐴0 )𝑙𝑛≈ 𝑙𝑛 ≈ −+[𝐴0 − ]
𝑘0
𝑅𝑇𝑒𝑓𝑓
𝑅𝑇𝑒𝑓𝑓

5.1

In equation 5.1, the apparent gas-phase acidity term, GAapp, is the analog of the apparent gas-phase
basicity introduced by Wesdemiotis and coworkers.24 At the x-intercept (where the ratio of ki/k0 equals 1),
GAapp(HA0) = ΔH°acid(HAi). The effective temperature (Teff) represents the internal energy of thecollisionally activated heterodimer. It is not a true thermodynamic temperature, but rather a measure of
the distribution of the energy available to the dimer.22, 29, 30 When ln[Ai–/A0–] is plotted against
ΔH°acid(HAi), a linear regression is obtained with slope equal to ‒1/RTeff. Each unique collision activation
condition provides a characteristic Teff and GAapp(HA0). Because the extended kinetic method requires theuse of multiple activation conditions, multiple Teff and GAapp(HA0) values are obtained. The GAapp(HA0)
values are then plotted versus Teff using the relationship:
𝐺𝐴𝑎𝑝𝑝 (𝐻𝐴0 ) = ∆𝐻° 𝑎𝑐𝑖𝑑 (𝐻𝐴0 ) − 𝑇𝑒𝑓𝑓 ∆∆𝑆 ° (𝐻𝐴0 , 𝐻𝐴𝑖 )5.2

For equation 5.2, ΔΔS°(HA0,HAi) is provided by the slope of the line, while determination of the yintercept gives ΔH°acid(HA0). This two plot technique for determining ΔH°acid(HA0) was first introduced
by Armentrout25 as the “alternative method.”More recently, another method for analyzing extended kinetic method data, called the ‘Orthogonal
Distance Regression (ODR)’ by Ervin and Armentrout26 or the ‘isoequilibrium (or isothermal) point’
method by Bouchoux et al.19 has been described. Instead of using two plots to find ΔH°acid(HA0), theisoequilibrium method determines the values of ΔH°acid(HA0) and ΔΔS°(HA0,HAi) from a single plot in
which the linear regressions of ln(ki/k0) vs. ΔH°acid(HAi) have converged to a single point, called theisoequilibrium point.19, 27 Problematically, several experiments have shown the existence of multiple
85isoequilibrium points.21 Therefore, a treatment has been developed which forces the convergence of an
isoequilibrium point.26, 28 However, it has been found that results for data analyzed using theisoequilibrium and alternative methods deviate by less than 0.5 kcal/mol when compared over the same
collision energy range.21
5.2

Experimental
The MS/MS experiments were performed on a QTrap 4000 (AB SCIEX, Toronto, Canada), 31 triple

quadrupole mass spectrometer; the QTrap is described in detail in Chapter 2. Using ESI allowed facile
introduction of the proton-bound hetero dimers into the gas-phase. The QTrap is used to perform collisioninduced dissociation (CID) on the mass-selected proton-bound heterodimers, Scheme 5.1. Collisionenergies from 5–22 eV (lab frame) are used. The collision energies are not corrected for perturbations of
the true zero of energy, nor are they corrected for multiple collision events. The intensity of ion fragments
and undissociated parent ions are monitored with Q3, which is coupled to an electron multiplier. Typical,
raw CID spectra from a kinetic method experiment are shown in Figure 5.2.

86Figure 5.2 CID spectra for the proton-bound heterodimer of 2-nitroimidazole and 3-nitrobenzoic
acid at four collision energies
m/z 279, 166, and 112 correspond to the proton-bound heterodimer, deprotonated 3-nitrobenzoic acid,and deprotonated 2-nitroimidazole, respectively. m/z 122 corresponds to the loss of CO2 fromdeprotonated 3-nitrobenzoic acid. The four spectra illustrate how the relative intensities of the CID
fragments will change as a function of collision energy. Note that these experiments are duplicated foreach reference acid when performing an extended kinetic method experiment.

The kinetic method data were analyzed using the alternative method.25 Statistical analysis of the data
was performed using Igor Pro (WaveMetrics, Inc.; Lake Oswego, Oregon, USA), which implements thefreely available software package, ODRPACK95.32 The first plots of ln[Ai–/A0–] vs. ΔH°acid(HAi) were
weighted using ordinary least squares with respect to the standard deviations of ln[Ai–/A0–] at the 95%
confidence interval. The experimental errors for the acidities of the reference acids, ΔH°acid(HAi), were
not included in the statistical analysis of the linear regressions of the first plots. For the second plots of87

GAapp vs. Teff, weighted ODR was implemented to determine ΔΔS° and ΔH°acid(HA0) using the curvefitting function ODR=2 in Igor Pro, also at the 95% confidence interval. Although we use ODR to weightthe linear regressions of the second plots, we are not using the isoequilibrium method described by Ervinand Bouchoux.19, 26-28
All of the chemicals used in this study were purchased from Sigma-Aldrich (St. Louis, Missouri,
USA) except for 2-nitropyrrole, 3-nitropyrrole, and 3-nitropyrazole which were purchased from Oxchem
Corporation (Irwindale, California, USA), Acella Pharmaceuticals (San Diego, California, USA), andMatrix Scientific (Columbia, South Carolina, USA), respectively.
5.3

Calculations
Theoretical calculations were carried out using the Gaussian 09 suite of programs.33 Hybrid

functional calculations at the B3LYP/aug-cc-pVTZ level of theory were used to determine theoretical
ΔG°acid, ΔH°acid, and ΔS°acid values for the unknown nitroazoles investigated in this study. Thermal
corrections to entropy and free energy were applied under the rigid-rotor harmonic-oscillator
approximation.
In addition to reporting theoretical ΔG°acid values, we also report ΔG°acid,experiment values determined by
a combination of theory and experiment where ΔG°acid,experiment = ΔH°acid,experiment ‒ TΔS°acid,theory.
5.4

Results
The reference acids used for this study are listed in Table 5.1. The ΔH°acid and ΔG°acid values werepreviously determined using ion-molecule reaction equilibrium methods.10-13

ΔH°acid for 7 (4-cyanobenzoic acid) was redetermined for use in this study with the extended kinetic method and thealternative method for data analysis.25 The new value of ΔH°acid(HAi) = 329.4±2.2 kcal/mol was used for
subsequent experiments in which 7 was employed as a reference; this value agrees with the previously
determined value, 327.8±2.1 kcal/mol,34 within experimental error. For its experimental determination, 788was referenced against the carboxylic acids 5, 6, 8, 9, 10, and 11. Because 7, like the reference acids it
was compared to, is a carboxylic acid, the second plot of GAapp vs. Teff is characterized by a line with asmall slope, where ΔΔS°=0.533±0.303 cal/(mol∙K). This result is expected for proton transfer between
functionally and structurally similar anions.
Table 5.1 Reference acids
Reference Acid
ΔHᵒacid

Literaturea
ΔGᵒacid

Theoryb
ΔHᵒacid ΔGᵒacid

1: phenylacetic acid
2: 3-toluic acid
3: benzoic acid
4: phenoxyacetic acid
5: chloroacetic acid
6: 4-(trifluoromethyl)benzoic acid341.5(±2.1)
340.6(±2.1)
340.2(±2.2)
338.0(±2.2)336.5(±2.2)
332.3(±2.1)334.5(±2.0)333.6(±2.0)333.1(±2.0)
331.0(±2.0)
329.0(±2.0)
325.3(±2.0)

339.8
340.8
340.0
335.8
332.7
331.9

332.7
332.9
332.3
328.4
325.6
324.3

7: 4-cyanobenzoic acidc
8: 3-nitrobenzoic acid
9: 4-nitrobenzoic acid
10: salicylic acid
11: 3,5-bis(trifluoromethyl)benzoic acid329.4(±2.2)
329.1(±2.1)
328.1(±2.2)325.5(±2.2)324.4(±2.1)321.7(±2.4)322.1(±2.0)
321.1(±2.0)317.8(±2.0)317.4(±2.0)

329.0
328.8
327.5
328.0
325.8

321.3
321.1
319.7
320.6
318.1

a

Values from the NIST Chemistry WebBook (Ref. 34). bTheoretical calculations were carried out
using Gaussian 09 at the B3LYP/aug-cc-pVTZ level of theory, and thermal corrections to entropy
and free energy are included under the rigid rotor-harmonic oscillator approximation for T = 298
K. cThe ΔH° and ΔG° values for 7 were redetermined before use in this experiment using theextended kinetic method with the alternative method for data analysis.

The nitrated azole species in this study were extensive enough in their ΔH°acid(HA0) values to warrant
their pairing with three separate groups of reference acids. 2- and 3-nitropyrrole were referenced against
1, 2, 3, 4, 5, and 6. 3- and 4-nitropyrazole were referenced against 5, 6, 7, 8, and 9. Lastly, 2- and 4nitroimidazole were referenced against 6, 7, 8, 9, 10, and 11. As shown, there is some overlap between
the groups of reference acids and unknown acid pairs.
We collected data for each of the azole acid systems at multiple lab-frame collision energies from ~522 eV. The limit on the collision energy range is determined for each experiment by analyzing howsmoothly the effective temperature (Teff) rises with respect to the collision energies set by the89instrument.21 Although the data can also be analyzed in the center-of-mass frame, Armentrout andcoworkers reported that converting to the center-of-mass frame does not change the character of theresults.21 For that reason, we have analyzed the data in the conventional laboratory frame.
The ratios of the product ions, Ai–/A0–, were determined by examining the peak heights of pertinent
ions as extracted from MS/MS data. Notably, several of the carboxylic reference acids exhibited a loss ofm/z 44 (CO2) upon using higher collision energies (≥14 eV).

Additionally, some of the species

containing a nitro group (–NO2) exhibited a loss of m/z 30 (NO). Chloride ion was observed as asecondary product ion when reference acid 5 was used. The intensities of the secondary product ions were
accounted for by adding their intensities to the respective primary ion intensities and no further
corrections were made.
Figure 5.3 shows five plots of ln[ki/k0] vs. ΔH°acid(HAi) for 4-nitropyrazole and its reference acids atthe collision energies 5, 7, 10, 14 and 18 eV. Most reports overlay all regressions onto a single plot, andin an ideal situation the linear regressions of the different collision energies will cross at a single point,
known as the isoequilibrium point. However, most extended kinetic method data deviate from a single
crossing point,21 and the data for this study are no exception. The slope of the linear regression decreases
as the collision energy increase from Figure 5.3a to Figure 5.3e. This decrease in slope corresponds to anincrease in Teff, because slope is proportional to 1/Teff. Simultaneously, the x-intercept, indicated by a blue
arrow, increases from Figure 5.3a to Figure 5.3e, corresponding to an increase in the GAapp term. The twovalues, GAapp and Teff, are plotted in a second graph, shown as Figure 5.4. Analysis of the y-intercept andslope of the second plot provides values for ΔH°acid(HA0) and ΔΔS°(HA0,HAi), as shown by the
relationship in equation 5.2.

90Figure 5.3 Plots of ln(ki/k0) vs. ΔH°acid (kcal/mol) for 4-nitropyrazole.
Plots a, b, c, d, and e correspond to data taken at collision energies 5, 7, 10, 14, and 18 eV (lab frame),
respectively. As the collision energy increases, the slope (–1/RTeff) becomes less negative, and the xintercept increases; at the x-intercept, GAapp = ΔH°acid. Note that the ΔH°acid values plotted correspondto the reference acids 5, 6, 7, 8, and 9, not 4-nitropyrazole.

91

Figure 5.4 Plot of GAapp vs. Teff values for 4-nitropyrazole.
The data in Figure 5.4 are provided by Figure 5.3. The points are weighted in both the dependent andindependent variables at the 95% confidence interval using ODR; the error bars represent two standard
deviations of the data. The slope and y-intercept of the weighted line gives values for ΔΔS° (3.7cal/(mol∙K)) and ΔH°acid (329.5 kcal/mol), respectively.

Table 5.2 gives a complete summary of the results of our measurements. The experimental results
ΔH°acid(HA0) were obtained by the alternative method of data analysis. Additionally, ΔH°acid(HA0) values
were theoretically calculated for each nitrated azole. The calculated ΔH°acid(HA0) values agree with theexperimental values within 1 kcal/mol, except for 3-nitropyrrole which deviates from the experimental
value by 1.7 kcal/mol. We also report ΔG°acid(HA0) values for the unknowns; these values were
determined by calculating TΔS°acid (T=298 K) and subtracting the result from the experimental ΔH°acid
values (ΔG°acid,experiment = ΔH°acid,experiment - TΔS°acid,theory). This approach provides reasonably accurate
ΔG°acid values because the entropic contribution to the free energy of deprotonation can be calculated with92

reasonable certainty. The absolute error reported for the ΔG°acid values is ±2.4 kcal/mol because of theadded uncertainty introduced by the calculation.
Table 5.2 Results of kinetic method and computational studies
Nitrated Azole

Ecoll(eV) Teff(K)cExperimenta

ΔHᵒacid(kcal/mol)

TheoryaΔΔSᵒ[cal/(mol∙K)]b

ΔGᵒacidd

ΔGᵒacid

ΔHᵒacid(kcal/mol)(kcal/mol) (kcal/mol) (kcal/mol)

TΔSᵒacid6‒18435‒645 337.0(±2.2) 2.3(±0.2)

329.4 (±2.4) 336.0

328.4

7.6

3-nitropyrrole 6‒18
3-nitropyrazole 6‒18

560‒880 335.8(±2.2) 3.2(±0.2)690‒1010 330.5(±2.2) 4.7(±0.3)

328.4(±2.4) 334.1
323.1(±2.4) 331.2

326.6
323.7

7.4
7.4

4-nitropyrazole 5‒18570‒930 329.5(±2.2) 3.7(±0.3)

322.0(±2.4) 330.0

322.5

7.5

2-nitroimidazole 6‒22
4-nitroimidazole 6‒22500‒735 327.4(±2.2) 1.9(±0.3)625‒975 325.0(±2.2) 3.0(±0.3)

319.7(±2.4) 328.2
317.6(±2.4) 325.8

320.5
318.4

7.7
7.4

2-nitropyrrolea

The values ΔHᵒacid and ΔΔSᵒ were determined using the alternative method for extended kinetic method data
analysis. ODR was used to weight the two values in the second plot. The errors for ΔH°acid arise from theuncertainties of the reference acids. The experimental precision from the ODR statistical treatment of the data at theb
95% confidence interval is much better (≤ ±0.4 kcal/mol) than the absolute error. ΔGᵒacid,experiment was determined byca combination of theory and experiment where (ΔG°acid,experiment = ΔH° acid,experiment - TΔS°acid,theory). Theoretical
calculations were carried out using Gaussian 09 at the B3LYP/aug-cc-pVTZ level of theory, and thermal correctionsdto entropy and free energy are included under the rigid rotor-harmonic oscillator approximation. T=298 K.

In addition to the experimental kinetic method measurements, we performed a theoreticalinvestigation to determine the ΔH°acid(HA), ΔG°acid(HA), and electron affinity (EA) values for 24 azolate
and nitrated-azolate species. These results supplement the experimental work by investigating higher
order aza compounds, as well investigating azoles with two nitro groups. The results are presented inTable 5.3,14-16, 34-38 which, in addition to theoretical values, compiles previously determined experimental
values. The theoretical anion proton affinities agree with the experimental measurements within 3
kcal/mol.

93

Table 5.3 Experimental and theoretical thermochemical properties of nitrated azoles
Azole

Experiment

Theoryb,cc,dΔH°acid(kcal/mol)eΔG°acid(kcal/mol)

a

b,c

c,d

e

ΔH°acid(kcal/mol)

ΔG°acid(kcal/mol)

EA
(eV)

360.1

352.6

2.104

336.0

328.4

3.436

334.1

326.6

3.416

2,3-dinitropyrrole

317.2

309.8

4.165

2,4-dinitropyrrole

315.0

307.4

4.487

2,5-dinitropyrrole

315.2

307.5

4.501

355.7

348.1

2.908

(N1)331.2(N2)331.7(N2)323.7(N2)324.0

4.154

330.0

322.5

4.220

(N1)313.1(N2)312.1(N1)306.4(N2)304.6

5.011

311.3

303.7

5.125

350.8

343.3

2.562

328.2

320.5

3.789

(N1)325.8(N3)326.6(N1)318.4(N3)319.0

3.865

2,4-dinitroimidazole(N1)307.9(N3)307.3(N1)300.1(N3)299.5

4.836

4,5-dinitroimidazole

309.6

302.5

4.851

(N1)343.2(N2)346.9(N1)335.6(N2)339.3

3.420

4-nitro-1,2,3-triazole(N1)319.4(N2)322.3(N3)318.5(N1)311.9(N2)314.8(N3)310.9

4.626

4,5-dinitro-1,2,3-triazole(N1)301.5(N2)304.9(N1)294.1(N2)297.9

5.377(N1)345.2(N4)339.4(N1)337.7(N4)331.9

3.367

3-nitro-1,2,4-triazole(N1)321.8(N2)322.6(N4)317.2(N1)314.6(N2)314.8(N4)309.5

4.503

3,5-dinitro-1,2,4-triazole(N1)303.1(N4)298.1(N1)295.4(N4)290.0

5.428

pyrrole359.54(±0.25)
35358.6(±2.2)2-nitropyrrole

337.0(±2.2)3-nitropyrrole

335.8(±2.2)

pyrazole

353.6(±2.4)3-nitropyrazole

330.5(±2.2)4-nitropyrazole

329.5(±2.2)

14350.9(±2.0)

EA
(eV)
14

2.145(±0.010)

14

f

f

16

346.4(±2.3)

162.938(±0.005)

16

f

f

3,4-dinitropyrazole

2,5-dinitropyrazole
imidazole

349.93(±0.72)
36
350.1(±2.1)2-nitroimidazole

327.4(±2.2)4-nitroimidazole

1,2,3-triazole

1,2,4-triazole

325.0(±2.2)

15342.60(±0.40)
36342.8(±2.0)

152.613(±0.006)

15

f

f

(N2)346.4(±2.1)(N1)344.2(±2.1)

3734(N2)339.1(±2.0)(N1)336.9(±2.0)

37

34

94

3.447(±0.004)

38

tetrazole(N1)333.7(±2.1)

34(N1)326.4(±2.0)34nitrotetrazolea(N1)331.1(N2)333.0(N1)323.5(N2)325.4

4.070(N1)308.4(N2)310.2(N1)300.6(N2)302.9

5.172

Theoretical values calculated using B3LYP/aug-cc-pVTZ methods. ΔH°acid and ΔG°acid are reported at 298 K using the rigid-rotor harmonicbcoscillator approximation. EA are reported at 0 K including zero point energies. Proton affinity of the anion. The nitrogen-atom acid site isdeindicated in parentheses for both experimental and theoretical anion proton affinities and gas-phase acidities. Gas-phase acidity. Electronfbinding energy of the anion (electron affinity of the corresponding neutral radical). Experimental values determined from this work.

5.5

Discussion

Recently,

Armentrout

andcoworkers

investigatedthegas-phase

aciditiesofrigid-planar-

monohalogenated phenols using the B3LYP/aug-cc-pVTZ level of theory. (It is noteworthy to mention
that they optimized the molecular geometries using a slightly smaller basis set, aug-cc-pVDZ.) Theirexperimental and theoretical 298 K acidity values present a fairly good correlation, with mean absolutedeviations of 2±1 kcal/mol.21 To further investigate the correlation between theory and experiment, we
calculated the 298 K acidity values for the reference acids used in these studies, and found mean absolutedeviations of 1±3 kcal/mol. However, when species containing halogens are removed, as well as
phenoxyacetic acid, which likely has anharmonic vibrations due to the floppy nature of the etheral,carboxylic arm, the correlation is improved to 0±3 kcal/mol. Therefore, computations at the B3LYP/augcc-pVTZ level of theory should appropriately predict the gas-phase acidities of the rigid-planar-organic
molecules investigated in Table 5.3 with a reasonable certainty.
The aza substituent effect was investigated previously by Taft et al.36 Upon finding that imidazole is astronger acid than pyrazole, they argued that the lone pair repulsion from the ring-nitrogen atoms on thepyrazole anion destabilizes it, making pyrazole a weaker acid than imidazole. Additionally, they report
the conjugate bases of pyrazole and imidazole to be more stable than pyrrole by 4.8 and 8.4 kcal/mol,
respectively. The work of Taft et al. exploits ΔG°acid(HA), while our work examines ΔH°acid(HA).
However, relative differences between these two thermodynamic properties are comparable because theentropic contribution is roughly constant with TΔS usually changing by less than 1 kcal/mol for any given95

pair of acids. Our work expands these observations and explores trends in nitration effects. Based on the
difference in acidity between 2-nitropyrrole, 3-nitropyrrole and pyrrole; 3-nitropyrazole, 4-nitropyrazole
and pyrazole; and 2-nitroimidazole, 4-nitroimidazole and imidazole, we report that when a nitro group issubstituted onto an azole, it increases the acidity by 23.5±1.9 kcal/mol at the 95% confidence interval.14-16
This is in perfect agreement with the experimentally determined differences in acidity between benzene
and nitrobenzene, 24.2 kcal/mol.18, 39
As stated in the results, the anions under investigation for this study were sufficiently different in
ΔH°acid(HA0) values to require their evaluation with three separate groups of reference acids. Although wereport the absolute error for each measured ΔH°acid(HA0) as ±2.2 kcal/mol, the relative difference inacidity between each pair is 1.2±0.3 kcal/mol for 2-nitropyrrole and 3-nitropyrrole, 1.0±0.7 kcal/mol for
3-nitropyrazole and 4-nitropyrazole, and 2.4±0.6 kcal/mol for 2-nitroimidazole and 4-nitroimidazole.
Relative uncertainties are reported at the 95% confidence interval (2σ). This ordering is reinforced
because the acidities of each pair were determined using the same reference acids. Furthermore, theexperiments are much more precise (±0.4 kcal/mol) than the absolute uncertainty (±2.2 kcal/mol). This
allows us to arrange with confidence the proposed azoles in order of increasing acidity as 2-nitropyrrole,
3-nitropyrrole, 3-nitropyrazole, 4-nitropyrazole, 2-nitroimidazole, and 4-nitroimidazole, which have
ΔH°acid (kcal/mol) values of 337.0, 335.8, 330.5, 329.5, 327.4, and 325.0, respectively. To provide somecontext, other common anions of ionic liquids including Cl–, Br–, CF3SO3–, and NTf2– have ΔH°acid(kcal/mol) values of 333.40, 323.54, 305.4, and 293.3, respectively.
The differences in acidity among the nitrated azoles can be attributed to nitrogen atom and nitro
group placement. Imidazole is more acidic than pyrazole because pyrazole is destabilized by the adjacentplacement of the two nitrogen atoms. A similar destabilization occurs when a nitro group is locatedadjacent to a ring-nitrogen atom. The optimized equilibrium geometries for each of the six nitrated azoles
experimentally investigated in this study yield planar molecules. Therefore, the oxygen atoms of the nitro
groups are in close proximity to the neighboring ring atoms. This orientation introduces a region of high96

electron density in proximity to ring atoms that are adjacent to the nitro group. Understanding thegeometries helps us explain the relative ordering between the sub groups. 2-nitropyrrole is a weaker acidthan 3-nitropyrrole because the electron dense region of the oxygen atom is repulsive towards the electron
dense region of the ring-nitrogen atom. This same argument holds for the difference in acidity between 3and 4-nitropyrazole; 4-nitropyrazole is more acidic than 3-nitropyrazole because 3-nitropyrazole faces ahigher degree of electron repulsion when the electron dense oxygen atom is in close proximity to theelectron dense ring-nitrogen atom. Lastly, the difference in acidity between 2- and 4-nitroimidazole ismore pronounced because the nitro group of 2-nitroimidazole is placed between two ring-nitrogen atoms,
resulting in greater electron density overlap. In summary, when a nitro group is placed on an azole ringadjacent to a ring nitrogen atom, it will destabilize the anion by about 1.1±0.3 kcal/mol as compared tothe isomer with the nitro group located away from the ring-nitrogen atom.
A few more trends are evident from the additional theoretical investigation, which are consistent withour chemical intuition. First, for higher order aza substitution, molecules of higher acidity are produced.
Second, as the number of nitro groups on the molecule increases, stronger acids are generated. The
electron withdrawing properties of the ring-nitrogen atoms and the nitro groups stabilize the additional
electronic charge of the anion, allowing the parent molecule to more easily lose a proton. Again, we foundwithin the dinitro compounds that the nitro group placement can affect the stability of the anion. Placing
the nitro groups adjacent to each other will decrease the anion’s stability by about 2 kcal/mol. (See thedifference between 2,3- and 2,4-dinitropyrrole, and 2,4- and 4,5-dinitroimidazole.) While this result is
consistent with an increase in electrostatic repulsion, spatial constraint issues are also contributing, i.e.
2,4-dinitropyrrole and 2,4-dinitroimidazole are planar while our calculations show that the nitro-group–
oxygen atoms of 2,3-dinitropyrrole and 4,5-dinitroimidazole are outside of the molecular plane. Finally,as the acidity of the molecule increases, the electron binding energy of the conjugate base, or anion, also
increases. Both of these trends indicate an increase in the stability of the anion.

97

5.6

Conclusion
We have employed the extended kinetic method with the alternative method of data analysis todetermine ΔH°acid for six nitrated-azole compounds. The success of this work is attributed to employing
reliable reference acids of similar function and structure. The reference acids used were similar in acidity(± 6 kcal/mol) to the acidity of the nitrated azoles, and the reference acids’ acidity range encompassed theacidity of the nitrated azoles. The results of this experiment elucidate trends in the mono-nitration ofpyrrole, pyrazole, and imidazole. We conclude that the pertinent anions of this study are very stable, andthey should be capable of generating robust ionic liquids. In addition, previous studies have shown thatthe basicity of an ionic liquid is dominated by the gas-phase basicity of the anion.8, 40 Therefore, theseexperimental measurements provide insight into the Lewis-base properties of ionic liquids in which theanions of this study are employed.
Furthermore, 4-nitroimidazole, the strongest acid of this study, has already found use as the anioniccomponent of ionic liquids in a recent study of azole based ionic liquids conducted by Rogers andcoworkers.9 For future work, it would be interesting to investigate the gas-phase acidity of the remaining
molecules of Rogers’ report including the nitrated triazoles, tetrazoles, benzoimidazoles, andbenzotriazoles. In addition to our experimental work, we have computationally determined the gas-phase
acidity of five other azolate anions found in Rogers’ study including 2,4-dinitroimidazole; 4,5dinitroimidazole; 4-nitro-1,2,3-triazole; 3,5-dinitro-1,2,4-triazole; and tetrazole (ΔH°acid = 307, 310, 319,
298, and, 331 kcal/mol, respectively, for the most acidic sites). These results suggest that most of theazolate based ionic liquids should have a reasonably low melting point while maintaining thermal
stability up to around 200 °C. From our investigation, no clear trend between acidity and ionic liquidmelting point is revealed. Conversely, the thermal stability of the ionic liquid pair increases with the gasphase stability of the anion.998

Finally, these studies indicate that the addition of a nitro group to an aromatic ring will increase thecompound’s acidity/anionic-stability by about 23.5±1.9 kcal/mol. Moreover, these molecules are
marginally destabilized by about 1.1±0.3 kcal/mol when the nitro group is placed adjacent to a ringnitrogen atom. By understanding these trends, researchers will have greater control over thethermodynamic properties of novel ionic liquids.
5.7

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Martin, K. Morokuma, V.G. Zakrzewski, G.A. Voth, P. Salvador, J.J. Dannenberg, S. Dapprich,
A.D. Daniels, Ö. Farkas, J.B. Foresman, J.V. Ortiz, J. Cioslowski, D.J. Fox, Gaussian 09, Gaussian,Inc., Wallingford CT, 2009.
[34] S.G. Lias, J.E. Bartmess, J.F. Liebman, J.L. Holmes, W.G. Mallard, NIST Chemistry WebBook, in:
P.J. Lindstrom, W.G. Mallard (Eds.) Ion Energetics Data, National Institute of Standards andTechnology, Gaithersburg, MD, 2014.
[35] J.E. Bartmess, J.A. Scott, R.T. McIver, Scale of Acidities in the Gas-Phase from Methanol to Phenol,
J. Am. Chem. Soc., 101 (1979) 6046-6056.
[36] R.W. Taft, F. Anvia, M. Taagepera, J. Catalan, J. Elguero, Electrostatic proximity effects in therelative basicities and acidities of pyrazole, imidazole, pyridazine, and pyrimidine, J. Am. Chem.
Soc., 108 (1986) 3237-3239.
[37] J. Catalan, R.M. Claramunt, J. Elguero, J. Laynez, M. Menendez, F. Anvia, J.H. Quian, M.
Taagepera, R.W. Taft, Basicity and Acidity of Azoles - the Annelation Effect in Azoles, J. Am. Chem.
Soc., 110 (1988) 4105-4111.
[38] T. Ichino, D.H. Andrews, G.J. Rathbone, F. Misaizu, R.M.D. Calvi, S.W. Wren, S. Kato, V.M.
Bierbaum, W.C. Lineberger, Ion chemistry of 1H-1,2,3-Triazole, J. Phys. Chem. B, 112 (2008) 545557.
[39] R.F. Gunion, M.K. Gilles, M.L. Polak, W.C. Lineberger, Ultraviolet photoelectron-spectroscopy ofthe phenide, benzyl and phenoxide anions, with ab initio calculations, Int. J. Mass Spectrom. Ion
Process., 117 (1992) 601-620.
[40] L. Crowhurst, P.R. Mawdsley, J.M. Perez-Arlandis, P.A. Salter, T. Welton, Solvent-soluteinteractions in ionic liquids, Phys. Chem. Chem. Phys., 5 (2003) 2790-2794.
101

6 Gas-Phase Investigations of the α-Effect

6.1

Introduction
About 50 years ago, Pearson and Edwards initiated comprehensive studies of the rates of nucleophilic

chemical reactions in solution. They discovered that some nucleophiles react more rapidly than expectedbased purely on the basicity of the nucleophile.1 These highly-reactive super-nucleophiles were named αnucleophiles, as they were observed to possess a lone pair of electrons on the atom adjacent to theattacking center, Figure 6.1. Their enhanced reactivity was named the α-effect.

Figure 6.1 α-nucleophiles
α-nucleophiles possess one or more lone pair of electrons (red) on the atom adjacent to the nucleophilic
atom. The α-nucleophile may be anionic (i.e., peroxide, hypochlorite) or neutral (i.e., hydrazine,hydroxylamine).

Since Pearson and Edward’s discovery, the origin of the α-effect has been discussed in great detail.
Some suggested reasons for enhanced nucleophilicity of α-nucleophiles include ground-state
destabilization, transition-state stabilization, and solvent effects. The first two reasons imply that theenhanced reactivity arises from intrinsic properties of the α-nucleophiles, while the latter arises fromdifferential solvation of the α-nucleophiles compared with the normal-nucleophiles.
The notion of ground-state destabilization assumes that the lone pairs of electrons on the α-atomexperience electrostatic repulsion with the lone pairs at the reaction center. The ground-state
destabilization effect would increase the energy of the ground-state of the α-nucleophiles,2-4 resulting in alower relative transition-state barrier therefore increasing the rate of reaction. Nonetheless, this argument
is problematic because ground-state destabilization would not only affect the rate of the reaction, but also102

affect the basicity of the nucleophile under investigation.5-7 Therefore, no deviation from the rate-energy
relationship portrayed by Brønsted correlations (vide infra) would be observed.
Transition-state stabilization has also been proposed as a reason for the increased reactivity of αnucleophiles vs. normal-nucleophiles. An increase in the stability of the transition state would result in a
reduction of the transition-state barrier height thereby increasing the rate of reaction. The best chemicalargument for transition-state stabilization follows the assumption that nucleophilic reactions occur
through a single-electron transfer process.8, 9 Because nucleophilic reactions occur through incipient bondformation, rationalizing the mechanism through single-electron transfer is certainly reasonable. If singleelectron transfer is occurring, the nucleophile will resemble a radical in the transition state, and it is wellknown that radicals are stabilized by neighboring lone pairs.10
Since its discovery, the α-effect has been observed in diverse solvents,5, 11-16 although the magnitudeof the α-effect (Effα-Nu/Effnormal-Nu) is highly solvent dependent.5, 17 The results suggest that α-nucleophiles
are more easily desolvated than normal nucleophiles in some environments. It is possible that the diffuse
charge of the α-nucleophiles assists desolvation at the reaction center when compared to normalnucleophiles. However, these results alone cannot dismiss the possibility of an intrinsic contribution to
the α-effect.
This chapter will summarize several gas-phase α-effect experiments executed using the selected ion
flow tube (SIFT). Thanks to context provided by decades of discussion and experimentation, the α-effecthas been observed in the gas-phase.18-24 The first section on Marcus theory will discuss experiments forobserving the α-effect in the gas-phase; in particular, the reaction must not be too exothermic or else
small changes in the intrinsic transition-state barrier between nucleophiles will be masked.20 The
following section examines the α-effect through SN2 and E2 reactions, and reports a minor intrinsic
contribution to the α-effect.21 Micro-solvation is the topic of the next section, and the micro-solvation
studies reveal that, even in the presence of a single solvent molecule, the magnitude of the α-effect is103

dramatically increased.23, 24 The final section reports α-effect studies through competing reaction
mechanisms with methyl formate. Again, a minor intrinsic component of the α-effect is revealed,21 andinvestigations of the micro-solvated nucleophiles show a large enhancement of the α-effect.22
6.2

Marcus Theory and the α-Effect in Gas-Phase SN2 Reactions
Marcus theory was developed to describe barriers associated with single electron transfer dynamics.

Moreover, the theory is generally relevant for describing gas-phase double-well potentials, such as thoseobserved in SN2 reactions.25 Gas-phase identity SN2 reactions (where the attacking and leaving groups are
identical) are inherently thermoneutral, but may have an intrinsic barrier, ΔH‡intrinsic. Conversely, theactivation barriers for non-identity SN2 reactions is influenced by the reaction thermodynamics. When thereaction is exothermic, the overall activation barrier, ΔH‡activation, is reduced by thermodynamic driving
forces. Figure 6.2 provides a diagram for this relationship.

104

Figure 6.2 Marcus Theory for a SN2 potential energy surface
The SN2 potential energy surface for an identity reaction (a) and a non-identity reaction (b). The
quantity ΔH‡intrinsic represents the central barrier of an identity reaction. The central barrier for a nonidentity reaction (ΔH‡activation) is modified by the thermodynamic driving force, ΔH‡driving force. The
central barrier is often described relative to the reactant energy, ΔH‡overall.

Marcus Theory models the magnitude of the reduction in ΔH‡activation for non-identity reactions; theresults match well with experimental findings.26 Marcus’s formulation follows equation 6.1.
ΔH‡activation = ΔH‡intrinsic + ΔH‡driving force

6.1

ΔH‡driving force is calculated using equation 6.2,
ΔH‡driving force = ΔHrxn/2 + (ΔHrxn)2/(16 × (ΔH‡intrinsic – ΔHwell))

105

6.2

where ΔHwell is the complexation energy of the reactant ion-dipole complex. When discussing a nonidentity reaction, ΔH‡intrinsic is determined by considering the intrinsic barrier for the identity reaction of
the anion (ΔH‡(X,X)) and the substrate (ΔH‡(Y,Y)), equation 6.3.
ΔH‡intrinsic = ΔH‡(X,Y) = [ΔH‡(X,X) + ΔH‡(Y,Y)]/2

6.3

Figure 6.3 provides a graphical representation of equations 6.1 and 6.2 as a function of reaction
exothermicity, ΔHrxn. ΔH‡activation correlates with ΔHrxn at lower exothermicities, maintaining the relative
differences for ΔH‡intrinsic(X,Y). However, at larger reaction exothermicities the ΔH‡activation values invert. In
this ‘inversion region’ it is possible for small differences in ΔH‡activation to be masked by strong
thermodynamic driving forces.
Applying Marcus Theory to the reagent selection can be helpful for observing an α-effect. The
Marcus inversion region illustrates that large exothermicities of reaction can mask small differences in theactivation barrier. Therefore, a successful investigation of the α-effect requires a multifaceted approach toexperimental design.
To test this theory, previous α-effect studies from this laboratory were revisited. The α-effect waspreviously demonstrated in gas-phase SN2 reactions from experiments involving the normal anions, HO–,CH3O–, C2H5O–, and the α-nucleophile HOO– with methyl fluoride, CH3F; anisole, CH3OC6H5; and 4fluoroanisole, CH3OC6H4F. The magnitude of the α-effect ranged from 2.3 to >50. The results from thisstudy, as well as the tabulated Marcus Barriers, are reported in Table 6.1.

106Figure 6.3 Marcus Theory relationship of ΔH‡activation and ΔH‡driving force to ΔHrxn
Table 6.1 α-Effect and Marcus barriers for SN2 reactions
Marcus Barriersd (relative to reactants)
Nu– + M

PAa

ΔHrxnaSN2 Effb

α-effectb

ΔH‡overallc

ΔH‡(X,X)

ΔH‡(Y,Y)ΔH‡(X,Y)ΔH‡driving

ΔH‡overallforce

HO– + CH3F

1633

–91

0.0042

0.62

–15.6

56.7

–11.0

22.9

–38.5

–15.6

–

CH3O + CH3F

1598

–70

0.0007

3.7

–14.1

43.6

–11.0

16.3

–30.4

–14.1

–

1585

–60

<0.00005

>50

–5.2

52.9

–11.0

21.0

–27.0

–6.0

HOO + CH3F

1575

–65

0.0026

–19.4

28.1

–11.0

8.6

–28.0

–19.4

HO– + CH3OC6H5

C2H5O + CH3F
–

1633

–162

0.13

0.69

–28.6

40.5

39.2

39.9

–68.4

–28.6

–

1598

–141

0.04

2.3

–21.6

39.4

39.2

39.3

–60.9

–21.6

–

C2H5O + CH3OC6H5

1585

–131

0.01

9

–11.2

53.8

39.2

46.5

–57.7

–11.2

HOO– + CH3OC6H5

1575

–135

0.09

–25.6

26.7

39.2

33.0

–58.6

–25.6

HO– + CH3OC6H4F

CH3O + CH3OC6H5

1633

–174

-

-

–36.6

38.1

35.2

36.7

–73.2

–36.6

–

1598

–153

0.10

2.3

–29.6

37.2

35.2

36.2

–65.8

–29.6

–

C2H5O + CH3OC6H4F

1585

–143

0.02

12

–19.2

51.8

35.2

43.5

–62.7

–19.2

HOO– + CH3OC6H4F

1575

–147

0.23

–33.9

23.8

35.2

29.5

–63.4

–33.9

CH3O + CH3OC6H4F

Thermodynamic data in units of kJ mol–1; Proton Affinities (PA) of the nucleophile, reaction exothermicities (ΔHrxn). bSN2 reaction
efficiencies (SN2 Eff), relative α-effects (Effα-Nu/Effnormal-Nu) of reactions involving CH3F, CH3OC6H5, and CH3OC6H4F. cTransition statebarriers (ΔH‡overall) calculated using G2(+) methods for methyl fluoride reactions, G3MP2 for anisoles. dMarcus barriers calculated using
equations 6.1, 6.2, and 6.3 where ΔHwell (CH3F=50, CH3OC6H5=92, and CH3OC6H4F=102 kJ mol–1)a107

Comparing HOO– to HO– does not reveal an α-effect with the neutral reagents listed, likely because
the SN2 reactions of HO– with anisole and 4-fluoroanisole are highly exothermic (–162 and –174 kJ/mol,
respectively). These reactions therefore occur in the Marcus inversion region. In addition, the proton
affinity (PA) of HO– (1633 kJ/mol) is much greater than that for HOO– (1575 kJ/mol) suggesting different
nucleophilic behavior. Conversely, the reactivity of HOO– with the remaining alkoxides does reveal an αeffect; the exothermicities and PAs of the normal-oxyanions better match those of HOO– than hydroxide.
Investigating the α-Effect in SN2 and E2 Reactions

6.3

Nucleophiles may undergo both substitution (SN2) and elimination (E2) reactions. Reagent anions
were selected to be similar in both structure and basicity to produce meaningful Brønsted correlations. In
this section, the reactions of the normal-oxyanions FCH2CH2O– and C6H5CH2O– are compared to the αoxyanion (CH3)3COO–. The neutral reagents CH3Cl and (CH3)3CCl were selected to investigate SN2 andE2 reactivity, respectively. Reactions with CH3Cl produce Cl– exclusively through an SN2 reaction; Cl–cannot be generated from CH3Cl through an E2 reaction because CH3Cl does not contain β-hydrogens.
The reactions with (CH3)3CCl will also produce Cl–, but due to steric crowding at the electrophilic carbon,
the reactions are assumed to produce Cl– strictly from an E2 mechanism. The results of our kinetic studiesof the six reactions are presented in Table 6.2.
Table 6.2 α-effect for SN2 and E2 reactivity
Thermodynamic dataa
–

ΔHrxn

ΔH overall

kexp (× 10-10)

Eff

Effα-Nu/Effnormal-Nu

1553(1544)–130(–149)–37.4

6.28±0.20

0.31

1.6

C6H5CH2O + CH3Cl

1548(1538)–125(–148)–31.8

4.46±0.09

0.24

2.1

(CH3)3COO– + CH3Cl

1552(1552)–141(–164)–44.9

9.74±0.20

0.51

FCH2CH2O– + (CH3)3CCl

1553(1544)–88.3(–80.0)–35.1

12.3±0.2

0.56

0.9

C6H5CH2O– + (CH3)3CCl

1548(1538)–83.3(–73.7)–32.2

5.17±0.10

0.27

1.9

–

FCH2CH2O + CH3Cl–

E2

α-EffectPA

Nu + MSN2

kinetic datab‡–(CH3)3COO + (CH3)3CCl
1552(1552) –87.3(–82.3) –39.6
10.2±0.2
0.51
Thermodynamic data in units of kJ mol–1; experimental and theoretical proton affinity (PA) and reaction enthalpy (ΔHrxn),
theoretical values are in parentheses and calculated using G3MP2; theoretical transition state barriers (ΔH‡overall) calculatedusing G3MP2. bOverall rate constant expressed in units of cm3 s–1, Error bars represent one standard deviation (1σ) of the data,
absolute uncertainty is ±30%. Reaction efficiency (Eff = kexp/kcol), kcol calculated from PTT [Ref. 19]
a

Su.27

108

The SN2 reactions showed a small α-effect. The reaction efficiency is slightly lower for C6H5CH2O–
compared with FCH2CH2O– in agreement with the computational barriers. The theoretical reaction barrierof C6H5CH2O– is higher than that of FCH2CH2O– by ~6 kJ mol–1. The barrier for the α-oxyanion(CH3)3COO– is the lowest. While the magnitude of the α-effect is small, it is consistent with previous
estimations concerning small differences in intrinsic reaction barriers. Although enhanced reactivity isobserved with the α-nucleophile relative to proton affinity, part of this effect may arise from the larger
exothermicity of the reaction (~10 kJ mol–1).
The α-effect has not been previously reported in experimental elimination reactions, but has beensuggested to exist in E2 reactions involving ethyl chloride from the computational work of Ren and
Yamataka.28 Ethyl chloride was not investigated due to the competing SN2 reaction pathway. The
experimental reaction efficiencies from investigating the E2 reactions with (CH3)3CCl correlate well withthe proton affinities and reaction exothermicities of the reagent anions. This is consistent with the E2
mechanism, which requires protonation of the anion. The strength of the correlation shows that the E2
mechanism is highly sensitive to the proton affinity of the anion. The α-oxyanion (CH3)3COO– pairs
better with FCH2CH2O– than C6H5CH2O– due to a more similar proton affinity. The proton affinity of
C6H5CH2O– is lower by approximately 5 kJ mol–1. Although the reaction efficiency of (CH3)3COO– isslightly lower than that for FCH2CH2O–, the reaction efficiencies are the same within experimental error.
Due to slightly larger entropic barriers within the SN2 mechanism, the rate constants for the reagent
anions are slightly higher for the E2 mechanism than for the SN2 mechanism even though thecomputational barriers for the E2 mechanism are slightly higher.
Experimental E2 reaction efficiencies involving (CH3)3CCl from this work were combined withpreviously measured reaction efficiencies to form the Brønsted correlation in Figure 6.4. Brønsted
correlations compare the reaction efficiency (Eff = kexp/kcol, y-axis) to the anion proton affinity (x-axis),
forming a rate-energy correlation. While anions 2 through 10 have an approximately linear-fit, the overallcorrelation is non-linear, and the curvature increases at lower anion proton affinities. Figure 6.4 illustrates
109that the efficiencies of E2 reactions are strongly correlated to the proton affinity of the anion, but the nonlinear behavior suggests significant contributions from factors other than the activation barrier. The originof the curvature has been previously discussed, and is attributed to minor differences in the transition state
structures.29 There is a high amount of variability in an E2 transition state due to the three components:
proton transfer (PT) from Cβ to the nucleophile, π-bond formation between Cα and Cβ, and the geometry
between the leaving group and Cα. On the other hand, some of the curvature at higher anion proton
affinity can be attributed to the reaction efficiencies approaching the collision rate.

Figure 6.4 Brønsted correlation of E2 reaction efficiencies of oxyanions with (CH3)3CCl
Reaction efficiency (kexp/kcol) vs the proton affinity (PA) for the E2 reactions of normal- and α-nucleophiles with
(CH3)3CCl. Previously measured rate constants from references [Ref. 22-25]. Efficiencies calculated from PTT [Ref.
19]. (■, normal-anions; ● α-anions): 1. HO– 2. CH3O– 3. C2H5O– 4. HOO– 5. CFH2CH2O– 6. (CH3)3COO– 7.
CF3CH2O– 8. ClO– 9. CF3CF2CH2O– 10. BrO– 11. HC(O)OO– 12. C6H5O– 13. HC(O)O– 14. CH3C(O)O– 15.
CH3C(O)OO–.
27, 30-33

Small deviations in the correlation between proton affinity and reaction efficiencies in Table 6.2 andFigure 6.4 suggest that a minor α-effect may occur in E2 transition states. However, a substantive intrinsic
origin does not seem to prevail for an enhanced reactivity of α-nucleophiles in the E2 mechanism.
110

6.4

The α-Effect within Micro-Solvated Systems
The investigation of physical-organic chemistry in the gas-phase is highly desirable because intrinsic

properties and reactivities may be observed in the absence of solvent. In addition, several experiments
attempt to bridge the gap between condensed- and gas-phase environments by studying gas-phasereactions through micro-solvation, where a controlled number of solvent molecules are added to areagent. The micro-solvated reactions investigated in this chapter illustrate that, while the α-effect might
have a minor intrinsic contribution, solvation clearly enhanced the magnitude.22-24
The gas-phase nucleophilic reactions with methyl chloride proceed exclusively through an S N2reaction mechanism. Since there are no competing mechanisms, reaction efficiencies of normal- and αnucleophiles with methyl chloride can be compared directly. A summary of the S N2 reactions of unsolvated HO–, CH3O–, C2H5O–, (CH3)2CHO– and HOO–;18, 34 water micro-solvated HO–[H2O], CH3O–[H2O], C2H5O–[H2O], (CH3)2CHO–[H2O] and HOO–[H2O];23 and methanol micro-solvated CH3O–[CH3OH], C2H5O–[CH3OH], (CH3)2CHO–[CH3OH] and HOO–[CH3OH]24 with CH3Cl are presented inTable 6.3. Methanol solvated hydroxide, HO–[CH3OH], cannot be investigated because hydroxide will
deprotonate methanol to form hydrated methoxide, CH3O–[H2O].

111

Table 6.3 Water- methanol- and un-solvated nucleophiles in reactions with methyl chloride
thermodynamic dataa

kinetic datab

PA

ΔHrxn

kexp (×10–11)

Eff

αeffect

product fractionc
Cl–[H2O] orCl–
Cl–[CH3OH]

HO–[H2O]

1538–101/–161

36

0.15

0.29

0.85

0.15

CH3O–[H2O]

1520–103/–164

10±1

0.049

0.90

0.78

0.22

C2H5O [H2O]

1511–95/–155

3.3±0.1

0.016

2.8

0.47

0.53

(CH3)2CHO–[H2O]

1505–82/–143

0.54±0.05

0.0028

16

0.42

0.58

HOO–[H2O]

1493–80/–1409.3±0.3

0.044

0.72

0.28

CH3O–[CH3OH]

1508–91/–157

2.6±0.1

0.013

0.30

0.68

0.32

C2H5O–[CH3OH]

1500–83/–150

0.61±0.04

0.0032

1.2

0.52

0.29d

(CH3)2CHO–[CH3OH]

1496–71/–138

0.14±0.01

0.00075

5.2

ee

HOO–[CH3OH]

1466–64/–1300.79±0.06

0.0039

0.67

0.33

HO–

1635 (1633)

–214

190

0.67

0.78

1.0

CH3O–

1601 (1598)

–202

160

0.60

0.87

1.0

1584 (1585)

–187

130

0.60

0.87

1.0

(CH3)2CHO

1574 (1576)

–171

110

0.54

0.96

1.0

HOO–

1574 (1575)

–186

123

0.52

Nu

‒

–

C2H5O––

1.0

Thermodynamic data in units of kJ mol–1 calculated using G3; proton affinities (PA) of the micro-solvated anion withexperimental values in parentheses when available (Ref. 27); reaction enthalpies (ΔH rxn) for the products Cl– and clustered Cl–[H2O] or Cl–[CH3OH], respectively (Cl–/Cl–[H2O]) or (Cl–/Cl–[CH3OH]). bOverall rate constant expressed in units of cm3 s–1(Ref. 15, 16, 27); Error bars represent one standard deviation (1σ) of the data, absolute uncertainty is ±30%. Reaction efficiency(Eff = kexp/kcol), kcol calculated from PTT (Ref. 19). cAbsolute uncertainty is ±50%. dAdditional 0.19 Cl–[C2H5OH] formation.
e
Because of the low reactivity, a product distribution could not be determined.
a

PT35T27[HO–,CH3O–, C2H5O–, HOO–]36 [Water Solvated]23 [CH3OH Solvated]24
Figure 6.5 provides a Brønsted rate-energy correlation between the reaction efficiencies and proton
affinities of the un- and micro-solvated nucleophilic anions. While an α-effect is not observed for the unsolvated nucleophiles, an enhanced reactivity is observed for both the water and methanol solvated HOO–.

112

Figure 6.5 Brønsted correlation of water-, methanol-, and un-solvated nucelophiles reacting with
CH3Cl
Reaction efficiency (kexp/kcol) vs proton affinity (PA) for the SN2 reaction of un-solvated (○), watersolvated (shaded circles), and methanol solvated (●) anions with methyl chloride. Normal- and αnucleophiles are represented in blue and red, respectively. The linear trend lines are fits to the normal
nucleophiles data set comprising CH3O–, C2H5O–, and (CH3)2CHO– with methanol- and water-solvation
respectively.

Figure 6.5 clearly illustrates an α-effect for the reactivity of HOO–[H2O]; an α-effect is not observed forthe un-solvated system. Considering the PA of HOO–[H2O], this anion reacts more than an order ofmagnitude faster than is expected based on the rate-energy relationship illustrated by the other water
solvated nucleophiles. Although the un-solvated nucleophiles have a large proton affinity range (~60 kJ
mol–1), their range of reaction efficiencies is quite small (0.52 to 0.67). In contrast, the water-solvatednucleophiles show reaction efficiencies that vary from 0.0028 to 0.15, a factor of ~50. The addition of asingle water molecule to the nucleophilic reagents dramatically affects their reactivity.
113

The methanol-solvated reactions with methyl chloride also clearly illustrate an α-effect. Adding amethanol molecule (ΔPAaverage,methanol = –85 kJ mol–1) to the normal nucleophiles results in a greater
decrease in the proton affinity than adding a water molecule (ΔPAaverage,water = –74 kJ mol–1), Table 6.3.
However, solvation of HOO– shows even larger changes in stability; the computational PA of HOO– withthe addition of methanol is ΔPA = –110 kJ mol–1 and with the addition of water is ΔPA = –81 kJ mol–1.
Considering the rate-energy correlation for the methanol-solvated–normal nucleophiles, the observed
reactivity of HOO–[CH3OH] is more than two orders of magnitude faster than expected.
Due to the larger decrease in the proton affinity of the methanol-solvated ions compared with thewater-solvated ions, the relative efficiencies of the methanol solvated system are lower than the water
solvated system with methyl chloride. Despite the reduction in efficiency, Figure 6.5 illustrates a similar
curvature for the water and methanol solvation rate-energy relationships; the linear regressions presented
are only intended to guide the eye. Furthermore, the linear regressions illustrate that both water andmethanol solvation have a similar consequence on the oxy reagent α-effect studies. While the α-effect is
not observed in the un-solvated reagent system, both solvent molecules allow observation of an α-effectfor the SN2 reactivity of HOO– with methyl chloride.
Studying the α-effect between the water- or un-solvated reagents with methyl bromide is notproductive because the high exothermicity of reaction will mask subtle differences in transition statebarriers and reaction efficiencies, as established in the discussion of section 6.2. However, methanol
solvation reduces the PA of the reagent nucleophiles, therefore reducing the exothermicity of the reaction
with methyl bromide. Consequently, in addition to the reactions with methyl chloride, an α-effect isobserved in the reactions of the methanol-solvated reagents with methyl bromide. The results arepresented in Table 6.4.

114

Table 6.4 Methanol Solvated Nucleophiles Reactivity with Methyl Bromide
thermodynamic dataa

kinetic datab

product fractionc

Nu–

PA

ΔHrxn

kexp (×10–11)

Eff

α-effect

Br–

Br–[CH3OH]

CH3O–[CH3OH]

1508–123/–181

57±2

0.33

0.8

0.95

0.05

C2H5O–[CH3OH]

1500–115/–174

38±2

0.23

1.1

0.89

0.08d

(CH3)2CHO–[CH3OH]

1496–103/–162

20±1

0.13

2.0

0.91

0.09

–

HOO [CH3OH]
1466–97/–154
45±1
0.26
0.92
0.08
a
Thermodynamic data in units of kJ mol–1 calculated using G3; proton affinity (PA) of the micro-solvated anion; reaction
enthalpy (ΔHrxn) displayed for Br– and clustered Br–[CH3OH] products, respectively (Br–/Br–[CH3OH]). bOverall rate
constant (kexp) in units of cm3 s–1. Error bars represent one standard deviation (1σ), absolute uncertainty is ±30%. Reaction
efficiency (Eff = kexp/kcol), kcol calculated from PTT [Ref. 19]. cAbsolute uncertainty is ±50%. dAdditional 0.04 Br–[C2H5OH]
formation.

Figure 6.6 shows the Brønsted rate-energy correlation between the methanol solvated system withmethyl bromide (upper panel) and methyl chloride (lower panel.) The methyl bromide reactions are more
efficient, likely due to the higher exothermicity of reaction and a strong thermodynamic driving force.
The results illustrate an enhanced reactivity of HOO–[CH3OH] with both methyl bromide and chloride as
the efficiencies of the reactions with HOO–[CH3OH] deviate from the rate-energy relationship displayedwith the normal nucleophiles. Studying the reactivity of methanol-solvated nucleophiles with methyl
bromide and chloride reveal an α-effect in two different efficiency ranges.

115

Figure 6.6 Brønsted correlation of methanol-solvated nucleophiles with CH3Br and CH3Cl
Reaction efficiency (kexp/kcol) vs proton affinity (PA) for the SN2 reaction of methanol-solvatednucleophiles with CH3Br (upper panel) and CH3Cl (lower panel). The linear trend lines are fits to thenormal nucleophile data sets.

The product distributions and reaction exothermicities for the reactions of the water- and methanolsolvated nucleophiles with methyl bromide and chloride are summarized in Table 6.3 and Table 6.4.
Although the reactions with methyl bromide and chloride may only proceed through an S N2 reaction
mechanism, two products are possible, the un-solvated (6.4) and solvated (6.5) halide ion.
Nu–[CH3OH] + CH3X → NuCH3 + X– + CH3OH

6.4→ NuCH3 + X–[CH3OH]

6.5

116

The reactions following equation 6.5 are more exothermic, and they could reasonably be expected todominate the product distribution. However, in all but two reactions studied, formation of the bare halideion (equation 6.4) dominates the distribution. The remaining two reactions involving the water-solvatednucleophiles C2H5O–[H2O] and (CH3)2CHO–[H2O] generate roughly equal amounts of Cl– and Cl–[H2O].
Previous studies of SN2 reactions with micro-solvated nucleophiles list similar product distributions.37, 38
Many factors may contribute to the preferred formation of bare product ions including steric hinderance, 39
product dynamics,40 and redistribution of excess energy.41 The micro-solvation study reveals similar
product distributions for the normal- and α-nucleophiles suggesting that the enhanced reactivity observedwith the micro-solvated HOO– is not reflected in the product distributions.
6.5

The α-Effect and Competing Mechanisms with Methyl Formate
Methyl formate provides an interesting substrate to investigate nucleophilic-reactivity due tocompeting mechanisms. PT, SN2, and BAC2 mechanisms are in competition when a nucleophile reacts
with methyl formate, Scheme 6.1. Preference for SN2 or BAC2 reaction channels suggests an increased
nucleophilicity of the anionic reagent compared to PT channels.

Scheme 6.1

Previously studied methyl formate reactions reported similar branching fractions (S N2:PT:BAC2) forHO– (5%:61%:34%) and HOO– (8%:64%:28%).42 Due to the similar reactivity between the prototypical
α-nucleophile HOO– compared with HO– the authors concluded that the α-effect is not intrinsic in
117

origin.42 This conclusion was later questioned by Patterson and Fountain. Patterson and Fountain argue
that, because the PA of HOO– is ~60 kJ mol–1 smaller than the PA of HO–, a similarity in reactivity
supports an intrinsic α-effect.9 The reactions of HO– and HOO– with methyl formate is reinvestigated andexpanded by adding the normal-nucleophiles CH3O–, C2H5O–, and (CH3)2CHO–. (Actual reactivity isinvestigated with deuterium and 18O labeled compounds to differentiate reaction pathways.) The results
are tabulated in Table 6.5.
Table 6.5 The α-effect for reactions of un-solvated anions with methyl formate
SN2
–

–9 b

BR (Eff)

ΔH

BR(Eff)cΔH‡dα-effect0.08(0.05)

–39.3

1

0.59(0.47)–83

0.33(0.20)

–53

0.9

1.33±0.03

0.03(0.02)

–31.6

2.5

0.79(0.48)–64

0.18(0.11)

–48

1.5

1585±31.25±0.02

<0.01

–21.8

~13

0.97(0.62)–53

0.03(0.02)

–45

8.5

1576±41.20±0.02

—

–13.1

—

1.0(0.64)

–47

—

–44

—

kexp (×10 )BF(Eff)1633±21.73±0.03

1598±2

CH3CD2O(CH3)2CHO–

H18O–
–

CD3O

–

BAC2

α-effect

PA

Nu

PT
‡d

s

c

ΔH

c

‡d

–

HOO
1575±21.34±0.03
0.08(0.05)–35.7
0.64(0.39)–76
0.28(0.17)–55aExperimental proton affinity (PA, Ref. 27). bOverall rate constants in units of cm3 s–1, error bars reported as 1σ, absolute uncertainty is ±30%.
c
Branching fractions (BF), Efficiency (Eff) compares the relative rate constant for the branching fraction to the theoretical collision rate (kcol)
calculated using PTT (Ref. 19). dBarrier heights and reaction exothermicity are relative to reactants and calculated with G3MP2.
35

All of the methyl formate reactions are very efficient. The theoretical reaction barriers support theexperimental branching fractions, and the experimental results can be applied to understand intrinsic
differences in each of the reaction channels. The reaction mechanisms involving methyl formate werecomputationally evaluated by Pliego and Riveros.43 Their results suggest that: the PT channel occurs
through a single-well potential; the SN2 channel occurs through a double-well potential, and the BAC2
channel occurs through a multi-well potential where the controlling barrier is for the formation of thetetrahedral intermediate.43 These controlling barriers were theoretically computed with the G3MP2
method, and reported in Table 6.5. Both the SN2 and BAC2 reaction efficiencies correlate nicely with thetheoretical barrier heights. However, the PT channel efficiencies seem a bit low when considering thereaction exothermicity. Pliego and Riveros modeled the PT mechanism as a direct formation of themethoxide ion bound to the neutral molecules CO and NuH.43 This approach creates a proton abstraction
process that is highly exothermic, but models a reaction that is more difficult than direct proton118

abstraction. The PT process occurs on the formyl end of the molecule, while the SN2 and BAC2
mechanisms require the nucleophile to approach the methyl end of the ester. Nucleophilic attack at the
methyl end of the ester by HO– has been statistically modeled by RRKM theory, and the results match ourobserved branching fractions suggesting that the behavior of the SN2 and BAC2 mechanisms are statistical.
The reported product fractions (SN2:BAC2) from RRKM theory (15%:85%) are in agreement withexperimental data (20%:80%).43
Comparing the relative efficiencies for the nucleophilic SN2 and BAC2 reactions certainly favors the
HOO– anion as the proton affinity of the normal-nucleophiles decreases. It is unfortunate that the reactionof (CH3)2CHO– (the normal-nucleophile closest in proton affinity to HOO–) only occurs through the PT
reaction pathway, masking the nucleophilic reaction channels. This prohibits a report of an upper limit to
the nucleophilic reaction channels for (CH3)2CHO–, and therefore the magnitude of the α-effect cannot bereported. Nonetheless, a slight enhancement in the nucleophilicity for HOO– is observed in comparison to
CH3CD2O– and CD3O–. While the magnitude of the α-effect is greater for the BAC2 reactions in solution,5
this does not appear to be the case in gas-phase reactions with methyl formate, as the SN2 reactions at thesp3 carbon are the same or slightly higher than BAC2 reactions at the sp2 carbon.
The same reactions with methyl formate were investigated in the presence of a single watermolecule. The hydrated-anionic nucleophiles were introduced into the reaction flow tube for reactions
with methyl formate. The micro-solvated system exhibited a much larger α-effect than was previously
observed with the solvent free reactant system. Figure 6.7 presents a Brønsted energy-rate correlation
diagram comparing the reactivity of the un- and micro-solvated nucleophiles with methyl formate. The
plot uses two linear regressions to exhibit the rate-energy relationship of the two respective systems (unand micro-solvated). The plot unambiguously illustrates the enhanced reactivity of the micro-solvated αnucleophile HOO–[H2O] as it deviates from the rate-energy relationship followed by the normalnucleophiles. Even the addition of a single water molecule seems to mimic the solvation environment,
particularly the enhanced reactivity of HOO– observed in solution.
119

Figure 6.7 Brønsted plot: Micro-solvated and un-solvated ions with methyl formate
The reaction efficiencies (kexp/kcol) are compared to the anion PA for the un-solvated and micro-solvatednucleophiles with methyl formate. The anion proton affinities are calculated using the G3 method. Vertical error
bars represent uncertainty in the rate constants (±30%), while the horizontal error bars represent the uncertainty inthe experimentally determined PA. Blue and red circles represent normal-and α-nucleophiles, respectively.

A complete report of the experimental results including kinetic measurements, branching fraction,
relative reaction channel efficiency, and transition state barriers is provided in Table 6.6. The reactivity ofthe micro-solvated nucleophile compared with the un-solvated anion mostly differs in the BAC2 channel,
where the efficiency rapidly decreases in going from the smaller alkoxy anions to the larger anions. The
larger normal-nucleophile, (CH3)2CHO–[H2O], does not exhibit any BAC2 reactivity. This is not surprisingbecause if (CH3)2CHO–[H2O] forms a tetrahedral complex as required in the BAC2 mechanism, theintermediate will dissociate back to reactants as (CH3)2CHO–[H2O] is a better leaving group than CH3O–[H2O].

120

Table 6.6 The α-effect for reactions of water-solvated anions with methyl formate
SN2
Nu–[H2O]PAa

kexp(×10-10)b

H18O–[H218O]

1538

13.9±0.3

CD3O–[H2O]

1527

CH3CD2O–[H2O](CH3)2CHO–[H2O]

PT

BAC2

ΔH‡d

BF(Eff)c

ΔH‡d

BF(Eff)cΔH‡dα-effect—–20.50(0.32)–51

0.50(0.32)

–48

1.4

6.4±0.2

0.40(0.13)–30.20(0.07)–47

0.40(0.13)

–44

3.5

1517

4.0±0.1

0.65(0.14)

30.25(0.06)–41

0.10(0.02)

–43

23

1505

1.8±0.1

0.60(0.06)

80.40(0.04)

–38

—

–40

BF(Eff)c

HOO–[H2O]
1493 8.7±0.1
—
1
—
–38
1.0(0.45)–45aExperimental proton affinity (PA) in units of kJ mol–1 calculated using G3. bOverall rate constants in units of cm3 s–1, error bars
reported as 1σ, absolute uncertainty is ±30%. cBranching fractions (BF), Efficiency (Eff) compares the relative rate constant for the
branching fraction to the theoretical collision rate (kcol) calculated using PTT (Ref. 19). dBarrier heights and reaction exothermicity
are relative to reactants and calculated using G3.

It is striking that, despite the high enthalpic SN2 transition state barriers presented in Table 6.6, theSN2 product channel can compete with the PT and BAC2 channels. Figure 6.8 illustrates the lowest energy
transition state structures. Examination of the transition states reveals that the PT and BAC2 reaction
channels proceed through highly ordered intermediate structures. Whereas the S N2 transition state
structures for the un-solvated reaction system require the most ordered structures, the micro-solvated SN2transition state structures are the least ordered, and therefore the PT and B AC2 channels must surmounthigh entropic barriers in the micro-solvated systems.

121

Figure 6.8 Lowest energy transition states for micro-solvated ions with methyl formate
Optimized at the MP2/6-311++G(d,p) level, the transition state geometries are illustrated for the PT,
SN2, and BAC2 reaction pathways for HOO–(H2O) and CH3O–(H2O).

Previous studies investigating substitution reactions at saturated carbon centers report that
incremental micro-solvation rapidly decreases the efficiency of reaction.23, 37, 38, 44-46 While the study
regarding SN2 reactions of micro-solvated ions with methyl chloride illustrated a decrease in the reaction
efficiency by a factor of 100 (section 6.4),23 the efficiency for reaction between the micro-solvated anions
and methyl formate is only decreased by a factor of at most six. Although it is not favored entropically,
the interaction of the water molecule with the carbonyl group for the PT and BAC2 transition states (Figure6.8) may stabilize and increase the reaction efficiency. The stabilization effect of the water molecule on
the PT and BAC2 channels is more obvious when Table 6.5 and Table 6.6 are compared. The transition
states of the PT and BAC2 channels are comparable between the un-solvated and micro-solvated systems,
while the enthalpic SN2 transition state barriers are much higher within the micro-solvated system.
Similar to the study with the un-solvated ions, the micro-solvated system illustrates an α-effect within
the BAC2 reaction pathway. The reaction of HOO–[H2O] with methyl formate only proceeds through the
BAC2 mechanism in contrast to the un-solvated system, where the major product of HOO– is PT. The
122

magnitude of the α-effect is largest when comparing the BAC2 reactions of HOO–[H2O] with CH3CD2O–[H2O], on the order of 23. This is similar to the largest α-effect for the un-solvated system, where the αeffect for BAC2 reaction efficiency between HOO– and CH3CD2O– is on the order of 8.5.
6.6

Conclusion
Observation of the α-effect in the gas-phase was possible thanks to decades of context provided bydiscussion and experiments. The rate-energy relationship portrayed by Brønsted correlations provide the“gold standard” for reporting an α-effect. This relationship was not considered in early investigations of
the α-effect in our lab, and resulted in a false conclusion that the α-effect is not intrinsic in origin.42
Furthermore, the application of Marcus theory has provided a series of guidelines for designing
experiments that investigate the intrinsic origin of the α-effect. Following these guidelines inexperimental design has assisted the discovery of an intrinsic component of the α-effect, although pure
intrinsic contributions are minor.
Furthermore, the micro-solvation studies have revealed that the large magnitudes of the α-effectreported in solution are indeed the result of differential solvation of the normal- vs. α-nucleophile. Whilea small intrinsic origin of the α-effect has been unveiled, the magnitude of the α-effect is clearly
influenced by solvation environments, and the influence of the solvent must be considered for a complete
understanding of the α-effect. Even in the presence of a single water or methanol molecule, the magnitudeof the α-effect is dramatically increased, supporting both intrinsic and solvent contributions to the αeffect.
In conclusion, the α-effect is a multifaceted, complex phenomenon where several factors—both
intrinsic and environmental—contribute to its magnitude. While these studies provide more context to theorigin of the α-effect, a complete understanding of the α-effect will likely continue to challenge future
researchers.

123

6.7

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127

7 Appendix A

Supplimental material for construction of the ion funnel.

Figure A1 Schematic for Ion-Funnel Balun
Designed by R. M. Moision and R. Jones. Revised 2007.
A) Balun consists of 2 stacked FT-290-77 toroid (www.amidon-inductive.com) wrapped as a trifilar with
18 gauge magnet wire. Previous versions used 14 gauge wire. 18 gauge has proven to be sufficient.
(www.amidon-inductive.com/aai_cost_toroidal.htm)
The FT-240-77, FT-290-J, or FT-240-J would also work in this circuit.
B) Two 100 Ω (carbon) resistors in parallel, resulting in a total of 50 Ω. These must not be wire wound.
They should be non-inductive, and should be able to dissipate the amount of power that the RF
amplifier produces. Surplus Sales of Nebraska has large carbon resistors (Type C). Our current setup
uses four 50 Ω, 100 Watt resistors, but these are more difficult to obtain. If these large carbon
resistors are used, then the appropriate resistor clips are needed for mounting. Over the frequency
range typically used with the ion funnel, the massively parallel 50 Ω resistor described in the original
version of this document works just fine.
C) RF Balancing posts. Placing capacitors from one post to ground can adjust the peak-to-peak voltageof the output. This can be used to equalize the peak-to-peak voltage of the two output channels. We
have these on all designs, but rarely find them useful. In later designs, the 10:1 monitoring posts havealso been removed.
In the circuit it is important to keep all wires equal in length within the balun as well as for each output
channel. This includes the coax cable going from the balun to the chamber and any wiring within thechamber as well.
To build the trifilar, take 3 lengths of magnet wire, ~10-15 feet in length each. Place one end of each into
the chuck of an electric drill. Stretch the other sides across the room and hold the strands together. Spinthe 3 strands together so that they effectively form a single strand. Wrap the coiled strand around thestacked toroid. It helps to tape the two toroids together before wrapping with the magnet wire.
128

Figure A2 Lens elements for ion funnel
Material: 0.020” (±0.001”) hard brass or stainless steel. All dimensions ±0.005” unless stated. Scale 1:1

129

Figure A3 Insulator/Spacer elements for ion funnel
Material: 0.020” (±0.001”) Teflon. All dimensions ±0.005” unless stated. Scale 1:1

130

8 Appendix B

Cartesian Coordinates (Å) and Energies (Hartree) for optimized geometries for chemical species in
Chapter 4
Figure 4.1
1A, 1B, Dicyanamide, NCNCN─Atom X
Y
Z
N
0.00000000
0.00000000
0.68821500
C0.00000000
1.15135100
0.05646900
N0.00000000
2.24305700-0.39250900
C0.00000000-1.15135100 0.05646900
N0.00000000-2.24305700 -0.39250900
ZPE (Zero Point Energy) Corrected Energy: -240.516411
1A, 1B, Nitric Acid, HNO3
Atom X
YN-0.14930900 0.03441500
O-1.03699000 -0.77287500
O1.14202700-0.53580400
O-0.19034500 1.24816800
H
1.72763000
0.24318700
ZPE Corrected Energy: -280.913951

Z-0.00023500
0.00013300
0.00002300
0.00004200
0.00005800

2A
Atom X
Y
N
2.91132400-0.00845300
O4.04329600-0.47278800
O
1.91729000-0.91304000
O2.61888800
1.18098600
H
0.96907700-0.42446100
N-2.80699200 0.72955700
C-3.87787400 -0.03807200
N-4.90160800 -0.61700100
C-1.59651200 0.27325500
N-0.45394500 -0.01094600
ZPE Corrected Energy: -521.470462

Z-0.00117900-0.00329200
0.00184400-0.00161500
0.00358100-0.00071000-0.00208200-0.00368000
0.00338400
0.00744200

2A/3A
Atom XN
1.55618000
O
2.66158700
O1.28764500
O0.65599800
H-0.41825800
N-1.84713100

Y-0.38275000-0.79635100-0.23147000-0.10537600
2.89077500
0.51939600

Z0.01460000-0.37730200
1.22398900-0.86877700
0.75823000-0.14181100
131C-2.01031100 -0.78268800
N-2.33586700 -1.90094500
C-0.79528000 1.21530600
N-0.17175900 2.27559700
ZPE Corrected Energy: -521.424903

0.00836200
0.12317400-0.08562500-0.01281100

3A
Atom XYN
1.63350600
0.56815100
C0.45086200
1.08712500
N-0.00097100 2.28852000
C
1.91730600-0.70878200
N2.29435700-1.82311100
H0.73459500
2.87762000
O-0.50834900 0.16204600
N-1.41748600 -0.42810100
O-1.30422400 -0.22938400
O-2.25110700 -1.10589700
ZPE Corrected Energy: -521.439271

Z-0.16865100
0.15600500
0.10283600-0.03657000
0.03401300-0.29315700
0.79629300-0.06962100-1.26083800
0.50035800

3A/4A
Atom X
Y
N-1.65616300 0.82860700
C-0.36523100 0.52550700
N0.63993800
1.40719500
C-2.57937600 -0.11453000
N-3.50331800 -0.84317600
H
0.46400000
2.24400500
O0.08657600-0.70543500
N1.76538600-0.26586500
O2.07391900-0.54649200
O
2.39984800-0.32272100
ZPE Corrected Energy: -521.415658

Z
0.07158500-0.06086500-0.03279800
0.01554600-0.01025500
0.52290900-0.26196900
0.03186500
1.16396500-0.98621700

4A
Atom X
YN
1.48237400-0.70231800
C0.50391200
0.22277200
N-0.74771200 -0.53582200
C
2.73739200-0.30856500
N3.88964400-0.09178800
O0.49869200
1.43470700
H-0.67769900 -1.53690900
N-2.01690300 -0.10507900
O-2.84432400 -1.00954500
O-2.28211200 1.08692800
ZPE Corrected Energy: -521.491439

Z-0.10019100-0.05963400-0.24641800
0.02034600
0.11593200
0.08351400-0.11043400-0.00252700
0.22534800-0.06153700

2B
Atom XN
5.63194400

Y-0.21163700

Z
0.00004900
132

O
6.70997800-0.78229300
O4.55172200-1.02672500
O5.44387000
0.99642800
H3.68532500-0.45052200
N0.00015600
1.16974600
C-1.15045500 0.56395400
N-2.24403100 0.13751700
C
1.15045800
0.56336400
N
2.24380300
0.13633900
H-3.68522200 -0.44997200
O-4.55143700 -1.02648700
N-5.63194300 -0.21179900
O-6.70978100 -0.78283200
O-5.44430700 0.99633700
ZPE Corrected Energy: -802.416924
2B/3B
Atom X
YN-2.79351200 1.26202100
O-3.26246200 2.36037700
O-2.66628600 0.91559900
O-2.41956000 0.44209300
H-3.94303400 -2.24948500
N-1.15840000 -1.78113200
C-0.14010700 -0.95609800
N0.87471900-0.38068200
C-2.39706000 -1.50766200
N-3.58030600 -1.81086800
H
2.31936100
0.32579200
O3.15139900
0.92736300
N
4.25349700
0.14611700
O
5.31289300
0.74686600
O4.09335200-1.06004100
ZPE Corrected Energy: -802.368530
3B
Atom
N
C
N
C
N
H
O
N
O
O
H
O
N
O

X
-1.25091600-2.46084000-3.61515000-0.15588800
0.90400100-3.50773300-2.47224900-2.76567300-2.84604600-2.89547700
2.26084700
3.11041700
4.23348700
5.27777800

Y-1.62712900-1.14340500-1.69140100-0.93940800-0.43131900-2.64000100
0.19837200
1.22460300
0.94781800
2.30419100
0.21994300
0.84403700
0.12152500
0.75149000-0.00074700-0.00247000
0.00285300-0.00145500
0.00049000
0.00071600
0.00100000-0.00011200-0.00068800
0.00148300
0.00245600-0.00036400
0.00074800-0.00372300

Z
0.04278000-0.28931300
1.23721300-0.88500100
0.83816400-0.11343000-0.09779500-0.08725100-0.07107200-0.00088400-0.07731100-0.09589500
0.03480000
0.02385200
0.14917500

Z0.11572800-0.20831000-0.17884500
0.02604100-0.01953300
0.18622700-0.79012200
0.11788300
1.29407800-0.41703800-0.07766000-0.16269700
0.00996400-0.07166900
133

O
4.12645400-1.07803600
ZPE Corrected Energy: -802.390688

0.23103300

3B/4B
Atom X
Y
N
0.85013900
1.32885900
C
2.03267600
0.68833400
N
3.21327700
1.29594300
C-0.27208400 0.67006500
N-1.35405500 0.21159300
H
3.25415000
2.22730200
O
2.14356100-0.61858300
N
3.93944900-0.57458500
O4.18994200-0.84315800
O
4.49673300-0.84079500
H-2.73168800 -0.38404100
O-3.60962100 -0.97594600
N-4.69969100 -0.19708300
O-5.77027800 -0.78684100
O-4.54156700 1.00948000
ZPE Corrected Energy: -802.367400

Z-0.02572000-0.08840100-0.09117300-0.04594100-0.05222900
0.31941900-0.17664500
0.06856300
1.21077900-0.95504000-0.06761300-0.10324000
0.02667200-0.00382000
0.16189800

4B
Atom X
Y
N
0.84668400-0.62854900
C
1.87786800
0.25303700
N
3.08641800-0.55340600
C-0.37639600 -0.18441800
N-1.51792600 0.08547300
O
1.91112800
1.46323300
H
2.98395200-1.55335100
N
4.37251000-0.15798600
O
5.16423500-1.08829600
O4.66675200
1.02485800
H-2.99495200 0.43490400
O-3.95571300 0.86406700
N-4.89190000 -0.10578500
O-6.04831000 0.28623700
O-4.52163500 -1.27153500
ZPE Corrected Energy: -802.442001

Z-0.15009800-0.08598900-0.20709800-0.08953100-0.04487900
0.03614500-0.08507200
0.04208100
0.27020900-0.00998500
0.00674700
0.04557400
0.04080400
0.08304700-0.00426900

5B
Atom
N
C
N
C
N
O
H
N

X
-0.76839300-1.74401900-1.23578700
0.29625500
1.29985500-2.88432900-0.25522200-2.03606900

Y
2.02931900
1.04633200-0.20471500
1.98773700
2.09450600
1.20855500-0.46586100-1.31964900

Z-0.43877600-0.39111500
0.09170100
0.23175900
0.87860200-0.76059400-0.08303000
0.29200400
134

O-1.45675300 -2.40178600
O-3.20818800 -1.15607600
H
2.13393700
1.38466300
O
3.58337700-1.50396900
N
2.78536000-0.61623200
O3.18226300
0.35064000
O
1.59526900-0.62309000
ZPE Corrected Energy: -802.436330

0.19885100
0.59170000
0.74978000-0.50026200-0.19345600
0.56027600-0.60511300

5B/6B
Atom X
Y
N-0.07544400 1.27578000
C-0.96266200 0.76628300
N-2.02175200 0.14900500
C
1.16129800
1.43114600
N2.25268100
1.99845200
O-0.96039700 0.83871100
H-1.76907000 -0.24714700
N-3.11329900 -0.50125900
O-3.60661800 -1.37012800
O-3.53710300 -0.15333200
H
2.20248300
3.00670700
O1.65637800-0.57493600
N2.73406900-1.00076000
O3.16532300-0.39020100
O3.27504100-2.02419600
ZPE Corrected Energy: -802.389883

Z-0.75040500
0.17111600-0.58952200-0.51668600-0.51718400
1.37687100-1.48802500-0.05446600-0.78399300
1.03371800-0.37768800-0.44090200
0.15510500
1.14244600-0.29883600

6B
Atom X
YN
0.10234000
0.96567500
C0.93826400-0.09007200
N
2.28667000
0.50784500
C-1.24071500 0.87101700
N-2.10807600 1.81886500
O0.80512500-1.29604200
H
2.34394900
1.50103500
N
3.48311000-0.09522800
O4.40876400
0.68074800
O3.59685900-1.30595600
H-1.60447400 2.70387700
O-1.75561600 -0.45855500
N-3.19534700 -0.48647000
O-3.74431500 -0.56374300
O-3.67402400 -0.53962500
ZPE Corrected Energy: -802.405451

Z-0.10904900-0.08701300-0.24756100-0.03636100-0.00740600
0.02798100-0.05796800
0.01260100
0.31995700-0.11028600-0.04282500
0.00156000
0.04839500-1.02316700
1.15422700

6B/7B
Atom XN
0.09708300
C
1.03067900

Y
0.79699200-0.18388400

Z0.01051600
0.01149800
135N
2.30316600
0.51637500
C-1.20967900 0.47665200
N-2.16891800 1.39458200
O
1.00039700-1.38688900
H
2.28109600
1.51989500
N3.56297000
0.03298800
O4.43273100
0.90203300
O3.77291600-1.16908000
H-1.95341000 2.32592600
O-1.70476700 -0.73980600
N-3.39958100 -0.15009700
O-3.75664800 -0.23144200
O-3.99671900 -0.34210500
ZPE Corrected Energy: -802.382534
7B
Atom X
Y
N-0.00004300 -0.22366700
C-1.16917200 0.46200400
N-2.20554200 -0.58414400
C
1.16912600
0.46191700
N
2.20550200-0.58421400
O-1.47112800 1.63499000
H-1.91791900 -1.51955500
N-3.54921000 -0.42905000
O-4.14948300 -1.45869700
O-4.07546000 0.65161500
H
1.91802100-1.51961500
O
1.47116300
1.63488900
N
3.54922300-0.42900300
O
4.14953200-1.45856900
O4.07545800
0.65154700
ZPE Corrected Energy: -802.455366-0.26224900
0.09955100
0.20843200
0.18681900-0.12725000-0.04763500
0.15109900-0.10535000-0.14512100
0.09001500-0.05823100-1.20554000
0.96423900

Z
0.00021800-0.05395500-0.22267300
0.05390500
0.22249100
0.01512200
0.03882800-0.01468000
0.34496800-0.23065900-0.03918400-0.01550800
0.01475900-0.34515600
0.23121400

Figure 4.2
A, Dicyanamide, NCNCN─Atom X
Y
N
0.00000000
0.00000000
C
0.00000000
1.13336200
N0.00000000
2.21606800
C0.00000000-1.13336200
N0.00000000-2.21606800
ZPE Corrected Energy: -239.913961

Z
0.74478100
0.06723500-0.43002100
0.06723500-0.43002100

B, Protonated Dicyanamide, HNCNCN
Atom X
Y
Z
N-0.05882300 0.65888400
0.01237200
C
1.04488700
0.09161800
0.03215500
N
2.18208800-0.30852700 -0.12907100
C-1.25192100 0.05069700
0.00247700
136

N-2.34582500 -0.39000700
H
2.80012500-0.57634700
ZPE Corrected Energy: -240.403638
C, Protonated Dicyanamide, HN(CN)2
Atom X
Y
N-0.04067400 0.72008800
C-0.00047100 0.04615800
N-0.00047100 -0.51923900
C-0.00047100 0.04615800
N-0.00047100 -0.51923900
H
0.29695800
1.67483600
ZPE Corrected Energy: -240.401951-0.00244500
0.62621800

Z
0.00000000
1.17508400
2.20601800-1.17508400-2.20601800
0.00000000

Figure 4.3
A, NCNCN─[HNO3]
Atom X
YN-2.89093300 -0.00620400
O-4.03404300 -0.43617300
O-1.92987000 -0.94897400
O-2.55047400 1.16899000
H-1.00298000 -0.47765000
N2.80700700
0.79202400
C
3.83851500-0.03743300
N4.83623500-0.67998700
C
1.60156300
0.30097300
N0.45877700-0.01645300
ZPE Corrected Energy: -520.253165

Z-0.02498000-0.06566600
0.05456200-0.04876900
0.08066100
0.01822200-0.04062900-0.09548400
0.07156900
0.13262600

B, NCNCN─[HNO3]2
Atom X
Y
N
0.71638400
1.31037500
O1.72179900
1.90444600
O-0.28716400 1.96510800
O
0.63360600
0.04719500
H-1.36336300 1.13937100
N-4.01124000 -1.09074700
C-5.20899300 -0.80382400
N-6.30125700 -0.66398600
C-3.13864700 -0.17809100
N-2.28141200 0.58313500
N3.78734600-0.84817400
O3.42032800-0.18161000
O
2.87034300-1.13433100
O4.89704100-1.30966000
H
2.03283500-0.61128400
ZPE Corrected Energy: -800.578177

Z-0.28430800-0.67210900
0.16018400-0.30772800
0.51360800
0.32360800-0.18095700-0.61560000
0.50842800
0.82043400
0.14897400
1.10348900-0.81457700-0.05115000-0.57506100

C, NCNCN─[HNO3]3
Atom XN
0.49955700

Y-0.03549200

Z-0.04577400
137O-0.45034800 0.55747200
O0.33808400-1.20657300
O
1.60743400
0.53506600
H-1.14694700 -1.72110000
N-4.37092900 -1.58158300
C-5.37286500 -1.15919600
N-6.33281700 -0.79914900
C-3.20628500 -1.78210700
N-2.11261400 -2.14638800
N
3.63362600-2.10439400
O
2.98321200-2.01309900
O
3.49401500-1.11107100
O4.41850500-2.97999500
H
2.74747900-0.52843000
N0.39014900
3.71246400
O
0.80660200
4.84934100
O
0.13077800
3.07797800
O0.17137900
3.09768900
H-0.07172900 2.12467700
ZPE Corrected Energy: -1080.904003-0.61611800
0.38827300
0.09189000
0.20587900-0.65139500
0.12076500
0.70943400-0.20193600
0.11324400-0.04686800-1.07446900
0.88082200
0.26526700
0.54313500
0.08783800-0.03039300-1.09508500
1.11773400-0.85070300

138

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