Core Ion Structures in CO2- and N2O-Based Cluster Anions Studied by Infrared
Photodissociation SpectroscopybyMichael Charles Thompson
B.S., Baylor University, 2013

A thesis submitted to the
Faculty of the Graduate School of the
University of Colorado in partial fulfillmentof the requirement for the degree ofDoctor of Philosophy
Department of Chemistry and Biochemistry2018This thesis entitled:
Core Ion Structures in CO2- and N2O-Based Cluster Anions Studied by Infrared
Photodissociation Spectroscopy
has been approved for the Department of Chemistry and Biochemistry

Professor J. Mathias Weber

Professor Veronica M. Bierbaum

Date

The final copy of this thesis has been examined by the signatories, and wefind that both the content and the form meet acceptable presentation standardsof scholarly work in the above mentioned discipline.

iii

Thompson, Michael Charles (Ph.D., Chemistry)
Core Ion Structures in CO2 and N2O based Cluster Anions Studied by Infrared Photodissociation
Spectroscopy
Thesis Directed by Prof. J. Mathias Weber

The interactions of anionic species with other molecules opens up new avenues to study
the process of bond formation or charge transfer. Gas phase clusters are a useful tool for studying
such systems in the absence of competing effects in condensed phase. Cluster ions can be studied
as model analogues to more complex condensed phase systems. The majority of this thesis focuseson gas phase cluster ions of the form [M(CO2)n] (M=Bi, Sn, Mn, Fe). These clusters consist of acharged molecular anion surrounded by weakly bound “solvent” CO2 species. IR
photodissociation spectroscopy is used to probe the infrared spectra of the molecular core ions.
[Bi(CO2)n]- and [Sn(CO2)n]- clusters are studied as model systems of the reduction of CO2
at a corner or edge site of a Bi or Sn electrode surface. The structures of the core ions for theseclusters give insight into potential docking motifs of the CO2 species. In both species, the formationof an ɳ1-C docked CO2 species results in a metal carboxylate complex. These structures havecharacteristic CO2 stretching frequencies that can be used to identify these species in the condensed
phase. These species also exhibit oxalate ligand formation, which is interesting since it requires
the formation of a C-C bond.
[Mn(CO2)n]- and [Fe(CO2)n]- clusters are studied to understand how a change in theelectron configuration of the metal affects the binding motifs of the CO2. As in other first row
transition metals, these species exhibit a rich collection of interaction motifs, among which the ɳ2

iv(C,O), where M-C and M-O bonds are formed, is most prevalent. In addition, oxalate ligands also
are present.
We also investigated the structure and charge distributions of neat and heterogeneous N2O
clusters [(N2O)n- and (N2O)nO-]. In heterogeneous N2O clusters, we find that initially the core ion
is an NNO2- molecular anion. We find that this core ion switches to an O- core ion at larger cluster
sizes. We also find an N2O- core ion for neat anionic clusters of N2O and report for the first time
frequencies of the N-N and N-O stretches of the N2O- anion.

v

Dedication

To my family and my teachers for always encouraging me to explore my curiosity.

vi

Acknowledgements

As I conclude this chapter of my life and move onto the next, it is appropriate to look backand acknowledge all those who have been instrumental throughout this journey. Any list of people
must begin with my advisor Prof. J. Mathias Weber. I will be eternally grateful for his thoughtful
counsel and excellent example. His constant jokes and puns always served to lighten the mood on
many occasions and made my graduate experience all the more rich and enjoyable. Mathias has
not only been a good mentor, but he has been an example of patience and thoughtfulness that I
will strive to emulate.
Throughout my time in the Weber group, I have had the opportunity to overlap with anumber of very talented group members. Initially, I overlapped with Dr. Benjamin Knurr and Dr.
Shuang Xu. Our conversations about science and philosophy helped grow my own critical thinking
and laboratory skills. I am immensely grateful especially to Ben, who showed me how to turn the
instrument on and how to make ions. His patience and guidance are one of the reasons I have beenas productive as I was during my graduate career. I would like to thank Curtis Beimborn for his
friendship and willingness to discuss anything. Recently, I have had the privilege of working with
Dr. Leah Dodson. Leah’s work ethic and curiosity are inspiring. I thank her for challenging me tothink beyond what I am doing in the moment, and constantly pushing me to be better. I have hadless time to interact with Wyatt Zagorec-Marks but, I have no doubt that he will be extremely
successful.

vii

I have been fortunate enough during my time to work at JILA and participate in the
Molecular and Ionic Clusters Supergroup consisting of the groups of Profs. Carl Lineberger,
Ronnie Bierbaum, and Barney Ellison. The talks and discussions afterwards have been insightful
and enlightening. Specifically, I thank Allan Maple and Dan Nelson for being willing to chat aboutproblems from time to time. I would also like to thank the members of the JILA Machine Shop,
the Computing Group, and Electronics Shop. Their knowledge and experience has been
indispensable, whether showing me how to use a lathe, helping to restart our calculation computers(for the thousandth time), or diagnosing broken power supplies. Their eagerness to teach anddiscuss problems have helped mold me into a better scientist.
I am privileged to have been supported by several funding sources during my graduate
career. I am thankful to the NSF JILA Physics Frontier Center (PHY-1125844 and PHY-1734006)
and the US Department of Education GAANN program for supporting my research. I also must
thank the Department of Chemistry and Biochemistry for support through Teaching
Assistantships.
Lastly, I would like to thank my friends and family. Their support has helped make mygraduate career much easier than it could have been. Specifically, I would like to thank themembers of the Magic Club for helping me escape the lab every weekend. To my family, I thankyou for your quiet support and willingness to listen. While it has been hard being so far away, yoursupport has meant the world to me.

viiiContents
1

Introduction and Motivation ............................................................................................1

2

Experimental and Theoretical Methods ..........................................................................6
2.1 Overview ..............................................................................................................................6
2.2 Experimental Methods .........................................................................................................6
2.2.1

Ion Generation .........................................................................................................6

2.2.1.1

Laser Vaporization Ion Source ....................................................................8

2.2.1.2

Entrainment of Volatile Species ................................................................10

2.2.2

Mass Spectrometry and Photodissociation Spectroscopy ......................................11

2.2.3

Light Source ...........................................................................................................15

2.2.4

Signal Acquisition and Processing.........................................................................17

2.2 Computational Methods .....................................................................................................18

3

Interactions of CO2 with Bismuth: Formation of Metalloformate andMetallooxalate Complexes...............................................................................................20
3.1 Introduction ........................................................................................................................20
3.2 Results and Discussion ......................................................................................................21
3.3 Conclusions ........................................................................................................................33

ix

4Interaction of CO2 with a Sn Atom in the Presence of an Excess Electron:
Formation of a Weak Sn‒C bond ...................................................................................34
4.1 Introduction ........................................................................................................................34
4.2 Results and Discussion ......................................................................................................35
4.3 Conclusions ........................................................................................................................51

5

Interactions of CO2 with Fe and Mn: Formation of Complex and DiverseCore Ions ...........................................................................................................................52
5.1 Introduction ........................................................................................................................52
5.2 [Fe(CO2)n]‒ Results and Discussion ..................................................................................53
5.3 [Mn(CO2)n]‒ Results and Discussion ................................................................................64
5.4 Conclusions .......................................................................................................................83

6

Enhancement of the Intensity of the Symmetric CO2 Stretching Transition in
MCO2‒ Complexes (M = H, Ag, and Bi) Compared to CO2‒ .......................................85
6.1 Introduction ........................................................................................................................85
6.2 Theoretical Methods .........................................................................................................86
6.3 Results and Discussion .....................................................................................................87
6.4 Conclusions .......................................................................................................................95

7

Core Ion Geometries in (N2O)nO‒ and (N2O)n‒ .............................................................96x

7.1 Introduction ........................................................................................................................96
7.2 Results and Discussion of (N2O)nO‒ Clusters ...................................................................99
7.3 Core ion structure determination of (N2O)n‒....................................................................106
7.4 Conclusions .....................................................................................................................110

8Bibliography ...................................................................................................................112

xi

List of Figures
Figure 1.1

Walsh Diagram of CO2 (Left).11 Reproduced from Ref. 12 with permission
from Taylor and Francis Group. The HOMOs of CO2 (lower) and CO2‒ (upper)
based on HF-SCF calculations are shown on the right. ...........................................2

Figure 2.1

Schematic of the ion source. In the laser vaporization method, a laser
propagates from right to left and is focused onto the edge of a metal target(see text). In the entrainment of volatile gases experiment, the target gas is
introduced with a general valve and entrained into the main expansion(see text), additionally ions can be produced using only the main expansion. ........9

Figure 2.2

Diagram of the experimental apparatus with paths outlined of thevaporization laser (blue), ion path (green), electron beam (red), and laser(purple). ..................................................................................................................11

Figure 2.3

Mass spectrum obtained via the laser vaporization ion source with bismuth
target and a pulsed CO2 expansion. .......................................................................13

Figure 2.4

Schematic of the Optical Parametric Oscillator/Amplifier system(/2, half-wave plate; OC, output coupler; BD, beam dump; Si-Pol., siliconpolarizer; Pol., polarizer). ......................................................................................15

Figure 3.1

IR photodissociation action spectra of [Bi(CO2)n]‒ (n=2‒9). Labels denotethe core ion geometry (F=metalloformate, O=oxalate) and the subscript
letters denote the type of oscillation (S=symmetric, A=asymmetric, M=metalbound, F=free). The dashed lines are meant to draw the eye to experimental
signatures of the metalloformate complex (magenta) and metallooxalate
complex (green). ....................................................................................................22

Figure 3.2

Calculated core ion geometries (Bi, blue; C, black; O, red). .................................23

Figure 3.3[Bi(CO2)2]‒ experimental spectrum and calculated spectra. The colors of the
traces indicate the spin of the structure (red, triplet; blue, singlet), and the coreion whose vibrational signatures are shown are pictured above each trace. The
top two structures have solvent species behind the ligand. ..................................26

xii

Figure 3.4

Calculated HOMOs (70% contour) of several MCOO‒ species [M=Ag (silver),Au (yellow), and Bi (blue)]. ...................................................................................27

Figure 3.5

Positions of the signatures in the F series as a function of cluster size for theasymmetric (red, top left) and symmetric (black, bottom left) as well as thedifference (blue, top right) and the average between FS and FA (magenta, bottom
right). ......................................................................................................................29

Figure 3.6[Bi(CO2)6]‒ experimental spectrum and calculated spectra (black, top). The colorsof the lower traces indicate the spin of the structure (red, triplet; blue, singlet).
Calculated vibrational frequencies of a metallooxalate based [Bi(CO2)6]‒ cluster(A) , a metalloformate based [Bi(CO2)6]‒ cluster (B), and a butterfly based[Bi(CO2)2]‒ cluster (C). ..........................................................................................30

Figure 3.7

Comparison of the relative energies of the triplet formate and triplet oxalate coreion clusters as a function of cluster size.................................................................32

Figure 4.1

Infrared photodissociation spectra of [Sn(CO2)n]‒ (n = 2‒6). Cluster size isdenoted by the numbers between the two panels. The red dashed line gives the
position of the asymmetric stretching mode of free CO2. Each trace is individuallynormalized to the highest intensity feature. ...........................................................36

Figure 4.2

Calculated cluster geometries of [Sn(CO2)2]‒. Calculations were performed basedon the methods described in Chapter 2. Clusters are labeled based on core ion
geometry (capital letter), number of solvent CO2 species (number) and solvent
configuration (lower case letter) according to increasing relative energy. ............39

Figure 4.3

A comparison of calculated vibrational spectra (red) with the experimental
spectrum of [Sn(CO2)2]‒. The intensity of the solvent signatures is approximately
6‒10 times weaker than the feature at 1850 cm-1. ...............................................40

Figure 4.4

Calculated potential energy curve connecting structures D(0) and C(1b) along the
Sn‒C bond distance. All energies are relative to the lowest energy calculated data
point and are corrected for zero point motion except for the Sn‒C stretchingmode.......................................................................................................................43

xiii

Figure 4.5

Calculated HOMO contours for different positions along the potential energy
curve described in Figure 4.4. Contours are 70% of maximum orbital
amplitude. ..............................................................................................................44

Figure 4.6

Calculated cluster structures for [Sn(CO2)3]‒. The cluster labeling procedure is
described in Figure 4.2, and calculations were performed as described in Chapter
2..............................................................................................................................46

Figure 4.7

Comparison between the experimental spectrum of [Sn(CO2)3]‒ (black) andcalculated vibrational spectra (red). Calculated traces are labeled by cluster
structure. The region from 1816‒1975 cm-1 is enhanced by a factor of 5 tohighlight two low intensity bands in that region. ...................................................47

Figure 5.1

Experimental spectra of [Fe(CO2)n]‒ cluster ions. The numbers along the righthand side indicate the number of CO2 species in the cluster (n). Spectra in theasymmetric stretching region of neutral CO2 were taken over a single day and areshown to demonstrate the presence of at least one solvent CO2 species. ..............54

Figure 5.2

Calculated structures of [Fe(CO2)3]‒ (left) and [Fe(CO2)4]‒ (right). Calculations
were performed as described in Chapter 2. Clusters are labeled based on core ion
structure (capital letter), number of solvent species (number in parentheses) andsolvent position (lower case letter in parentheses). All calculated structures are in
the quartet spin state...............................................................................................56

Figure 5.3

Comparison of the experimental spectrum of [Fe(CO2)3]‒ (black, top) to
calculated IR spectra (red). Each trace is labeled according to the corresponding
calculated structure (see Figure 5.2). .....................................................................58

Figure 5.4

Comparison of the experimental spectrum of [Fe(CO2)4]‒ (black, top) to
calculated IR spectra (red). Each trace is labeled according to the corresponding
calculated structure (see Figure 5.2). .....................................................................61

Figure 5.5

A composite of all calculated infrared spectra for [Fe(CO2)4]‒ clusters (red)compared to the experimental spectrum (black). ...................................................63

Figure 5.6

Experimental spectra of [Mn(CO2)n]‒ cluster ions. The numbers along the righthand side indicate the number of CO2 species in the cluster (n). ........................65

xiv

Figure 5.7

Comparison of the experimental spectrum of [Mn(CO2)2]‒ (top, black) to thecalculated IR spectra of several isomers (red). Isomers are labeled based on thesame method described in Figure 5.2. Energies are relative to the lowest energy
calculated isomer. ..................................................................................................66

Figure 5.8

Potential energy curve along the C-Mn-O bond angle. The Born-Oppenheimer(black), vibrationally adiabatic ground state (red), and the vibrationally adiabatic
excited state (blue) surfaces are shown, along with relevant structural figures.....68

Figure 5.9

Calculated structures of [Mn(CO2)3]-. Calculations were performed as describedin Chapter 2. All calculated structures are in the quintet spin state. ......................70

Figure 5.10

Comparison of the experimental spectrum of [Mn(CO2)3]‒ (top, black) to thecalculated IR spectra of several isomers (red). Isomers are labeled based on thesame method described in Figure 5.2. Energies are relative to the lowest energy
calculated isomer. ..................................................................................................71

Figure 5.11

A comparison of the calculated IR spectra of different spin states (superscript
number) of cluster geometry E(1b) [See Figure 5.7] with the experimental
spectrum of [Mn(CO2)3]-........................................................................................74

Figure 5.12

Calculated structures of [Mn(CO2)4]‒. Calculations were performed as describedin Chapter 2. All calculated structures are in the quintet spin state. ......................75

Figure 5.13

Comparison of the experimental spectrum of [Mn(CO2)4]‒ (black, top) to
calculated IR spectra (red). ....................................................................................77

Figure 5.14

Calculated structures of [Mn(CO2)5]-. Calculations were performed as describedin Chapter 2. All calculated structures are in the quintet spin state. ......................80

Figure 5.15

Comparison of the experimental spectrum of [Mn(CO2)5]‒ (black, top) to
calculated IR spectra (red). ....................................................................................81

Figure 5.16

Exploratory Structure and calculated spectra for [Mn(CO2)9]‒. ............................83

xv

Figure 6.1

Structures and HOMOs for CO2‒, HCOO‒, AgCOO‒, and BiCOO‒ (presented left
to right in the order listed). Carbon atoms are shown in black, oxygen atoms areshown in red, hydrogen in white, silver in grey, and bismuth in blue. ..................87

Figure 6.2

Photodissociation spectra of CO2‒∙(CO2)7 (top, black), HCOO-·Ar (second trace,
magenta), AgCOO‒ ∙ (CO2)4 (third trace, red), and BiCOO‒ ∙ (CO2)3 (bottom trace,blue). Transitions are identified as symmetric (ωS) or asymmetric C‒O stretchingmotions (ωA). Assignments of the spectra are discussed below. ...........................89

Figure 6.3

Combined plot of the change in charge of the COO group, Δq, plotted along thesymmetric C‒O stretching coordinate. The change in charge of HCOO‒ is shownin magenta squares, AgCOO‒ shown in red circles, and BiCOO‒ shown in blue
triangles. The classical turning points along the normal mode coordinate are at ±1,
and the equilibrium geometry is shown at 0. ........................................................93

Figure 7.1

Walsh diagram of N2O, showing molecular orbital (MO) energies as a function of
NNO angle. The Walsh diagram for N2O was calculated using the RI-MP2
method (second-order Møller-Plesset perturbation theory applying the resolutionof-identity for evaluation of two electron integrals)125,126 with def2-TZVPP basis
sets for all atoms.58 Labels for MO curves in linear and bent geometries are shown
above the respective curve. In cases where more than one curve is shown, thelabels are listed with decreasing energy. The LUMO of neutral N2O is the 3π
orbital. ....................................................................................................................97

Figure 7.2

Mass spectrum of ions generated by electron injection into a supersonic expansionof neat N2O. ...........................................................................................................99

Figure 7.3

Infrared photodissociation spectra of (N2O)nO‒ (n = 1‒12). The numbers on theside indicate the number of N2O in the cluster. ...................................................100

Figure 7.4

Comparison of the experimental spectra of (N2O)5O‒ (top) and (N2O)5O‒∙Ar2(bottom)................................................................................................................101

Figure 7.5

Calculated isomers of N2O2‒. Calculations were performed as described in
Chapter 2. .............................................................................................................102

xvi

Figure 7.6

Comparison of the experimental spectrum of N2O2‒ ∙ 2Ar (top trace) withsimulated spectra of different isomers of N2O2‒ (bottom three traces). ...............103

Figure 7.7

Top panel: ratio of integrated peak areas (square points) compared with thepredicted values from the model explained in the text (dashed line). Error bars are
10% of the first data point and 50% for all subsequent data points. Bottom panel:
percentage of clusters with a NNO2‒ core ion according to the model. The model
assumes 100% NNO2‒ core ion population for n = 1...........................................105

Figure 7.8

Photodissociation spectra of (N2O)n- cluster ions (n = 7-11,14,15). ....................107

xvii

List of Tables
Table 3.1

Spin states of calculated core ion geometries. ......................................................24

Table 4.1

Spectral assignments of the experimental bands at n = 2 compared to predictedvibrational frequencies. ..........................................................................................41

Table 4.2

Spectral assignments for n = 3 compared to predicted frequencies. ......................48

Table 6.1

A summary of the calculated M‒C bond distances, charge on the COO groups, theintensities of ωS and ωA (IS and IA) and the calculated and experimental ratio(IS/IA). The ratio labeled DFT refers to the calculated infrared intensities from thecalculated vibrational spectrum of each complex. .................................................92

Table 7.1

Experimental transitions and calculated (harmonic approximation) (in cm-1) for
neutral and anionic N2O. Experimental data for neutral N2O have been taken from
Ref. 27. Experimental data for N2O– are from the present work. .......................108

1

Chapter 1
Introduction and Motivation
The steadily increasing concentration of CO2 in the atmosphere has been one of the primary
drivers of anthropogenic climate change.1 One promising avenue toward reducing the amount of
CO2 released is to replace combustion of hydrocarbons as the primary source of electricity.
However, new technologies involving solar or wind energy will require modernizing current
infrastructure and changing many technology approaches we currently rely on. In the interim, thedevelopment of a carbon neutral fuel cycle would allow current technology to be maintained, while
decreasing the amount of CO2 released from industrial sources such as power plants.2,3
The chemical conversion of CO2 into a usable feedstock chemical for the generation ofhydrocarbon fuels occurs in the simplest variant via a two-electron reduction process to CO:
CO2 + 2H+ + 2e‒→ CO + H2O (-0.53 V).4
The first step in this conversion in the one electron reduction of CO2 to CO2‒,5 although a hydride
transfer to CO2 has also been proposed as the initial step.6 In order to discuss this process in thecontext of catalysis, a thorough understanding of how CO2 interacts with an excess electron isnecessary.
In its neutral ground state, CO2 is a linear molecule with D∞h symmetry and CO bondlengths of 117 pm.7 The highest occupied molecular orbital (HOMO) of CO2 is the 1πg orbital.
Examination of the Walsh Diagram of CO2 shows that an excess electron will be accommodated
into the 2πu orbital. This orbital is stabilized by the deformation of the OCO bond angle away from180˚. In its ground state, the CO2‒ has a calculated OCO bond angle of 134˚,8,9 C2v molecularsymmetry, and the singly occupied molecular orbital has a1 symmetry. Occupation of this orbital

2

also results in an increase of the CO bond distance from 117 pm in the neutral7 to 124 pm in theanion.8 This is not surprising since the HOMO of CO2 is antibonding.10

Figure 1.1. Walsh Diagram of CO2 (Left).11 Reproduced from Ref. 12 with permission from
Taylor and Francis Group. The HOMO of CO2 (lower) and CO2‒ (upper) based on HF-SCF
calculations are shown on the right.

Formation of this radical anion is complicated by the negative adiabatic electron affinity(AEA) of CO2. Several experimental studies estimate the AEA to be -0.6 eV.13 Although the CO2‒
anion is unstable against spontaneous electron emission, several observations of CO2‒ have been
claimed in mass spectrometry experiments with an estimated lifetime of the ion of up to amillisecond.14-16 Observations of the anion are possible due to its large vertical detachment energy,
1.4 eV.17 The CO2‒ anion can be stabilized through interactions with a matrix18-20 or throughsolvation of the ion in the condensed phase or in gas phase cluster ions.21-25 For more details about

3

how CO2 interacts with negative charge, the author directs the reader to the following review
articles and citations therein.12,26
The population of the antibonding HOMO in CO2‒ causes changes in the vibrationalfrequencies of the anion compared to neutral CO2. As an example, the antisymmetric CO2
stretching frequency is red shifted from 2349 cm-1 (neutral) to 1665 cm-1 (anion).18,27 Experimental
work by Weber and coworkers on X‒∙CO2 (X=Cl, Br, I) have shown a correlation between the
charge on a CO2 species and its asymmetric CO2 stretching frequency.28 Further work on [M-CO2]‒
species (M = Au, Ag, and Cu) have corroborated and expanded upon this relationship, establishingthat the asymmetric CO2 stretching frequency can be used to estimate the amount of charge on a
CO2 species.29-31
Energetically, it is very costly to add an excess electron to CO2 (0.6 eV).14 Industrially thesolution to an energetically costly reaction is often to employ a catalyst to both reduce high reaction
barriers and as a result increase the reaction rate without the input of large amounts of energy into
the reactor. The reductive activation of CO2 is often employed over atomically rough electrodesurfaces in order to increase the accessible surface area. Work by Kanan and coworkers has shownthat electrochemically roughened catalyst surfaces are superior to their more atomically flat
counterparts.32,33 This observation has corroborated the hypothesis that the reduction of CO2 occurs
at corner or edge sites of an electrode surface. Ideally, the process of refining a reaction forindustrial application uses precise understanding of the fundamental physics and chemistry
governing the individual chemical reactions. In the context of CO2 reduction, one of the firstinteractions that must be understood is that of CO2 with these catalytically active corner and edge
sites. For more information about catalytic reduction of CO2 the author recommends the following
review articles and references therein.34-40

4Experimental determination of the binding motifs of CO2 to corner sites in situ iscomplicated by speciation, low concentrations, and short lifetimes of reactive intermediates in thecondensed phase. In order to limit the number of complicating factors, gas phase cluster ions canbe used as a model system to study these interactions in the absence of speciation. Corner sites ona charged electrode surface can be approximated as a single atomic anion, since charge tends toaccumulate at these corner sites, and the CO2 species may interact with a single atom at the site.
With this in mind, gas phase cluster ions of a metal atom and CO2 in the presence of an excess
electron can be used to study the binding motifs of CO2 molecules to strongly charged corner sites.
At the very least, these model systems provide insight into the fundamental interactions between
CO2 and undercoordinated metal atoms in the presence of an excess electron.
The structure of the molecular core ions in clusters, and how this structure is perturbed bysolvation are the main topics of this thesis. In Chapter 2, the experimental apparatus andexperimental and computational methods are discussed. The low concentration of cluster species
precludes the use of traditional linear IR absorption spectroscopy. Instead, infrared
photodissociation spectroscopy was employed to study the cluster species described in this thesis.
Chapters 3 and 4 will compare and contrast the interaction of CO2 with bismuth and tin anions asa method of understanding the activation of CO2 at bismuth or tin based electrocatalysts. Chapter
5 examines the interaction of CO2 with an iron atom and with a manganese atom. While notnecessarily applicable to the initial question of CO2 reduction catalysis, these clusters serve as away to probe the structure of novel metal-organic compounds containing several carbon dioxide
ligands. The final two chapters deviate from the overall theme discussed above. In Chapter 6 thefundamental physics behind the infrared intensity ratios of the asymmetric and symmetric CO 2
stretching modes of MCOO− (M = H, Ag, Bi) is examined. Chapter 7 focuses on the fundamental

5interaction of the solvation environment with the core ions of N2O cluster anions. N2O is a fairly
interesting analogue to CO2 because they are isoelectronic. However, electron attachment to N2O
is complicated by a low lying dissociative electron attachment process that results in the formationof (N2O)nO−, providing access to fundamentally different types of core ions.

6

Chapter 2
Experimental and Theoretical Methods
2.1

Overview
Experiments were performed with a custom photodissociation action spectrometer. Ions

were generated with one of three methods, entrainment of laser-ablated material into a supersonic
expansion41 (CO2), entrainment of volatile species into a supersonic expansion (Ar or N2O)
coupled with electron bombardment,42 or electron bombardment of a neat supersonic expansion(N2O and CO2). Each method generates a large number of anionic, cationic and neutral species;
the anionic species are injected into a time-of-flight mass spectrometer. Once mass separation isachieved (within the first 10 µs after injection), the ions are mass selected by a pulsed mass gate.43
The ions are irradiated with tunable light (600 cm-1-4000 cm-1) generated by an infrared opticalparametric converter (IR-OPC) system. Absorption of a photon of sufficient energy will result inthe loss of a weakly bound component of the cluster. A reflectron is used as a secondary mass
analysis step, in order to separate daughter ions from the heavier parents, and the formation ofdaughter ions is monitored as a function of the laser wavenumber. In the following sections, theexperimental apparatus and methods will be discussed in more detail.
2.2

Experimental Methods

2.2.1 Ion Generation
There are several methods for generating gas phase molecular ions, many of which dependon the stability of the ion with respect to the neutral. For fairly large ions, electrospray ionization44
or matrix assisted laser desorption and ionization (MALDI)45 are appropriate ionization methods.
However, in the case of unstable anionic species (such as CO2-), special care must be taken to7stabilize these species against electron auto-detachment.13 Unstable anions are often formed via
electron attachment to neutral cluster species.14 Lineberger and coworkers developed a supersonicion source to generate a beam of molecular anions by intersecting a supersonic expansion with theoutput of an electron gun.46 Entrainment, described by Campargue in 1970, is a process by whichmolecules penetrate the first shock boundary of a supersonic expansion and are drawn into thesupersonic flow.47 This phenomenon is useful in preparing reactive gas mixtures without directly
seeding the carrier gas prior to expansion. By entraining volatile gases into a supersonic expansion,
the exact conditions (i.e. concentrations, backing pressures, etc.) can be easily adjusted to optimize
ion production.42 The supersonic entrainment ion source operates by generating an electron impact
plasma in the high-density region of the supersonic expansion. The plasma contains slow
secondary electrons that combine with neutral clusters (A•mB) forming “hot” anionic clusters(A•mB‒*) (Equation 2.1).
A•mB + e‒ (slow) → A•mB‒*(2.1)

The hot anionic clusters will then evaporate weakly bound constituents in order to dissipate theattachment energy of the electron (Equation 2.2).
A•mB-* → A•(m-n)B- +nB,

(2.2)

The newly formed clusters exist as a molecular core ion surrounded by one or more weakly boundneutral molecular species that are solvating the ion, and can possess a significant amount of internalenergy, in some cases much more than the binding energy of one of the solvating molecules. The
process of losing excess energy can be described by an evaporative ensemble,48 where the cluster
species dissipate excess internal energy by evaporating one or more weakly bound molecularspecies until the ions are stable over the observation time of the experiment. The internal energy

8of the clusters can be estimated to be on the order of the binding energy of the last species to beevaporated. The assumption places an upper limit on the internal energy content of a cluster species
that possesses at least one weakly bound solvent. The experiments described in this thesis were
performed using two ion source configurations, a laser vaporization ion source and an electron
impact entrainment source.
2.2.1.1 Laser Vaporization Ion Source
Many of the metal-CO2 cluster species discussed in the later chapters were generated usinga laser vaporization ion source. Initially developed independently by the group of Smalley and by
Bondybey and English as a method of generating cold molecular beams of charged metal atoms
and clusters,49,50 laser vaporization ion sources are used to study materials whose low vapor
pressure would normally preclude study in a molecular beam experiment. The laser vaporization
source used here (Figure 2.1) features a long channel through which a laser is allowed to propagate
and impinge on a rotating metal target. The resulting metal vapor is entrained into a pulsed
expansion of CO2 generated from an Even-Lavie valve (stagnation pressure 550 kPa).

9

Figure 2.1. Schematic of the ion source. In the laser vaporization method, a laser propagatesfrom right to left and is focused onto the edge of a metal target (see text). In the entrainment ofvolatile gases experiment, the target gas is introduced with a general valve and entrained into themain expansion (see text), additionally ions can be produced using only the main expansion.

The metal disk targets were 50 mm in diameter with a thickness of 6 mm, and were eitherpurchased (Fe, Mn, Sn) or manufactured in JILA (Bi). Metal from the rim of the target wasvaporized with the third harmonic of an Nd:YAG laser (InnoLas Spitlight 300, 355 nm, 7 ns pulse
duration, 40 mJ/pulse) approximately 200 µs before the pulsed expansion. A lens, placed at theentrance to the entrainment channel, focuses the laser beam at the target. The disk is rotated toallow the laser to hit a fresh portion of the surface with each experimental cycle. Adiabatic coolingof the entrained metal vapor occurs in the expansion. The laser impact plasma creates a large
number of cationic species as well as a sufficient number of slow secondary electrons to allow for
the formation of [M(CO2)n]- (n=0-15) in sufficient abundance to perform our experiments. The gas
expands into a vacuum chamber typically held at several 10-6 mbar, pumped by a 400 Ls-1
turbomolecular pump (Leybold Turbovac 361) and a 2000 Ls-1 diffusion pump (BOC Edwards
Diffstak 250/2000). The current ion source is capable of generating large abundances of clusters

10

containing a single metal atom and several CO2 molecules, with small abundances of metal dimer
centered structures, depending on the metal in question and source conditions. The experiment is
operated at a repetition rate of 20 Hz.
2.2.1.2 Entrainment of Volatile Species
The entrainment of gases into pulsed supersonic expansions was initially developed as analternative to using seeded buffer gas sources in electron beam ionized supersonic expansion
sources.42 One of the constant struggles with using seeded gas sources in these experiments is thedifficulty of changing concentration of the seeded gas in the expansion. Johnson and coworkers
describe an ion source for the generation of cold ion-molecule complexes via entrainment of avolatile gas into a supersonic expansion.42 Entrainment of volatile gases into a supersonic
expansion was performed with the same ion source configuration described above using a different
ionization technique (see Figure 2.1). For this experiment a Series 9 General Valve (Parker) isused to introduce a secondary gas (B) into the experiment several microseconds before the main
expansion. Ions are produced by coupling an electron impact plasma to the supersonic expansion
by focusing the output of an electron gun (EKE=800 eV, 10‒50 µA) into the high density region of
the expansion. The target species is entrained into the main expansion of inert gas (Ar). The
resulting electron impact plasma generates slow secondary electrons that then interact with themolecule of interest, forming metastable [B- • mAr] cluster species. The excess energy of the B
anion can be dissipated by evaporating the weakly bound (500 cm-1)51 Ar atoms. A pulse delay
generator (Quantum Composers 9614+) controls the gas pulse timing.

11

2.2.2 Mass Spectrometry and Photodissociation Spectroscopy
Each of the ion sources described above results in the formation of cationic, neutral, andanionic cluster species. In order to separate the anionic species from this mixture, and allow the
study of mass selected anions, the ion source is coupled to a Wiley-McLaren time-of-flight mass
spectrometer. A schematic of the experimental set up is shown in Figure 2.2. The anionic portionof the molecular beam is accelerated, orthogonal to the expansion axis, into the field-free time-offlight region by two accelerator plates pulsed from ground to -4000 kV (repeller) and -3000 to 3400 kV (extractor) in a Wiley-McLaren arrangement.52 Two sets of electrostatic ion deflectors
are used to steer the ion beam, and an Einzel lens is employed in order to focus the ion beam. The
time-of-flight tube is typically held at a pressure of 10-8 mbar and is pumped by a 520 Ls-1
turbomolecular pump (Pfeiffer TC600). A delay generator (Quantum Composers 9614+) controls
the timing of all triggering pulses.

Figure 2.2. Diagram of the experimental apparatus with paths outlined of the vaporization laser(blue), ion path (green), electron beam (red), and laser (purple).

12

In an idealized mass spectrometer with a point source of ions, each ion receives the sameamount of initial kinetic energy in the direction of flight. The relationship between the masses of
two ions and their flight times can be written as:
m1 t1
=
√
m2 t 2(2.3)

The mass spectra can therefore in principle be assigned by knowing the mass of a single peak inthe spectrum. The initial assignment of a mass spectrum can be aided by the formation of simple
cluster ions (i.e. (CO2)n-) which have a known intensity profile as a function of size.53,54 A typicalmass spectrum can be found in Figure 2.3.
In reality, the ions arrive at the acceleration region with a distribution of positions along
the time-of-flight coordinate. Ions that are closer to the ion detector will spend less time in theaccelerating field and as such will be slower than ions (of the same m/z) that spend more time in
the accelerating field. This difference in velocity results in the faster ions eventually overtaking
the ions that spent less time in the field. The exact position at which this occurs, as described by
Wiley and McLaren,52 can be adjusted by tuning the biases of the two accelerator electrodes. This
position is referred to as the space focus of the mass spectrometer.

13

Figure 2.3. Mass spectrum obtained via the laser vaporization ion source with bismuth target anda pulsed CO2 expansion.

A mass gate, placed at the space focus of the spectrometer, is used to mass select ions. 43
The mass gate consists of a series of parallel shim electrodes that can be biased to ±100 V withneighboring plates alternating in polarity. When the plates are biased, the ions encounter an electricfield, which deflects them into the walls of the flight tube. If the plates are grounded, the ionscontinue on their original path. This device is used to mass select particular mass-to-charge ratios
by pulsing the mass gate to ground for a short time (800 ns). Ions that are allowed to pass throughthe mass gate are irradiated by the output of an IR-OPC system (see section 2.2.3). The ionscontinue to travel before encountering a two-stage reflectron (Jordon TOF) which reflects the ions.
The reflectron serves several purposes in the experiment, primarily the reflectron is used in concert
with photodissociation spectroscopy techniques to separate laser generated daughter ions from theparent ions. Secondly, the reflectron is used to account for any differences in velocity of the ionsalong the time-of-flight axis. Thirdly, the reflectron doubles the flight distance of the ions. This

14

will result in an increase of the relative flight times of each ion, since there is a larger distance inwhich to separate the ions. The ions are detected by a microchannel plate detector (MCP) installedat the second space focus of the spectrometer, whose position is defined by the ratio of the electricfields in the two stages of the reflectron.
Photodissociation spectroscopy is a method by which absorption of a photon results in the
dissociation of the target molecule. This technique is paired with a tandem mass spectrometer. The
first mass spectrometry step serves to mass-select the target ions, which is achieved in the firstspace focus of the Wiley-McLaren setup in the current experiment. The second mass spectrometry
step, occurring after irradiation, is used to monitor dissociation events, since the subsequent lossof mass from the target cluster can be detected as a change of flight time. For the metal-containing
cluster systems that will be discussed later in this thesis, the cluster ions in question will typicallyundergo the following process:[M(CO2)n ]‒+ hν → [M(CO2)n-1 ]- + CO2(2.4)

If a cluster absorbs a photon, the absorbed energy is distributed among all energetically accessiblemodes through intramolecular vibrational redistribution. The now “hot” molecular cluster candissipate excess energy by evaporating one or more weakly bound solvent species (Equation 2.4).
In the case of CO2 containing cluster anions, an infrared photon, in the wavenumber range 10002150 cm-1, typically has enough energy to cause the evaporation of a single CO 2 solvent species.
If the photon energy is higher than the binding energy of the most weakly bound solvent species,
this process occurs with near unity efficiency during the observation time of the experiment. Thisis the time between irradiation of the ions and when they reach the reflectron (10-30 µs). When aphoton with lower than the binding energy is absorbed, then the cluster must utilize its internalenergy prior to photon absorption to evaporate the solvent species. Often this is the case with very

15

low photon wavenumber (<1300 cm-1), and results in a suppression of the observed intensity ofsignatures in this energy range.
2.2.3 Light Source
Tunable infrared light in the wavenumber range of 600 cm-1 to 4500 cm-1 is generated using
an IR-OPC system (Laservision, see Figure 2.4).

Figure 2.4. Schematic of the Optical Parametric Oscillator/Amplifier system (/2, half-wave plate;OC, output coupler; BD, beam dump; Si-Pol., silicon polarizer; Pol., polarizer).

16

The optical parametric converter system is pumped by a Nd:YAG laser (InnoLas Spitlight 600, 7
ns pulse, 500 mJ/pulse, 1064 nm, ν1064) operating at a 20 Hz repetition rate. The output of the
Nd:YAG laser is split into two beams. One of the beams is frequency doubled into 532 nm lightand is used to pump an optical parametric oscillator (OPO). The OPO consists of two potassium
titanyl phosphate crystals (KTP) inside an optical cavity. The 532 nm pump wave is split into a
signal wave (νsignal) and an idler wave (νidler). The signal and idler wavenumbers are connected byconservation of energy and momentum such that the sum of the energy of these two waves must
equal the energy of the 532 nm pump (see equation 2.5). Since the signal and idler waves haveperpendicular polarizations, the signal wavelength is separated out of the OPO output by apolarizer. The idler beam profile is rotated 90˚ by a dove prism to minimize the divergence of thebeam in the vertical direction in the next stage. Any residual 532 nm light is removed with a Sipolarizer, which acts as a filter for light below 1127 nm. The second 1064 nm beam is routed intoa delay stage and directed into an optical parametric amplifier (OPA), which consists of four
potassium titanyl arsenate crystals (KTA). The delayed 1064 nm beam is difference-frequency
mixed with the idler output of the OPO (see equation 2.6). This generates mid-infrared light (νmidIR)in the range of 2100 cm-1 to 4500 cm-1 (pulse energy 5 mJ to 20 mJ). Any residual 1064 nm lightis reflected by a 1064 mirror and dumped.
The mid-IR and idler beams co-propagate out of the OPO/OPA and onto a second table.
Depending on the frequency of light needed, the idler beam is removed via a polarizer or the midIR and idler beams are allowed to undergo difference frequency mixing in a single silver gallium
selenide (AgGaSe2) crystal. Mixing in the crystal generates a far-IR (νfar IR) beam with an energy
range of 600 cm-1 to 2150 cm-1 (see equation 2.7; pulse energy 10 µJ to 100 µJ). Any remaining
mid-IR and idler radiation are removed from the far-IR beam by a germanium filter, which only

17

allows far-IR light to pass (transmission 5-14 µm). A summary of the various frequency
relationships can be found below in equations 2.5 to 2.7.
2 × ν1064 = νsignal + νidler(2.5)ν1064 ‒ νidler = νmidIR(2.6)

νidler‒ νmid IR = νfar IR(2.7)

The OPO is calibrated by measuring the signal wavelength (719 – 820 nm) after the OPO
stage. A combination of two spectrometers is used to measure the signal wavelength, an Ocean
Optics fiber-optic spectrometer (820 to 750 nm) and an Avantes fiber optic spectrometer (750 to719 nm). We estimate the bandwidth of the optical parametric converter system to be around2 cm-1. Given that H2O and CO2 exhibit very strong lines in the tuning range of the OPO/OPA, theentire beam line is purged with dry nitrogen, in order to eliminate the loss of power due toabsorption by atmospheric trace gases.
2.2.4 Signal Acquisition and Processing
In order to monitor the mass spectrum, the signal from the microchannel plate detector isvisualized using two digital oscilloscopes (Tektronix TDS 2022). The IR photodissociation spectra
are measured by monitoring the formation of lighter daughter ions as a function of laser
wavenumber. One of the oscilloscopes is set to display the entire width of the daughter peak (~100ns). The intensity of the ion signal is averaged over 64 shots per data point. A second oscilloscope
monitors a section of the mass spectrum with no ion signal. This oscilloscope captures any changes
in the baseline of the mass spectrum (background) as the fragment intensity is monitored. The IROPO/OPA is scanned along the tuning range by a LabVIEW program, which also collects theoutput data of the oscilloscopes. Laser power is monitored by placing a KBr window into the18beamline and measuring the back reflection with a pyroelectric joule meter. Since the transmission
efficiency of KBr is constant over the tuning range of the OPO/OPA the pulse energy of the IR
beam can be monitored using this back reflection. This also allows for simultaneous measurementof both laser power and photodissociation signal. The raw fragment formation data are further
processed by subtracting the background from the laser induced fragment signal. The resulting
infrared spectrum is then corrected for photon fluence by dividing the signal intensity by thephoton number as a function of photon wavenumber. Several spectra collected over multiple days
are averaged in order to increase the signal-to-noise ratio of the average spectrum. In addition,collecting data over several days allows us to confirm the reproducibility of the infrared spectra.
2.3 Computational Methods
In this section, common computational methods will be briefly discussed. In some cases,
different systems required other computational methods. These calculations will be discussed in
the appropriate chapter as necessary. Generally, in order to aid in the interpretation of thephotodissociation spectra, quantum chemical calculations are employed in order to determine thestructure of the clusters in question.
Calculations were performed on the TURBOMOLE suite of programs55 (ver. 5.9.1 and ver.
6.2) using density functional theory with a dispersion corrected B3LYP functional,56,57 and def2TZVPP basis sets for all atoms.58 Harmonic frequency calculations were performed using the
AOFORCE program.59 For CO2 based systems, harmonic frequencies were corrected foranharmonicity using a correction factor of 0.975, which was found by comparing the harmonicfrequency of free CO2 with experimental values. Some vibrations were corrected using othervalues, these motions will be described in the appropriate section. Approximations of the charges
on each individual atom were calculated using Natural Population Analysis.60

19

Initially, cluster structures are created based on chemical intuition and then optimized using
the above methods. The calculated vibrational spectra are then convoluted with a 12 cm-1 full widthat half-maximum Gaussian to generate a calculated vibrational spectrum that is then compared tothe experimental spectrum of a given cluster. Assignments are made based on the agreement
between these calculated spectra and the experimental spectrum.

20

Chapter 3
Interactions of CO2 with Bismuth: Formation of Metalloformate and Metallooxalate
Complexes
This chapter has been adapted with permission from:
Thompson, M. C., Ramsay, J., and Weber, J. M. “Solvent-Driven Reductive Activationof CO2 by Bismuth: Switching from Metalloformate Complexes to Oxalate Products.”
Angewandte Chemie International Edition, 2016. 55: p. 15171-15174. DOI:
10.1002/anie.201607445. Copyright John Wiley and Sons. License Number: 4310361273540

3.1 Introduction
The chemical conversion of CO2 into hydrocarbon fuels in a carbon neutral fuel cycle is apromising avenue toward reducing worldwide dependence on fossil-based fuels. Many syntheticroutes to hydrocarbon fuels from CO2 begin with the reduction of CO2 to CO.39 Researchers suchas Hori and coworkers during the late 1980s and 1990s studied the product distributions andenergetics of CO2 electroreduction over many different electrode materials ranging from transition
metals to p-block metals.61-63 These studies showed that the product of electroreduction of CO2 at
p-block metal electrodes was primarily formic acid, while reduction over noble metal electrodes(Au, Ag) produced carbon monoxide as a principal product.4
Several years ago, Rosenthal and coworkers were successful in developing a device for theformation of CO from CO2 by electrodepositing bismuth onto glassy carbon. The resulting surface

21

was incorporated into an electrochemical device with ionic liquid saturated with CO2.64 The resultof this experiment was a working device capable of forming CO at Faradaic efficiencies
comparable to devices made of electrochemically roughened or sputtered gold surfaces. Further
experiments using deposited bismuth nanoparticles as electrode material resulted in Faradaic
efficiencies comparable to those of bismuth foil and electrodeposited bismuth surfaces.65
An open question regarding these catalytic devices is the exact mechanism of CO
formation. While several theories exist for homogeneous CO2 reduction catalysts,66,67 the sheer
number of morphologies seen on electrode surfaces makes hypothesizing an exact mechanism
difficult. A consensus is that the first step in the reduction of CO2 to CO is the single electron
reduction of the CO2 to CO2•-.5,39,68 At the same time, this is believed to be the rate-limiting stepin the reduction of carbon dioxide. As such, an important first process to examine would be thedocking of a CO2 molecule onto the electrode surface and the transfer of a single electron to themolecule. Gas phase clusters have been used as effective model systems to study the behavior of
CO2 interacting with under-coordinated atoms on a metal surface.29-31 In this chapter, theinteraction of CO2 with a bismuth atom in the presence of an excess electron is discussed.
3.2 Results and Discussion
In order to understand how a CO2 molecule interacts with a bismuth atom in the presenceof an excess electron, we measured the IR photodissociation spectra of [Bi(CO)2)n]‒ (Figure 3.1).
The experimental spectra are displayed as a series of individual traces, where the numbers indicate
the size of the cluster, n. Each individual spectrum is scaled to the highest intensity peak. In anionic
metal-CO2 clusters, the CO2 molecules can take on one of two roles; the CO2 molecule is directly
bound to the metal atom as a ligand, or the CO2 molecule solvates the molecular core ion.
Signatures that appear in the spectral region shown in Figure 3.1 are CO stretching vibrations of22

CO2 species that are directly bound to the metal as ligands. These ligands exhibit binding energies
in excess of 4000 cm-1,69,70 such bonds cannot be broken with a single photon in the energy ranges
used in this experiment. The CO2 species that solvate the core ion are less tightly bound than theligands (around 1500 cm-1). The presence of one or more solvent CO2 molecules is necessary forclusters to be observable in this experiment. CO vibrational frequencies of CO2 ligands are red
shifted compared to the analogous motions in free CO2. This is due to charge sharing that occurs
between the metal center and the ligand(s). For an in-depth discussion of how CO2 interacts withnegative charge, see Chapter 1.

Figure 3.1. IR photodissociation action spectra of [Bi(CO2)n]‒ (n=2‒9). Labels denote the core ion
geometry (F=metalloformate, O=oxalate) and the subscript letters denote the type of oscillation(S=symmetric, A=asymmetric, M=metal bound, F=free). The dashed lines are meant to draw theeye to experimental signatures of the metalloformate complex (magenta) and metallooxalate
complex (green).

The spectra of the smaller cluster ions (n=2‒4) share a set of spectral features, labeled F in
Figure 3.1. These spectral features appear to blue-shift (FS) and red-shift (FA) as cluster size

23

increases. In the intermediate cluster size regime (n=5‒7) another set of spectral features appear(OM, OF) and grow in intensity relative to FS and FA. At larger cluster sizes (n=8‒9), the second
set of spectral features become the only observed signatures. Unlike FS and FA, the positions of OF
and OM do not change with increasing cluster size. The appearance of the second set of peaks (OFand OM) and the disappearance of the first set (FA and FS) suggests a change in core ion geometry
as cluster size increases. Such changes have been observed previously, and often can be associatedwith the formation of a first solvation shell, or the stabilization of a core ion structure due toincreased solvation.23,25,71
The structure of cluster species was determined by comparing calculated vibrationalfrequencies of a variety of cluster geometries with the experimental spectra. The harmonic
calculated frequencies are corrected by a factor of 0.9398 for the asymmetric CO stretching motion
and 0.9655 for the symmetric CO stretching motion of the metalloformate core ion as found bycomparing known experimental values18,29 with calculated values.72 All other motions werecorrected as described in Chapter 2. A summary of the stable core ion geometries is shown in
Figure 3.2.

Figure 3.2. Calculated core ion geometries (Bi, blue; C, black; O, red).

24

The metalloformate ion is a simple species where the CO2 ligand forms a single bond from its
carbon atom to the metal center. The geometry is the result of an attraction of the partially positive
carbon atom to the negatively charged metal center, accompanied by a repulsion of the partiallynegative oxygen atoms. The result is a formate-like complex where the hydrogen atom has been
replaced with a bismuth atom. The (C,O)-bidentate ion is a species where the CO2 molecule forms
two bonds to the bismuth center through the carbon atom and one of the oxygen atoms. Thisgeometry is only computed in the singlet spin configuration, while the metalloformate geometry
is found only in the triplet spin configuration.
The larger core ions feature two CO2 species acting as ligands. The butterfly structure
consists of two CO2 ligands each bound in a (C,O)-bidentate motif. In the oxalato ligand, two CO2
molecules are connected to one another through the formation of a C-C bond, and the ligand binds
to the metal atom through two oxygen atoms belonging to different CO2 subunits. This geometryis stable in both singlet and triplet configurations, while the butterfly structure is stable only in the
singlet spin state. A summary of which calculated core ions are observed in either triplet or singlet
spin states is shown in Table 3.1.

Table 3.1. Spin states of calculated core ion geometries
Core Ion Geometry

Stable in singlet spin state?

Stable in triplet spin state?

Metalloformate

No

Yes

(C,O)-Bidentate

Yes

No

Butterfly

YesNo

Oxalate

Yes

Yes

25

Each set of spectral features (F and O) can be assigned by comparing calculated vibrational
spectra to the experimental spectra (Figure 3.3). Beginning with [Bi(CO2)2]‒, two spectral features
are present at 1707 cm-1 (FA) and 1186 cm-1 (FS). These experimental signatures are compatible
with either (C,O) bidentate or metalloformate based clusters. We note that clusters with (C,O)
bidentate core ions lie approximately 400 meV higher in energy than the metalloformate based
clusters. This suggests that the spectral series FA and FS are due to clusters with metalloformate
core ions, and that clusters with a (C,O) bidentate core ion geometry are not present in our
experiment. Therefore, we assign FS and FA to the symmetric and asymmetric CO stretchingmotions of the carboxylate ligand. It is interesting to observe the contrast in relative intensities in
the calculated vibrational spectrum for metalloformate versus the experimental spectrum.

26

Figure 3.3. [Bi(CO2)2]‒ experimental spectrum and calculated spectra. The colors of the tracesindicate the spin of the structure (red, triplet; blue, singlet), and the core ion whose vibrationalsignatures are shown are pictured above each trace. The top two structures have solvent species
behind the ligand.

The formation of a metalloformate complex in this study compares well with previous
work on anionic clusters of CO2 with Ag and Au atoms.29,30,72 In those cases, the metal andcarbon atoms form a σ bond between the s-orbital on the metal atom and the π* orbital of theCO2 ligand. Similarly, the Bi-C bond is also a σ bond. However, the bond is formed between the
p-orbital of the Bi atom and the π*-orbital of the CO2 ligand (Figure 3.4). This suggests that the
charge donation from the metal atom to the CO2 ligand occurs from the orbital on the metal that27

the excess electron occupies (the 5s- or 6s-orbital in the case of Ag or Au, and the 6pz-orbital ofBi).

Figure 3.4. Calculated HOMOs (70% contour) of several MCOO‒ species [M=Ag (silver), Au(yellow), and Bi (blue)].

The spectral shifts of the F series contain important information regarding the overall
structure of the solvation shell around the metalloformate core ion. As has been seen previously,
the red shift of the FA signature is due to increased solvation around the ligand side of the coreion.29-31 Solvation around the carboxylate ligand polarizes excess electron density onto the ligand.
As discussed in Chapter 1, excess electron density is accommodated in an antibonding orbital of
CO2. This weakens the CO bonds and should result in a red shift of the CO oscillators. This
expectation is consistent with the behavior of the FA signature, which strongly red shifts as cluster
size increases. The blue shift of the FS signature is directly contrary to this expectation, as FS should
also red shift, in the absence of other effects. There is, however, an additional structural change in
CO2 upon accommodation of an excess electron, as the OCO bond angle decreases from 180˚ to
134˚.73 The same occurs in the bismuthoformate core ion, as the OCO bond angle decreases as afunction of cluster size from a calculated 140˚ in the bare complex to 133˚ in the [Bi(CO2)6]‒. The

28

decrease of the OCO bond angle results in a decrease of the kinetic coupling of the CO oscillators.
As the coupling decreases, the energy difference between the symmetric and asymmetric CO
stretching modes decrease. Experimentally, this would present itself as a red shift of theasymmetric stretching frequency and a blue shift of the symmetric stretching frequency. This effectshould result in a shift of the same magnitude and opposite sign for the symmetric and asymmetric
CO2 stretching frequencies, in the absence of the red shift associated with a weakening of the CObonds. This means that the average position of these two frequencies and its shift should show the
weakening of the CO bonds, while the decrease of the kinetic coupling between the two CO
oscillators is encoded in the diminishing splitting between the two modes. Figure 3.5 shows thatthe average position of the two signatures redshifts as cluster size increases. This demonstrates
that the weakening of the CO bonds is contributing to the observed spectral shifts. Also shown isthe splitting between FS and FA as a function of cluster size; the decrease of this splitting indicatesthat the OCO bond angle is decreasing with increasing cluster size.

29

Figure 3.5. Positions of the signatures in the F series as a function of cluster size for theasymmetric (red, top left) and symmetric (black, bottom left) as well as the difference (blue, top
right) and the average between FS and FA (magenta, bottom right).

A second set of spectral features appears beginning with [Bi(CO2)6]‒, and becomes thedominant set of signatures seen in the spectra of the larger clusters. The two signatures appear at1665 (OF) and 1304 cm-1 (OM). A comparison of calculated vibrational spectra and theexperimental spectrum (Figure 3.6) shows that the new spectral features are clearly due to theformation of a bismuth oxalate complex.

30

Figure 3.6. [Bi(CO2)6]‒ experimental spectrum and calculated spectra (black, top). The colors of
the lower traces indicate the spin of the structure (red, triplet; blue, singlet). Calculated
vibrational frequencies of a metallooxalate based [Bi(CO2)6]‒ cluster (A) , a metalloformate
based [Bi(CO2)6]‒ cluster (B), and a butterfly based [Bi(CO2)2]‒ cluster (C).

Based on these calculations the new spectral feature at 1665 cm-1 is assigned to the symmetriccombination of free CO stretches in the oxalate ligand. We note that the asymmetric combinationof free CO stretches is predicted to have infrared intensity, but its intensity is calculated to be 30%of that of the more intense symmetric free CO stretch. At cluster size n = 6, this feature is buried
underneath FA. The spectral signature at 1304 cm-1 is assigned to the symmetric combination of
metal bound CO stretches in the oxalate ligand. Similar to the free CO stretches, the asymmetric31

combination of metal bound CO stretches is also predicted to be much lower in intensity comparedto the symmetric combination.
The absence of any clear oxalate based vibrational signatures in the photodissociation
spectra of smaller clusters is something that warrants further discussion. Although clusters withsinglet oxalate core ions are calculated to be the lowest energy isomer for all cluster sizes, oxalatefeatures are only observed for intermediate and larger cluster sizes (n=6‒9). The presence of triplet
based metalloformate containing cluster species in small cluster sizes suggests that a comparison
between the relative energies of triplet oxalate and metalloformate based clusters is appropriate. A
comparison of the relative energies of clusters containing these two core ion geometries as afunction of cluster size (Figure 3.7) shows that the triplet oxalate based clusters are stabilized ascluster size increases. This trend continues until the two types of clusters are isoenergetic at n = 5
and then the oxalate based cluster is lower in energy at n = 6.

32

Figure 3.7. Comparison of the relative energies of the triplet formate and triplet oxalate core ionclusters as a function of cluster size.

The calculated cluster size where the two types of clusters are isoenergetic is identical to theexperimental observation of oxalate based signatures at cluster size n=5. Since the appearance ofoxalate based signatures seems to be tied to the relative stability of the triplet core ion geometry,clusters are most likely formed in the triplet electronic configuration. This suggests that the ionsare formed via association reactions of CO2 with Bi‒ in its 3P2 ground state during supersonic
expansion. We note that there are of course many different solvation positons around a given coreion. We have calculated several different solvation positions and have generally found that thedifference in energy between different solvation positions is small (±50 meV) for these cluster
species, and this value is used as the error bars in Figure 3.7.
The increasing stability of triplet oxalate based clusters can be explained through ananalysis of the amount of charge on this ligand compared to metalloformate-based clusters. Natural
population analysis assigns a maximum charge of -0.9 e on the carboxylate ligand of large

33

metalloformate based clusters (n=6). The oxalate ligand contains a calculated charge ofapproximately -1.5 e. The differential solvation energy gain demonstrated in Figure 3.7 is due tosolvating a larger amount of charge, which is more favorable than solvating a smaller amount ofcharge. This proposed mechanism is slightly different from other observed core ion changes incluster ions. In the case of (N2O)nO‒, changes in the core ion structure are driven by core ion size
rather than purely by charge as it is in the case of [Bi(CO2)n]‒.23,25
3.3 Conclusions
The interaction of CO2 with bismuth in the presence of an excess electron was studied
using photodissociation spectroscopy of mass selected [Bi(CO2)n]‒ (n=2‒9). Small anionic clusters(n=2‒4) have a metalloformate core ion structure. Intermediate clusters (n=5‒6) have both
metalloformate and metallooxalate core ion structures. Larger cluster sizes only exhibit ametallooxalate core ion structure. Solvation around the ligand of a metalloformate core ion
polarizes charge onto the ligand. The increased charge can be detected as a red shift of theasymmetric CO stretching frequency. Increased charge also results in a decrease of the OCO bondangle. The decrease of this angle decreases the kinetic coupling of the individual CO oscillators.
This is observed as a blue shift of the symmetric CO stretching frequency. The change frommetalloformate to metallooxalate core ion geometries is due to the larger amount of charge on theoxalate ligand, which has a larger solvation energy compared to the metalloformate core ion.

34

Chapter 4Interaction of CO2 with a Sn Atom in the Presence of an Excess Electron: Formation of a
Weak Sn‒C bond.
This chapter has been adapted with permission from:
Thompson, M. C., Ramsay, J., Weber, J. M., “Infrared Photodissociation Spectraof [Sn(CO2)n]− Cluster Ions.” Journal of Physical Chemistry A, 2018. ASAP DOI:
10.1021/acs.jpca.8b00362. Copyright 2018 American Chemical Society.

4.1 Introduction
As stated previously, one of the limitations in developing economically viable catalysts forCO2 electroreduction to liquid fuels is the development of cheaper catalytic materials. Some of themore promising materials explored in recent years have been inexpensive post transition metals,
such as Sn and Bi.65,74 Hori and coworkers have studied the catalytic activity of Sn toward CO2
previously.4,61 Additional work has shown that electroreduction of CO2 at Sn electrode surfaces
resulted in formic acid formation.61,63,75 Sn nanoparticle based electrodes have shown formate
formation to be the dominant reaction pathway under turnover conditions.76
An electrochemically roughened Sn based electrode surface has shown very high Faradaic
efficiencies for carbon monoxide formation.77 The efficiency of this device rivals that of traditional
catalytic materials for CO2 reduction (Ag and Au). Kanan and coworkers first reported thatelectrochemically roughened surfaces exhibited increased catalytic activity, compared to more
atomically flat surfaces.32,33,78 In principle, such roughened surfaces should have a larger
concentration of corner and edge sites, which are thought to be the catalytically active sites on thesurface. Experimental measurements of the binding motifs of CO2 at catalytically active sites onan electrode surface in situ is often hindered by the chemical environment in the condensed phase.

35This problem is further enhanced under turnover conditions where reaction intermediates may notsurvive long enough to be investigated. In this chapter, the IR photodissociation action spectra of[Sn(CO2)n]‒ are reported. The structure of the core ions will be discussed along with the effects
the solvent molecules have on the core ion’s structure and charge distribution.
4.2 Results and Discussion
The experimental spectra of [Sn(CO2)n]‒ (n = 2‒6) are shown in Figure 4.1. As mentioned
previously, vibrational signatures in the region from 1000 cm-1 to 2150 cm-1 are due to reduced
CO2 stretching motions of the molecular core ions. Signatures in the region from 2175 cm-1 to2450 cm-1 can be assigned to asymmetric CO2 stretching modes of solvent CO2 species.
Vibrational bands in this region blue shift by approximately 3 cm-1 as cluster size increases, similar
to the behavior observed in other M-CO2‒ cluster species (M = Au, Ag, Cu, Co, Ni, Bi, Mn, Fe,Ti).29-31,69,70,79-82 The nature of the blue shift is due to a decreasing influence of the negative coreion on the solvent species as cluster size increases.

36

Figure 4.1. Infrared photodissociation spectra of [Sn(CO2)n]‒ (n = 2‒6). Cluster size is denotedby the numbers between the two panels. The red dashed line gives the position of the asymmetricstretching mode of free CO2. Each trace is individually normalized to the highest intensity
feature.

The vibrational modes of the core ion cluster can be described as asymmetric andsymmetric stretching motions of negatively charged CO2 ligands bound to the metal atom. The
magnitude of the red shift of the asymmetric CO2 stretching frequency of metal bound CO2 species
compared to neutral CO2 species can be related to the amount of charge shared from the metal ion
to the CO2.29-31 As described in Chapter 3, the presence of excess charge on the CO2 species results
in a decrease in the OCO bond angle away from 180˚.11,12,26
An initial inspection of the experimental spectra will immediately reveal the large change
in spectroscopic signatures that occurs between cluster size n = 2 and n = 3. The appearance ofmany new experimental signatures in the region from 1650 – 1750 cm-1 with increasing cluster

37

size is often due to the experiment becoming sensitive to larger core ion geometries. As describedin Chapter 2, IR photodissociation spectroscopy requires the presence of a weakly bound solvent
species in these metal-CO2 cluster ions in order for an absorption event to be detected on the timescale of the experiment. There is not a large difference between the experimental spectrum ofcluster size n = 3 and the spectra for larger cluster sizes besides some changes in the relative
intensities of the signatures. This suggests that the largest core ion geometry present in themolecular beam has a chemical formula of [Sn(CO2)3]‒.
As described in Chapter 3, the intensity of infrared signatures below 1500 cm-1 are
artificially suppressed due to the energy of the photon being lower than the binding energy of a
weakly bound CO2 species. When the energy of an absorbed photon is lower than the binding
energy of a solvent species, only the cluster ions with internal energy above a critical level will
evaporate a solvent species on the timescale of the experiment. The internal energy of a cluster canbe estimated using an evaporative ensemble (see Chapter 2). Briefly, the internal energy content
of the cluster is estimated to be on the order of the binding energy of the last solvent species toevaporate from it during cluster formation.48 Calculations indicate that the binding energy of asolvent molecule to the molecular core ion is dependent on solvation position, particularly forsmall clusters. On average, the binding energy of a CO2 solvent species is around 1500 cm-1. For
small cluster sizes, the intensity of features below 1500 cm-1 will therefore be suppressed in theexperimental spectrum.
Just as in the case of Bi-CO2 cluster anions, Sn-CO2 clusters can exist in two different spinstates. The 5s24d105p3 valence electron configuration of a Sn anion implies that Sn-CO2 complexes
can exist in either a doublet or quartet spin state. Photoelectron spectroscopy studies of the Sn ‒anion confirm the expected ground electronic state of the anion as a 4S state.83 Quantum chemical

38

calculations for [Sn(CO2)n]‒ consistently show that doublet spin structures are 400 meV lower inenergy than similar clusters with quartet spin. In addition, several stable cluster geometries withdoublet spin dissociate into a Sn‒ solvated by multiple neutral CO2 species when calculated withquartet spin. It is likely that the discrepancy between the ground state spin of the Sn anion and thespin state predicted by our calculations for [Sn(CO2)n]‒ clusters is due to the interaction between
the Sn atom and the CO2 ligands breaking the symmetry of the 5p orbitals on the metal. This would
allow the doublet state to be stabilized in the complex. This behavior is analogous to Snn‒ clusters,
where already the dimer, Sn2‒, has a doublet ground state.84,85
A summary of the calculated structures with chemical formula [Sn(CO2)2]- along withrelative energies is shown in Figure 4.2. Inspection of the experimental spectrum for [Sn(CO2)2]‒
reveals two high intensity features at 1640 cm-1 and 1840 cm-1 with lower intensity features at1269 cm-1 and 1178 cm-1. The asymmetric stretching region of neutral CO2 also features a very
low intensity signature at 2333 cm-1 which signals that some of the clusters at this size must becomposed of a [SnCO2]‒ core ion and a solvent CO2 species.

39

Figure 4.2. Calculated cluster geometries of [Sn(CO2)2]‒. Calculations were performed based on
the methods described in Chapter 2. Clusters are labeled based on core ion geometry (capitalletter), number of solvent CO2 species (number) and solvent configuration (lower case letter)according to increasing relative energy.

The lowest energy structures at this cluster size feature a Sn metal center bound to a planar
oxalate ion [A(0)] similar to the structures observed in [Bi(CO2)n]‒ cluster anions (Chapter 3).81
The next highest energy cluster configuration also features a C2O4 ligand [B(0)], except that it isbound to the metal in a nonplanar configuration. Since we require the presence of at least one
weakly bound CO2 molecule in order to detect absorption events, structures A(0) and B(0) can besafely excluded from contributing to the experimental spectrum of [Sn(CO2)2]‒. We can do thisbecause the binding energy of CO2 ligands to the metal center are much higher (of the order of

40

104 cm-1) than the photon energies used in the experiment. It is highly likely that these two isomers
are present in the molecular beam at this cluster size. The calculated vibrational spectra of theremaining structures are compared to the experimental spectrum in Figure 4.3.

Figure 4.3. A comparison of calculated vibrational spectra (red) with the experimental spectrumof [Sn(CO2)2]‒. The intensity of the solvent signatures is approximately 6‒10 times weaker than
the feature at 1850 cm-1.

The remaining two cluster structures feature bidentate ɳ2-(C,O) CO2 ligands with either a solvent
CO2 species at the metal or behind the ligand [C(1a) or C(1b)] or accompanied by a monodentate
ɳ1-C bound CO2 ligand similar to that discussed in Chapter 3 [D(0)]. The presence of these highenergy isomers at this cluster size suggests that they are kinetically trapped species, indicating thatsince the collisional cooling rate of the ions is faster than the isomerization rate. Based on thecalculated vibrational frequencies, the experimental signatures are assigned as follows (see Table
4.1 for a summary).

41

Table 4.1. Spectral assignments of the experimental bands at n = 2 compared to predictedvibrational frequencies.
observed [cm-1]amainly contributing core ion and vibrational modepredicted [cm-1]

1178

D, symmetric CO stretch of 2-(C,O) ligand

1092

1269

D, symmetric carboxylate CO stretch

1196

1640

D, asymmetric CO stretch of 2-(C,O) ligand

1698

1850

D, antisymmetric carboxylate CO stretch

1898

2333

C, antisymmetric CO stretch of solvent

2326a

Average value of isomers C(1a) and C(1b).

The ɳ2-(C,O) binding motif is common to core ion structures C and D. The asymmetric
CO2 stretching motion of this ligand is predicted to be close in frequency to the experimental
signature at 1640 cm-1. We assign this band to the asymmetric CO2 stretching mode of a ɳ2-(C,O)
CO2 ligand. The symmetric CO2 stretching mode of this ligand is predicted to appear around 1100
cm-1 and is likely responsible for the weak spectral signature at 1178 cm-1. It is evident thatvibrational motions of the ɳ2-(C,O) ligand do not account for the experimental signatures at 1850
cm-1 and 1269 cm-1.
Comparison of the calculated vibrational frequencies of structure D(0) with theexperimental spectrum shows that one of the vibrations of this structure recovers the experimental
signature at 1850 cm-1. We assign this experimental feature to the asymmetric CO2 stretchingmotion of the carboxylate ligand. The symmetric CO2 stretching mode is assigned to theexperimental band at 1269 cm-1. Based on the assignments discussed, it is clear that structure D(0)42

is the main contributor to the experimental signatures at this cluster size. This is contrary to thepresence of the vibrational signature at 2333 cm-1, which indicates the presence of a weakly boundCO2 solvent species. Since structure D is not compatible with this signature, there must be someamount of structures C(1a) or C(1b) also present in the molecular beam. The low signal-to-noise
ratio of the peak at 2333 cm-1, especially compared to analogous transitions at the same cluster
size in Chapters 3 and 5, suggests that a very small fraction of the experimentally observable ion
population have a solvent CO2 molecule. The determination of a ratio between clusters with core
ion C or D present in the molecular beam is not straightforward since there is no spectral overlap
between the two spectral regions examined. Based on the raw ion intensities and typical laser pulse
energies, we estimate that the band at 2333 cm-1 is roughly a factor of 6-10 times weaker than theband at 1850 cm-1. This makes it very unlikely that we would be able to observe vibrationalsignatures of the corresponding core ions in the experimental spectrum.
The ability to observe photodissociation of cluster structure D(0) is a testament to how
weakly bound the carboxylate ligand is to the metal center. The calculated bond dissociationenergy of the Sn-C bond of the carboxylate ligand is around 1500 cm-1. This suggests that a single
infrared photon in the wavenumber range of this experiment should be able to dissociate that bond.
The intensities of vibrational signatures with lower wavenumber will be suppressed as discussed
above.
Similar to the case of Bi-CO2, the interaction of the ɳ1-C bound CO2 ligand with the metal
results in charge donation from the metal center to the ligand. The presence of excess electronic
charge lengthens the C‒O bonds and results in a red shift of the asymmetric CO2 stretching modecompared to the same mode in neutral CO2. Compared to other ɳ1-C bound CO2 ligands, the Sn-C
bond is about a factor of two weaker than the bonds observed in Au-CO2 cluster anions.72,86 This43

is likely due to the presence of the ɳ2-(C,O) ligand. Natural population analysis shows that amajority of the excess charge (around -1.0 e) is localized onto the ɳ2-(C,O) ligand in structure
D(0). The carboxylate ligand has -0.41 e of charge localized onto it. This leaves an overall cationic
Sn atom. In structure C(1a) the Sn atom is effectively neutral with nearly all of the electron density
residing on the ɳ2-(C,O) ligand.
There is a remarkable similarity between structures C(1b) and D(0), where the only
difference is the distance between the Sn and the carbon atom of the weakly bound CO2 species.
In order to understand the barrier toward interconversion between these two structures, we
calculated a potential energy curve along the Sn-C bond distance (Figure 4.4).

Figure 4.4. Calculated potential energy curve connecting structures D(0) and C(1b) along the
Sn‒C bond distance. All energies are relative to the lowest energy calculated data point and arecorrected for zero point motion except for the Sn‒C stretching mode.

It is clear that structures C(1b) and D(0) represent two minima along this potential energy surface.
Structures C(1b) and D(0) represent CO2 that is physisorbed and chemisorbed onto the SnCO2
complex anion, respectively. The existence of physisorbed and chemisorbed CO2 onto a Au anion
has been observed by Bowen and coworkers using photoelectron spectroscopy.87 The calculated

44

barrier to interconversion from D(0) to C(1b) is approximately 6 meV. This value is very low andwould most likely correspond to rapid interconversion between structure D(0) and structure C(1b).
One would therefore expect for the experimental spectrum to exhibit a broad spectral featurecorresponding to species in the process of converting from structures D(0) and C(1b). However,
the experimental spectrum has fairly narrow vibrational signatures corresponding only to structure
D(0). A plausible explanation is that the barrier height for the interconversion of these two isomers
is drastically underestimated by the calculation. Another explanation is that structure C(1b) simply
does not exist, and the spectral feature at 2333 cm-1 is due solely to structure C(1a).
We note that the calculated values of the C‒O bond lengths and the OCO bond angle varysmoothly along the Sn‒C bond length, suggesting a gradual conversion of structure D(0) tostructure C(1b). Examination of the highest occupied molecular orbitals (HOMOs) of structures
D(0), C(1b) and the transition state show a gradual elongation of one of the Sn 5p orbitals as theCO2 species moves closer to the metal center. Upon bond formation, there is strong mixing
between the 5p orbital of the Sn with one of the π orbitals of the ligand CO2. (See Figure 4.5).

Figure 4.5. Calculated HOMO contours for different positions along the potential energy curve
described in Figure 4.4. Contours are 70% of maximum orbital amplitude.

45

It is likely that the experimental spectrum of [Sn(CO2)2]‒ does not represent the overall distributionof different cluster structures, since the lowest energy isomers will not dissociate after absorbinga single infrared photon. The clusters that do dissociate at this cluster size provide insight into how
CO2 molecules interact with an undercoordinated Sn atom in the presence of an excess electron,i.e., under reducing conditions. As stated previously, gas phase cluster ions are interesting model
systems for the interaction of CO2 with strongly charged corner sites. The presence of isomer D(0)in the molecular beam demonstrates binding behavior that is unique to the Sn-CO2 system and may
be related to the efficiency of CO2 reduction at Sn based electrodes.
Adding another CO2 molecule to the clusters results in a large change in the experimentallyobserved infrared spectrum. The experimental spectrum contains three high intensity bands at
1625, 1678, and 1721 cm-1 with lower intensity bands at 1299 and 1323 cm-1. The loss of theexperimental signature at 1850 cm-1 suggests that clusters with core ion structure D are no longer
major contributors to the experimentally observed signatures. The overall increase in the signalto-noise ratio of the CO2 solvent signature compared to cluster size n = 2, implies that most if not
all of the structures contributing to the experimental spectrum contain at least one weakly boundCO2 species. A summary of the calculated structures and relative energies at this cluster size isshown in Figure 4.6. A comparison of the calculated vibrational spectra to the experimental
spectrum is provided in Figure 4.7, and a summary of the spectral assignments can be found in
Table 4.2.

46

Figure 4.6. Calculated cluster structures for [Sn(CO2)3]‒. The cluster labeling procedure is
described in Figure 4.2, and calculations were performed as described in Chapter 2.

47

Figure 4.7. Comparison between the experimental spectrum of [Sn(CO2)3]‒ (black) andcalculated vibrational spectra (red). Calculated traces are labeled by cluster structure. The region
from 1816‒1975 cm-1 is enhanced by a factor of 5 to highlight two low intensity bands in that
region.

48

Table 4.2. Spectral assignments for n = 3 compared to predicted frequencies.
observed [cm-1]amainly contributing core ion and vibrational modepredicteda [cm-1]

1299

A, in-phase combination of metal-bound CO stretches

1286

1323

B, out-of-phase combination of symmetric CO2 stretches

1297

1625

C, asymmetric CO2 stretch, structure C(2a)

1616

1678

B, antisymmetric stretch of free CO2 subunit

1664

1721

A, in-phase combination of free CO stretches

1698

1850

D, antisymmetric carboxylate CO2 stretch, structure D(1b)

1855

1945

D, antisymmetric carboxylate CO2 stretch, structure D(1a)1929

Unless assigned to specific solvent isomers, average values of different solvent isomers arereported

As stated above, the loss of the spectral feature at 1850 cm-1 in the spectrum of [Sn(CO2)3]‒
would seem to suggest that structures D(1a) and D(1b) are not present at this cluster size. However,a close examination of the experimental spectrum of [Sn(CO2)3]‒ reveals two very weak featuresat 1850 and 1945 cm-1 that correspond to the asymmetric CO2 stretching mode of the carboxylate
ligand of core ion D. The position of this band is directly related to the position of the solvent
species relative to the carboxylate ligand, and results in shifts similar to other metalloformate
complexes.29-31,81 The loss of intensity of these modes informs on the stability of structures D(1a)and D(1b) compared to structure C(1b). Solvation of structure D(0) at the ɳ2-(C,O) ligand resultsin charge polarization out of the carboxylate ligand. This further weakens the already very weakSn‒C σ bond of the carboxylate ligand, making it more likely that isomerization to structure C(1b)occurs.

49

The experimental signature at 1625 cm-1 can in principle be due to clusters with core ion
C or core ion D. Given the discussion above and the intensity of the signatures at 1850 and 1945
cm-1, it is unlikely that structures D(1a) or D(1b) contribute significantly to the intensity of thefeature at 1625 cm-1. Structures C(2a) and C(2b) each represent different solvation positions
around core ion C, where both solvents are behind the metal bound CO2 ligand, or there is one
solvent each at the ligand position and beside the metal atom. Based on the calculated vibrational
spectra we assign the feature at 1625 cm-1 to the asymmetric stretching motion of the ɳ2-(C,O)
ligand of core ion C. Specifically, cluster structure C(2a) matches the position of this band best.
The same vibration in structure C(2b) most likely contributes to the experimental signature at 1678
cm-1.
The addition of a third CO2 molecule to the cluster allows for larger core ion structures tobe solvated. This means that core ion structures A and B are observable. The individual CO2
subunits of the C2O4 ligand in core ion structure A are equivalent. This equivalence results incoupling of free and metal bound C‒O stretching motions into in-phase and out-of-phasecombinations. We assign the vibrational signature at 1721 cm-1 to the in-phase combination of freeC‒O stretching motions of structures A(1a) and A(1b). Similar to the case of Bi-oxalate
complexes, the out-of-phase combination of free C‒O stretching motions is not predicted to havesignificant infrared intensity. We assign the in-phase combination of metal-bound C‒O stretchingmotions to the experimental signature at 1299 cm-1.
In core ion structure B, rotation about the C‒C bond is not hindered by the presence of themetal atom. This results in a twisted configuration where the partially negative oxygen atoms repel
each other, resulting in a rotation of the free CO2 moiety of 90˚. This core ion geometry is similar
to the observed structure of MgCO2Cl‒ examined by Asmis and coworkers.88 In that structure, the50

CO2 ligand binds to the Mg atom through both of its oxygen atoms; this results in a carbene typestructure featuring a highly nucleophilic carbon atom. Structure B likely comes about as a resultof a nucleophilic attack by this carbon atom onto the carbon atom of a second CO2 molecule (thecarbon atom in neutral CO2 is slightly positively charged). Unlike in core ion structure A, the two
CO2 subunits that make up the C2O4 ligand are not equivalent. In fact, there is a significant
difference in the OCO bond angle of the free CO2 subunit (132˚) and the metal bound CO2 subunit(116˚). This difference in OCO bond angle results in a predicted splitting of the asymmetric CO2
stretching modes of 279 cm-1. This is remarkable, considering that each CO2 subunit has
approximately the same amount of charge associated with it. We assign the experimental signature
at 1678 cm-1 to the asymmetric CO2 stretching motion of the free CO2 subunit. The signature of
the asymmetric CO2 stretching motion of the metal bound CO2 subunit is predicted to occur around1385 cm-1. The experimental signature corresponding to this band is likely suppressed due to thebinding energy of a CO2 solvent species to this core ion. However, a compatible experimental
feature does appear for larger cluster sizes. Surprisingly, unlike the asymmetric stretching motions,the symmetric CO2 stretching motions are strongly coupled between the two CO2 subunits. We
assign the experimental signature at 1323 cm-1 to the out-of-phase combination of symmetric CO2
stretching motions. The in-phase combination of these motions is predicted to have no infraredintensity and is not observed in the experimental spectrum.
Unlike the behavior observed for [Bi(CO2)n]‒ cluster ions, there does not appear to be anychange in position of any of the signatures as cluster size increases.81 There is a reduction of theintensity of the band at 1625 cm-1 that could correspond to isomerization of core ion structure C
into either core ion A or B.

51

The observation of high-energy isomers is an indication that these isomers are kineticallytrapped due to significant barriers toward isomerization. This behavior has been seen previously
in other metal-CO2 cluster ions.31,69,70,82 In order to estimate the barrier toward isomerization forstructure B to structure A, exploratory calculations were performed along the Sn‒C‒O bond angle.
Based on these calculations, the barrier toward isomerization of core ion structure B to structure
A is predicted to be approximately 725 meV.
4.3 Conclusions
The structures of core ions found in [Sn(CO2)n]‒ cluster ions were determined using
photodissociation spectroscopy in concert with quantum chemical calculations. For [Sn(CO2)2]‒
clusters, the dominant core ion geometry that contributes to the experimental spectrum is a uniquestructure featuring two CO2 units bound to the metal center with ɳ2-(C,O) and an ɳ1-C binding
motifs. The ɳ1-C CO2 ligand is very weakly bound to the metal atom (bond dissociation energy
predicted to be approximately 1500 cm-1). For clusters with n ≥ 3, a plurality of clusters arecomprised of planar and nonplanar Sn-oxalate core ions and a ɳ2-(C,O) Sn-CO2 core ion. For large
cluster sizes, it is likely that the ɳ2-(C,O) Sn-CO2 core ion isomerizes into either of the Sn-oxalate
structures. Unlike Bi-CO2 cluster ions, Sn-CO2 core ions do not readily form metalloformate
structures, despite both Bi and Sn being p-block metals.81

52

Chapter 5Interactions of CO2 with Fe and Mn: Formation of Complex and Diverse Core Ions

This chapter has been adapted with permission from the following previously published articles:
Thompson, M. C., Dodson, L. G., Weber, J. M., “Structural Motifs of [Fe(CO2)n]− Clusters(n = 3–7).” Journal of Physical Chemistry A, 2017. 121: p.4132‒4138. DOI:
10.1021/acs.jpca.7b02742. Copyright 2017 American Chemical Society.
Thompson, M. C., Ramsay, J., Weber, J. M., “Interaction of CO2 with Atomic Manganesein the Presence of an Excess Negative Charge Probed by Infrared Spectroscopy of [Mn(CO 2)n]−Clusters.” Journal of Physical Chemistry A, 2017. 121: p.7534‒7542. DOI:
10.1021/acs.jpca.7b06870. Copyright 2017 American Chemical Society.

5.1 Introduction
The interaction of CO2 with transition metals is important both in the context of CO2
reduction catalysis as well as from a fundamental perspective.4,89,90 The electrochemical reductionof CO2 has many applications towards the sequestration or recycling of CO2 in a carbon neutral
fuel cycle.5,91,92 While the coinage metals have shown promise as materials for heterogeneous
catalytic conversion of CO2 to CO, transition metals outside of group 9 have not displayed theselectivity of these materials for CO formation. In many cases, homogeneous catalysts containing
first row transition metals have higher selectivities for CO formation compared to heterogeneous
analogues.4
The continued interest in first row transition metal based catalytic materials is due to theirrelatively large abundances compared to traditional electrode materials. Infrared studies of cationic
metal‒CO2 systems have shown primarily metal oxygen interactions between the metal cation andCO2 neutral species.93-100 These interactions are largely electrostatic in nature and often result in
negligible charge transfer to the CO2 species. Anionic clusters of a metal atom and CO2 exhibit a53much richer collection of interactions. In many cases, strong M‒C and M‒O bonds are formed
resulting in a wide variety of structural motifs.31,79,80 In contrast to earlier work by Knurr et al.,79,80
the systems discussed in the following sections exhibit many different ligand binding motifs
between metal and CO2 ligands.
5.2 [Fe(CO2)n]‒ Results and Discussion
In an effort to characterize the interactions of CO2 with an Fe atom in the presence of anexcess electron, photodissociation spectroscopy was performed on [Fe(CO2)n]- cluster ions . The
experimental spectra are presented in two spectral regions (990‒2150 and 2175‒2450 cm-1, see
Figure 5.1); signatures in the higher wavenumber region are due to CO2 species that are solvating
the core ion, while signatures in the lower wavenumber region are primarily due to reduced CO
stretching motions of CO2 ligands in the core ion. As discussed previously, the presence of a
weakly bound solvent species is necessary for our technique to detect laser-induced
photodissociation on the timescale of the experiment. As observed in [M(CO2)n]‒ (M=Au, Ag, Cu,
Co, Ni, Sn, and Bi), CO2 ligands are bound from anywhere between 4000 cm-1 to as high as 104
cm-1.29-31,79-81 This precludes observation of laser-induced dissociation of CO2 ligand species at thephoton energies used in this experiment. There was no observed photodissociation action fromcluster sizes smaller than n = 3. This suggests that the smallest stable core ion present in themolecular beam has stoichiometry [Fe(CO2)2]-.

54

Figure 5.1. Experimental spectra of [Fe(CO2)n]‒ cluster ions. The numbers along the right handside indicate the number of CO2 species in the cluster (n). Spectra in the asymmetric stretchingregion of neutral CO2 were taken over a single day and are shown to demonstrate the presence ofat least one solvent CO2 species.

55

As before, the interpretation of the experimental spectra is aided by quantum chemical
calculations. A series of cluster geometries were generated based on chemical intuition and thenrelaxed. Assignments were made by comparison of calculated vibrational frequencies to theexperimental spectrum. A set of viable calculated cluster structures is shown in Figure 5.2.

56

Figure 5.2. Calculated structures of [Fe(CO2)3]‒ (left) and [Fe(CO2)4]‒ (right). Calculations were
performed as described in Chapter 2. Clusters are labeled based on core ion structure (capitalletter), number of solvent species (number in parentheses) and solvent position (lower case letterin parentheses). All calculated structures are in the quartet spin state.

57

Based on the electron configuration of the Fe‒ (4s23d7), Fe‒CO2 complex ions can have two
possible spin states, doublet or quartet. Calculations show that doublet spin complexes lie
approximately 1 eV higher in energy than quartet spin structures with the same geometry. Thissuggests that doublet spin complexes can relax through a spin-flip to a quartet state without large
geometry changes. Similar transitions occur in other metal-organic complexes with lifetimes up toa microsecond.101 An alternative is that the ions are formed via an Fe‒ anion in its 4F groundstate.102 The relatively short spin flip transition lifetime suggests that doublet Fe‒CO2 complexes
formed in the ion source will not survive between formation and irradiation (milliseconds).
The smallest cluster that displays photodissociation action is [Fe(CO2)3]‒. The
experimental spectrum (Figure 5.3) for this cluster features a group of peaks in the region from1600‒1750 cm-1, a signature around 1900 cm-1, and several features below 1300 cm-1. The twolarge features at 1699 and 1749 cm-1 are assigned to the in-phase and out-of-phase combination ofasymmetric CO stretching vibrations in structures B(1a) and B(1b). Core ion geometry B features
two CO2 ligands each bound to the metal atom through the carbon atom and one of the oxygen
atoms [ɳ2-(C,O)]. This core ion has C2v symmetry, and its “butterfly” geometry has been seen
previously in studies of CO2 interactions with Co, Cu, and Ni.31,79,80 As discussed in Chapter 1,
when CO2 accommodates excess charge, there is a large red shift in the asymmetric stretchingfrequency. Since the assigned CO stretches are red shifted by about 700 cm-1 compared to theneutral asymmetric CO stretching motion, the CO2 ligands must be partially reduced. Natural
population analysis confirms this assumption, finding that each CO2 ligand contains about -0.89 eof charge (in the unsolvated complex). Kinetic coupling between individual CO2 oscillations
results in in-phase and out-of-phase combinations of CO2 stretching motions in the complex. The
higher energy feature (1749 cm-1) is assigned to the in-phase combination of asymmetric CO2

58

stretches, and the lower energy feature (1699 cm-1) is due to the out-of-phase combination ofasymmetric CO2 stretches. Calculations show a third high intensity feature should be present
around 1042 cm-1, corresponding to the in-phase combination of symmetric CO2 stretchingmotions. We assign this motion to the experimental signature at 993 cm-1. As discussed in Chapter
3, features with energy far below the binding energy of a solvent CO2 species (1400 cm-1) do notreadily dissociate a solvent species after absorption of a single photon in the time scale of theexperiment. Only a small portion of the ion population will contain enough internal energy (after
photon absorption) to dissociate on the timescale of the experiment. The suppression of theintensities of these features should diminish with increasing cluster size.

Figure 5.3. Comparison of the experimental spectrum of [Fe(CO2)3]‒ (black, top) to calculated IR
spectra (red). Each trace is labeled according to the corresponding calculated structure (seeFigure 5.2).

59

The feature at 1900 cm-1 is assigned to the CO stretching motion of structure A(1). The
core ion consists of a carbonate ligand bound to the metal along with a carbonyl ligand on the
opposite side. This geometry is only possible via insertion of the metal atom into a CO bond of
CO2. The CO2 molecule has a CO bond energy of 525.9 kJ/mol,103 so one would expect any
insertion of a metal atom into one of the CO bonds would feature a fairly high energetic barrier.
However, work by Dodson et al. has demonstrated that insertions of a Ti atom into a C‒O bond of
CO2 occurs with a negligible barrier.82 Unlike other transition metal-CO2 systems studied, thisinsertion signature has significant intensity, indicating that structure A(1) is not a minor contributor
to the experimental spectrum. There is an experimental signature at 1646 cm-1, which is consistent
with the calculated frequency of the asymmetric CO2 stretching motion of the carbonate ligand of
structure A(1).
A third possible core ion goemetry is an oxalate complex similar to others described
previously in this thesis. Structures C(1a) and C(1b) represent two different clusters with thesolvent species being to the side of the core ion or above the core ion, respectively. Calculated
vibrational frequencies suggest that one of the two weak shoulders at 1653 and 1676 cm-1 could
account for the symmetric combination of free CO stretches in the oxalate complex. Just as before,
the asymmetric combination of free CO stretches is predicted to have negligible infrared intensity.
The absence of a signature around 1350 cm-1 (corresponding to the symmetric combination of
metal bound CO stretches) could be due to a suppression of this feature due to the binding energyof a solvent CO2 species to this core ion. The author cannot rule out the presence of structures
C(1a) or C(1b) in the molecular beam at this cluster size.
The coexistence of structural families A and B merits further discussion. In the absence ofhigh barriers to isomerization, one would expect, on the timescale of the experiment, for all of the60clusters to relax to the global minimum geometry configuration, followed by CO2 solvent
evaporation as necessary to accommodate the resulting heat of reorganization. An examination of
the relative binding energies of the bare complexes reveals that structure A(0) is roughly 1 eVlower in energy than structures B(0) and C(0). An evaluation of the bond insertion process with asmaller test system reveals a barrier to metal insertion of around 1.4 eV.82 It is not surprising tofind higher energy isomers populated for this system due to kinetic trapping. If the collisionalcooling rates in the cluster source are faster than the isomerization rate, higher energy isomers canbecome kinetically trapped in higher lying potential wells.
Due to our detection scheme, the clusters need to have at least one weakly bound solvent
species in order for absorption events to be detectable. For [Fe(CO2)3]‒ clusters, our experiment is
not sensitive to core ions that have molecular formula [Fe(CO2)3]‒. This means that if these coreions are present in the molecular beam, they would only be detectable in the experimental spectraof larger cluster ions (n ≥ 4). Examination of the experimental spectrum of [Fe(CO2)4]‒ reveals theappearance of several new features that could be due to larger core ion structures becoming
detectable at this cluster size (Figure 5.4). Starting with signatures already identified as being
vibrational modes of smaller core ion geometries, butterfly based clusters (B) and insertion based
structures (A) are clearly still present. In these cases the positions of these signatures are slightly
shifted due to solvent position affecting the charge distributions of the core ions.

61

Figure 5.4. Comparison of the experimental spectrum of [Fe(CO2)4]‒ (black, top) to calculated IR
spectra (red). Each trace is labeled according to the corresponding calculated structure (seeFigure 5.2).

The viable calculated cluster geometries for this cluster size are shown in Figure 5.2. The lowest
energy isomer is structure D(1). This geometry is very similar to structural family C, except for athird CO2 species interacting with the metal atom. This CO2 forms two bonds to the metal center,through the C atom and one of the O atoms. The ɳ2 (C,O) ligand ends up with a smaller proportionof charge on it compared to the same binding motif seen in the butterfly complexes. This resultsin structural family D being identifiable from family C by the presence of an asymmetric CO
stretching signature around 1800 cm-1. This is compatible with the experimental signature foundat 1799 cm-1. The other prominent new experimental signature at 1313 cm-1 can be assigned to the62symmetric combination of metal bound CO stretches in the oxalate ring, and is compatible withstructure D(1) or possibly structures C(2a) and C(2b). The presence of the bidentate CO 2 ligand
in structure D(1) results in a calculated blue shift of the oxalate signatures (as shown in thecalculated spectra). This shift is due to charge migration out of the oxalate ring and into thebidentate CO2 ligand. The bidentate ligand contains a smaller amount of charge (-0.68 e) comparedto similar ligands in other complexes (-0.89 e). The relative position of the bidentate CO2
asymmetric stretching signature of structure D(1) compared to comparable signatures in structural
family B reflect this.
Several smaller features in the region from 1800‒1900 cm-1 are seen in the experimental
spectrum for [Fe(CO2)4]‒. These features can be attributed to combinations of asymmetric CO
stretching motions of structure E(1). This core ion geometry is a trisbidentate species with C3v
symmetry where three CO2 species are bound directly to the metal atom in an ɳ2 (C,O) motif. In
the bare complex, each CO2 ligand on average contains -0.64 e of charge. Upon solvation, theentire molecule skews toward the solvated side of the complex, breaking the C3v symmetry andresulting in the solvated CO2 ligand containing -0.68 e of charge. As was the case with the smallercluster size, the sheer number of populated isomers suggests that many of them are kineticallytrapped species.
Given the large number of spectral features and the variety of calculated core ion structures,it was useful to combine our calculated spectra to evaluate the overall fit of our calculations to theexperimental spectrum. The combined calculated curve was generated by scaling the individual
calculated traces using the experimental intensities of positively identified unique experimental
signatures assigned to each core ion. As an example, the calculated spectrum of structure D(1) was63

scaled to the intensity of the experimental signature at 1799 cm-1. The calculated spectrum for thiscluster size is compared to the experimental spectrum in Figure 5.5.

Figure 5.5. A composite of all calculated infrared spectra for [Fe(CO2)4]‒ clusters (red) comparedto the experimental spectrum (black).

Surprisingly, this method reproduces the experimental signatures and intensities very well. Thissuggests that it is unlikely that any other core ion geometry is significantly populated at this cluster
size. This also holds for larger cluster sizes, since no new signatures appear as cluster size
increases.
Unlike other metal-CO2 clusters discussed thus far, Fe-CO2 clusters exhibit a large varietyof core ion geometries. This is similar to previous studies of [Cu(CO2)n]‒, where the clusters
exhibited many diverse core ion geometries.31 An interesting note is that similar to other transition
metal-CO2 systems there were no observed core ions smaller than [M(CO2)2]‒.69,79

64

5.3 [Mn(CO2)n]‒ Results and Discussion
Laser vaporization of a manganese target entrained into a supersonic expansion of CO2
generated a series of [Mn(CO2)n]‒ (n=2‒10) cluster ions. Unlike other metals discussed in thisthesis, manganese does not exist as a stable anion (its electron affinity has never been measured
but is predicted to be strongly negative).104 If Mn‒ were to be stable, it would have a valence
electron configuration of 4s23d6, isoelectronic with neutral atomic iron. This electron configurationis compatible with three spin configurations: quintet, triplet and singlet. Calculated cluster
geometries consistently demonstrate that complexes with quintet spin configuration are the lowest
energy isomers, in agreement with Hund’s rules that would suggest 5D4 as the ground state of theatomic anion.
Infrared photodissociation action spectra of [Mn(CO2)n]‒ (n = 2‒10) are presented in Figure5.6. For the smallest cluster size, the region from 1600‒1700 cm-1 is relatively bare compared tothe spectra for larger cluster sizes. As cluster size increases, the number of features in this region
also dramatically increases for clusters larger than [Mn(CO2)3]‒. Complexity of the spectra of[M(CO2)n]‒ in this spectral range is not unusual, especially for clusters with first row transition
metals, and can be indicative of either highly polarizable core ions, many different core ion
geometries populated, or both.

65

Figure 5.6. Experimental spectra of [Mn(CO2)n]‒ cluster ions. The numbers along the right handside indicate the number of CO2 species in the cluster (n).

Spectra are assigned based on comparison of calculated vibrational frequencies of proposed
structures as described in Chapter 2. Starting with the smallest cluster size there are two features,
one at 1592 cm-1 and another at 1710 cm-1, with a broad feature between them. A summary of threestable cluster geometries, along with calculated vibrational spectra are shown in Figure 5.7.

66

Figure 5.7. Comparison of the experimental spectrum of [Mn(CO2)2]‒ (top, black) to thecalculated IR spectra of several isomers (red). Isomers are labeled based on the same method
described in Figure 5.2. Energies are relative to the lowest energy calculated isomer.

Structure A(1) consists of a single CO2 ligand bound in the ɳ2 (C,O) motif. In this particular cluster
geometry, the CO2 solvent species is directly behind the ligand. Structure B(1) consists of acarboxylate CO2 ligand similar to what has been discussed previously. As in the case of structure
A(1), the solvent species is located behind the CO2 ligand. There is no spectroscopic evidence for
either core ion species in any particular spin state. On energetic grounds, the quintet spin states are

67

roughly 1 eV lower in energy than the corresponding triplet spin analogue and around 2 eV lower
in energy than the corresponding singlet spin analogue.
Based on vibrational frequency calculations of structures A(1) and B(1), the feature at 1710
cm-1 was assigned to the asymmetric CO stretching motion of structure A(1) and we assigned thefeature at 1590 cm-1 to the asymmetric CO stretching motion of structure B(1). Structure C(1)could in principle also contribute to the peak at 1590 cm-1. Structure C(1) features a CO2 ligand
bound to the Mn atom through one of its O atoms. The ligand is effectively a CO2 anion, with amajority of the excess charge residing on the ligand. This structure is in contrast to the similarbinding method of CO2 in cationic clusters with metal atoms.93-100,105 While in theory this species
could contribute to the feature at 1590 cm-1, further discussion will show that this species does notcontribute to the experimental spectrum.
The nature of the broad signature between 1600 and 1700 cm-1 merits further discussion.
Since no individual structure seems to recover this broad feature, a potential energy scan was
performed along the C-Mn-O bond angle, calculating a potential energy curve along thiscoordinate (Figure 5.8).

68

Figure 5.8. Potential energy curve along the C-Mn-O bond angle. The Born-Oppenheimer(black), vibrationally adiabatic ground state (red), and the vibrationally adiabatic excited state(blue) surfaces are shown, along with relevant structural figures.

By freezing the C-Mn-O bond angle and allowing all other degrees of freedom to relax, we cangenerate a Born-Oppenheimer potential energy surface in order to examine possible barriers tointerconversion between the different quintet structures at this cluster size. Immediately apparent
is the flat nature of the potential between 80 and 140 degrees. This suggests that interconversion
between structures A(1) and B(1) will occur readily with no barrier. Structure C(1) has small
barrier (25 meV) toward isomerization, suggesting that it will readily isomerize into either A(1) or
B(1). This eliminates structure C(1) as a contributor to the experimentally observed signatures.
Vibrationally adiabatic surfaces were calculated according to Equation 5.1:86,106
1

Uad,0 (θ) = UBO (θ) + 2 ∑J ħωJ .

(5.1)

69

The vibrationally adiabatic ground and excited surfaces are very similar to the Born-Oppenheimer
surface. Both adiabatic surfaces were generated by summing over the two asymmetric CO
stretching modes, since they are faster than the other vibrational modes of these species. The
similarity between the three curves suggests that unresolved sequence bands could be responsible
for the broad transition seen in the experimental spectrum.
It is interesting to note that in the absence of the solvent species, structure B(0) lies at the
top of a 196 meV barrier along the C-Mn-O bond angle. Structure A(0) also lies 68 meV higher inenergy than isomer C(0). This rearrangement of the relative energy of these three core ion isomers
highlights the role of solvation in determining which core ions are observed in the experimental
spectrum.
There are several differences between the experimental spectra for [Mn(CO2)2]‒ and[Mn(CO2)3]‒. Most notably is the relative absence of the high intensity feature at 1590 cm-1 in thesmaller cluster size. Instead the spectrum of [Mn(CO2)3]‒ has two high intensity features at 1680
and 1724 cm-1, and two weaker features around 1575 and 1594 cm-1. There are also low intensity
features at 1003 cm-1, 1076 cm-1 and 1330 cm-1. At this cluster size, vibrational signatures ofmolecular core ions that consist of two CO2 ligands can now be detected. This drastically increases
the number of possible geometries that need to be examined. A selection of these geometries aswell as a comparison between calculated vibrational spectra and the experimental spectrum areshown in Figure 5.9 and Figure 5.10.

70

Figure 5.9. Calculated structures of [Mn(CO2)3]-. Calculations were performed as described in
Chapter 2. All calculated structures are in the quintet spin state.

71

Figure 5.10. Comparison of the experimental spectrum of [Mn(CO2)3]‒ (top, black) to thecalculated IR spectra of several isomers (red). Isomers are labeled based on the same method
described in Figure 5.2. Energies are relative to the lowest energy calculated isomer.

The lowest energy calculated isomer at this cluster size is structural family D. This coreion has two CO2 units forming a C-C bond and bonding to the metal atom through two oxygen
atoms. This planar core ion geometry is a commonality in metal CO2 chemistry, having been
observed in several metal-CO2 cluster systems (Fe, Ni, Cu, Bi, Sn).31,70,80,81 Natural population
analysis reveals that this core ion features a cationic Mn center (+0.5 e) with an almost dianionic

72

oxalate ligand (-1.5 e). The next higher energy structure is family E. The core ion in this case is abutterfly structure similar to what has been described in the previous section on [Fe(CO2)n]‒clusters. Structural family E appears to recover the two main experimental features at 1680 and
1724 cm-1. Also of note is that these structures also recover the two lowest energy features at 1003
and 1076 cm-1. As a reminder, due to the binding energy of solvent CO2 species to the core ionsbeing higher than the energy of these transitions, absorption of a single photon may not result in
evaporation on the timescale of the experiment for each ion. We can only detect the population ofions that have enough internal energy (combined with the absorbed photon energy) to evaporate a
CO2 species fast enough to be detected in our experiment. The low intensity feature observed at1330 cm-1 can be assigned to the symmetric combination of metal bound CO stretches in theoxalate ligand of structural family D. This signature is also suppressed due to the binding energyof the solvent CO2 species. Structures D(1a and b) also contribute to the intensity of the feature at1680 cm-1. Structure F(1) is an inverted butterfly structure with the same connectivity as core ion
E. Given the potential energy surface calculated in Figure 5.8, it is likely that this high-energy
isomer would readily isomerize to core ion structure E. Similarly, core ion structure G is also likely
too high in energy to be populated. Given the narrow breadth of the spectral features it is unlikelythat isomers E and F are coexisting in the beam. Similar to other first row transition metals, it does
not appear that different solvation positions around the core ion geometries significantly affect thecalculated vibrational frequencies. The exception is structural family D, whose asymmetriccombination of free CO stretches gain infrared intensity upon solvation of the core ion from theside [D(1b)] compared to the top of the complex [D(1a)].
The nature of the peaks at 1575 and 1594 cm-1 is ambiguous. One could assume that thesesignatures could be due to structural families B or H, however; unlike in other cluster systems

73

where these structures have been seen (Bi and Sn), there is no trace of other vibrational signaturescharacteristic of these species. While such peaks would almost certainly be suppressed by the same
effect that suppresses the other spectral features discussed above, the signatures should be apparent
given the high signal-to-noise ratio of the experimental spectrum at this cluster size. As such, theorigin of these two signatures remains ambiguous.
As mentioned earlier, Mn-CO2 anionic complexes have three compatible spin
configurations: singlet, triplet, and quintet, and unlike in the case of Fe and Bi, bare atomic
manganese does not support an excess electron. This means that beyond relative energies, there isno intuitive way to determine which spin states are likely to be populated in our molecular beam.
In the cases of structural families A, B, C, and D, there is no significant difference in the calculated
vibrational frequencies of a cluster geometry as a function of spin state. Structural family E
presents different calculated spectra as the spin state is changed. This is due to changes in the
structure of the butterfly anion, primarily in the C-Mn-C bond angle, as a function of spin. As aresult, the spin state of clusters with butterfly core ions can be determined by comparing the
calculated vibrational frequencies of complexes in different spin states with the experimental
spectrum (Figure 5.11). Direct comparison of the calculated spectra with the experimental
spectrum shows that the quintet spin structure of the butterfly complex best matches theexperimental spectrum, in both position and relative intensities. It is unlikely that other core ion
geometries will exist with different spin states since each core ion geometry has approximately thesame relative energy differences for each spin state.

74

Figure 5.11. A comparison of the calculated IR spectra of different spin states (superscript
number) of cluster geometry E(1b) [See Figure 5.7] with the experimental spectrum of[Mn(CO2)3]-.

The addition of a fourth CO2 molecule in the cluster results in an increase of the complexity
seen in the experimental spectrum. New spectral features appear at 1321 cm-1 and 1777 cm-1. The
fourth CO2 species allows for core ion structures of molecular formula [Mn(CO2)3]‒ to be
detectable. Calculated cluster geometries, along with relative energies are shown in Figure 5.12.

75

Figure 5.12. Calculated structures of [Mn(CO2)4]‒. Calculations were performed as described in
Chapter 2. All calculated structures are in the quintet spin state.

The lowest energy cluster geometry at this cluster size features an oxalate bidentate core ion(structural family I). This core ion structure is very similar to structural family D except that a third
CO2 species has bonded to the metal atom through a carbon and one of its oxygen atoms.
Comparison of calculated vibrational frequencies of clusters with this core ion, along with othercalculated structures (see Figure 5.13), reveal that the new spectral feature at 1777 cm-1 could be
due to the asymmetric stretching motion of the bidentate CO2 ligand of core ion family I. There is

76

also excellent agreement with other experimental signatures from 1600‒1750, due to in-phase andout-of-phase combinations of free CO stretching motions of the oxalate ligand. The signature at
1321 cm-1 is due to the in-phase combination of metal bound CO oscillations of the oxalate ligand.
In principle, this transition could also be due to structural family D, since there is only a slight red
shift of this signature in structures D(2a) and D(2b). The author notes that solvation of core ion
family D at the metal position results in the formation of core ion family I. This suggests that coreion family D will not be present in larger cluster sizes, since the probability of solvation at themetal center increases with cluster size, as solvation positions around the oxalate ligand are filled.

77

Figure 5.13 Comparison of the experimental spectrum of [Mn(CO2)4]‒ (black, top) to calculated
IR spectra (red).

Clusters with the butterfly (structural family E) core ion geometry are responsible for muchof the intensity in the region from 1600‒1700 cm-1. The low intensity features at 1007 and 1080
cm-1 are assigned to in-phase and out-of-phase combinations of symmetric CO stretching motions.
Interestingly, the solvation environment around this core ion results in large change in the positionof the two modes from 1600-1700. Asymmetric solvation around a single CO2 ligand localizes alarger portion of the excess charge onto this ligand. This breaks the symmetry of the charge

78

distribution and changes the coupling of the oscillators. This results in the localization ofoscillatory motion onto single CO2 units, and is accompanied by an increase of the splitting
between the two modes. Despite the fact that clusters with a butterfly core ion geometry are higher
in energy than structural families I and D, clusters with butterfly core ion geometries contributesignificantly to the experimental spectrum at this cluster size and are likely kinetically trapped.
The assumption that structural family E is kinetically trapped is based on the fact that isomerizationout of this core ion would require the Mn-C bond to break. Based on the calculated potential energy
curve described above (Figure 5.8), this process is estimated to have a barrier of at least 352 meV.
The low intensity signatures at 1562 and 1604 cm-1 could in principle be due to clusters with coreion geometry B (metalloformate) or could be due to structure H. However, since the symmetric
CO2 modes of these species are absent from the experimental spectrum it is unlikely that clusters
with these two core ion geometries are present in the experiment.
Differences in the experimental spectra of larger Mn-CO2 clusters are more subtle
compared to the changes described for smaller clusters. In addition to small changes in peak
position is the increase of the intensity of the features below 1400 cm-1. This is due to decreasing
solvent binding energies as cluster size increases. This will affect the spectra of smaller core ionsmore readily than for larger core ions due to the smaller ions having more solvent molecules at agiven cluster size compared to the larger core ions. Calculations of clusters with five CO2 species
present show the formation of a new core ion structure with chemical formula [Mn(CO2)4]- (seeFigure 5.14). Core ion J has four strongly bound CO2 ligands, two pairs of which each form a CC bond and bind into the metal center as planar C2O4 ligands. The symmetry of the core ion results
in a simple calculated vibrational spectrum, with only two signatures which are compatible withthe experimental signatures at 1330 and any of the number of peaks in the region from 1650-1750

79

cm-1 (See Figure 5.15). Two other isomers with similar structure have been calculated, with one
or both of the oxalate ligands adopting a twisted configuration (Structures K(1) and L(1)). These
structures lie at a minimum over 1 eV higher in energy than structure J and do not appear to match
any unique experimental signature. The individual oxalate ligands of core ion J each have acalculated charge of -1.29 e with the Mn having a charge of +1.60 e. While the individual oxalate
ligands are weakly charged compared to the oxalate ligand in core ion D (around -1.5 e), thecombined presence of two of them results in significant oxidation of the Mn atom. In this structure,
the charge of the Mn atom is considerably larger than the largest charge calculated for an Fe atom
coordinated to multiple CO2 species (+1.2 e). The author notes that other experimental signaturesat this cluster size are due to different solvation positions around smaller core ions previously
discussed.

80

Figure 5.14. Calculated structures of [Mn(CO2)5]-. Calculations were performed as described in
Chapter 2. All calculated structures are in the quintet spin state.

81

Figure 5.15. Comparison of the experimental spectrum of [Mn(CO2)5]‒ (black, top) to calculated
IR spectra (red).

While the asymmetric CO2 stretching region of neutral CO2 is a vitally important metric toconfirm the presence of weakly bound CO2 species in the clusters, normally this region is
uninteresting for the purpose of structural identification. Occasionally, this region can exhibit
interesting signatures that can report on the nature of the core ion’s interaction with solvent
species.29,30 Clusters smaller than [Mn(CO2)6]‒ have signatures in this region that show thepresence of at least one neutral CO2 molecule in the cluster. This signature blue shifts as cluster
size increases, and is consistent with the behavior of all other metal-CO2 anionic cluster species
studied thus far. For clusters larger than this, there is a splitting that results in two different
signatures, a signature that is red shifted from the asymmetric CO stretching frequency of 2349

82

cm-1 and a signature that is blue shifted from this value. A blue shift of the asymmetric CO
stretching frequency is commonly seen in cationic metal CO2 complexes due to complexation of
the carbon dioxide molecule and the metal cation.93-100,105 The overall shift is much weaker thanthat observed for metal bound CO2 species in anionic metal-CO2 systems because the binding is
primarily a charge quadrupole interaction between the metal cation and the CO2 with little to no
charge sharing between the two.93,107 It is theorized that the CO2 species encounters a large
repulsive potential barrier on the metal side of the complex, resulting in a blue shift of theasymmetric stretching frequency.96 With these observations in mind, it is not unlikely that a CO2
solvent species could weakly interact with a positive metal center of one of the core ions (theprobability of this increases as solvation positions around the negatively charged ligands areoccupied at larger cluster sizes). To test this theory, exploratory calculations of a [Mn(CO2)9]cluster anion with a double oxalate core ion were performed. The converged configuration has asingle solvent CO2 species out of the plane of the ligands interacting with the metal cation center
through one of the oxygen atoms (see Figure 5.16).

83

Figure 5.16. Exploratory structure and calculated spectrum for [Mn(CO2)9]‒.

5.4 Conclusions
The interactions of CO2 with first row transition metals in the presence of an excess
electron have proven to yield much richer chemistry than similar interactions with heavier group
11 or post transition metal atoms29,30,81 While the sheer number of structures can be intimidating,the key interactions can be broken down into combinations of either ɳ2 (C,O) or oxalate binding
motifs. As was the case for clusters with heavier first row transition metals,31,79,80 the predominantcore ion structure is the butterfly motif, featuring two ɳ2 (C,O) bound CO2 ligands especially atcluster sizes less that n = 4. For [Fe(CO2)n]‒ clusters the butterfly structure is expanded to accept
a third ɳ2 (C,O) bound CO2 ligand; this results in a complex with C3v symmetry that is unique to84

the Fe-CO2 cluster anion. The second most common ligand seen in first row transition metalclusters with CO2 (and more generally in other metal-CO2 anionic systems) is the C2O4 motif
especially at larger cluster sizes. In many cases, the C2O4 ligand is paired with an ɳ2 (C,O) boundCO2 ligand generating a characteristic vibrational signature seen in the spectra of [M(CO 2)n]‒(where n ≥ 4) clusters. For [Mn(CO2)n]‒ clusters the C2O4 ligand is observed to be paired with asecond oxalate ligand.

85

Chapter 6
Enhancement of the Intensity of the Symmetric CO2 Stretching Transition in MCO2‒
Complexes (M = H, Ag, and Bi) Compared to CO2This chapter has been adapted with permission from the following previously publishedarticle:Thompson, M. C. and Weber, J. M., “Enhancement of Infrared Activity by Moving
Electrons through Bonds ‒ The Case of CO2 anion and Carboxylate.” Chemical Physics Letters.
2017 683: p. 586-590. DOI: 10.1016/j.cplett.2017.01.060. Copyright Elsevier.

6.1 Introduction
One of the most basic measurements in spectroscopy is the intensity of a vibrationaltransition. Within this value are encoded many molecular properties, the most important of which
is how the charge distribution of the species varies with structural changes during vibration. This
⃗ (𝑄)
𝜕𝜇

principle is enshrined in the selection rule for infrared activity, ( 𝜕𝑄 ) ≠ 0.
0

The driver of infrared intensity is the vibrational transition dipole moment ⟨𝑣 ′ |𝜇 (𝑄)|𝑣′′⟩
coupling vibrational levels v ' and v" of a normal mode Q. This is often understood with a fixedcharge approximation, which assumes that the atomic charges are constant over motion along thenormal mode coordinate. This approximation, however, does not hold true in all situations; in somecases, the intensity of a particular infrared transition defies what can be reasonably expected basedon the fixed charge approximation. One prominent example is seen in the intensity ratio of thesymmetric and asymmetric CH2 stretching motion in CH2X radicals (X = F, Cl) as measured by

86

Nesbitt and coworkers.108,109 They found that the CH stretching motion was strongly coupled tothe charge distribution of the radical species. The flow of charge along the normal mode
coordinates runs counter to the direction of the transition dipole moment approximated with thefixed charge model. This reduces the magnitude and even reverses the direction of the overall
transition dipole moment. Nesbitt and coworkers found that the intensity of the symmetric CH2
stretching motion was enhanced by “charge sloshing” along the mode, while the asymmetric CH2
stretching motion is suppressed by the same effect.108,109 This results in a reversal of the expectedratio of the intensities of these two modes compared to what is expected based on a fixed charge
model.
While the term charge sloshing certainly evokes the correct picture of the physical process
occurring, for the remainder of this chapter this process of vibrationally mediated electron density
modulation will be referred to as charge oscillation. Charge oscillation has been observed inseveral ionic systems.110-112 In this chapter, charge oscillation in several MCOO- species (M = H,Ag, and Bi) with be contrasted with the photodissociation spectrum of solvated CO2-.
6.2 Theoretical Methods
Many of the calculations were performed with the same procedure as described in Chapter
2. Briefly, calculations were performed using density function theory as implemented in the
TURBOMOLE suite.55 Calculations were done with a dispersion corrected B3LYP functional56
with def2-TZVPP basis set58 with effective core potentials for the Bi atom.113 In order tocharacterize how the charge distribution of a species changes along the symmetric CO2 stretching
coordinate, geometries along the symmetric and asymmetric CO2 stretching coordinates weregenerated by distorting the molecular geometries along each coordinate using the VIBRATION
program. A single point calculation was performed for each geometry, and the charge of each atom

87

was assigned using several methods. Charge analysis using Mulliken, Loewdin, Natural
Population Analysis, and population analysis based on orbital occupation numbers withmulticenter correction (PABOON114) were compared, and the PABOON method was found toyield the best agreement with the overall dipole moment calculated obtained by density functional
theory. The charges on the COO‒ moieties were calculated by adding the charges on the carbon
and both oxygen atoms.
6.3 Results and Discussion
We set out to compare and contrast the infrared intensity of symmetric and asymmetric
CO2 stretching modes in CO2‒ and several MCOO‒ species (M = H, Ag, Bi). In each of the MCOO‒
species, a single σ bond forms between the M atom and the carbon atom of the COO‒ group. In thecase of both CO2‒ and MCOO‒, the excess electron density is accommodated in what would be an
antibonding orbital of neutral CO2 (See Chapter 1). Structures and highest occupied molecularorbitals of CO2‒ and MCOO‒ are shown in Figure 6.1.

Figure 6.1. Structures and HOMOs for CO2‒, HCOO‒, AgCOO‒, and BiCOO‒ (presented left toright in the order listed). Carbon atoms are shown in black, oxygen atoms are shown in red,hydrogen in white, silver in grey, and bismuth in blue.

88

Unlike the other species discussed here, CO2‒ is unstable with respect to electron
autodetachment. The species can be stabilized in a matrix18 or a (CO2)n- cluster.21,23,24 We
performed photodissociation spectroscopy on (CO2)8‒, which has been confirmed to have CO2‒ as
the core ion by Johnson and coworkers, and can therefore be characterized as having thecomposition CO2‒∙(CO2)7 .23 The MCOO‒ species studied (M = Ag, Bi) each form metalloformate
type complexes (see Chapter 3 for an in depth discussion of the BiCOO‒ core ion). We report thephotodissociation action spectra of HCOO‒∙Ar, CO2‒∙(CO2)7, BiCOO‒∙(CO2)3, and AgCOO‒∙(CO2)4 in Figure 6.2.

89

Figure 6.2. Photodissociation spectra of CO2‒∙(CO2)7 (top, black), HCOO-·Ar (second trace,
magenta), AgCOO‒ ∙ (CO2)4 (third trace, red), and BiCOO‒ ∙ (CO2)3 (bottom trace, blue).
Transitions are identified as symmetric (ωS) or asymmetric C‒O stretching motions (ωA).
Assignments of the spectra are discussed below.
The spectrum of CO2‒∙(CO2)7 was collected by monitoring the formation of CO2‒∙(CO2)6
resulting from the loss of a single CO2 molecule. The spectrum is dominated by the high intensity
feature at 1661 cm-1, which is assigned to the asymmetric C‒O stretching mode of CO2‒ (ωA). Thisvalue is in good agreement with matrix isolation measurements perfomed by Jacox andcoworkers.18 The two lower intensity features are each due to symmetric C‒O stretching motionsof the core ion (1274 cm-1, ωS) and of the solvent molecules (1379 cm-1). The interaction of solvent

90

CO2 molecules with the core ion slightly breaks their symmetry, which allows for this weaktransition to be observable.
The photodissociation spectrum of HCOO‒∙Ar is collected by monitoring the formation of
HCOO‒ as a result of the loss of the Ar atom. This spectrum does not feature the same suppressionof low energy transitions described in previous chapters of this thesis. The binding energy of an
Ar atom to a negatively charged molecule is approximately 500 cm-1, which means that a single
infrared photon in the energy range of the experiment will readily cause the evaporation of the Ar
atom on the timescale of the experiment. The spectrum has two spectral signatures at 1617 and1314 cm-1. These transitions are assigned to ωA and ωS, respectively. These values are consistent
with photodissociation work performed by Johnson and coworkers.115 The presence of thehydrogen atom results in a significant red shift of ωA compared to the value for CO2‒ accompanied
by a blue shift of ωS. There is also a decrease in the calculated OCO bond angle from 135˚ (CO2‒)
to 131˚ (HCOO‒). As described in Chapter 3, the spectral shifts of ωA and ωS can be traced to a
weakening of the kinetic coupling of the individual CO oscillators as the OCO bond angle
decreases. Unlike in the behavior described in [Bi(CO2)n]‒, this effect is combined with andecreasing amount of charge on the COO‒ group compared to CO2‒. A decreasing amount ofcharge would result in a strengthening of the C-O bonds and a blue shift of both CO2 stretchingfrequencies. This highlights the strong effect that the more acute bond angle has on the stretchingfrequencies of HCOO‒.
The spectrum of AgCOO‒∙(CO2)4 was collected by monitoring the formation of AgCOO‒ ∙(CO2)3 as a result of the loss of a single CO2 molecule. We assign the experimental features at 1756
and 1187 cm-1 to ωA and ωS of the AgCOO‒ core ion, respectively. The feature at 1667 cm-1 isassigned to ωA of a CO2‒ core ion. These values are consistent with observations by Weber and91

coworkers.30 The Ag atom is solvating the CO2‒ core ion and the loss of Ag as a result of vibrational
excitation can be monitored. The presence of the Ag atom shifts the position of ω A to slightly
higher wavenumber compared to the value for CO2‒∙(CO2)7. For more information about thissystem the author directs the reader to the following article.30 The spectrum of BiCOO‒∙(CO2)3 has
two spectral features that can be assigned to ωA (1663 cm-1) and ωS (1211 cm-1). For more
information about this system, please see Chapter 3.
Of interest is the change in the intensity ration of ωS and ωA (IS/IA) for the systems understudy. Consistently for all MCOO‒ species discussed, the intensity ratio IS/IA is enhanced comparedto the intensity ratio of the same transitions in CO2‒. In the case of BiCOO‒, this ratio is enhanced
by a factor of 40. An even greater enhancement of a factor of 50 is observed experimentally `for
AgCOO‒. The suppression of low-energy transitions described in previous chapters will result ina slight reduction of the intensity of ωS, meaning that the experimentally observed enhancement
may prove to be smaller than the actual change of this ratio for infrared absorption.
In the typical description of electric dipole allowed vibrational transitions, the transition
dipole moment of a transition between two vibrational states v' and v" can be written as
⃗ (𝑄)
𝜕𝜇
𝜕𝑄

| ⟨𝑣 ′ |𝑄|𝑣"⟩. Taking the derivative of the dipole moment operator, 𝜇 (𝑄) = ∑𝑗(𝑞𝑗 ∙ 𝑟𝑗 ), we0arrive at the equation,
⃗ (𝑄)
𝜕𝜇

𝜕𝑟𝑗

𝜕𝑄

𝜕𝑄 0

| = ∑𝑗 (𝑞𝑗0| + 𝑟𝑗

𝜕𝑞𝑗

| ).

𝜕𝑄 0

(6.1)

The first term in equation 6.1 describes the derivative of the dipole moment under the fixed charge
approximation. In this approximation, the atomic charges, qj, are held constant while the nuclear
coordinates change along the normal mode. The second term quantifies the derivative of the dipole

92

moment due to only the changes in the charge distribution along the normal mode coordinate,while the nuclear coordinates themselves remain fixed. The sum of both of these components willenter into the expression for the overall transition dipole moment.
The large differences in the intensity ratio IS/IA for the different ions cannot be explainedwith the fixed charge approximation (Table 6.1). Using equation 6.1, we can separate thecontributions of the fixed charge approximation and charge oscillation. The charge oscillation term
can be calculated by varying the atomic charges of each atom, while maintaining the equilibriumgeometry, in center of mass coordinates, of the complex as a whole. Since all of the molecules inquestion have C2v symmetry, only the component of the transition dipole moment along the σ bondin MCOO‒ species will contribute to the overall transition dipole for the symmetric stretch. For
the asymmetric stretch, the overall transition dipole lies perpendicular to the axis, in the plane of
the molecule.

Table 6.1. A summary of the calculated M‒C bond distances, charge on the COO groups, theintensities of ωS and ωA (IS and IA) and the calculated and experimental ratio (IS/IA). The ratio
labeled DFT refers to the calculated infrared intensities from the calculated vibrational spectrumof each complex.
Species

M-C
distance
116

Chargeon COO[e]

IS[km·mol-1]
(DFT)

IA
[km·mol-1]
(DFT)

IS/IA
(DFT)

IS/IA
(fixedcharge)

IS/IA(chargeoscillation)

IS/IA(experimental)

CO2-

N/A

-1.00

21

642

0.033

0.45

0.004

0.021

HCOO-

1.14

-0.90

146

671

0.22

0.57

0.11

0.22

AgCOO-

1.51

-0.37

592

342

1.7

0.12

0.42

1.4

BiCOO-

2.46

-0.53

851

569

1.5

0.30

0.71

0.65

93

The ratio IS/IA was calculated for both the fixed charge approximation and the charge
oscillation approximation, and compared to the calculated ratios based on the predicted vibrational
spectra from DFT calculations and the experimentally observed ratios. In general, the fixed charge
approximation overestimates the experimentally observed intensity ratio for CO2‒ and HCOO‒
while underestimating the experimentally observed ratio for BiCOO‒ and AgCOO‒. In contrast,
the ratio calculated only considering charge oscillation underestimates the experimentallyobserved ratio for CO2‒ and HCOO‒, but overestimates the ratio for BiCOO‒ and AgCOO‒ (Table
6.1). This can be understood by considering just how much charge does oscillate along thesymmetric stretching coordinate in these systems. Figure 6.3 shows the portion of charge assigned
to the COO group as a function of the position along the normal mode coordinate Q. This
description is of course only meaningful for MCOO-, since CO2‒ does not exhibit any charge
movement along this coordinate, because there is no bonding partner in which to move the charge.

Figure 6.3. Combined plot of the change in charge of the COO group, Δq, plotted along thesymmetric C‒O stretching coordinate. The change in charge of HCOO‒ is shown in magenta
squares, AgCOO‒ shown in red circles, and BiCOO‒ shown in blue triangles. The classicalturning points along the normal mode coordinate are at ±1, and the equilibrium geometry is
shown at 0.

94

The metalloformate complexes exhibit a much higher proportion of the excess charge moving
through the complex compared to the formate anion (HCOO‒). We can understand this byconsidering that the COO group in formate anion has about -0.9 e of charge assigned to it in theequilibrium geometry. This results in an upper limit of the amount of charge flow into the COO
moiety since most of the charge already resides on the COO group. Charge flow from the H atom
to the COO group accounts for approximately 0.1 e of charge. Conversely, the metalloformate
species have between -0.5 e and -0.7 e on the COO group in their equilibrium geometry, which
naturally allows for the binding partner to be able to push more charge into the COO groups.
Charge flow for these species amounts to approximately 0.2 e of charge moving between the M
atom and the COO group.
It is clear that the binding partner M acts as a charge reservoir, allowing for excess charge
to be moved into and out of the bonding partner along the symmetric stretching coordinate. During
CO elongation charge moves out of the binding partner M and into the COO group, while during
CO contraction charge migrates out of the COO group and into the binding partner M. The identityof the binding partner is a key role in the amount of charge that is moved during this process.
We note that the intensity of the asymmetric stretching infrared signature is also enhanced
by charge oscillation. In this case, charge oscillation occurs from one C−O bond into the other as
each bond is contracted and extended, respectively. The magnitude of this enhancement is notconnected to the identity of the binding partner M, since the charge oscillation is limited only to
the COO− group.

95

6.4 Conclusions
The intensity ratios IS/IA for MCOO‒ complexes (M = H, Ag, Bi) are not recovered using
the fixed charge model, which can often be used successfully for calculating the transition dipole
moment. The reason for the failure of this model is that a significant amount of charge moves
between the M atom and the COO group along the symmetric CO2 stretching coordinate.
Comparison of the intensity ratio IS/IA in MCOO‒ to the same ratio in CO2‒ shows enhancementof over 50 times the value for CO2‒. We find that the intensities of both the symmetric and theasymmetric C‒O stretching mode are enhanced by charge oscillation along the normal mode
coordinate, although the enhancement of the intensity of the symmetric stretching mode shows agreater dependence on the binding partner, M, compared to the asymmetric stretching mode.

96

Chapter 7
Core Ion Geometries in (N2O)nO‒ and (N2O)n‒This work has been adapted with permission from the following previously publishedarticle:Thompson, M. C., and Weber, J. M. “Infrared spectroscopic studies on the cluster size
dependence of charge carrier structure in nitrous oxide cluster anions.” Journal of Chemical
Physics, 2016. 144:10. DOI: 10.1063/1.4943189. Copyright AIP Publishing. License Number:
4310371324667
7.1 Introduction
Since its detection by mass spectroscopy in the 1960’s, the molecular properties of the
N2O‒ anion have largely remained elusive.14,117-121 The lack of experimental values of many of theproperties of this molecular anion lie in the fact that it is very difficult to produce. Early collisionexperiments to determine the adiabatic electron affinity (AEA) of this species yielded wildly
different results (not even agreeing on the sign of the AEA).122,123
Nitrous oxide is isoelectronic with CO2. Like CO2, N2O adopts a bent geometry upon
accepting an electron. The Walsh diagram of N2O (Figure 7.1), calculated based on a method
discussed in section 7.2, shows that an excess electron will occupy one of the empty 3π orbitals of
N2O. This orbital is stabilized by bending the NNO bond angle away from the linear neutral
geometry to approximately 135˚. High-level ab initio calculations124,125 predict a negative AEA
for N2O, which would mean that the anion is unstable against electron autodetachment. This would
explain the difficulties in the generation of the N2O‒ anion. This change in geometry results in very

97small Frank-Condon overlap between anionic and neutral states, rendering the ground state of theneutral inaccessible from the ground state of the anion.

Figure 7.1. Walsh diagram of N2O, showing molecular orbital (MO) energies as a function of NNO
angle. The Walsh diagram for N2O was calculated using the RI-MP2 method (second-order
Møller-Plesset perturbation theory applying the resolution-of-identity for evaluation of twoelectron integrals)126,127 with def2-TZVPP basis sets for all atoms.58 Labels for MO curves inlinear and bent geometries are shown above the respective curve. In cases where more than one
curve is shown, the labels are listed with decreasing energy. The LUMO of neutral N2O is the 3π
orbital.

98

Another difficulty in generating N2O− comes from the fact that direct electron attachment
to N2O results only in dissociation, with products primarily being O- and N2.14,120,128-130 The
ground state geometry of N2O‒ is predicted to lie very close in energy to a dissociative
curve.121,124,125,131 It is theorized that the formation of bare N2O‒ is likely the result of low energy
charge transfer processes occurring with other anionic species (NO‒ or O‒).117 Unlike CO2‒, theanion of N2O has not been isolated in a matrix, and consequently the vibrational frequencies ofthis species have not been experimentally measured.
Previous work looking at electron attachment to neutral (N2O)n clusters has found that themain product channel is a dissociative attachment to N2O resulting in the formation of (N2O)n-1O‒
clusters. Additionally it was reported that smaller abundances of (N2O)n‒ and (N2O)nNO‒ clusters
can be produced at low electron kinetic energies.14,130,132-138 These anionic clusters have primarily
been studied using photoelectron spectroscopy. Johnson and coworkers were able to characterizedifferent isomers of N2O2‒ using photoelectron spectroscopy by introducing different precursor
species into their experiment.139 Sanov and coworkers collected photoelectron spectra of (N2O)nO‒
cluster ions.71,140,141 They report that for small clusters the core ion is a molecular anion withformula N2O2‒, and between cluster size n = 3 and 4 there is a change in the photoelectron signaturefrom a molecular anion to an atomic anion. This most likely is due to a core ion change from an
N2O2‒ to an O‒ anion. As discussed earlier in Chapter 3, such core ion switching events can be due
to a balance between electronic stability and solvation energy.
Despite the wealth of experimental work on (N2O)nO‒ clusters, there has been no infrared
work to determine the exact nature of the charge carrier. In the case of (N2O)n‒ there has been nowork characterizing the nature of the core ion. In this chapter, the infrared photodissociation

99

spectra of (N2O)nO‒ and (N2O)n‒ are reported and the nature of the core ion in both cluster systems
is discussed.
7.2 Results and Discussion of (N2O)nO‒ Clusters
Electron bombardment of a supersonic expansion of N2O (See Chapter 2) resulted in theformation of (N2O)nO‒, (N2O)n‒, and (N2O)nNO‒ clusters (Figure 7.2). The most abundant species
were (N2O)nO‒ clusters, which is consistent with previous work studying the electron attachment
to neutral N2O clusters.14,130,132-138

Figure 7.2. Mass spectrum of ions generated by electron injection into a supersonic expansion ofneat N2O.
The infrared photodissociation spectra of (N2O)nO‒ are shown in Figure 7.3. For smaller ions (n≤ 5) the spectra were obtained by monitoring the loss of both Ar tags from (N2O)nO‒ ∙ 2Ar. Allother spectra were obtained by monitoring the loss of N2O from (N2O)nO‒. We note the spectrum
for n = 5 was obtained with and without the presence of Ar tags and there was no significant
difference in either peak position or relative intensities (Figure 7.4).

100

Figure 7.3. Infrared photodissociation spectra of (N2O)nO‒ (n = 1‒12). The numbers on the side
indicate the number of N2O in the cluster.

101

Figure 7.4. Comparison of the experimental spectra of (N2O)5O‒ (top) and (N2O)5O‒∙Ar2(bottom).

The experimental spectrum of (N2O)O‒∙Ar2 contains two features at 1197 and 1352 cm-1.
This indicates the presence of a molecular core ion, since O‒ does not have any vibrationalfrequencies. Also, if O‒ were to be the core ion, we would expect to see the N-O stretching motionof a solvent N2O species around 1285 cm-1. With the signature not being present, we conclude thecore ion must be a molecular ion with formula N2O2‒. This observation is consistent with previous
work by Sanov and coworkers who observed molecular photoelectron signatures from N2O2‒.71
Quantum chemical calculations reveal three different isomers of the N2O2‒ (Figure 7.5). Two of
the isomers are cis and trans configurations of ONNO‒ with the cis isomer lying 116 meV higher
in energy than the global minimum trans configuration. A third isomer, which lies 291 meV above

102

the global minimum trans configuration, is an NNO2‒ configuration where the central nitrogenbonds to a second oxygen atom. This structure was found to be the predominant core ion structurein (N2O)nO‒ generated by electron injection into a supersonic expansion of N2O studied by Johnson
and coworkers using photoelectron spectroscopy.139

Figure 7.5. Calculated isomers of N2O2‒. Calculations were performed as described in Chapter 2.

Comparison of the calculated vibrational frequencies of these three isomers with the experimental
spectrum reveal that the core ion is the NNO2‒ isomer (Figure 7.6). We assign the experimental
signature at 1352 cm-1 to the N‒N stretching motion and the lower energy signature to theasymmetric NO2 stretching mode. The narrow spectral width of the two signatures and the absenceof any lower intensity features preclude the presence of any other isomer. It is interesting to notethat the highest energy isomer is the only one that is populated. We believe that this geometry canbe explained by understanding the charge distribution of neutral N2O. Charge analysis assigns acharge of +0.51 e on the central N atom, while the terminal N atom has a small negative charge (‒0.07 e) and the O atom has a charge of ‒0.44 e. We concur with the argument made by Johnson
and coworkers, that an O‒ ion, most likely generated as a result of dissociative electron attachment
to neutral N2O, performs a nucleophilic attack at the central nitrogen of a neutral N2O species toform NNO2‒.139

103

Figure 7.6. Comparison of the experimental spectrum of N2O2‒ ∙ 2Ar (top trace) with simulated
spectra of different isomers of N2O2‒ (bottom three traces).

Larger cluster sizes show the presence of a new feature at 1292 cm-1 that red shifts withincreasing cluster size and reaches a stable position at 1282 cm-1 in the limit of the largest cluster
sizes under study. We assign this band to the N‒O stretching motion of neutral N2O molecules
solvating the core ion. For smaller cluster sizes, this feature has some substructure that is likely
due to small differences in the solvation position of the solvent around the core ion. The intensityof the solvent signature increases compared to the signatures of the NNO 2‒ ion with increasingcluster size. This is to be expected since the number of solvent species increases with cluster size.

104

Sanov and coworkers observed a core ion switch from NNO2‒ to O‒ at cluster size n=4. Inspectionof the spectra for larger cluster sizes shows that no new spectral features appear with increasingcluster size. This suggests that if a core ion switch were to occur it would be to a core ion geometry
without observable vibrational signatures. The most likely scenario would be a core ion change
from NNO2‒ to O‒. In the present experiment, signatures of the NNO2‒ persist for all cluster sizes
studied, suggesting that the core ion change remains incomplete with our source conditions. We
note, however, that if a core ion change were to occur, there would be a disproportional change inthe ratios of the peak areas of NNO2‒ vibrational signatures to the peak areas of the solvent
signature. We analyzed the integrated peak areas of signatures up to (N2O)7O‒. If there is no coreion change, then the ratio of NNO2‒ vibrational signatures to the solvent signatures would change
according to
I NNO2I solv1n 1(n ≥ 2),(7.1)

with the onset of solvent signatures beginning at n = 2. Since our infrared spectra do not support
the presence of either cis or trans ONNO‒, we assume that the only molecular core ion species
present is NNO2‒. This is in line with work by Sanov and coworkers who only observed a single
molecular signature in their photoelectron images.71 The assumption that we only have an NNO2‒
or an O‒ molecular core ion allows us to directly compare our experimental results to the model
discussed above (see Figure 7.7).

105

Figure 7.7. Top panel: ratio of integrated peak areas (square points) compared with the predictedvalues from the model explained in the text (dashed line). Error bars are 10% of the first data
point and 50% for all subsequent data points. Bottom panel: percentage of clusters with a NNO2‒
core ion according to the model. The model assumes 100% NNO2‒ core ion population for n = 1.

Clusters with NNO2‒ core ion geometries drop to around 30% of the overall population by n = 3
and then drop to around 15% around n = 6. This is in conflict with behavior observed by Sanovand coworkers where the core ion change occurred abruptly at n = 4. We believe that different
source conditions between the two experiments are responsible for this discrepancy.
In conclusion, the core ion geometry for small (N2O)nO‒ clusters is the NNO2‒ molecularanion. For larger cluster ions there is a core ion change to a species that does not have any
detectable vibrational signatures (presumably O‒). These clusters were generated by electron

106

injection into a supersonic expansion of neat N2O. Previous work by Johnson and coworkers has
shown that the geometry of the core ion is somewhat dependent on the method used to create theions.139 As such, the currently reported results may only be applicable to clusters generated in asimilar manner with the same precursor gases.71,139-141
7.3 Core ion structure determination of (N2O)n‒The photodissociation action spectra of (N2O)n‒ (n = 7‒11, 14,15) are shown in Figure 7.8.
The spectra were generated by monitoring the loss of a single N2O species. No other loss channels
were observed resulting from the infrared absorption. The main spectral feature is a high intensity
band at 1282 cm-1, which corresponds to the N‒O stretching mode of neutral N2O species in thecluster. These N2O species are acting as solvent species, solvating the core ion in the cluster. The
other signatures in the experimental spectra lie in regions where neutral N2O does not have avibrational feature. We conclude that these signatures must be due to vibrations of a molecularcore ion. At the smallest cluster size examined (n =7), there is a low intensity feature at 1643 cm1. At n = 8 there are two spectral features in this region one at 1665 and another at 1583 cm-1. For

clusters larger than n = 9, there is only one low intensity spectral feature in this region, its position
varying from 1592‒1598 cm-1 as cluster size increases.

107

Figure 7.8. Photodissociation spectra of (N2O)n- cluster ions (n = 7-11,14,15).

The presence of only one unknown vibrational signature in the region from 1400‒2000 cm1

suggests that the core ion is a single N2O‒ species. If the excess electron were to be shared

between multiple N2O species, one would expect to see many different signatures in this region,
corresponding to local N‒N or N‒O oscillations of the core ion. The signature at 1595 cm-1 can be
assigned to the N‒N stretching mode of an N2O‒ core ion. This assignment is confirmed by bothdensity functional theory and ab initio calculations (see Table 7.1). Anharmonic correction scaling
factors for N2O‒ were calculated by comparing calculated vibrational frequencies toexperimentally measured frequencies. Using an anharmonic correction scaling factor of 0.949, our108

calculated vibrational frequency for the N‒N stretching mode in N2O‒ is exactly 1595 cm-1. Thisconfirms the assignment of this mode to the N‒N stretch, although the exact agreement is likely
fortuitous.

TABLE 7.1. Experimental transitions and calculated (harmonic approximation) (in cm-1) for
neutral and anionic N2O. Experimental data for neutral N2O have been taken from Ref. 27
Experimental data for N2O– are from the present work.
mode
method
N2O
N2O–
bend

N-O stretch

N-N stretch

experimental

589

N/A

calc. (this work, DFT)

619

654

calc. (Ref. 142)

549

560

calc. (Ref. 143).

N/A

657

experimental

1285

894

calc. (this work, DFT)

1329

993

calc. (Ref. 142)

959

863

calc. (Ref. 143)

N/A

998

experimental

2224

1595

calc. (this work, DFT)

2343

1681

calc. (Ref. 142)

2434

1610

calc. (Ref. 143)

N/A

1684

As mentioned above, the exact position of the N‒N stretching mode depends on the size ofthe cluster, with two different signatures separated by around 80 cm-1. It is not unreasonable to109suggest that this could be due to the N2O‒ experiencing two different solvation environments. In
the limit of a small cluster, solvation will occur preferentially at the portion of the core ion
containing more charge. Preferential solvation around the oxygen side of the anion will result incharge polarization of the excess electron toward that side of the molecule. This will result in astronger attraction of additional solvent species to this side of the anion. As the cluster grows,
eventually there will be no more solvation sites on this side of the anion and solvation will occur
on the other side. As the solvation shell around the anion begins to grow, the charge willredistribute more evenly.
When N2O accepts an excess electron, it is accommodated in an antibonding orbital. This
weakens the N-N and N-O bonds and results in a red shift of both stretching frequencies. Whenthe charge distribution shifts toward the N‒N bond this will weaken it resulting in a red shift ofthe signature. This is similar to the behavior observed in the transition from cluster size n = 7‒9.
At these cluster sizes the N‒N stretching signature red shifts over 80 cm-1 from 1665 cm-1 for small
clusters to 1595 cm-1 for larger clusters. One hypothesis consistent with this shift is that a half shell
is formed around the oxygen atom of the N2O‒, which is completed at n = 8. The solvation shellthen continues to grow to a more symmetric structure around the core ion at larger cluster sizes.
The behavior is similar to IBr‒∙(CO2)n clusters, where solvation preferentially occurs at thebromine atom of the core ion until a half shell is formed at n ≈ 8.144 Electronic structure calculationsof clusters at this size are difficult, and due to the large number of possible structures, we wereunable to confirm half shell formation at large cluster sizes by computational chemistry. We wereable to calculate the energy difference between an O solvated cluster and a more symmetricsolvation environment for cluster sizes n = 3 and 4. At these cluster sizes, asymmetric solvation
structures around the O atom of the core ion are 50-60 meV lower in energy than structures110

symmetrically solvating the core ion. This result is consistent with our present hypothesis.
However, there was only a negligible difference in the calculated vibrational frequency of the N‒
N stretching mode (less than 10 cm-1). While the predicted red shift of the experimentally observed
signature is consistent with our calculations, the magnitude of the shift is not recovered.
Nevertheless, the overall theory that the experimentally observed red shift is due to the formationof a half shell around the core ion is still a plausible explanation.
In principle, the N‒O stretching motion of N2O‒, predicted to be between 863‒998 cm-1
depending on the level of theory, is within the wavenumber range of our experiment. To attempt
to observe this signature, we extended our scanning range to 800 cm-1 for (N2O)10‒ clusters, which
was the most abundant cluster size, and detected a low intensity feature at 894 cm-1. Transitions
below around 1000 cm-1 are suppressed due to the binding energy of N2O solvent molecules,
leading to kinetic shift effects, as described in the previous chapters of this thesis. Calculations
performed for the current experiment using an anharmonic correction factor of 0.967 predict theN‒O stretching mode to be at 960 cm-1. This does not agree well with the experimental observation,and likely is due to an underestimation of the anharmonicity of this vibrational mode.
7.4 Conclusions
We report the photodissociation spectra of (N2O)nO‒ and (N2O)n‒ generated by electron
injection into a supersonic expansion of N2O. In (N2O)nO‒, we identify the core ion for smallercluster sizes as NNO2‒. There is a core ion change that occurs around n = 3 to predominantly anO‒ core ion, although the molecular core ion persists in small abundances.
For (N2O)n‒, the core ion is N2O‒. IR measurement of the vibrational frequencies of thision show the N‒N stretching vibrational frequency to be 1595 cm-1 and the N‒O vibrational

111

frequency to be 894 cm-1. The observation of an abrupt red shift of the N‒N stretching mode is
attributed to the formation of a half solvation shell around the oxygen side of the core ion at cluster
size n = 8.

112

Chapter 8
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