Structures and Charge Distributions of Cluster Anions
Studied by Infrared Photodissociation Spectroscopy
byB. J. Knurr
B.A., Macalester College, 2009

A thesis submitted to the
Faculty of the Graduate School of the
University of Colorado in partial fulfillmentof the requirements for the degree ofDoctor of Philosophy
Department of Chemistry and Biochemistry2015This thesis entitled:
Structures and Charge Distributions of Cluster Anions Studied by Infrared Photodissociation
Spectroscopy
written by B. J. Knurr
has been approved for the Department of Chemistry and BiochemistryProf.
J. Mathias
MathiasWeber
Weber
Prof. J.

Prof.
Veronica M. Bierbaum
Prof. Veronica

Date

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

Knurr, B. J. (Ph.D., Chemical Physics)
Structures and Charge Distributions of Cluster Anions Studied by Infrared Photodissociation Spectroscopy
Thesis directed by Prof. Prof. J. Mathias Weber

Anions are relevant in many di↵erent facets of chemistry. The study of anions, their interaction with other atoms and molecules and the e↵ect of solvation on those interactions are of interest
to further the understanding of chemical processes. Studies on cluster ions in vacuo allow for thegeneration of and excellent control over chemically interesting species. Experiments of this nature
are ideal for studying the fundamental interactions of molecular partners by eliminating many of
the complications of obtaining spectroscopic data on an individual species in the condensed phase.
The clusters studied can also be viewed as model systems for more complex chemical processes thatare otherwise difficult to understand at a molecular level.
In one part of this work, infrared spectra were obtained for [M(CO2 )n ] (M = Au, Ag, Co, Ni
and Cu) clusters. These studies were performed in an e↵ort to understand the interactions resulting
in the reduction of CO2 . These cluster systems are viewed as model single atom catalysts for the
purpose of CO2 reduction. Au, Ag and Cu were all found to form ⌘ 1 (monodentate) [MCO2 ]complexes of the metal and one CO2 molecule with a structure akin to the formate anion. Study ofthe solvation environment of these ⌘ 1 complexes reveals that the excess charge on the core can bepolarized onto the CO2 moiety with increasing efficiency from Au to Cu. Co, Ni and Cu all exhibit
core ions of the form [CO2 MCO2 ]where the CO2 ligands are bound in a bidentate (⌘ 2 ) fashion.

The polarizability of the core species increases from Co to Cu. Cu is a member of both the firstrow transition metals and the coinage metals, and it shows interaction motifs characteristic of bothgroups of metals. Additionally, two studies were performed on [CoO(CO2 )n ]and [NiO(CO2 )n ]clusters. It is shown that similar core structures are present in both CoO and NiO clusters andthat the general behavior of the clusters is comparable to the analogous non-oxide clusters.

iv
Two studies were performed on non-metal containing systems. A charge transfer reaction
between a nitromethane anion and iodomethane was vibrationally induced. It is found that vibrational excitation of C-H stretches on either of the species in the complex leads to charge transferfrom the nitomethane to the iodomethane and results in the formation of I

product anions. Ar

tagging is used to study the energetics of the reaction and achieve quenching of the reactive channel.
The reaction is discussed in the framework of a vibrationally induced SN 2 reaction.
The infrared spectra of a naphthalene anion clustered with one to six water molecules were
measured. The ⇡-system of the naphthalene accommodates the excess electron with a network ofwater molecules evolving on one side of the naphthalene. It was found that water-water H-bonds are
generally more favorable than water-ion ⇡-H-bonds. With three or more water molecules present,
the spectra indicate that there is more than one conformation of the water network present.

Dedication

To my parents and my teachers for always pushing me to be a good scientist and an evenbetter person.

vi

Acknowledgements

There have been many people who have helped me on my journey through graduate school.
I must begin by acknowledging my advisor, Professor J. Mathias Weber for not only thinking I
would be a good fit in the graduate program but also in his group. Without his support, teaching
and awful puns (at which I have not only laughed but also contributed) I don’t think I would havebeen nearly as productive or enjoyed my graduate experience as much as I have. Mathias has been
not only an excellent mentor but also a good friend and there is no way I can thank him enoughfor all he has done for me.
I must also thank my group members both former and present. Jesse Marcum, Chris Adams
and Sydney Kaufman were all present when I joined the lab and were instrumental in making me
feel welcome as a group member and as a scientist. They were all patient and helpful while I waslearning the ropes of being a graduate student. I learned some of my best lab problem solving andcritical thinking skills from them. I must also acknowledge Chris specifically for teaching me how
to turn the instrument on and what everything does with respect to “making and shooting ions”.
Without his guidance and patience with me, even while writing his own thesis, I would not havebeen as productive as I have been. I have to thank Casey Christopher for his unwavering friendshipand willingness to discuss just about everything. It was a pleasure to be his lab mate for three years
and friend for longer. I must thank the four undergraduate researchers, Eric Pozzi, Grant Gogul,
Caitlin Majlinger and Luke Long, letting me vicariously participate in their research. I wish them
all well in their current studies and endeavors. In the last two years I have gotten to know the two
newest members of the group, Shuang “Xavier” Xu and Michael Thompson, and in the past year I

vii
have handed over the reins of “my” experiment to Michael. Michael is already proficient with the
instrument and, with likely little luck, will easily be more productive than I was. I wish them boththe best in their graduate careers and know that there is much success in their futures.
I was fortunate enough to work in within the JILA community. The resources and, inparticular, the people I have had the pleasure of meeting and learning from are many and are
all appreciated. From learning about electronics to machining techniques to how the purchasingsystem works and all the time they took to talk to me about various (and potentially o↵ topic)
subjects, I am grateful. I also received a lot of support from the Ion Super Group. Professors
Veronica Bierbaum, Barney Ellision and Carl Lineberger and their respective groups have given meso much valuable feedback on my research and professional growth. It has been a pleasure meetingand engaging with such a strong group of scientists.
I gratefully acknowledge the grants that have funded this research and my education. All of
this research has been funded by grants from the National Science Foundation through either the
Career Foundation (CHE-0845618) or the JILA AMO Physics Frontier Center (PHY-1125844 andPHY-0551010). I was further supported by the University of Colorado Department of Chemistry
and Biochemistry through Teaching Assistantships.
Lastly I must acknowledge my family and friends. My gratitude for the support and encouragement everyone has given me during my time as a graduate student can not be put into words.
In particular all the time my parents spent listening to me rant about research and teaching. Most
importantly I have to thank my wife, Anne, for her support and encouragement through this process
and for sticking with me through everything.

viiiContents

Chapter
1 Introduction and Motivation

1

2 The Experiment3

2.1

Overview of Experiment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

3

2.2

Experimental Details . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

3

2.2.1

Cluster Generation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

3

2.2.2

Time of Flight Mass Spectrometer . . . . . . . . . . . . . . . . . . . . . . . .

7

2.2.3

Laser System . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

2.2.4

Data Collection and Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

3 Infrared Studies of Metal-(CO2 )n Anion Clusters

16

3.1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17

3.2

[Au(CO2 )n ] and [Ag(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . . . . . . 20

3.3

3.2.1

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22

3.2.2

Infrared Spectra of [Au(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . 23

3.2.3

Infrared Spectra of [Ag(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . 30

[Co(CO2 )n ] and [Ni(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . . . . . . 41
3.3.1

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43

3.3.2

Infrared Spectra of [Co(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . 43

3.3.3

Infrared Spectra of [Ni(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . 55

ix
3.4

3.5

Infrared Studies of [Cu(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . . . . . 65
3.4.1

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65

3.4.2

Infrared Spectra of [Cu(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . . 66

Comparison of all Metal-(CO2 )n Anion Clusters . . . . . . . . . . . . . . . . . . . . . 74

4 Infrared Studies of [CoO(CO2 )n ] and [NiO(CO2 )n ] Clusters

79

4.1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 79

4.2

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

4.3

Spectra and Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 80

4.4

4.3.1

Infrared Spectra of [CoO(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . 80

4.3.2

Infrared Spectra of [NiO(CO2 )n ] Clusters . . . . . . . . . . . . . . . . . . . 89

Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95

5 Observation of a Vibrationally Induced Electron Transfer Reaction in a Binary MolecularComplex97

5.1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97

5.2

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100

5.3

Results and Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 101

5.4

Conclusion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111

6 Infrared Spectra of Naphthalene Water Clusters

112

6.1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 112

6.2

Computational Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114

6.3

Results and Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 115

6.4

6.3.1

OH Stretching Vibrations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 115

6.3.2

CH Stretching Vibrations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 130

Conclusion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131

xBibliography

132

Appendix
A Supplementary Material for Chapter 3.2.2

147

B Supplementary Material for Chapter 3.3.3

150

C Supplementary Material for Chapter 6

154

C.1 [C10 H8 ·H2 O]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 154

C.2 [C10 H8 ·(H2 O)2 ]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 157

C.3 [C10 H8 ·(H2 O)3 ]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 160

C.4 [C10 H8 ·(H2 O)4 ]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 163

C.5 [C10 H8 ·(H2 O)5 ]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 167

C.6 [C10 H8 ·(H2 O)6 ]. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 173

xi

TablesTable
3.1

Calculated partial charges for the solvation isomers of [Co(CO2 )4 ]presented in

Figure 3.19. Isomers 1 - 4 refer to the top to bottom structures presented in Figure
3.19. The column labeled Solvent indicates the amount of partial charge delocalized
into the solvent environment. The partial charges are reported in units of e. . . . . . 49
3.2

Calculated partial charges of the core isomers for [Ni(CO2 )n ] . See Figure 3.23 anddiscussion in the text for additional information on the core structures. The column
labeled Solvent indicates the amount of partial charge delocalized into the solvent
environment. The partial charges are reported in units of e. . . . . . . . . . . . . . . 63

4.1

Calculated partial charges for the predicted core structures of [CoO(CO2 )n ]pre-sented in Figures 4.2 and 4.4. The partial charges are reported in units of e . . . . . 87
4.2

Calculated partial charges for the predicted core structures of [NiO(CO2 )n ]pre-sented in Figure 4.7. The partial charges are reported in units of e . . . . . . . . . . 95
5.1

Experimentally determined and calculated peak positions in cm 1 for the dominantvibrational features observed in Figure 5.2. The experimental frequencies are reported for CH3 NO2 ·CH3 I with values for CH3 NO2 ·CD3 I given in parentheses whereobserved. Calculated frequencies (see above for discussion) are given for the nondeuterated complex only and are reported in km·mol 1 . The vibrational assignments
given are based on structure I in Figure 5.3. . . . . . . . . . . . . . . . . . . . . . . . 103

xii

FiguresFigure2.1

Schematic for the entrainment setup. X and Y represent di↵erent molecular species
and are pulsed into the entrainment region using series 9 General Valves. . . . . . .

2.2

Schematic diagram of the laser vaporization setup used to generate metal containing
clusters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

2.3

8Schematic of the experimental apparatus including the vaporization source with theelectron gun shown for completeness. . . . . . . . . . . . . . . . . . . . . . . . . . . .

2.4

6

9

Mass spectrum observed using the vaporization source with a cobalt disc and CO2
as the carrier gas. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10

2.5

Schematic diagram of the Laservision optical parametric converter setup used forthe generation of tunable infrared radiation in these experiments. . . . . . . . . . . . 13

2.6

Schematic diagram of the multi-pass cell ring reflector used to increase the numberof passes of IR light on the mass selected ion packet. . . . . . . . . . . . . . . . . . . 14

3.1

The Walsh diagram for CO2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19

3.2

The observed redshift of the antisymmetric CO stretch in CO2 when interactingwith Auand various halides. As is indicated on the graph, the redshift of theantisymmetric stretching frequency of CO2 interacting with Au is calculated while
the other redshifts have been experimentally measured. The redshifts are all withrespect to 2349 cm 1 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21

xiii
3.3

Experimental spectra of [Au(CO2 )n ]clusters. The numbers denote the numberof CO2 molecules in the cluster. The right and left panes are individually normalized so they are on di↵erent scales. The arrow indicates the predicted frequency of
the AuCO2 binary complex from the calculations of Boese et al. All spectra wereobtained monitoring the loss of a single solvent CO2 . . . . . . . . . . . . . . . . . . . 24
3.4

Comparison of the three calculated lowest energy structures for [Au(CO2 )2 ] clusters
to the experimental spectrum for [Au(CO2 )2 ] . Relative energies are given in meV. . 26

3.5

Comparison of the lowest energy structures for [Au(CO2 )n ] clusters (n = 3, 4). Relative energies are given in meV and are with respect to the lowest energy conformerof a given cluster size. Since these structures are all within 60 meV of each other,
they are considered to be isoenergetic. . . . . . . . . . . . . . . . . . . . . . . . . . . 28

3.6

Experimental spectra of [Ag(CO2 )n ]clusters. The numbers denote the number of

CO2 molecules in the cluster (n). The right and left panes are individually normalized
so they are on di↵erent scales. All of these spectra were obtained monitoring theloss of a single solvent CO2 . The dashed lines are meant to guide the eye. . . . . . . 31
3.7

Comparison of the lowest energy structures for [Ag(CO2 )n ]clusters (n = 2

4).

All reported structures are within 60 meV relative energy of each other for a given
cluster size and are considered to be isoenergetic. . . . . . . . . . . . . . . . . . . . . 33
3.8

Peak positions in the lower energy region (left panel in Figure 3.6) plotted againstcluster size. Solid squares represent structures where the terminal solvation positionis filled. The open squares represent structures where the terminal solvation positionis vacant. Solid circles indicate when the charge carrier is CO2 . . . . . . . . . . . . 34

3.9

Comparison of structures exhibiting symmetric solvation around a Ag
n=2anion for4. The green dotted line is meant to show the red shift trend observed andpredicted. The magenta dot-dash line shows the nearly unshifted solvent CO2 peakcorresponding to the existence of a formate type core. . . . . . . . . . . . . . . . . . 36

xiv
3.10 Comparison of Ag loss channels for [Ag(CO2 )4 ] (D) and [Ag(CO2 )5 ] (C) to (CO2 )nclusters exhibiting di↵erent core structures (A and B). The feature at 1660 cm 1
corresponds to a CO2 core while the set of peaks between 1800 cm 1 and 2000
cm 1 indicates a C2 O4 core. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37
3.11 Potential energy curve for [Ag(CO2 )4 ] generated by varying the OCO bond angle ofthe partially reduced CO2 moiety. The red curve represents the structures as shown.
The black curve contains a dielectric continuum corresponding to CO2 around theshown structures to approximate the e↵ect of increased solvation (see Section 3.2.1 foradditional information). A robust second minimum is recovered with the formationof a C2 O4 species, which is additionally stabilized by the dielectric continuum. The
relative energies presented here are not zero-point corrected. . . . . . . . . . . . . . . 38
3.12 Potential energy curve for [Au(CO2 )4 ] generated by varying the OCO bond angle ofthe partially reduced CO2 moiety. No second minimum is recovered corresponding toa C2 O4 core unlike in [Ag(CO2 )4 ] (see Figure 3.11). The relative energies presented
here are not zero-point corrected. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 40
3.13 Plot of the antisymmetric CO stretching frequency in a formate core against thecalculated partial charge localized on the activated CO2 moiety for an assigned
structure (see text). The blue triangles represent [Ag(CO2 )n ]clusters while theorange squares represent [Au(CO2 )n ] clusters. The black circle is the antisymmetric stretching frequency of a CO2 in a Ne matrix. The solid line is a linear fit of
all the points in the plot. The dashed lines represent the maximum charge transfer
based on the most red shifted frequency observed. . . . . . . . . . . . . . . . . . . . 41
3.14 Experimental spectra of [Co(CO2 )n ] clusters monitoring the loss of one CO2 . The
two regions (1000 cm 1 to 2150 cm 1 and 2270 cm 1 to 2400 cm 1 ) are individuallynormalized and are therefore on di↵erent scales. Data was not acquired down to1000 cm 1 for larger clusters due to low parent cluster intensity and correspondingly
poor daughter anion signal intensity. . . . . . . . . . . . . . . . . . . . . . . . . . . . 44

xv
3.15 Comparison of the peak positions for [Au(CO2 )n ] (top), [Ag(CO2 )n ] (middle) and[Co(CO2 )n ] (bottom) clusters. Note that fragmentation is first observed for Au andAg at n = 2 while it not observed until n = 3 for Co. . . . . . . . . . . . . . . . . . . 45
3.16 Structures of CoCO2 for singlet and triplet occupations. The capital letter denotes
the core isomer while the subscript indicates a singlet or triplet conformation. Relative energies are given in units of meV. . . . . . . . . . . . . . . . . . . . . . . . . . 46
3.17 Lowest energy calculated structures for [Co(CO2 )2 ] . The letter denotes the core
structure while the number in parentheses indicates the number of solvent molecules
and the superscript denotes singlet or triplet configuration. Relative energies are
given in meV. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 48
3.18 Comparison of the calculated lowest energy structures for [Co(CO2 )2 ]to the ex-perimental spectrum. The letter denotes the core structure while the number inparentheses indicates the number of solvent molecules and the superscript denotes
singlet or triplet configuration. Relative energies are given in meV. . . . . . . . . . . 50
3.19 Comparison of calculated di↵erent solvation isomers around structure 1 E to theexperimental spectrum for [Co(CO2 )4 ] . Relative energies are given in meV. . . . . 51
3.20 Comparison of calculated likely minor isomers to the experimental spectrum for[Co(CO2 )4 ] . Relative energies are given in meV. . . . . . . . . . . . . . . . . . . . . 53
3.21 Orbitals in the d-block of the [CO2 CoCO2 ] core ion expected for a “butterfly” ML4
configuration (left) and a bent ML2 configuration (center), compared to the calculated molecular orbitals for the singlet core ion structure 1 E (right). Note that the
HOMO-3 orbital is not part of the d-block but is shown for completeness. The occupied orbitals are shown as bold horizontal lines for each configuration. Calculated
orbital energies are given in eV. The dotted lines show how the dz 2 orbital changes
from ML4 to ML2 and correlation between calculations and ligand field theory. Notethat the definition of the z-axis is not along the symmetry axis of the complex but
rather is defined by the dz 2 character of the HOMO. . . . . . . . . . . . . . . . . . . 54

xvi
3.22 Experimental spectra of [Ni(CO2 )n ]clusters. The numbers indicate the numberof CO2 molecules (n) present in each cluster. The traces on the left and right are
individually normalized and are on di↵erent scales. . . . . . . . . . . . . . . . . . . . 56
3.23 Calculated lowest energy structures for [Ni(CO2 )n ]clusters. The capital letter

indicates the core structure and the number denotes the number of solvent molecules
present. Plain text indicates n = 2 structures while bold text denotes n = 3
structures. Relative energies are reported in meV and are given with respect to thelowest energy structure for a given cluster size. Relative energies for isomers B and
G are not reported as there are no structures to compare them with and are shownin italic text. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57
3.24 Comparison of calculated structures for [Ni(CO2 )3 ]to the experimental spectrum

for n = 3. See Figure 3.23 and text for discussion on the structures. Relative energies
are given in meV. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59
3.25 Comparison of various solvation positions around the core isomer E. Note that thepredicted spectrum shifts only slightly with varying solvation position. The calculated spectrum of isomer E* is shown on the bottom for completeness. . . . . . . . . 60
3.26 Calculated vibrational frequencies as a function of cluster size for assigned core
structures compared to the experimental peak positions as a function of cluster
size (see text for discussion). The letters indicate the core structure giving rise toeach set of peaks (see Figures 3.23). The experimental spectrum for [Ni(CO2 )5 ] isshown on top for reference. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 62
3.27 A comparison of the combined calculated spectra for isomers E and F to the experimental spectrum of [Ni(CO2 )2 ] . Note that while the fit is not ideal, it is reasonable
that a multiphoton process could recover the observed experimental spectrum. . . . 64
3.28 Experimental spectra of [Cu(CO2 )n ] . The numbers denote the number of CO2
molecules in the cluster. All spectra are individually normalized so that the left andright traces are on di↵erent scales. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 67

xvii
3.29 Experimental spectra of [Cu(CO2 )n ]with the three groups of spectral features

highlighted (see text for discussion). Numbers denote the number of CO2 moleculesin the cluster. All spectra are individually normalized so that the left and righttraces are on di↵erent scales. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68
3.30 Calculated lowest energy structures for [Cu(CO2 )n ]chosen to be consistent with that for [Co(CO2 )n ](n = 1 to 4). The notation isand [Ni(CO2 )n ]clusters. The

letter denotes the core structure and the number/letter combination in parentheses
indicates the number of solvent molecules and their positions. The purple box indicates n = 1 species, the blue boxes indicate n = 2 species, the green boxes indicate
n = 3 species and the orange box indicates n = 4 species. Bare core structures areshown on the top row while solvation isomers are shown in the descending columns.
Relative energies are given in meV and are with respect to the lowest energy isomer
for a given cluster size. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69
3.31 Comparison of calculated structures B(1a), B(1b) and A(1) to the experimental
spectrum of [Cu(CO2 )2 ] . See Figure 3.30 for relative energies. . . . . . . . . . . . . 70
3.32 Plot of the antisymmetric CO stretching frequency in a formate core against the calculated partial charge localized on the activated CO2 moiety for an assigned structure (see text). The blue triangles represent [Ag(CO2 )n ] clusters while the orange
squares represent [Au(CO2 )4 ] clusters and the red squares represent [Cu(CO2 )n ] .
The black circle is the antisymmetric stretching frequency of a CO2 in a Ne matrix.
The solid line is a linear fit of all the points in the plot. The dashed lines represent the maximum charge transfer observed based on the most red shifted frequency
observed. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 71
3.33 Comparison of calculated structures with a [Cu(CO2 )2 ]core to the experimental

spectrum for n = 3. See Figure 3.30 for description and relative energies of thestructures presented. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73

xviii3.34 Comparison of the calculated spectra for di↵erent solvation positions around isomers
E and G to the experimental spectrum for [Cu(CO2 )6 ] . Note that since isomer G
is a minor contributor to the experimental spectrum, the features around 1700 cm 1
will likely be obscured by the formate signature. See Figure 3.30 for additional
information on the structures presented here and text for additional discussion on
symmetric and antisymmetric solvation. . . . . . . . . . . . . . . . . . . . . . . . . . 75
3.35 Spectral shifts of the experimental and simulated features for [Ni(CO2 )n ](n = 3to 6). The top trace show the spectrum for n = 6 for comparison. The lines aremeant to guide the eye and are not a fit of any of the data, the letters indicate theassignments of the core isomers as shown in Figure 3.30. . . . . . . . . . . . . . . . . 76
4.1

Experimental spectra of [CoO(CO2 )n ]for (n = 3

8). The numbers denote thenumber of CO2 molecules present in the cluster. The exploratory experimental data
for n = 8 was truncated after determining that the spectrum was not appreciably
evolving as a function of cluster size. The left and right traces are individuallynormalized and are not on the same scale. . . . . . . . . . . . . . . . . . . . . . . . . 81
4.2

Calculated lowest energy structures for [CoO(CO2 )n ] . The capital letter denotes
the core structure and superscripts 1 and 3 denote singlet and triplet, respectively.
Relative energies are given in meV and are all with respect to isomer 3 A. . . . . . . 83

4.3

Comparison of calculated core isomers for [CoO(CO2 )2 ] to the experimental spectrum of [CoO(CO2 )3 ] . Isomer labels and relative energies are given for each structure. See text for discussion. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85

4.4

Calculated lowest energy structures for [CoO(CO2 )3 ] . The capital letter denotes thecore structure, the superscript 1 and 3 indicate singlet and triplet configurations, andthe number/letter combination denote the number and configuration of the solvent
molecules. Relative energies are reported in meV. . . . . . . . . . . . . . . . . . . . . 86

xix
4.5

Comparison of the calculated singlet structures in Figure 4.4 to the experimental
spectrum for [CoO(CO2 )3 ] . Isomer labels and relative energies are given to the right. 88

4.6

Experimental spectra of [NiO(CO2 )n ] . The numbers indicate the number of CO2
molecules present. The left and right traces are individually normalized and are ondi↵erent scales. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

4.7

Calculated lowest energy structures for [NiO(CO2 )n ] . Structures in the leftmost
column are bare core structures while the structures to the right are solvation isomer(s) of the core structures. The letters denote the core structure while the letter/number combination indicate the number of solvent molecules and their orientation around the core. Relative energies for [NiO(CO2 )]structures are shown initalics, [NiO(CO2 )2 ] are shown in plain text and [NiO(CO2 )3 ] are shown in bold.
No relative energy is presented for F(1) since there is no other structure of the form
[NiO(CO2 )4 ] to compare it with. All relative energies are reported in meV.
4.8

. . . . 91

Comparison of calculated structures G(1), H(1) and A to the experimental spectrumof [NiO(CO2 )2 ] . See Figure 4.7 for relative energies and text for discussion. . . . . 92

4.9

Comparison of calculated structures for [NiO(CO2 )3 ] to the experimental spectrumof [NiO(CO2 )3 ] . See Figure 4.7 for relative energies and text for discussion. . . . . 94

5.1

Schematic representations of the potential results of vibrational excitation of the barebinary complex (top) and the Ar solvated complex (bottom). R1 refers to Equation
5.1 while R2 refers to Equation 5.2. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100

5.2

Photodissociation action spectra of CH3 NO2 ·CH3 I (top trace) and CH3 NO2 ·CD3 I(bottom trace) acquired by monitoring the formation of I

photoproducts. The

dashed lines indicate the CH stretching modes attributed to the CH3 I molecule
within the complex. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102
5.3

Calculated lowest energy structures for [CH3 NO2 ,CH3 ,I]complexes. The relative

energy for conformation II is with respect to conformation I.

. . . . . . . . . . . . . 104

xx
5.4

Comparison of the experimental spectrum of CH3 NO2 ·CH3 I·Ar to the simulated
spectra of calculated conformers I (middle panel) and II (bottom panel). The spectrum of the argon tagged ion is used because it is better resolved compared to thespectrum of the bare ion. Note that the peak at ⇠2950 cm 1 for conformer I contains
two close lying transitions (see Table 5.1) which account for the splitting observedin the experimental spectrum. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 105

5.5

Experimental spectra of CH3 NO2 ·CH3 I·Arn (n = 03). The di↵erent colors of thespectra indicate a change in the photofragmentation channel. The red curves aremonitoring charge transfer and the subsequent formation of I

photoproducts (seeR1 in Figure 5.1) while the blue curves are monitoring Ar evaporation (see R2 in
Figure 5.1). Note that both channels are observed for n = 2.
5.6

. . . . . . . . . . . . . 107

Schematic representation of the potential energy surface along the reaction coordinate of dissociative electron transfer from CH3 NO2 to CH3 I (energies are to scale).
(a) Adiabatic electron affinity of CH3 NO2 (172 meV). (b) Binding energy of the
CH3 NO2 ·CH3 I entrance channel complex (490 meV, calculated for conformation I).
(c) Estimated lower limit of the barrier height for the electron transfer reaction (230
meV, experimentally determined from Ar solvation). (d) Relative energy of the exit
channel complex for structure III (500 meV, solid curve) and structure IV (⇠1150
meV, dashed curve). The dashed line represents the SN 2 product. (e) Exit channel
asymptote for the formation of I

and neutral products based on the two possible

exit channel complexes. The solid curve is based on the heat of reaction for dissociative electron attachment to CH3 I (620 meV). The dashed curve represents the I+ CH3 NO2 CH·3 asymptote (⇠1550 meV). . . . . . . . . . . . . . . . . . . . . . . . . 108

xxi
5.7

Calculated potential energy curve (top) and charge on the iodine atom (bottom) as
a function of the CI bond distance. The full squares in the top trace are calculatedwith only the CI bond distance fixed, letting all other molecular coordinates relax.
The open diamonds represent calculations where the out-of-plane angle of the methylgroup of the CH3 I molecule (✓oop ) was fixed at 10.0 , corresponding to the geometry
at a CI bond distance of ⇠2.575 Å in the unrestricted calculation. The open squares
are data points calculated with the same angle (✓oop ) fixed at 5.5 , corresponding to
the geometry at 2.625 Å CI bond distance in the unrestricted calculation. We notethat ✓oop changes very rapidly as the transition state is approached. The resulting
approximated diabatic potential energy curves are shown as dashed lines in the insetof the top panel; the structures shown as insets from left to right represent theentrance channel, transition state, and exit channel geometries, respectively. . . . . . 110

6.1

Experimental spectrum and calculated spectra of [Np·(H2 O)] . Top trace is theexperimental spectrum of [Np·(H2 O)] ·Ar8 acquired by monitoring the loss of 6 Ar
atoms. Bottom traces are calculated conformers and their predicted infrared spectra.
The dashed lines are meant to show ionic ⇡-H-bonds. . . . . . . . . . . . . . . . . . . 116

6.2

Observation of vibrational autodetachment in [Np·(H2 O)] ·Arm clusters. Top trace:
m = 4, monitoring the loss of all 4 Ar atoms. Center trace: m = 5, monitoring theloss of all 5 Ar atoms. Bottom trace: m = 8, monitoring the loss of 5 Ar atoms.
Note that the loss of 6 Ar atoms (shown in Figure 6.1) yields the same spectrum as
the loss of 5 Ar atoms. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 119

6.3

Experimental spectrum and calculated spectra of [Np·(H2 O)2 ] . Top trace is theexperimental spectrum of [Np·(H2 O)2 ] ·Ar2 acquired by monitoring the loss of 2 Ar
atoms. Bottom traces are calculated conformers and their predicted infrared spectra.
The dashed lines show the hydrogen bonds. . . . . . . . . . . . . . . . . . . . . . . . 121

xxii
6.4

Experimental spectrum and calculated spectra of [Np·(H2 O)3 ] . Top trace is theexperimental spectrum of [Np·(H2 O)3 ] ·Ar8 acquired by monitoring the loss of 5
Ar atoms. Bottom traces are calculated conformers and their predicted infrared
spectra. The second trace from the top is a combination of conformers 3-1 and 3-2.
The dashed lines show the hydrogen bonds. . . . . . . . . . . . . . . . . . . . . . . . 123

6.5

Experimental spectrum and calculated spectra of [Np·(H2 O)4 ] . Top trace is theexperimental spectrum of [Np·(H2 O)4 ] ·Ar2 acquired by monitoring the loss of 2 Ar
atoms. Bottom traces are calculated conformers and their predicted infrared spectra.
The dashed lines show the hydrogen bonds. . . . . . . . . . . . . . . . . . . . . . . . 126

6.6

Experimental spectrum and calculated spectra of [Np·(H2 O)5 ] . Top trace is theexperimental spectrum of [Np·(H2 O)5 ] ·Ar2 acquired by monitoring the loss of 2 Ar
atoms. Bottom traces are calculated conformers and their predicted infrared spectra.
The dashed lines show the hydrogen bonds. . . . . . . . . . . . . . . . . . . . . . . . 127

6.7

Experimental spectrum and calculated spectra of [Np·(H2 O)6 ] . Top trace is theexperimental spectrum of [Np·(H2 O)6 ] ·Ar2 acquired by monitoring the loss of 2 Ar
atoms. Bottom traces are calculated conformers and their predicted infrared spectra.
The dashed lines show the hydrogen bonds. . . . . . . . . . . . . . . . . . . . . . . . 129

A.1 Calculated infrared spectra in the region of the partially reduced antisymmetric CO
stretch of the formate core for [Au(CO2 )4 ]clusters. The experimental spectrum

for [Au(CO2 )4 ] is shown in the top panel. The calculated spectra are shown in thelower panels with their corresponding structures and relative energies given to theright. Calculated OCO bond angles range from 139.3 to 142.6 for the bonded CO2
moiety and 175.2 to 178.0 for the solvent CO2 molecules. . . . . . . . . . . . . . . 148

xxiii
A.2 Calculated infrared spectra in the region of the partially reduced antisymmetric CO
stretch of the formate core for [Au(CO2 )5 ]clusters. The experimental spectrum

for [Au(CO2 )5 ] is shown in the top panel. The calculated spectra are shown in thelower panels with their corresponding structures and relative energies given to theright. Calculated OCO bond angles range from 139.3 to 142.6 for the bonded CO2
moiety and 175.2 to 178.0 for the solvent CO2 molecules. . . . . . . . . . . . . . . 149
B.1 Calculated spectra of bare core isomers for [Ni(CO2 )2 ] clusters. All of the isomers
presented here are discussed in Chapter 3.3.3. . . . . . . . . . . . . . . . . . . . . . . 151
B.2 Additional calculated spectra of bare core isomers for [Ni(CO2 )n ] clusters. Isomers
B and F are discussed in Chapter 3.3.3. Isomers H and I were determined not tocontribute to the experimental spectra and were not included in the discussion. . . . 152
B.3 Estimated barrier height for the conversion of isomer E* to D. Due to the barrierheight, we do not expect this isomer conversion process to occur in our experiment. . 153

Chapter 1
Introduction and Motivation

Ions are ubiquitous to chemistry. They are part of many chemical systems from crystalline
lattices to interstellar species to chemical reaction intermediates. [1–3] Because of the charged
nature of ions, they exhibit di↵erent behavior than their neutral analogs in many chemical settings.
One area of chemistry where the di↵erences between ions and neutrals is most noticeable is whenstudying solvation e↵ects. Solvation is present in many chemical reactions where the products andthe efficiency of the reaction can be dependent on the identity of the solvent species. Solvent cansometimes interact with ions in predictable ways like theanions and the+ H atoms orienting toward the Cl

O atoms to the Na+ cations when NaCl is dissolved in water. Other times thesolvent-ion interactions are less predictable. For example, when solvating SF6 with H2 O, the firstwater molecule makes a hydrogen bond to one of the F atoms localizing the excess charge on theinteracting F atom. When a second water molecule is added, it hydrogen bonds to the same F
atom rather than solvating any of the F atoms of the SF6 anion or forming a water-water hydrogenbond. [4] By studying the structure and charge distribution of gas phase anionic complexes, we
can understand the molecular level interactions. By also investigating solvation e↵ects on anioniccomplexes, understanding the role of solvent on a molecular level can be improved.
The study of molecular level properties of ionic systems in situ is difficult due to complications such as speciation, short lifetimes and low concentrations of chemically relevant transient
species. These issues can be avoided by performing experiments on cluster anions in vacuo rather
than in solution. There are a number of techniques that can be used for the generation of gas

2phase clusters. These techniques allow for control over the composition of the cluster species andthe solvation environment. These advantages grant the ability to generate and study spectroscopically model chemical systems. An attractive method of study is infrared spectroscopy. Within
the molecular vibrations are encoded the structure and charge distribution of the species being
probed. The e↵ects of solvation on these properties are included within the evolution of the spectra as the solvent environment changes. If the solvation environment can be made large enough,measurements of or extrapolation to behavior in bulk solution is possible. [5, 6] Using computational chemistry to simulate the structures, vibrations and charge distributions of clusters can be a
powerful complement to infrared spectroscopy. These two techniques used in tandem can identify
structures and the e↵ects of solvation on the core charged species.
This thesis focuses on the interactions of two di↵erent species in the presence of an excess
electron and how solvation by various solvents a↵ects those interactions. Chapter 2 details theexperimental apparatus and techniques used to study gas phase anionic clusters. Due to the low
concentrations of anions in mass selected ion beams, direct absorption infrared spectroscopy is not
feasible. Instead infrared photodissociation spectroscopy is used to take spectra of the clusters.
Chapter 3 presents the study of five di↵erent transition metal anions interacting with multiple
CO2 molecules. These results are generalized into trends of the observed charge carrier and thee↵ects of solvation by CO2 on the charged species. These clusters are treated as model systems
for single atom catalysts for the reduction of CO2 . Chapter 4 is an extension of Chapter 3 where
the structural motifs of two di↵erent transition-metal oxide anions solvated by CO2 are presented.
The last two Chapters (5 and 6) deviate from metal containing clusters. Chapter 5 presents a
vibrationally induced charge transfer reaction in a bimolecular system, while Chapter 6 looks at
the interaction of 1 - 6 H2 O molecules with a naphthalene anion and can be viewed as a model
system for the wetting of graphene.
All of the studies discussed in this thesis were accompanied by density functional theory
calculations. However, since the parameters of the calculations are di↵erent for each study, thecomputational details will be presented and discussed in the corresponding Chapter or Section.

Chapter 2
The Experiment2.1

Overview of Experiment
In general, the experimental setup generates gas phase cluster anions that are mass selected

and probed with tunable infrared radiation. The setup is based on a design by Lineberger andcoworkers. [7] Clusters are generated by entrainment of either laser vaporization or pulsed gaseous
species into a supersonic expansion of a carrier gas (here either Ar or CO2 ). The anionic clusters arethen accelerated perpendicular to the propagation direction of the expansion into a Wiley McLarentime of flight mass spectrometer where they are mass selected based on their mass to charge ratio.
A single cluster size is then irradiated by the tunable output of a optical parametric converter. The
generation of daughter fragment ions is monitored as a function of photon wavelength resulting in
an infrared photodissociation action spectrum. The experiment operates at a repetition rate of 20
Hz.

2.2

Experimental Details

2.2.1

Cluster Generation

There are a number of ways to form gas phase molecular clusters. One of the most common
ways is to use pulsed supersonic expansions of gas phase species to create cold molecular clusters
in vacuo. Depending on the expansion gases used, their stagnation pressures and the mechanicalshape of the nozzle as well as di↵erent expansion conditions, a variety of cluster species can be4formed. Electron bombardment of these clusters with high energy electrons (EKE ⇠800 eV) can
ionize molecules in the expansion creating slow secondary electrons. These slow secondary electrons
can then attach themselves to a molecular cluster forming (through evaporative cooling) anionicmolecular clusters (see eqn. 2.1).

[Xn Ym ] + e ! [Xn p Ym q ] + pX + qY

(2.1)

Alternatively, the electrons could attach themselves to an individual molecule generating ananion that can then go on to collide with other molecular species in three-body collisions. These
species can then bind to the anion creating clusters. A cluster such as [Xn Ym ]can further besolvated in a noble gas such as Ar, Ne or Kr forming a species of the form [Xn Ym ] ·Ark . If
a molecular cluster is formed prior to electron attachment, the electron binding energy must be
dissipated upon electron attachment. This can occur in a number of ways. If the cluster constituentsare bound by more than the binding energy of the electron, then the cluster will remain intact but
will be considered a “hot” cluster. Here the energy released will be stored in vibrational or rotationalmotion or electronic excitation. Without a way for the energy to dissipate from the cluster, it islikely that the excess electron will be ejected again and the anion will not survive. If there arecomponents bound to the cluster by less than the binding energy of the excess electron, they will
likely be evaporated to dissipate the energy of attachment. This has been described before as anevaporative ensemble, [8] where the energy will dissipate through evaporative cooling until there
is not enough internal energy remaining to “boil o↵” the most weakly bound remaining cluster
constituent. The internal energy of the cluster can be approximated to be roughly of the order ofthe binding energy of the last molecule that was evaporated. Since the internal cluster temperature
is proportional to the binding energy of the weakest bound constituent, solvation with noble gas
atoms generally leads to colder clusters, as they have lower binding energies. In turn, colder clusters
tend to exhibit more well resolved spectra, simplifying data analysis and potentially allowing forthe detection of previously unobservable features. The process of electron attachment to an Ar

5solvated molecular cluster is given in equation. 2.2.

[Xn Ym ] · Ark + e ! [Xn Ym ] · Ark l + lAr(2.2)

The general experimental setup used in all of the experiments described here is based on adesign by Lineberger and co-workers [7] and has been described extensively by Weber. [9] For theseexperiments two di↵erent source setups have been used to generate the clusters of interest.

2.2.1.1

Pulsed Valve Entrainment

In the first configuration of the ion source, two Series 9 General Valves are used to introduce
a molecular species of interest into an entrainment block (see Figure 2.1). After the entrained gases
are allowed to mix for 1 to 10 msec, a supersonic expansion of a carrier gas (Ar or CO2 ) generated by
an Even-Lavie pulsed valve [10] creates a supersonic expansion through the entrained gas mixture.
During the entrainment process, the species X and Y are pulled (entrained) into the expansion ofcarrier gas (this is shown by the red arrows in Figure 2.1). [7, 11] The expansion propagates out of
the entrainment block and into the source chamber. As the expansion propagates, collisions of thespecies within the expansion result in collisional cooling and aggregation into molecular or ionicclusters. The propagation distance is approximately 5 cm and can be varied by a few millimeters.
Changes to this distance a↵ect the amount of time the expansion can propagate, which a↵ects thecomposition of the clusters formed. The expansion is bombarded with a high energy (800 - 1000eV) beam of electrons generated from a resistively heated tungsten filament. The beam of electrons
is focused into the high density region of the expansion using an Einzel lens, electrostatic deflectors
and a ring magnet. The resulting electron impact plasma contains slow secondary electrons which
can attach to individual molecules or molecular clusters generated by the expansion. The timingof the gas pulses is controlled by a Quantum Composers delay pulse generator. Pressures in thesource chamber are typically on the order of low 10 6 mbar to 1 ⇥ 10 4 mbar. This pressure range
is measured by a nude Bayard Alpert ion gauge where the chamber is being evacuated by a 1000
L·s 1 TURBOVAC 1000c turbomolecular pump and a ⇠2500 L·s 1 Edwards di↵usion pump.

6

Figure 2.1: Schematic for the entrainment setup. X and Y represent di↵erent molecular species
and are pulsed into the entrainment region using series 9 General Valves. The red arrows denotethe general trajectories of the entrained species (here X and Y) into the carrier gas expansion. The
carrier gas (argon) is pulsed out using an Even-Lavie pulsed valve. The black arrow represents the
trajectory of the high energy electron beam.

This ion source setup was used to generate Ar-tagged nitromethane-iodomethane clusters((CH3 NO2 ) ·CH3 I·Arn ) and naphthalene-water clusters ([Np(H2 O)n ] ·Arm ) discussed in Chapters
5 and 6, respectively.

7
2.2.1.2

Laser Vaporization

The previously described entrainment setup is unable to easily generate molecular clusters
containing metal atom(s) due to the exceptionally low vapor pressure of most metals. To circumvent
this issue, laser vaporization is employed to create metal vapor. In this case, one of the General
Valves present in the entrainment setup is replaced with a rotating metal rod driven by a DC motor.
A metal disc is then mounted to the end of the rod as shown in Figure 2.2. The third harmonicof a nanosecond pulsed Nd:YAG laser (355 nm) is used to vaporize material from the edge of therotating metal disc. By rotating the disc, generally even wear across the edge of the disc is achieved
resulting in a more stable ion beam. The vaporization occurs about 1.5 cm o↵ to the side of the
Even-Lavie valve nozzle (see Figure 2.2). Within 300 to 500 µseconds of the vaporization laser
firing, a supersonic expansion of CO2 is generated from the Even-Lavie valve (stagnation pressure
typically 5.5 bar). In this source configuration, the vaporization process already creates a sufficientquantity of slow secondary electrons so the electron gun is not required for anion cluster formation.
The vaporization setup typically results in the formation of metal-(CO2 )n clusters as well as pure
(CO2 )n clusters, metaloxide-(CO2 )n clusters and some dimetal-(CO2 )n clusters for cobalt andnickel. The timing of the gas pulse and the vaporization laser pulse are controlled by synchronized
Quantum Composer pulse delay generators.

2.2.2

Time of Flight Mass Spectrometer

In both laser vaporization and entrainment setups, clusters with anionic, cationic and neutral
charges are produced. To separate out only the anionic species, the source chamber is set up to
work in tandem with a time of flight mass spectrometer. After the molecular beam has propagatedinto the acceleration region of the time-of-flight mass spectrometer, it is accelerated perpendicular
to the direction of propagation by two electrostatic reflectors operating in a Wiley-McLaren field
configuration (see Figure 2.3). This is done by pulsing two plates to -4000 V (repeller) and -2800
V to -3300 V (extractor). This accelerates the anionic species into the flight tube with the same

8

Figure 2.2: Schematic for the laser vaporization entrainment source. Here, one of the General
Valves in Figure 2.1 is replaced with the rotating metal disc (see text). The laser enters on theleft hand side of the setup and is focused onto the edge of the metal disc using a small lens. The
metal vapor is expanded in a supersonic expansion of CO2 . No electron beam was required in theseexperiments.

average kinetic energy and allows for the spatial focusing of ions with the same mass to charge
ratio at the first space focus, whose position is controlled by the voltage on the extractor plate.
The first spatial focus is at the laser interaction region (see Figure 2.3). The ion beam is steered
through the flight tube by two sets of electrostatic reflectors and an Einzel lens. The ion beam isthen reversed and refocused onto a multichannel plate (MCP) detector using a reflectron secondary
mass analyzer. Typical pressures in the flight tube are on the order of 10 8 mbar yielding a mean
free path on the order of 104 meters. The total ion trajectory length is approximately 3 meters.
Since all of the anionic species generated in the source are given the same kinetic energy
in the same direction from nearly the same starting point and they all have the same distance totravel to the detector, their flight time and mass to charge ratio are related as follows (for species

9Figure 2.3: Schematic of the experimental apparatus including the vaporization source with theelectron gun shown for completeness. The green line indicates the ion beam trajectory and all
essential parts of the apparatus are labeled and discussed further in the text.

of the same charge):

t1
=
t2

r

m1
.
m2

(2.3)

Assuming that all ions are singly charged, the mass of all the constituents of the ion beam
can be identified. With the masses known and the precursor species used in the expansion known,
the chemical composition of the mass peaks can be determined. As an example, Figure 2.4 showsa portion of the mass spectrum acquired using a cobalt disc as a vaporization target. We areable to identify progressions owing to (CO2 )n clusters, [Co(CO2 )n ]ping [CoO(CO2 )n ]and [Co2 (CO2 )n 1 ]clusters and nearly overlap-

clusters. Mass peaks corresponding to the formationof dianions were not observed in any of the studies presented. The mass resolution of the mass
spectrometer ( mm ) is on the order of 500.
Following identification of the mass peaks in the ion beam, a mass gate can be used to
select a small mass window ( mm ⇡ 15) for spectroscopic interrogation. The mass gate operates

10
Co- 1

3

2
3

0

0

1

100

4
5

4
2

5

3

200

6

6
4

7
5

300

8

7

9

10 11 12 13 14 15 16 17 18 198

9 10 11 12 13 14 15 16 17 18 19 20

6

7

8

[Co(CO2)n](CO2)-n[Co2(CO2)n-1]-

9 10[CoO(CO2)n]-

400

500

600

700

800

900

1000

mass/charge [m/z]
Figure 2.4: Mass spectrum observed using the vaporization source with a cobalt disc and CO2 as
the carrier gas.

by generating an electric field on a charged set of plates that disrupts the flight path of the ion
beam. By switching the field to zero for an appropriate time window (⇠1 µsec), a selected cluster
size range is allowed to continue on the normal ion beam trajectory. The plates are then recharged
and the remainder of the ion beam is deflected. The timing of the mass gate is controlled by a
Quantum Composer delay pulse generator.
Absorption of an infrared photon results in excitation of a vibrational mode in the cluster.

11
Through anharmonic coupling and the resulting intramolecular vibrational relaxation, the energyof the initially excited vibrational state is redistributed over the cluster constituents. If enoughenergy is absorbed, a species possessing a lower binding energy than that of the absorbed photon
can dissociate. If dissociation occurs prior to the ions reaching the reflectron, the lighter daughterions will be turned around and impinge on the detector earlier than the undissociated parent ions.
The reflectron voltage can then be scaled to the ratio of the daughter mass divided by the parent
mass to allow unambiguous assignment of the daughter species. The generation of daughter ions canbe detected as a function of the laser wavelength yielding an infrared photodissociation spectrum of
the parent cluster (see equation 2.4 as an example of photodissociation of a metal-(CO2 )n cluster
ion).

[M(CO2 )n ] + h⌫ ! [M(CO2 )n m ] + mCO2

(2.4)

In summary, the order of events for cluster generation followed by interrogation by tunable
IR light is as follows: 1.) General Valves emit gas into the entrainment block/Vaporization laser
fires creating metal vapor 2.) Carrier gas is expanded out of the Even-Lavie valve 3.) Anionicclusters are accelerated into the Time-of-Flight mass spectrometer 4.) A small mass range of ionsis transmitted using a mass gate 5.) The selected anions are irradiated by tunable IR light 6.)
Remaining parent and newly formed daughter anions are turned around by the Reflectron andimpinge upon a MCP detector

2.2.3

Laser System

The mass selected ions are irradiated by the tunable output of a LaserVision OPO/OPA(optical parametric oscillator/optical parametric amplifier) system with an additional di↵erence
frequency mixing stage, operating in the range of 600 cm 1 to 4400 cm 1 . The OPO/OPA system
is pumped by approximately 500 mJ of 1064 nm light generated from an Innolas Spitlight 600
Nd:YAG laser operating at 20 Hz with a pulse duration of 7 ns. This laser system has two scanning
regions, mid IR and far IR. The mid IR region is from 2260 cm 1 to 4400 cm 1 . Light in this12
wavenumber range is generated by doubling part of the 1064 nm pump radiation, generating 532
nm light. The 532 nm light is then converted into signal and idler waves by two KTP crystals
in the oscillator part of the OPO (see Figure 2.5 and equation 2.5). Since signal and idler haveorthogonal polarizations, the signal wave is rejected by a polarizer while the idler wave is allowed
to oscillate. The idler wave then exits the OPO resonator, its beam profile rotated 90 by a doveprism, and overlapped with the rest of the 1064 nm light in the OPA, which consists of a chainof four KTA crystals. Di↵erence frequency mixing of the idler wave and the 1064 nm pump light
results in the generation of mid IR light in a wavenumber range of 2260 cm 1 to 4400 cm 1 (seeequation 2.6).

2 ⇥ ⌫1064 = ⌫signal + ⌫idler

(2.5)

⌫1064

⌫idler = ⌫mid IR(2.6)⌫idler⌫mid IR = ⌫far IR(2.7)

The far IR tuning range is from 600 cm 1 to 2150 cm 1 . Light in this range is generated bydi↵erence frequency mixing of the mid IR output with the amplified idler wave (see equation 2.7).
The output power of the OPO/OPA is between 5 mJ and 20 mJ of mid IR light and 10 µJ to 100µJ of far IR light depending on the wavelength of the photons being generated. The bandwidth ofthe laser system is 2 cm 1 .
Irradiation occurs in a multi-pass cell based on a design by Liu and coworkers (see Figure2.6). [12] The multi-pass cell is made out of OFHC copper for high IR reflectivity and allows
adjustment of the number of passes of the laser across the ion beam from 1 to 14 (estimated). The
multi-pass cell can be rotated, which allows the user to choose the number of passes. Irradiating
the ion packet with multiple passes of the IR light generates additional fragment ions compared
to only one pass of the IR light and improves the signal to noise ratio of single scans by up to afactor of four. The ability of the cell to rotate provides a convenient way to test for multiphoton
processes by reducing the number of passes such that the intensity of a peak attributed to a single

131064 nm Pump Laser

High
Reflector

OPO (KTP)

Dove
Prism

Polarizer

OPA (KTA)

OutputCouplerBBO
AgGaSeGe Filter
Beam Splitter

Polarizer

To Multipass Cell

Figure 2.5: Schematic diagram of the Laservision optical parametric converter setup used for thegeneration of tunable infrared radiation in these experiments.

photon process is reduced by half. We can then monitor the intensities of all the other peaks inour spectra. All other single photon process will also reduce by half while any multiphoton process
will reduce by N12 where N is the number of photons in the process. This test was performed forall of the metal-(CO2 )n and metal-oxide-(CO2 )n cluster studies.
2.2.4

Data Collection and Analysis

Spectra are collected by centering a Tektronics TDS 2022 digital oscilloscope on a fragment
peak. The time window of the oscilloscope is adjusted so that it is nearly the same width as themass peak (⇠100 nsec). The average signal in the active window is recorded by the oscilloscope andaveraged for a di↵erent number of laser shots depending on the scanning rate (discussed below).
A LabVIEW program reads out the averaged signal and records it while the laser is scanned.
There are on average between 100 and 1000 parent ions per shot for a given cluster species with

14

Figure 2.6: Schematic diagram of the multi-pass cell ring reflector used to increase the number ofpasses of IR light on the mass selected ion packet. See text for discussion.

larger clusters and metal containing clusters having lower parent ion intensities compared to otherspecies. Photodissociation efficiency is around 5% meaning that there are at most ⇠50 daughterions generated. Due to this low number of fragment anions, the individual spectra acquired can
exhibit a poor signal to noise ratio. The spectra are collected in ⇠1 cm 1 steps averaging 16 shots(fast mode) or 64 shots (slow mode) of the experiment over a specified tuning range. The slow
scanning rate yields a higher signal to noise ratio on individual scans but leads to data acquisition
times that are four times longer compared to the fast scanning rate.
A number of spectra are taken for a given cluster species and fragmentation channel over

15
multiple days and averaged into one spectrum. This is done to improve the signal to noise ratio of
the average spectrum and ensure reproducibility from day to day. On average, 10 to 20 individual
spectra are averaged together. Each individual spectrum acquired by monitoring the generationof daughter ions is corrected for background signal. The final average spectrum is both corrected
for laser fluence by measuring the output power of infrared light as a function of frequency andcalibrated using an Ocean Optics HR2000+ spectrometer to measure the wavelength of the signal
wave. All data analysis is carried out within the OriginLab program.

Chapter 3
Infrared Studies of Metal-(CO2 )n Anion Clusters

This Chapter has been reproduced in part with permission from the following previouslypublished articles:
Knurr, B. J., Weber, J. M., “Solvent-Driven Reductive Activation of Carbon Dioxide byGold Anions.” Journal of the American Chemical Society, 2012. 134: p. 18804 - 18808. DOI:
10.1021/ja308991a. Copyright 2014 American Chemical Society.
Knurr, B. J., Weber, J. M., “Solvent-Mediated Reduction of Carbon Dioxide in AnionicComplexes with Silver Atoms.” Journal of Physical Chemistry A, 2013. 117: p. 10764 - 10771.
DOI: 10.1021/jp407646t. Copyright 2014 American Chemical Society.
Knurr, B. J., Weber, J. M., “Infrared Spectra and Structures of Anionic Complexes Cobalt
with Carbon Dioxide Ligands.” Journal of Physical Chemistry A, 2014. 118: p. 4056 - 4062. DOI:
10.1021/jp503194v. Copyright 2014 American Chemical Society.
Knurr, B. J., Weber, J. M., “Interaction of Nickel with Carbon Dioxide in [Ni(CO2 )n ]Clusters Studied by Infrared Spectroscopy.” Journal of Physical Chemistry A, 2014. 118: p. 8753
- 8757. DOI: 10.1021/jp507149u. Copyright 2014 American Chemical Society.
Knurr, B. J., Weber, J. M., “Structural Diversity of Cu-CO2 Complexes - Infrared Spectra
and Structures of [Cu(CO2 )n ] Clusters.” Journal of Physical Chemistry A, 2014. 118: p. 10246 10251. DOI:10.1021/jp508219y. Copyright 2014 American Chemical Society

17

3.1

Introduction
The creation of sustainable energy sources has been one of the most pressing recent scientificendeavors. On the assumption that our fuel economy will depend on carbon-based fuels for theforeseeable future, development of carbon neutral fuel cycles are essential for the environmental
and economic health of society. Carbon dioxide is a byproduct of many industrial processes as wellas combustion reactions involving fossil fuels. CO2 has also been shown to be a potent greenhousegas with a number of negative environmental e↵ects attributed to its increased emission. [13–15]
It follows that research and the development of techniques aiding in the creation of carbon neutral
processes are of importance to global sustainability.
One major avenue being explored is the sequestration and recycling of CO2 generated in
combustion reactions. Recovered CO2 can be converted into chemically usable fuels or feedstock
for the generation of fuels. Virtually any scheme for the conversion of CO2 involves reduction of
the CO2 , where the most simple case ends in the formation of formate (HCOO ). [16] This simple
scheme requires the transfer of two electrons and one proton to a CO2 molecule. A more recently
developed and promising avenue toward CO2 conversion involves a six-electron process whereby
CO2 is converted to methanol in the presence of pyridine based homogeneous catalysts. [17, 18]
A number of other approaches, typically involving metal based catalysts, have been used tomediate CO2 reduction. [19–26] Of particular interest are systems where the catalyst is composedof a single metal atom or a small metal cluster. [27] In these cases, the catalysts are supported on
surfaces [28–31] or embedded into supramolecular assemblies [32] or multiligand systems. [19,20,33]
A number of these studies have shown conversion from CO2 into CO or small organic species (i.e.
CH3 OH, CH4 , etc.). However, the mechanistic, molecular level details of these processes are notwell understood.
In the processes mentioned above, a CO2 molecule will undergo reduction, where the simplest
starting point is one-electron reduction. One-electron reduction of an isolated CO2 molecule isenergetically very expensive with a cost of ca. 0.6 eV. [34] In addition to the large energetic cost, a18free CO2 anion is only metastable where the longest observed lifetime is 90 µs. [34] However, theformation of a complex between the CO2 molecule and a catalyst can aid in both charge transfer
and stabilization of the activated species. In this case, a fractional amount of the excess charge
would be transfered to the CO2 and further interaction with reaction partners could complete thereduction process and subsequent chemical reactions.
Addition of an excess electron to CO2 in a linear geometry leads to a population the 2⇡uorbital (see Figure 3.1). However, by following the curves of the Walsh diagram [35] to minimize
the energy of the system, it is clear that the anion will rearrange from a linear geometry with
(1⇡g )4 (2⇡u )1 occupation to a bent geometry with a (1a2 )2 (4b2 )2 (6a1 )1 configuration. This weakens
the ⇡ interactions present in the neutral CO2 molecule, resulting in a lengthening of the CO bonds
from 1.162 Å in the neutral species [36] to ⇠1.24 Å in the anion. In complexes containing CO2 , theOCO bond angle depends on the amount of excess charge localized on the CO2 unit and can varyfrom 180 for the neutral to ⇠135 for the anion. These changes in bond strength and geometry
are reflected in the frequency of the antisymmetric CO stretching mode of CO2 . In neutral CO2 ,the antisymmetric stretch is observed at 2349 cm 1 . [37] In a CO2 anion, the frequency of thesame vibrational motion has been observed in a Ne matrix at 1658.3 cm 1 . [38] An antisymmetricstretch observed between these two extremes encodes the structure and charge distribution on theCO2 and, as will be shown in this chapter, can be used as a spectroscopic probe of the reduction
process.
Reduction of CO2 to convert it into usable chemicals usually occurs in solution. The condensed phase environment complicates the interpretation of spectroscopic probes of reaction intermediates and short lived transient species that are characteristic of the reaction mechanisms.
Due to speciation and the short lifetime of many reaction intermediates, it is difficult to definitively
isolate the spectral signatures of relevant species in situ. Experiments in vacuo a↵ord spectroscopic
characterization of well defined molecular targets. This allows for the identification of potential
reaction intermediates and can aid condensed phase research by elucidating the chemical processes
underlying their reactions.

19

Figure 3.1: The Walsh diagram for CO2 . [35, 39] Reproduced from [40] with permission of The
Royal Society of Chemistry. See text for discussion about the CO2 anion and the CO2 neutral.

While the presence of a complicated chemical environment makes spectroscopy in situ difficult, it is necessary for CO2 reduction reactions and can have profound e↵ects on the relevant
species and chemistry. In the context of the present work we are interested in the nature of thecharge carrier in metal-CO2 of the form [M(CO2 )n ]clusters (M = Au, Ag, Co, Ni, Cu), and in

the e↵ects of solvation (here the solvent is CO2 ) on the structure and charge distribution of theionic “core” species. We hope that this body of work will aid in rational catalyst design and agreater understanding of the chemical processes underlying CO2 reduction using metal containing
catalysts.

20

3.2

[Au(CO2 )n ] and [Ag(CO2 )n ] Clusters
There has been a substantial amount of interest in Au as a catalytic material. Gold clusters

supported on surfaces have been shown to be catalytically active toward the oxidation of CO to
CO2 as well as the reduction of a single CO2 molecule. [27, 41–44] In modeling the binary Au-CO2
complex in the context of the oxidation of CO to CO2 mediated by AuO, two types of interaction
were predicted. [45] First, the two could interact as an ion-molecule pair where the CO2 moleculesolvates a Au

anion due to the partial positive ( + ) charge on the carbon atom being attracted

to the Au . Additionally, the repulsive force of the partial negative (atoms would push them away from the Au) charge on the oxygenanion. These interactions would lead to relatively

small deformation of the CO2 molecule and the CO antisymmetric stretch would be only weakly
perturbed. Alternatively, the carbon atom of the CO2 molecule could form a covalent bond with
the Au anion. This interaction would create a y-shaped complex reminiscent of a formate anion(HCOO ) where the hydrogen has been replaced by the Au atom. This structure will be referred
to as the formate motif. The formate motif is structurally similar to the proposed key intermediate
for CO2 reduction by a pyridine anion where a C-N bond is formed between the CO2 and the N
atom (C5 H5 N-CO2 ). In this configuration, the geometry change of the CO2 ligand towards thatof the anion (see Figure 3.1) suggests that the formation of a Au-C bond would facilitate chargetransfer to the CO2 . As a consequence of this uptake of excess charge, there would be a substantial
redshift of the CO antisymmetric stretching frequency providing a distinct spectroscopic signature
to observe.
Spectroscopic studies have been performed on the binary [AuCO2 ] complex to characterize
the type(s) of interaction present. [46] Even though the ion molecule complex had been predicted
to be slightly lower in energy [45], only the covalent interaction was observed though multiphoton
photodissociation spectroscopy. [46] This result contrasts with those obtained for the interactionsof most halide anions with CO2 (with the exception of F [47]) where the CO2 was always found tosolvate the halide anion with minimal charge delocalization onto the CO2 molecules (see Figure 3.2).

21

600

-

Au
(calc)-1

red shift [cm ]

500
400
300
200
100

Br

-

I

Cl

-

-

0
0

100

200

300

400

500

600

binding energy [meV]Figure 3.2: The observed redshift of the antisymmetric CO stretch in CO2 when interacting with
Au and various halides. [46, 48] As is indicated on the graph, the redshift of the antisymmetricstretching frequency of CO2 interacting with Au is calculated while the other redshifts have been
experimentally measured. The redshifts are all with respect to 2349 cm 1 . [37][48] Both Au and the halides possess a closed outer sub-shell with comparable electronegativities(specifically 2.54 for Au and 2.66 for I). [49] However, the interaction of Au
a reactivity observed only in [F(CO2 )]with CO2 showsclusters while the other halides remain closed sub-shell

species solvated by CO2 molecules.
While an understanding of the binary complex is useful, the addition of solvent molecules is anecessary next step since any industrial process is likely to be carried out in a solvent environment(e.g. supercritical CO2 , H2 O, alcohols, etc.). By building clusters one solvent molecule at a time,
we can study the cumulative e↵ect of solvation on the metal-CO2 anion core.

22
Silver surfaces and nanoparticales have been tested as catalysts in the electrochemical reduction of CO2 . [50–55] In these studies, a number of species, including CO, CH3 OH and HCOO ,were formed. However, there is a lack of mechanistic understanding of the underlying chemical
processes leading to formation of these species. There have been e↵orts to characterize the surface
composition of Ag coated substrates [52] but in general there is a deficiency of spectroscopic studies
on the catalytic processes that facilitate CO2 reduction.
The electron configuration of Ag is [Kr]4d10 5s1 which is very similar to that of Au whichis [Xe]4f14 5d10 6s1 . However, many of the properties of these two atoms are quite di↵erent. For
example, the electron affinity of Au is 2.309 eV [56] while the electron affinity of Ag is 1.303 eV. [57]Similarly, Au has a greater electronegativity and ionization energy than Ag but a similar atomic
radius. This and other di↵erences can be attributed to larger relativistic e↵ects in Au compared to
Ag. Superficially, one might expect that the interactions of CO2 with Ag will be similar in nature
to those of Au but, due to the di↵erences in atomic properties, the charge distributions and e↵ectof solvent on the core species are likely to di↵er.

3.2.1

Computational Methods

For [Au(CO2 )n ] and [Ag(CO2 )n ] clusters, a variety of levels of theory and basis sets were
employed. The TURBOMOLE v. 5.9.1 and 6.2 suite of programs were used for all calculations. [58]For all calculations where vibrational spectra are reported, we employed density functional theory
using the B3-LYP functional [59, 60] and the def2-TZVPP basis sets for all atoms. [61] We also
performed exploratory calculations employing the PBE0 functional [62] and RI-MP2 level [63]of theory. Comparison of calculations performed with the def2-TZVPP basis sets [61] to def2TZVP [64] and aug-cc-pVTZ basis sets [65] with di↵erent functionals or levels of theory yieldedthat the B3-LYP functional with def2-TZVPP basis sets for all atoms was the best compromise
between computational cost and accuracy for this work. All vibrational frequencies were calculatedusing the AOFORCE program. [66, 67] The calculated vibrational frequencies corresponding to
CO2 molecules bonded to a metal atom were scaled by 0.9380 to account for anharmonicity. This

23
scaling factor as obtained by comparing the calculated value of the antisymmetric CO stretch in
AuCO2 with the high level calculations of Boese et al. [46] Frequencies corresponding to weakly
perturbed CO2 molecules were scaled by 0.97537. This factor was generated by comparing the
calculated frequency for a free CO2 to the measured antisymmetric stretch of CO2 reported by
Shimaniouchi. [37] All partial charges were calculated using a natural population analysis. [68] Allof the energies reported are zero-point corrected.
To model the e↵ects of solvation beyond the first solvation shell in both [Au(CO2 )4 ][Ag(CO2 )n ]and

(n = 2 and 4), we employed a conductor-like screening model (COSMO). [69–72]

We used a dielectric constant of 1.6 (corresponding to saturated liquid CO2 at 291 K) [73] and amolecular-shaped cavity generated using the default radii for carbon (2.0 Å), oxygen (1.72 Å) andsilver (2.223 Å). The results from the COSMO simulation do not include zero-point corrections.

3.2.2

Infrared Spectra of [Au(CO2 )n ]Clusters

To explore the e↵ect of solvation on the binary AuCO2 complex, we probed [Au(CO2 )n ]clusters by infrared photodissociation spectroscopy. Figure 3.3 shows the experimental spectra
obtained for [Au(CO2 )n ]clusters (n = 2 - 13) by observing the loss of one CO2 molecule. The

panel on the right side in the wavenumber range 2250 cm 1 to 2450 cm 1 contains signaturesof solvent type CO2 molecules. These CO2 molecules mostly retain their identity as free CO2
molecules and still exhibit an antisymmetric stretching frequency near 2349 cm 1 with only minor
deviation due to the presence of the negative charge. On the other hand, the left panel of Figure
3.3 shows signatures of partially reduced CO2 species. These CO2 units have lost their identity as
free CO2 molecules and their vibrational frequencies encode information about their structure androle in the charge carrying species.
An initial inspection of the experimental spectra reveals three distinct trends. First, thesignature of the partially reduced CO2 shifts to the red with increasing cluster size from n = 2- 7, 9. Second, the same partially reduced CO2 signature shifts back to the blue as cluster size
increases from n = 10 - 13. Third, the peaks corresponding to solvent CO2 molecules show multiple

24

2
3

intensity [arb. units]

4
5
6
7
8
9
10
11
12
13
1600 1700 1800 1900 2000 2100 2200 2300 2400 2500

wavenumber [cm-1]
Figure 3.3: Experimental spectra of [Au(CO2 )n ] clusters. The numbers denote the number of
CO2 molecules in the cluster. The right and left panes are individually normalized so they are ondi↵erent scales. The arrow indicates the predicted frequency of the AuCO2 binary complex fromthe calculations of Boese et al. [46] All spectra were obtained monitoring the loss of a single solvent
CO2 .

components at cluster sizes n

7. Additionally, [Au(CO2 )8 ]exhibits no signature of a partiallyreduced CO2 molecule and a rather complicated solvent signature compared to other cluster sizes.
These trends will be discussed one at a time, beginning with the observed redshift from n =2 - 9. As was discussed in the introduction to this chapter, strongly redshifted CO antisymmetricstretching frequencies indicate partial reduction of a CO2 molecule. Assuming that the core structure is of the form [AuCO2 ] assigned by Boese et al. [46], the redshift observed in the left panel of25
Figure 3.3 indicates that solvation is polarizing the excess charge density onto the partially reduced
CO2 molecule. Stated another way, initial solvation is increasing the degree of reduction of thepartially reduced CO2 unit.
To check the validity of this hypothesis we begin by considering the e↵ect of CO2 as a solvent
species. It has been observed before in halide and O2 ·CO2 cluster species that individual CO2
molecules interacting closely with the charged species can result in a small amount of the excess
charge becoming delocalized onto the solvent CO2 molecules. [48, 74] In short, the CO2 molecules
polarize the excess charge towards themselves and can take up a small amount of the charge without
forming covalent bonds. The analogy to this case would be that solvation around the CO2 moietyof the formate complex should result in polarization of the excess charge onto the covalently boundCO2 species and o↵ of the metal atom.
To test this theory, we performed DFT calculations at a variety of di↵erent levels of theory(see Section 3.2.1 for details). From converged geometric structures, we generated vibrationalfrequencies and compared them to the experimental spectra. Figure 3.4 shows a comparison of
the three lowest energy structures calculated for [Au(CO2 )2 ] . The structure shown in the bottom
panel of Figure 3.4 is termed symmetric solvation of the Au anion and is the result of solvation of
the Au atom (opposite the bonded CO2 ligand) of the binary core. Here the two CO2 molecules are
positioned on opposite sides of the Au atom and, as consequence, cancel the electron withdrawing
capabilities of each other. This results in a small amount of charge transfer onto both CO2 molecules
resulting in a ⇠20 cm 1 estimated redshift of the antisymmetric stretching mode and no activated
CO2 species. However, it is clear that this solvation motif does not recover the experimental featurearound 1810 cm 1 .
The other two structures presented in Figure 3.4 are examples of solvation around the [AuCO2 ]formate motif. It is obvious that the top two structures both recover the experimental

spectrum reasonably well but there are di↵erences. For the rest of the chapter, we refer to the top
structure as an example of a “side” solvation position while the structure shown in the middle panel
shows occupation of the “terminal” solvation position. If only side solvation is present, a smallerintensity [arb. units]

26

0 meV

6 meV

78 meV

1800

1900

2000

2100

2200

2300

2400

2500

wavenumber [cm-1]
Figure 3.4: Comparison of the three calculated lowest energy structures for [Au(CO2 )2 ]to the experimental spectrum for [Au(CO2 )2 ] . Relative energies are given in meV.

clusters

redshift of the bonded CO2 antisymmetric stretching frequency should be expected than when theterminal solvation position is occupied since the ability of the solvent CO2 to polarize the charge
will be greater when situated behind the bonded CO2 and along the Au-C bond axis. Calculations
corroborate this prediction and show that side solvation results in a lesser amount of the excess
charge being localized on the activated CO2 moiety than terminal solvation (-0.48 e vs. -0.55 e).

27This indicates, consistent with chemical intuition, that not only can CO2 act as polarizing solvent,but its location in the cluster alters its e↵ect on the solute. Here, we tentatively assign the sidesolvation isomer as the dominant solvation motif for [Au(CO2 )2 ] .
Based on our observations for [Au(CO2 )2 ] , solvation must be occurring around the CO2
moiety for cluster sizes n = 3

9. Interestingly, there are two peaks that begin to appear at n = 3,

both are very pronounced in n = 4 and only one peak is recovered at n = 5. The higher-energy
feature in n = 3 and 4 follows a constant 10 cm 1 shift with each additional CO2 from n = 2

4.

The lower-energy shoulder in n = 3 also shifts by 10 cm 1 to the lower-energy peak for n = 4 andagain by 10 cm 1 to the single peak in n = 5. No peak is recovered in n = 5 corresponding to a 10
cm 1 shift of the high-energy feature in n = 4. These observations suggest that there are uniquesolvation isomers populated at small cluster sizes that give rise to these di↵erent features.
Figure 3.5 shows a comparison of the calculated spectra for di↵erent solvation isomers of[Au(CO2 )n ] for (n = 3, 4). As can be seen, solvation in side positions (see middle panel on the left
side of Figure 3.5) results in a smaller predicted redshift than solvation in the terminal position (seebottom panel on left side of Figure 3.5). Comparing the theoretical and experimental spectra, it
becomes apparent that when the terminal solvation position is filled, the lower energy set of peaks
are recovered. When the terminal solvation position is left vacant, the higher energy peaks in the
spectra of n = 3 and 4 are predicted. At n = 5, the terminal position is always filled leaving only
the lower-energy peak. For the remaining cluster sizes, the terminal position is already filled soanother e↵ect must give rise to the multiple features observed at large cluster size (n7 discussed

below).
Calculations corroborate the idea that solvation around the CO2 moiety causes charge polarization onto the activated CO2 ligand and a concomitant redshift of the antisymmetric CO stretch.
If a blue shift is observed with increased solvation, then the additional CO2 molecules must be
occupying positions that polarize charge density o↵ of the CO2 moiety. At n = 9, it is likelythat all of the solvation positions around the bonded CO2 are filled. Subsequent addition of CO2
molecules would force them to solvate the Au atom as the solvation energy of occupying the firstintensity [arb. units]

28

1800

1900

2000
-1

wavenumber [cm ]

0 meV

0 meV

7 meV

16 meV

14 meV

48 meV

2100

1800

1900

2000

2100

-1

wavenumber [cm ]Figure 3.5: Comparison of the lowest energy structures for [Au(CO2 )n ] clusters (n = 3, 4). Relative energies are given in meV and are with respect to the lowest energy conformer of a given
cluster size. Since these structures are all within 60 meV of each other, they are considered to be
isoenergetic.

solvation shell of the Au moiety will be more favorable than the solvation energy of opening asecond solvation shell around the CO2 moiety. It is reasonable to assume that solvation around the
Au atom would polarize charge onto the Au and away from the CO2 unit based on what has already been seen for solvation around the bonded CO2 moiety. Additionally, calculations for higher
energy conformers of smaller complexes have already predicted this behavior (see bottom panel of
Figure 3.4) as solvation around the Au atom resulted in a deactivation of the bonded CO2 . Forn10, solvation is occurring around the Au moiety resulting in an observed blue shift as cluster

29
size increases because additional solvent can not be accommodated around the partially reduced
CO2 moiety. It is difficult to survey all of the possible solvation isomers for clusters larger thann = 5. Appendix A shows additional calculated structures for n = 4 and 5.
The next observation to discuss is the lack of a signature of a partially reduced CO2 in n = 8
clusters. In a photoelectron study, it was observed that Br ·(CO2 )n clusters (Brsize to Au ) exhibit closure of the first solvation shell around Br

is similar in

at n = 8. [47] Additionally,(CO2 )n clusters exhibit an abrupt change in the charge carrier as a function of cluster size. [75]Due to the electronic instability of the CO2 anion, (CO2 )n clusters initially possess a C2 O4 core.
At n = 6, the solvation energy gained by solvating a more compact CO2 core becomes greaterthan the electronic stability of the C2 O4 core and the charge carrier switches to CO2 . At n = 14
the first solvation shell of CO2 is filled and rather than expand into the second solvation shell,
it is energetically more favorable to switch back to the C2 O4 core so that the solvation energy ismaximized by the larger first solvation shell of C2 O4 compared to opening the second solvation
shell of CO2 . In [Au(CO2 )n ] clusters, a similar switch occurs but happens over only one cluster
size. At n = 8 the solvation energy of solvating the Au is greater than the electronic stability of
the formate core. For all other cluster sizes, the presence of the formate core is energetically morefavorable than solvation of a compact charge carrier.
Lastly, at n = 7 we observed the first splitting of the solvent feature and we retain at least
two solvent CO2 peaks for all larger cluster sizes. This is likely due to the position of the solvent
molecule with respect to the structure of the charge carrier. At n = 7, most of the space aroundthe activated CO2 moiety is probably filled since the blue shift is first observed two cluster sizes
later. At this cluster size, additional solvent molecules may begin to position themselves closer tothe Au atom. This could result in a di↵erent amount of charge uptake by various solvent molecules
depending on their position around the core resulting in multiple solvent signatures. While it isdifficult to definitively say, it is likely that for n

7, the first major change as to which part of theformate core is being solvated is occurring resulting in the multiple solvent CO2 peaks observed.
To briefly summarize, in [Au(CO2 )n ]clusters solvation occurs around the CO2 moiety at

30
n = 2 7, 9 with the terminal position(s) always being filled for n

5. This causes a polarization of

the excess charge onto the covalently bonded CO2 moiety. From n = 10

13, additional solvation

occurs around the Au atom resulting in a deactivation of the partially reduced CO2 . At n = 8, it ismore energetically favorable to solvate a Au anion rather than form the AuCO2 formate complex.
3.2.3

Infrared Spectra of [Ag(CO2 )n ]The experimental spectra of [Ag(CO2 )n ]Clusters
are shown in Figure 3.6 and were obtained bymonitoring the loss of one CO2 molecule. Again, there are two regions of the experimental spectra.
The higher energy region (2250 cm 1 - 2400 cm 1 ) shows signatures of weakly perturbed CO2
molecules. These features, like in [Au(CO2 )n ]clusters, are due to solvent CO2 molecules thatare not part of the charge carrier. The lower energy region (1600 cm 1 to 2150 cm 1 ) contains
signatures of the charge carrying species. There are three primary features in the spectra toconsider. First, there is a set of peaks in Figure 3.6, highlighted by the red dashed line, that shift
to the red with increasing number of CO2 molecules. Second, within the first few cluster sizes
probed the solvent feature changes shape and shifts to higher energy. Lastly, there is a relatively
stationary peak highlighted by the dashed blue line at ca. 1660 cm 1 that was not observed in[Au(CO2 )n ] clusters. Again, we will interpret these features one at a time.
We will first analyze the red shifting feature. Since fragmentation is first observed at n = 2,
it is reasonable to assume that the core species is of the form [AgCO2 ]since the binding energyof a CO2 bonded to the metal atom is calculated to be on the order of 0.5 eV. This is outside
the photon energy range of the experiment and fragmentation would only be observable throughmulti-photon processes (like were observed by Boese et al. [46]). Since Ag is electronically very
similar to Au, it is also likely that the structural motifs for both the charge carrier and solvation
observed in [Au(CO2 )n ] clusters are at play here.
We performed DFT calculations on [Ag(CO2 )n ](n = 24) clusters, employing the sameformate core with the side/terminal solvation model discussed above in Section 3.2.2. If we compare
the calculated vibrational frequencies for the predicted [Ag(CO2 )n ](n = 24) structures with

31

2
3

intensity [arb. units]

4
5
6
7
8
9
10
11
1600 1700 1800 1900 2000 2100 2200 2300 2400

wavenumber [cm-1]
Figure 3.6: Experimental spectra of [Ag(CO2 )n ] clusters. The numbers denote the number of
CO2 molecules in the cluster (n). The right and left panes are individually normalized so they are
on di↵erent scales. All of these spectra were obtained monitoring the loss of a single solvent CO2 .
The dashed lines are meant to guide the eye.

32
their corresponding experimental spectra, we observe very good agreement (see Figure 3.7). The
single peak in n = 2 is recovered by the side solvation isomer (see top trace in the left panelof Figure 3.7). The two peaks observed in n = 3 and 4 are predicted by di↵erent solvationisomers of the formate core. The higher energy feature is best reproduced by leaving the terminal
position vacant while structures where the terminal position is occupied recover the lower energy
peak. At n = 5, only a single peak remains in this region indicating that the terminal position
is now filled for all larger clusters. This is identical behavior to that observed in similarly sized[Au(CO2 )n ]clusters and is summarized in Figure 3.8. However, in contrast to the behavior oflarge [Au(CO2 )n ]clusters (i.e. n7), no blue shift is observed in [Ag(CO2 )n ]clusters ascluster size increases. The di↵erential red shift decreases in magnitude as the number of CO2
molecules increases but the peak corresponding to the activated CO2 molecule continues to shift tolower energies. Additionally the peak position for this set of features appears to be asymptoticallyapproaching the antisymmetric CO stretching frequency of CO2 at approximately 1660 cm 1 . Thisobservation will receive additional consideration below.
We next consider the structure of the solvent peaks in small [Ag(CO2 )n ] clusters (right sideof Figure 3.6). In [Ag(CO2 )2 ] , the solvent peak is centered at 2325 cm 1 rather than 2345 cm 1
as has been common in the spectra of [Au(CO2 )n ] . It also appears fairly broad compared to thesolvent peaks for cluster sizes n

5. In [Ag(CO2 )3 ] , there are two distinct solvent signatures, amore intense peak at 2330 cm 1 and a smaller peak at 2345 cm 1 . In [Ag(CO2 )4 ] , the intensity
ratio of the two peaks is reversed and the lower energy peak continues to shift to the blue now
centered at 2335 cm 1 . At [Ag(CO2 )5 ]only one solvent feature, at 2345 cm 1 , is observed.

The shifts observed here are akin to what was seen before in halide CO2 interactions for di↵erent
halides. [48] Based on the interpretation o↵ered for halide-CO2 interactions, there is likely only asmall amount of the excess charge transfered onto the CO2 molecules. Of the starting geometries
used, nearly all structures converged to a formate core where there is a significant amount of charge
localized on one CO2 molecule. The only family of structures that did not converge to the formate
are those exhibiting symmetric solvation of a Aganion. In these structures the CO2 molecules

33

n=3

n=4

intensity [arb. units]

n=2

1800

1900

1800

1900

1800

1900

wavenumber [cm-1]
Figure 3.7: Comparison of the lowest energy structures for [Ag(CO2 )n ] clusters (n = 2 4). All
reported structures are within 60 meV relative energy of each other for a given cluster size and areconsidered to be isoenergetic.

solvate the Aganion in such a way that they prevent any one CO2 molecule from creating acovalent bond with the Ag atom.
Calculations of structures exhibiting symmetric solvation around a Aganion are shown in

Figure 3.9. As can be clearly seen, this family of structures recovers the position and red shift

34

peak position(s) [cm-1]

1900

1800

1700

1

2

3

4

5

6

7

8

9

10

11

12

cluster size [Ag(CO2)n]Figure 3.8: Peak positions in the lower energy region (left panel in Figure 3.6) plotted againstcluster size. Solid squares represent structures where the terminal solvation position is filled. The
open squares represent structures where the terminal solvation position is vacant. Solid circles
indicate when the charge carrier is CO2 (see text).

of the solvent peak in all cases where it is observed. The reason symmetric solvation is no longer
observed at clusters larger than n = 4 is likely due to a lack of available solvation positions thatwould not preferentially enable one of the CO2 molecules to form a covalent bond with the Aganion. Stated another way, solvation behind an already symmetrically solvating CO2 moleculewould polarize charge onto it and prompt it to form a covalent bond with the Aganion. Basedon the signal to noise of the far-IR region and lack of a peak at 2345 cm 1 in [Ag(CO2 )2 ] clusters,
this cluster shows nearly exclusively symmetric solvation with only a small amount of the formate

35complex being formed. There likely exists a peak at 2345 cm 1 for n = 2 but it is obscured by thebroad feature centered at 2325 cm 1 . At n = 3, symmetric solvation is still preferred but there ismore of the formate complex being formed. By n = 4, the argentoformate complex [AgCO2 ]isthe dominant species.
The last trend to discuss is the feature that appears in n = 3

9 and is centered around1660 cm 1 . This frequency is near to that of the antisymmetric stretch of CO2 at 1658.3 cm 1 [38]suggesting that the charge carrier is no longer the formate like complex but rather a CO2 anion. If
this is true, then the Ag atom would no longer be part of the charge carrier, rather playing the roleof a solvent and consequently should be a possible fragment species. We performed experiments
monitoring the loss of a neutral Ag0 atom and the resulting spectra are shown by traces C and D in
Figure 3.10. Comparing the features observed in the loss of Ag channel to those in the loss of CO2channel, it is clear that the peak belonging to the formate core is not observed when monitoring theloss of Ag while the 1660 cm 1 feature is. Additionally, a set of three peaks is observed for n = 4 inthe loss of Ag spectrum. This set of peaks is indicative of a C2 O4 charge carrier rather than a CO2
charge carrier, which is the expected core species for small (CO2 )n clusters (n = 25). [75] Thisobservation indicates that the Ag atom can and does indeed act as a solvent species rather thanpart of the charge carrier. In addition to the formate core structure, the peak at 1660 cm 1 showsthat signatures similar to those of neat (CO2 )n clusters are observed in [Ag(CO2 )n ] clusters.
To further validate the observation that Ag can act as a “solvent” we performed calculations
on [Ag(CO2 )4 ]clusters in an attempt to isolate a structure where the Ag is not part of thecharge carrier. We varied the OCO bond angle of the covalently bonded CO2 moiety in the lowest
energy isomer of [Ag(CO2 )4 ] (see Figure 3.7 top structure right panel) to try and isolate a second
minimum energy structure where Ag is not part of the charge carrier. A robust local minimum
is recovered corresponding to the formation of a C2 O4 charge carrier rather than the formate-likespecies, consistent with spectroscopic observations.
We also tested the e↵ects of a continuum dielectric corresponding to room temperature CO2around the calculated structures. The black curve in Figure 3.11 represents this scenario. It362intensity [arb. units]

3

4

2

3

4

2250

2300

2350

2400

wavenumber [cm-1]
Figure 3.9: Comparison of structures exhibiting symmetric solvation around a Ag anion for n =2 4. The green dotted line is meant to show the red shift trend observed and predicted. The
magenta dot-dash line shows the nearly unshifted solvent CO2 peak corresponding to the existenceof a formate type core.

37

intensity [arb. units]A(CO2)7_

B(CO2)4_C

Ag / [Ag(CO2)5]_D

Ag / [Ag(CO2)4]_

1600

1700

1800

1900

2000

2100

wavenumber [cm-1]
Figure 3.10: Comparison of Ag loss channels for [Ag(CO2 )4 ] (D) and [Ag(CO2 )5 ] (C) to (CO2 )nclusters exhibiting di↵erent core structures (A and B). The feature at 1660 cm 1 corresponds to a
CO2 core while the set of peaks between 1800 cm 1 and 2000 cm 1 indicates a C2 O4 core. See
text for discussion.

is apparent that, while an oversimplification of solvent e↵ects, the addition of a CO2 dielectric
continuum is stabilizing to the C2 O4 core. However, we stress that solvent is not merely a dielectricas can be seen in Figure 3.7. At this point, it is clear that CO2 -based charge carriers can exist in
the presence of a Ag atom.
We return briefly to the comment made above concerning the red shift of the peak corresponding to the formate core. Since behavior akin to (CO2 )n clusters is observed, it is reasonable to38[Ag(CO2)4]- with CO2 bath2000

relative energy [meV][Ag(CO2)4]-

1500

1000

500

0
60

40

20

0

-20

-40

-60

-80

OCO bond angle deviation from 180° [°]Figure 3.11: Potential energy curve for [Ag(CO2 )4 ] generated by varying the OCO bond angleof the partially reduced CO2 moiety. The red curve represents the structures as shown. The
black curve contains a dielectric continuum corresponding to CO2 around the shown structures
to approximate the e↵ect of increased solvation (see Section 3.2.1 for additional information). A
robust second minimum is recovered with the formation of a C2 O4 species, which is additionallystabilized by the dielectric continuum. The relative energies presented here are not zero-point
corrected.

hypothesize that the peak corresponding to the formate complex is indeed approaching a frequency
similar to the antisymmetric stretch of a CO2 anion. However, without spectra of larger clusters(i.e. n12) it is impossible to say with certainty that this is true. Figure 3.8 summarizes thisasymptotic behavior. This section will finish with a comparison of [Ag(CO2 )n ] and [Au(CO2 )n ]clusters.
The primary core structure formed in both [Au(CO2 )n ] and [Ag(CO2 )n ] clusters is a binary

39
complex of the metal and one CO2 in a structure resembling the formate anion. The behavior ofthese cores in the presence of solvent CO2 molecules is similar with preferential solvation around thecovalently bound CO2 resulting in increased reduction of that CO2 ligand initially being observed.
Solvation around the metal atom observed in [Au(CO2 )n ]clusters (n10) resulted in a blue

shift of the antisymmetric CO stretching frequency corresponding with deactivation of the bonded
CO2 ligand. No blue shift was observed in [Ag(CO2 )n ]clusters. This can be attributed to thedi↵erence in electronegativity between Au and Ag (2.54 vs. 1.93, respectively). [49] This strongdi↵erence results in it being more energetically favorable for the Au atom to retain part of theexcess charge and thus hamper charge transfer to the bonded CO2 ligand. This is also consistent
with the behavior observed when attempting to obtain a second minimum corresponding to a C2 O4
charge carrier in [Au(CO2 )4 ]clusters (see Figure 3.12). No second minimum is recovered due tothe high electronegativity of Au.
Because the formate structure exists in both [Ag(CO2 )4 ]and [Au(CO2 )4 ]clusters it isinformative to compare their results and draw a connection to the observed CO antisymmetricstretching frequency and the degree of reduction. Figure 3.13 shows the experimentally observed
frequency of the partially reduced CO2 ligand plotted against the calculated charge on that CO2
based on the assigned structure for that frequency. Additionally, the frequency of a CO2 anion
is also shown. [38] The data indicate an empirical linear correlation between the antisymmetricstretching frequency and the partial charge on the CO2 moiety in this wavenumber region. We
caution that this linear trend breaks down at higher vibrational frequencies and appears to be cubicin nature. [76, 77] Based on this relationship, we can extrapolate the amount of charge localizedon the activated CO2 ligand. For [Au(CO2 )4 ] clusters, a maximum of -0.79 e is donated to thebonded CO2 . For [Ag(CO2 )4 ]clusters, we observe -0.92 e

donated with no indication that thetrend of increased charge donation with increasing solvation will reverse.
The results in this section indicate that Ag is better suited to reductive activation of CO2
than Au. This is owing to the observation of lower antisymmetric stretching frequencies of theactivated CO2 ligand in [Ag(CO2 )n ]corresponding to increased reductive activation compared40

relative energy [meV]

3000
2500
2000
1500
1000
500
0
60

40

20

0

-20

-40

-60

-80

OCO bond angle deviation from 180o [o]Figure 3.12: Potential energy curve for [Au(CO2 )4 ] generated by varying the OCO bond angle ofthe partially reduced CO2 moiety. No second minimum is recovered corresponding to a C2 O4 coreunlike in [Ag(CO2 )4 ] (see Figure 3.11). The relative energies presented here are not-zero point
corrected.

to [Au(CO2 )n ]clusters. Additionally, the ability of Ag to take the role of a solvent rather thanpart of the charge carrier indicates that it is likely easier to abstract reduced CO2 from Ag rather
than from Au. These results suggest an initial trend that lower electronegativity of a 9th column
transition metal corresponds to an increased ability to reduce CO2 .

41

partial charge on CO2 [e-]

-0.4
-0.5
-0.6
-0.7
[Au(CO2)9]-

-0.8
-0.9

[Ag(CO2)11] -

-1.0
1650

1700

1750

1800

1850

1900

observed redshift [cm-1]
Figure 3.13: Plot of the antisymmetric CO stretching frequency in a formate core against thecalculated partial charge localized on the activated CO2 moiety for an assigned structure (see text).
The blue triangles represent [Ag(CO2 )n ] clusters while the orange squares represent [Au(CO2 )n ]clusters. The black circle is the antisymmetric stretching frequency of a CO2 in a Ne matrix. [38]The solid line is a linear fit of all the points in the plot. The dashed lines represent the maximum
charge transfer based on the most red shifted frequency observed.

3.3[Co(CO2 )n ] and [Ni(CO2 )n ] Clusters
Organometallic complexes containing transition metals have found interest and application in

many chemical contexts. In particular, transition metal complexes have been observed as important
catalysts and support materials for catalysts. Unfortunately, there is little detailed information on
the molecular level functionality of these catalytically relevant materials. In a deviation from the
9th column transition metals (or coinage metals Au and Ag), we shall consider the interactions

42of CO2 with cobalt and nickel anions. Since the electron configurations of these two metals aredi↵erent from those of Au and Ag and from each other ([Ar]4s2 3d7 for Co and [Ar]4s2 3d8 for Ni),
one would expect that di↵erent chemistry will likely be observed compared to what was reported
in Section 3.2.
Cobalt has garnered interest as a catalyst in several prototypical reactions and in a numberof metal-organic compounds. [78, 79] However, few spectroscopic studies have been performed invacuo on cobalt containing species that involve a net negative charge. Some work has been carriedout in the context of Fischer-Tropsch synthesis by Gerhards and coworkers, where the interaction of
alcohols with atomic and cluster anions of Co was probed by IR spectroscopy. [78, 79] The FischerTropsch process has received attention as a scheme by which alternative fuels can be generated byconverting CO and H2 to hydrocarbons. [80] This makes the conversion CO2 to CO a process ofpotential industrial importance. Cobalt has in fact seen application in reduction schemes where
CO2 is converted to CO, a feedstock for Fischer-Tropsch processes. [81]Nickel has also been utilized in a wide variety of catalytic applications. Some of these applications have involved being a direct component of the catalyst to being a part of the substrates
and other support materials. [82–84] However, there are again few spectroscopic studies of mass
selected species leaving the molecular level details of these materials and processes relatively uncharacterized. In the context of catalytically active materials, nickel has been shown to be a usefulactive component in dry reforming reactions between CH4 and CO2 where syngas (CO and H2 ) is
generated. [85] Additionally, nickel has also seen application in a number of electroplating processes
utilizing supercritical CO2 . [86, 87] Kubiak and coworkers have shown that Ni ions embedded in an
organic framework are active in the conversion of CO2 to CO23 and CO. [88] With the breadth ofcatalytic applications of Ni in processes involving CO2 , the properties of [Ni(CO2 )n ] clusters areof interest. In an e↵ort to elucidate bonding trends, charge carrier structures, and solvation e↵ects,
this section will discuss the structures, charge distributions, and bonding trends of [Co(CO2 )n ]and [Ni(CO2 )n ] clusters.

43
3.3.1

Computational Methods

We employed density functional theory to aid in the interpretation of the experimental spectra. We used the TURBOMOLE v. 5.9.1 and 6.2 suite of programs [58] to perform calculations
for [Co(CO2 )n ]and [Ni(CO2 )n ]clusters, respectively. The B3-LYP functional [59, 60] with adispersion correction was used in both cases. For [Co(CO2 )n ] clusters an older dispersion correction was used [89] while a more recent version was employed for [Ni(CO2 )n ]clusters. [90] It has

generally been found in this work that the form of the dispersion correction used does not change
the outcome of the calculations appreciably when compared to each other. However, the use of adispersion correction does lead to, on average, better results compared to calculations where thecorrection is not used. The def2-TZVPP basis sets were assigned to all atoms. [61] Vibrationalfrequencies were calculated using the AOFORCE program. [66, 67] The calculated vibrational frequencies were all scaled by 0.97537 to correct for anharmonicity. This factor was generated bycomparing the calculated frequency for a free CO2 to the measured antisymmetric stretch of CO2reported by Shimaniouchi. [37] All reported energies in this section are zero-point corrected and allpartial charges were calculated using a natural population analysis. [68]3.3.2

Infrared Spectra of [Co(CO2 )n ]Clusters

Figure 3.14 shows the experimental spectra obtained monitoring the loss of one CO2 moleculefrom [Co(CO2 )n ] clusters. A number of stark di↵erences are present when compared to the spectra
observed for [Au(CO2 )n ] and [Ag(CO2 )n ] clusters. First, the onset of fragmentation is observed
at n = 3 for Co rather than at n = 2 for Au and Ag. Second, there are two distinct and dominant
spectral signatures present in the far IR region rather than only one for Au and Ag (not including
the feature attributed to a CO2 charge carrier in [Ag(CO2 )n ]of the peaks observed in [Co(CO2 )n ]clusters). Third, the frequencies

clusters are relatively invariant with respect to increasing

solvation unlike in Au and Ag where ⇠10 cm 1 red and blue shifts were observed with each
additional solvent molecule. These three initial observations likely indicate that the core structure44

3intensity [arb. units]

4
5
6
7
8
9
10
11
1000

1200

1400

1600

1800

2000

2200

2400

wavenumber [cm-1]
Figure 3.14: Experimental spectra of [Co(CO2 )n ] clusters monitoring the loss of one CO2 . The
two regions (1000 cm 1 to 2150 cm 1 and 2270 cm 1 to 2400 cm 1 ) are individually normalized
and are therefore on di↵erent scales. Data was not acquired down to 1000 cm 1 for larger clusters
due to low parent cluster intensity and correspondingly poor daughter anion signal intensity.

is significantly di↵erent from the formate motif observed in [Au(CO2 )n ] and [Ag(CO2 )n ] clusters(see Figure 3.15).
Complicating matters further is the possibility of both singlet and triplet electron configurations being present due to the [Ar]4s2 3d8 electron configuration of Co . As can be seen in Figure
3.16, calculations predict that di↵erent electron configurations result in unique core structures. For[Co(CO2 )] , the formate structure is predicted for a triplet configuration while a bidentate interaction of the CO2 to the Co, where the CO2 molecules make both a carbon-Co and an oxygen-Co

45

1850
1800
1750
1700

peak position [cm-1]

1650
1850
1800
1750
1700
1650
1850
1800
1750
1700
1650

2

4

6

8

10

12

14

number of CO2 molecules
Figure 3.15: Comparison of the peak positions for [Au(CO2 )n ] (top), [Ag(CO2 )n ] (middle) and[Co(CO2 )n ] (bottom) clusters. Note that fragmentation is first observed for Au and Ag at n = 2
while it not observed until n = 3 for Co. See text for discussion.

46

Figure 3.16: Structures of CoCO2 for singlet and triplet occupations. The capital letter denotes
the core isomer while the subscript indicates a singlet or triplet conformation. Relative energies
are given in units of meV.

bond, is predicted for a singlet configuration. The bidentate interaction shown in Figure 3.16 (1 C)
is referred to in the metal-organic literature as an ⌘ 2 interaction. [91] The lowest energy structure
for [Co(CO2 )]is isomer A where the Co atom is inserted across a C-O bond. It is already clear

by looking at the predicted structures for [Co(CO2 )] that there are a number of isomers di↵erent
from the formate motif that will necessitate consideration to interpret the experimental spectra.
Since fragmentation is first observed at n = 3, it is likely that the core charge carrier is of
the form [Co(CO2 )2 ] where the third CO2 unit is the first solvent molecule and can be dissociated(this will be revisited below). Performing calculations on [Co(CO2 )2 ]species reveals a numberof di↵erent structures (see Figure 3.17). The predicted global minimum structure is an insertion
structure (see isomer D in Figure 3.17). In this family of structures the metal atom, here Co, is47
inserted across a C-O double bond resulting in CO and O subunits bonded to the metal atom(similar to isomer A in Figure 3.16). The second CO2 unit can then add to the lone O atom
forming a CO3 subunit. While the insertion structure is predicted to be the global minimum, we
calculated a substantial barrier (1.5 eV) to formation of this family of structures. This barrier islargely a result of the breaking of the C-O double bond. Due to this barrier to formation, only asmall population, if any, of this structure is expected to be found in the experiment.
The next higher energy family of structures are what we have termed butterfly structures.
In this family of structures, two CO2 molecules bind in a bidentate (⌘ 2 ) fashion to the metal atom
(here Co). There are a variety of orientations in which the two CO2 molecules can be (see 1 E, 3 E,
and 3 E* in Figure 3.17). Singlet and triplet occupations result in nearly perpendicular (1 E) andplanar (3 E) orientations of the two CO2 molecules, respectively. Additionally, the CO2 ligands canalso be oriented such that the [CO2 CoCO2 ]core exhibits C2h symmetry (3 E*). For all of thesestructures, the cobalt is predicted to be in a +1 oxidation state (Co(I)). The calculated binding
energy of a CO2 bound in a bidentate fashion to the Co atom is on the order of (8000 cm 1 or 1eV) which is well above the photon energy of our experiment. This explains why no fragmentation
is observed until n = 3, i.e. when a solvent CO2 exists.
The last family of structures that were calculated are those possessing a formate core similar
to the core structure observed in [Au(CO2 )n ]and [Ag(CO2 )n ]clusters (see structure 3 B(1) in

Figure 3.17). The formate structure is only recovered for a triplet electron configuration andis predicted to be substantially higher in energy than any of the other structures calculated for[Co(CO2 )2 ](see Figure 3.17). It is also unlikely that the formate core exists in any detectableamount since no fragmentation was observed for n = 2 where one should be able to fragment isomerB(1).
At n = 3, the core structures presented in Figure 3.17 can be solvated. A comparison of thecalculated spectra of these core structures with the [Co(CO2 )3 ]experimental spectrum reveals

that the butterfly family of structures most accurately recovers the experimental spectrum (seeFigure 3.18). The insertion (3 D), formate (3 B(1)) and C2h symmetric butterfly (3 E*) structures do

48

Figure 3.17: Lowest energy calculated structures for [Co(CO2 )2 ] . The letter denotes the core
structure while the number in parentheses indicates the number of solvent molecules and the superscript denotes singlet or triplet configuration. Relative energies are given in meV.

49
not recover the major features of the experimental spectrum. Unfortunately, there is not a clear
di↵erence between the singlet and triplet butterfly structures (1 E and 3 E) that would unambiguously
distinguish the two from each other. At this stage, the dominant peaks in the experimental spectrum
are tentatively assigned to the butterfly isomers 1 E and 3 E. It should be noted that the singlet
butterfly isomer has slightly better agreement with the experimental spectrum both positionallyand with respect to the predicted intensity ratio of the two dominant peaks. However, the possibilityof a mixture of singlet and triplet butterfly structures cannot be discounted.
To further corroborate the assignment of the dominant features in the experimental spectrato those of the butterfly structures (1 E and 3 E), it is necessary to calculate the e↵ect of solvationon the calculated frequencies. The experimental peak positions do not shift appreciably with increasing solvation. Figure 3.19 shows a comparison of the calculated spectra for di↵erent solvation
positions around the butterfly core to the experimental spectrum of [Co(CO2 )4 ] . Clearly, calculations predict little variation in the vibrational frequencies as a function of solvation position. Thisobservation is also corroborated by the almost negligible di↵erences in the calculated charge distributions on the butterfly core of the various solvation isomers (see Table 3.1) indicating that thebutterfly cores are indeed robust core species. At this point the dominant peaks in the [Co(CO2 )n ]spectra are assigned to both of the butterfly isomers (1 E and 3 E).
Table 3.1: Calculated partial charges for the solvation isomers of [Co(CO2 )4 ] presented in Figure
3.19. Isomers 1 - 4 refer to the top to bottom structures presented in Figure 3.19. The column
labeled Solvent indicates the amount of partial charge delocalized into the solvent environment.
The partial charges are reported in units of e.
Solvation Isomer
Isomer 1
Isomer 2
Isomer 3
Isomer 4

Co
+0.93
+0.92
+0.92
+0.94

CO2 (1)
-1.0
-1.0
-1.0
-0.96

CO2 (2)
-0.91
-0.91
-0.91
-0.97

Solvent
-0.02
-0.01
-0.01
-0.01

The last set of spectroscopic features to discuss are the minor peaks around 1635 cm 1 , 1685
cm 1 , and 1950 cm 1 (the 1950 cm 1 feature first appears at n = 4). All of these features can

50intensity [arb. units]

3

3

D

1

E

3

E

3

E*

B(1)

1000 1100 1200 1300 1400 1500 1600 1700 1800 1900 2000 2100

wavenumber [cm-1]
Figure 3.18: Comparison of the calculated lowest energy structures for [Co(CO2 )2 ] to the experimental spectrum. The letter denotes the core structure while the number in parentheses indicatesthe number of solvent molecules and the superscript denotes singlet or triplet configuration. Relative energies are given in meV.

51

intensity [arb. units]

0 meV

isomer 1

12 meV

isomer 2

26 meV

isomer 3

29 meV

1600

1700

1800

1900

2000

isomer 4

2100

wavenumber [cm-1]
Figure 3.19: Comparison of calculated di↵erent solvation isomers around structure 1 E to the experimental spectrum for [Co(CO2 )4 ] . Relative energies are given in meV.

52
be generally accounted for by the insertion structure (3 D), the C2h symmetry butterfly (3 E*) anda species where the cluster consists of the cobalt atom, a CO2 ligand and a C2 O4 unit (G) (seeFigure 3.20). Structure G will be discussed further in the next Section but for now it worth notingthat G has a [Co(CO2 )3 ]core and would be first observable at n = 4. It is difficult to make adefinite assignment of these weak features as their signal to noise ratio is poor compared to that ofthe dominant features. However, due to the position of the calculated and measured frequencies,
it is likely that there is a small population of the insertion structure (3 D) and possibly a mixtureof structures 3 E* and G being generated.
The structure of the butterfly core is reminiscent of many inorganic species. [91] It is informative to compare the calculated molecular orbitals of the singlet butterfly structure (1 E) to thoseof general metal ligand complexes (MLx ). For an octahedral ML6 complex, the five d-orbitals are
split into the t2g (lower energy) and the eg (higher energy) orbitals. This is due to overlap of theligands with the dx2 y2 and dz 2 orbitals on the metal while there is little overlap with the dxy , dxz
and dyz orbitals. If two of the equatorial ligands are removed, a ML4 complex is formed and the
dx2 y2 orbital is stabilized relative to the dz 2 orbital. This stabilization is a result of decreased
overlap of ligands with the dx2 y2 orbital (see left side of Figure 3.21). [92]The ML4 complex has a structural appearance close to that of the butterfly core, so one mightassume the molecular orbitals are similar and that the [CO2 CoCO2 ]core could be described assuch. However, the bidentate interaction creates a three membered ring which keeps the two
metal-ligand bonds very close to each other. This interaction could also be thought of as pseudomonodentate where even though two bonds are being formed, they are essentially bonding to thesame d-orbital on the metal. If this were the case, one would not expect ML4 behavior but rather
ML2 (see center column of Figure 3.21) where now the dz 2 orbital will be stabilized more than thedx2 y2 orbital. [92]The top five occupied molecular orbitals (HOMOs) were calculated in addition to the next
two lowest lying unoccupied molecular orbitals (LUMOs) for the isomer 1 E (see right side of Figure3.21). The structure of the HOMO is similar to that of the dz 2 orbital while the LUMO bears

53

intensity [arb. units]

3

3

3

E*

B(1)

1

1000

1200

1400

1600

1800

D

G

2000

wavenumber [cm-1]
Figure 3.20: Comparison of calculated likely minor isomers to the experimental spectrum for[Co(CO2 )4 ] . Relative energies are given in meV. See text for additional discussion of the isomers.

54

Figure 3.21: Orbitals in the d-block of the [CO2 CoCO2 ] core ion expected for a “butterfly” ML4
configuration (left) and a bent ML2 configuration (center), compared to the calculated molecularorbitals for the singlet core ion structure 1 E (right). Note that the HOMO-3 orbital is not part ofthe d-block but is shown for completeness. The occupied orbitals are shown as bold horizontal lines
for each configuration. Calculated orbital energies are given in eV. The dotted lines show how thedz 2 orbital changes from ML4 to ML2 and correlation between calculations and ligand field theory.
Note that the definition of the z-axis is not along the symmetry axis of the complex but rather isdefined by the dz 2 character of the HOMO.

resemblance to the half interacting dx2 y2 orbital in both ML2 and ML4 complexes. The energeticordering of the HOMO and LUMO suggest that the butterfly complex bares more resemblance toa ML2 complex than to a ML4 complex suggesting that the core CO2 ligands can be approximatedas monodentate.
In the organometallic literature the bidentate CO2 bonding is referred to as an ⌘ 2 complex,
since there are two bonds from the ligand to the metal center. The butterfly structures are unique
as most experiments have only observed one ligand ⌘ 2 complexes. [91] Here, there are two CO2

55
ligands bonding to the metal atom creating a diligand ⌘ 2 complex. It is also interesting to notethat while inorganic chemistry has characterized this bonding motif with two metal-ligand bonds,
the MO picture appears to treat each CO2 as a monodentate ligand. In addition to the unique
number of ligands, the oxidation state of the Co in these clusters is also uncommon. Cobalt tends
to form complexes where it is in an oxidation state +2. In these experiments, only Co(I) complexes
were observed. Co(I) has been seen before in disproportionation reactions and in complexes withbipyridine used for water splitting and CO2 reduction to CO but, again, is not a common species.
[93,94] These two observations of the interaction of CO2 with Co in the presence of an excess electron
suggest that there are less explored synthetic routes and unknown novel chemistry involving Co
and CO2 .

3.3.3

Infrared Spectra of [Ni(CO2 )n ]Clusters

Figure 3.22 shows the experimental spectra obtained by monitoring the loss of one CO2
molecule from [Ni(CO2 )n ] clusters (n = 2

8). The spectra, while clearly di↵erent from those of[Co(CO2 )n ] clusters (see Figure 3.14), exhibit similar trends to those present in Co. The spectra
tend to not shift as the number of CO2 molecules increases and there are multiple dominantpeaks present. These observations and the fact that the first solvent peak is observed at n = 3
likely indicates that, as is the case for Co, the dominant core species consists of two CO2 units
in addition to the Ni atom. There are, however, two primary di↵erences. First, even though
there is no signature of a solvent CO2 molecule, photodissociation is observed for n = 2 where
broad unresolved peaks are recovered. Second, rather than two dominant peaks as was observed in[Co(CO2 )n ] clusters, three dominant features are present in [Ni(CO2 )n ] .
Since the experimental spectra bear some resemblance to those of [Co(CO2 )n ] , it is reasonable to compare a similar set of calculated structures. To be consistent with Section 3.3.2, the same
designations for core structures and convention for indicating solvation will be used here. Figure
3.23 shows the energetically lowest lying calculated structures. Since there is only one unparied
electron for Ni , there is only one spin multiplicity to be concerned with here. The insertion

56

intensity [arb. units]

2
3
4
5
6
7
8
1000

1200

1400

1600

1800

2000

2200

2400

wavenumber [cm-1]
Figure 3.22: Experimental spectra of [Ni(CO2 )n ] clusters. The numbers indicate the numberof CO2 molecules (n) present in each cluster. The traces on the left and right are individuallynormalized and are on di↵erent scales.

structure (D) is again predicted to be the global minimum and (like in [Co(CO2 )n ]clusters) ispredicted to have a roughly 1.5 eV barrier to formation (see Appendix B). The butterfly familyof structures are also predicted to be low in energy. However, di↵erent from in [Co(CO2 )n ] clusters, the C2h butterfly (E*) is calculated to be closer in energy to the asymmetric butterfly (E).
The other major di↵erence is that a structure containing a C2 O4 subunit (F) is calculated to beenergetically relevant for [Ni(CO2 )n ] clusters.
A comparison of the core structures discussed above possessing one solvent CO2 to the ex-

57

Figure 3.23: Calculated lowest energy structures for [Ni(CO2 )n ] clusters. The capital letter indicates the core structure and the number denotes the number of solvent molecules present. Plaintext indicates n = 2 structures while bold text denotes n = 3 structures. Relative energies arereported in meV and are given with respect to the lowest energy structure for a given cluster size.
Relative energies for isomers B and G are not reported as there are no structures to compare themwith and are shown in italic text.

58
perimental spectrum of [Ni(CO2 )3 ] is shown in Figure 3.24. It is clear that the insertion structure(D(1)), the formate motif (B(2)) and the isomer containing the C2 O4 subunit (F(1)) do not recover
the dominant peaks in the experimental spectrum. Only the two butterfly structures accurately
recover the experimental spectrum. Because of the symmetry of isomer E* the antisymmetricstretches of the two CO2 moieties are coupled. This results in only the out of phase combination having an appreciable oscillator strength since the in-phase combination of the antisymmetricstretches results in a negligible change of the dipole moment of the core species. The single peak
recovered from structure E* is predicted to lie between the two frequencies calculated for isomer
E. These two structures are close in energy and the combination of their calculated frequencies
accounts for the three dominant peaks observed in the experimental spectrum. Two additional
structures were predicted but are high in energy and do not recover the experimental spectra (seeAppendix B).
It is necessary to check the e↵ect of solvation on the calculated vibrational frequencies andcore charge distributions for isomers E and E*. The experimental spectra suggest that the e↵ect ofsolvation is slightly more pronounced than it is in [Co(CO2 )n ] clusters since the dominant peaks
shift by 3 cm 1 to 5 cm 1 with increasing cluster size and bifurcate at large cluster size (i.e. n

6).

Figure 3.25 shows a comparison of various solvation positions around isomer E. There is a larger
variation in the calculated vibrational frequencies as a function of solvation position in [Ni(CO2 )4 ]compared to [Co(CO2 )4 ]clusters (see Figure 3.19). However, the variations in peak position as

a function of solvation are still not as large as they are in [Au(CO2 )n ]and [Ag(CO2 )n ] . These

calculations and comparisons suggest that the core charge distribution for [Ni(CO2 )n ]clusters(specifically isomer E*) is slightly more polarizable than comparable structures in [Co(CO2 )n ]clusters.
Like in [Co(CO2 )n ] , there are a number of minor peaks observed in the [Ni(CO2 )n ] spectra.
In particular, there are peaks observed at 1350 cm 1 , 1920 cm 1 and 2000 cm 1 . Referring back to
Figure 3.24, it appears as though the insertion structure could account for the peak observed near
2000 cm 1 . The predicted vibration at 1675 cm 1 for isomer D is likely present in the experimental

intensity [arb. units]

59

1000

1200

1400

1600

1800

D(1)

0 meV

F(1)

886 meVE(1)985 meV

E*(1)1161 meVB(2)1637 meV

2000

wavenumber [cm-1]
Figure 3.24: Comparison of calculated structures for [Ni(CO2 )3 ] to the experimental spectrum
for n = 3. See Figure 3.23 and text for discussion on the structures. Relative energies are given in
meV.

60

intensity [arb. units]

0 meV

9 meV

57 meV154 meV

1000

1200

1400

1600

1800

2000

wavenumber [cm-1]
Figure 3.25: Comparison of various solvation positions around the core isomer E. Note that thepredicted spectrum shifts only slightly with varying solvation position. The calculated spectrum ofisomer E* is shown on the bottom for completeness.

61
spectra but is obscured by the low signal to noise ratio below the dominant features. The other two
peaks are observed for the first time at n = 4, indicating that the core species responsible is likelyof the form [Ni(CO2 )3 ] where the fourth CO2 would be the first solvent molecule. Surveying thepossible core species of the form [Ni(CO2 )3 ] reveals only one reasonable structure. That structureis shown in Figure 3.23 (isomer G) and consists of CO2 and C2 O4 subunits (a similar structure wasalso found in [Co(CO2 )n ] clusters).
To definitively assign the experimental spectra, exploratory calculations on increased solvation of isomers D, E, E* and G were performed to simulate the evolution of the experimental
spectra. The results of these exploratory calculations is shown in Figure 3.26. It is clear thatthese four structures do account for the observed experimental peaks and recover the red and blue
shifts occurring with increasing solvation. In particular, note that structure G does indeed clearlyrecover the two peaks at 1350 cm 1 and 1920 cm 1 and their solvation mediated blue shift. Basedon the agreement of calculations with experiment, the three dominant peaks are assigned to thetwo butterfly structures (E and E*), the peak at 2000 cm 1 to the insertion structure (D), and thetwo peaks at 1350 cm 1 and 1920 cm 1 to structure G.
It is worth noting that the calculated charge distribution on the dominant core species indicates two negatively charged CO2 ligands and a positively charged Ni atom. However, isomer G
is the only structure predicted to contain a Ni(I) atom. Instead, the charge on the Ni atom is ⇠+0.5 e for the butterfly structures and ⇠ +0.7 e for the insertion isomer (see Table 3.2). This is incontrast to [Co(CO2 )n ] clusters where the Co atom was exclusively in an oxidation state +1.
The last spectrum to discuss is that of [Ni(CO2 )2 ] (see Figure 3.22). The peaks observed at
this cluster size are broad and unresolved. Since there is no solvent peak observed, the fragmented
CO2 must come from the core species. Since the binding energy of a core CO2 is on the order of
10000 cm 1 , it is likely that this spectrum in due to multiphoton excitation of a hot bare core species(internal energy ⇠5000 cm 1 ). The n = 2 cluster size was found in high abundance compared to
the rest of the [Ni(CO2 )n ]clusters. It is likely that due to the high abundance of parent ionsat this cluster size and the lack of a one photon dissociation channel, multiphoton dissociationintensity[arb. units]

62

6cluster size [n]543experimental

6
5
4
3
1000

G

G

calculated

E,E*

1200

1400

1600

E,E*

wavenumber [cm-1]

1800

D

2000

Figure 3.26: Calculated vibrational frequencies as a function of cluster size for assigned core structures compared to the experimental peak positions as a function of cluster size (see text for discussion). The letters indicate the core structure giving rise to each set of peaks (see Figures 3.23).
The experimental spectrum for [Ni(CO2 )5 ] is shown on top for reference.

is observable for this cluster size. Figure 3.27 shows a comparison of the calculated spectra forstructures E and F with the experimental spectrum for n = 2. The calculated peaks are in thesame general frequency range as the broad experimental peaks and overlap with the envelope of thefeatures in the experimental spectrum. Even though a combination of isomers E and F appears toreasonably recover the experimental spectrum, the fact that isomer F does not account for any ofthe dominant peaks at larger cluster sizes (i.e. n3) puts into question the likelihood that isomer

63
Table 3.2: Calculated partial charges of the core isomers for [Ni(CO2 )n ] . See Figure 3.23 anddiscussion in the text for additional information on the core structures. The column labeled Solvent
indicates the amount of partial charge delocalized into the solvent environment. The partial charges
are reported in units of e.
Core Structure
D
F
E
E*
B
G

Ni
+0.67
+0.66
+0.49
+0.41
-0.35
+0.99

CO2

CO3
-1.48

CO
-0.19

C2 O 4
-1.66

-0.75
-0.71
-0.65
-0.42

-1.57

Solvent
0.00
0.00
+0.01
+0.01
0.00
0.00

F exists at n = 2. However, the assignment of the n = 2 spectrum is nearly impossible to confirm
due to the low signal to noise ratio and the lack of a reference to discern multiphoton from single
photon e↵ects.
For both [Co(CO2 )n ] and [Ni(CO2 )n ] clusters, a “butterfly” core is the dominant structural
motif. For Co, the two CO2 ligands are nearly perpendicular to each other in isomer E while there
is only a 42 angle between the two CO2 moieties for the same isomer in Ni. For both Co andNi there are signatures of the insertion structure and other minor isomers. The existence of thesestructures indicates that there are multiple ways in which CO2 can bind to both Co and Ni atoms
unlike what was observed in Au and Ag.
The calculated polarizability of the butterfly core increases from Co to Ni. At the sametime, the electronegativity of the metal increases and the calculated charge on the metal becomes
more negative. All three of these observations are connected. As the electronegativity of the metal
increases, the metal will localize more of the excess charge onto itself. This results in the lesspositive partial charge on Ni compared to Co. Additionally, since Ni is now able to accommodate
more of the excess charge, the polarizability of the core increases as electrons can now move more
freely across the butterfly core.
Ni and Co form similar structures with CO2 in the presence of an excess electron. Neithernormalized intensity [arb. units]

64

1000

1200

1400

1600

1800

2000

wavenumber [cm-1]
Figure 3.27: A comparison of the combined calculated spectra for isomers E and F to the experimental spectrum of [Ni(CO2 )2 ] . Note that while the fit is not ideal, it is reasonable that a
multiphoton process could recover the observed experimental spectrum. See text for discussion.

appear particularly suitable for applications as single atom catalysts in CO2 reduction/conversion
schemes since, in both cases, strong covalent bonds are formed with the reduced CO2 ligands.
However, the structures observed here suggest that there may be unique applications of CO2 within
inorganic and organometallic chemistry.

65

3.4

Infrared Studies of [Cu(CO2 )n ] Clusters
There are a number of reasons why Cu carries a lot of interest in the context of these studies.

First, there has been work showing the application of copper in catalytic systems and fuel cells
where CO2 has been converted to CO, [95–97] methane, [97–99], alcohols and other small organic
species. [97, 100, 101] Unfortunately, there again has been little e↵ort to study the active species
and catalyst complexes spectroscopically to gain understanding of the molecular level functionality.
Copper is also the final member of the coinage metals. For anionic complexes with CO2 , one would
expect that the structure of the charge carrier and the solvation behavior will be similar to thatof Au and Ag. This would manifest itself in spectral signatures characteristic of the formate motif
with solvation occurring around the CO2 moiety, resulting in increased activation of the bonded
CO2 . Cu has a similar atomic radius and electronegativity to Ag, and one might expect moresimilarities to [Ag(CO2 )n ]clusters than to [Au(CO2 )n ]clusters. Cu is also immediately to theright of Ni continuing the study of first row transition metal anions interacting with CO2 . Theinteractions of Co and Ni with CO2 are markedly di↵erent than those of Au and Ag. Since Cu isalso a first row transition metal, one may expect to see butterfly structures like those observed in
Co and Ni. Studying [Cu(CO2 )n ] clusters allows us to study two periodic groups simultaneously
and discover if their behavior is interconnected and/or which bonding motifs are present.

3.4.1

Computational Methods

Density functional theory was performed using the TURBOMOLE v. 6.2 suite of programs.
[58] For all calculations performed in this section, the B3-LYP functional [59, 60] with dispersion
correction [90] was used. The def2-TZVPP basis sets were assigned to all atoms in these calculations.
[61] Harmonic frequencies were calculated using the AOFORCE program. [66, 67] The calculated
harmonic frequencies were scaled by 0.97537 to account for anharmonicity. This factor was obtainedby comparing the antisymmetric stretch of CO2 as reported by Shimanouchi [37] to the calculated
value for the same molecular motion under this level of theory. Charge distributions were calculated

66
using a natural population analysis. [68]3.4.2

Infrared Spectra of [Cu(CO2 )n ]Clusters

Figure 3.28 shows the experimental spectra for [Cu(CO2 )n ]for n = 2

9. These spectra

are clearly far more complicated than any of the spectra observed previously for the other metals
described in this chapter! Fragmentation is first observed at n = 2 where one dominant peak isobserved in the region of the charge carrier at 1750 cm 1 with a shoulder at 1720 cm 1 . At n = 3,three dominant peaks are observed in the region of the charge carrier with a small peak at lower
energies. As cluster size increases, the peaks begin to shift and split resulting in a large numberof peaks for n

6. While the prospect of assigning all of these peaks is at first glance daunting,identifying trends and using the information garnered from the other metal CO2 studies will proveinvaluable for interpreting these spectra.
Upon detailed inspection, there are three dominant trends within these spectra (shown in
Figure 3.29). We will discuss these trends one at a time. First, the trend highlighted in red showsa red shifting peak that changes its role from the dominant feature in n = 2 to a minor feature
at all larger cluster sizes. This set of peaks is similar to those observed in both [Au(CO2 )n ][Ag(CO2 )n ]andclusters. These features were attributed to a formate like structure where solvation

around the CO2 results in polarization of the excess charge onto the bonding CO2 moiety and aconcomitant red shift of the CO2 antisymmetric stretch for the bonded species. Since Cu, likeAu and Ag, is a coinage metal, it is likely to exhibit the formate structural motif present in both[Au(CO2 )n ] and [Ag(CO2 )n ] clusters.
Performing calculations on a CuCO2 species yields only two possible core structures (seeFigure 3.30). The formate structure (B) is predicted to be the lowest energy isomer. This is thesame core structure assigned in [Au(CO2 )n ]and [Ag(CO2 )n ]clusters. The insertion structure(A) is higher in energy. Unlike for Co and Ni, the insertion structure is not predicted to be thelowest energy isomer for n = 1 in [Cu(CO2 )n ] clusters. The monosolvated isomers are also shownin Figure 3.30. For isomer B these are the previously described side and terminal solvation isomers.

67

2

intensity [arb. units]

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4
5
6
7
8
9
1600

1700

1800

1900

2000

2100

2200

2300

2400

wavenumber [cm-1]
Figure 3.28: Experimental spectra of [Cu(CO2 )n ] . The numbers denote the number of CO2
molecules in the cluster. All spectra are individually normalized so that the left and right traces
are on di↵erent scales.

Isomer A can only be solvated around the CO moiety since solvation of the lone O atom would
likely result in the formation of isomer D (discussed below).
Figure 3.31 shows a comparison of the formate core with side or terminal solvation positions
occupied and the insertion structure to the experimental spectrum for [Cu(CO2 )2 ] . It is clearthat the insertion structure A(1) does not recover the experimental spectrum while the side solvated formate structure (B(1b)) does, and the terminal solvation isomer (B(1a)) recovers the small
shoulder at 1720 cm 1 . This behavior has been observed before in [Au(CO2 )n ] and [Ag(CO2 )n ]clusters, where occupation of the terminal solvation position results in lower energy peaks compared68

Figure 3.29: Experimental spectra of [Cu(CO2 )n ] with the three groups of spectral features highlighted (see text for discussion). Numbers denote the number of CO2 molecules in the cluster. All
spectra are individually normalized so that the left and right traces are on di↵erent scales.

to when the terminal position is left vacant. Interestingly, while the formate motif is present, it isclearly not the dominant structural motif as cluster size increases but rather persists as a minor
contributor to the experimental spectra.
We can compare the experimentally measured frequency of the red highlighted trend to thecalculated charge on the bonded CO2 ligand in the formate structures as was done for [Au(CO2 )n ]and [Ag(CO2 )n ][Cu(CO2 )n ]clusters (see Figure 3.13). Figure 3.32 shows the formate structure data forclusters together with the linear correlation generated from the Au and Ag data.

All data sets are consistent with an empirical linear correlation. This further corroborates the

69

Figure 3.30: Calculated lowest energy structures for [Cu(CO2 )n ] (n = 1 to 4). The notation ischosen to be consistent with that for [Co(CO2 )n ] and [Ni(CO2 )n ] clusters. The letter denotes thecore structure and the number/letter combination in parentheses indicates the number of solvent
molecules and their positions. The purple box indicates n = 1 species, the blue boxes indicate
n = 2 species, the green boxes indicate n = 3 species and the orange box indicates n = 4 species.
Bare core structures are shown on the top row while solvation isomers are shown in the descending
columns. Relative energies are given in meV and are with respect to the lowest energy isomer fora given cluster size.

assignment of the formate structure and provides convincing evidence that the linear correlation in
Figure 3.32 is a spectroscopic tool for assessing the degree of reduction of an ⌘ 1 bonded CO2 to ametal.
The second set of features to consider is those highlighted in blue (see Figure 3.29). Thesepeaks first appear at n = 3 indicating that the core species responsible is of the form [Cu(CO2 )2 ] .
The [M(CO2 )2 ]core composition was assigned primarily to butterfly structures in [Co(CO2 )n ]and [Ni(CO2 )n ]clusters. However, with the multitude of peaks present in the experimental

spectra, all reasonable structural variations of the form [Cu(CO2 )2 ] should be considered. Figure
3.30 shows all of the calculated structures for [Cu(CO2 )2 ] . As was predicted for Co and Ni, theintensity [arb. units]

70

1600

1700

1800

1900

2000

2100

B(1a)

0 meV

B(1b)

21 meVA(1)1388 meV

2200

wavenumber [cm-1]
Figure 3.31: Comparison of calculated structures B(1a), B(1b) and A(1) to the experimental spectrum of [Cu(CO2 )2 ] . See Figure 3.30 for relative energies and the text for discussion.

insertion structure (D) is the predicted lowest energy structure. The C2v butterfly (E), C2 O4 (F)and inversion butterfly (E*) are also predicted to be possible isomers. Figure 3.30 also shows avariety of di↵erent solvation positions possible for each of the core species in columns descending

71

Figure 3.32: Plot of the antisymmetric CO stretching frequency in a formate core against thecalculated partial charge localized on the activated CO2 moiety for an assigned structure (see text).
The blue triangles represent [Ag(CO2 )n ] clusters while the orange squares represent [Au(CO2 )4 ]clusters and the red squares represent [Cu(CO2 )n ] . The black circle is the antisymmetric stretchingfrequency of a CO2 in a Ne matrix. [38] The solid line is a linear fit of all the points in the plot.
The dashed lines represent the maximum charge transfer observed based on the most red shiftedfrequency observed.

from the bare cores.
A comparison of the calculated spectra for the singly solvated [Cu(CO2 )2 ]cores to theexperimental spectrum for n = 3 is presented in Figure 3.33. The butterfly core, with two di↵erent
solvation environments, most accurately recovers the experimental spectrum. These two motifs are
termed symmetric (E(1a) and E(1c)) and asymmetric (E(1b)) solvation. In the case of symmetric

72
solvation, the two core CO2 ligands are solvated equally. Stated another way, symmetric solvation
is where the solvation positions behind each of the core CO2 ligands are either both filled or vacant.
This behavior results in an even distribution of the excess charge between the two CO2 ligands. As
a consequence, the vibrational modes of the CO2 ligands are coupled, resulting in in-phase (higher
energy) and out of phase (lower energy) combinations of the antisymmetric stretches (see third
panel in Figure 3.33).
Asymmetric solvation occurs when one of the two core CO2 ligands is preferentially solvated(see Figure 3.30 isomer E(1b)). In this case, the solvated CO2 has additional charge density
polarized onto it relative to the unsolvated CO2 ligand. Since there is a di↵erence between the
two ligands, the C2v symmetry is broken and the antisymmetric CO2 stretching frequencies are nolonger coupled like they are in symmetric solvation. The solvated CO2 moiety, due to the increased
partial charge being polarized onto it, exhibits a lower energy stretching mode compared to thatof the unsolvated CO2 moiety (see E(1b) in Figure 3.33). These two solvation behaviors accountfor the three peaks initially observed in [Cu(CO2 )3 ] (see Figure 3.34).
As solvation increases, one can envision that these two solvation motifs, symmetric andasymmetric, will persist but become more convoluted. The solvation environment will likely besomewhere between these two extremes as more and more CO2 molecules are added, resulting inthe spectral congestion observed in the blue region. Since both solvation motifs recover the peakat approximately 1750 cm 1 , it is not surprising that this is the first feature to split and to havethe most congestion associated with it. However, asymmetric solvation generally accounts for the
peaks centered at 1810 cm 1 while symmetric solvation will generally recover the features centeredat 1770 cm 1 .
The final region of peaks to discuss is that highlighted in green in Figure 3.29. These peaks
first appear at n = 4, indicating the presence of a [Cu(CO2 )3 ] core species. Based on the calculated
structures of a Cu atom and three CO2 molecules with an excess electron, only one energetically
relevant structure is predicted. This core consists of CO2 and C2 O4 subunits bonded to the Cu
atom (see Figure 3.30, structure G) and is similar to the core structure assigned to analogous

intensity [arb. units]

73

1600

1700

1800

1900

2000

2100

D(1a)

0 meV

E(1a)

724 meVE(1b)770 meVB(2)853 meVE*(1a)

1718 meV

2200

wavenmuber [cm-1]
Figure 3.33: Comparison of calculated structures with a [Cu(CO2 )2 ] core to the experimental
spectrum for n = 3. See Figure 3.30 for description and relative energies of the structures presented.

74
features observed in [Ni(CO2 )n ] clusters.
While solvation position does not dramatically a↵ect the charge distribution of structure G in
[Ni(CO2 )n ] , di↵erent spectral features are predicted when di↵erent solvation positions are occupied
in the Cu analog. Specifically, solvation around the CO2 moiety results in increased localizationof the excess charge onto that ligand, causing lengthening of the CO bonds and a predicted lower
vibrational frequency compared to solvation around the C2 O4 moiety. Interestingly, a comparisonof these two di↵erent solvation motifs (CO2 solvation position filled or vacant) to the experimental
spectrum reveals that both are observed in the experiment. For the two peaks observed in the green
region of the experimental spectra (see Figure 3.29), solvation around the CO2 moiety recovers thelower energy feature while solvation around the C2 O4 recovers the higher energy peak. It should
be noted that while there is a signature predicted around 1700 cm 1 for structure G, it is likely
obscured by the dominant features of the butterfly core. At n = 8, only the lower energy featureappears in the green region of the experimental spectra. This is expected since as the number ofsolvent molecules increases, there is a higher likelihood that the solvation position behind the CO2
subunit will be filled. For n8, the solvation position behind the CO2 subunit is always filled andresults in only the lower energy feature being recovered.
As a final check of the assignment of the experimental spectra, we performed exploratory
calculations on the three primary core structures up to n = 6 to verify that the structures andsolvation motifs discussed above continue to recover the experimental spectra for larger cluster
sizes. Figure 3.35 shows the results of these exploratory calculations compared to the experimental
peak positions of [Cu(CO2 )n ](n = 36). As can be seen, the calculations recover the generalshape and trends of the experimental spectra even though a definitive assignment is not possible.

3.5

Comparison of all Metal-(CO2 )n Anion Clusters
It is clear from the assignment of the experimental spectra that [Cu(CO2 )n ] clusters exhibit

structures and properties that have been observed in other metal-CO2 clusters. Specifically, thebutterfly structure observed as the dominant structural motif in [Co(CO2 )n ]and [Ni(CO2 )n ]

75

intensity [arb. units]

E symmetric

E asymmetricG CO2 solvated

G C2O4 solvated

1600

1700

1800

1900

2000

2100

2200

wavenumber [cm-1]
Figure 3.34: Comparison of the calculated spectra for di↵erent solvation positions around isomers
E and G to the experimental spectrum for [Cu(CO2 )6 ] . Note that since isomer G is a minor
contributor to the experimental spectrum, the features around 1700 cm 1 will likely be obscuredby the formate signature. See Figure 3.30 for additional information on the structures presented
here and text for additional discussion on symmetric and antisymmetric solvation.

intensity[arb. units]

76

6

cluster size [n]543experimental

6
5
4
G

3

B

1600

calculated

E

1800

2000

wavenumber [cm-1]

Figure 3.35: Spectral shifts of the experimental and simulated features for [Ni(CO2 )n ] (n = 3 to6). The top trace show the spectrum for n = 6 for comparison. The lines are meant to guide theeye and are not a fit of any of the data, the letters indicate the assignments of the core isomers asshown in Figure 3.30. See text for further discussion.

clusters is the primary core charged species in [Cu(CO2 )n ]clusters. Additionally, the formate

motif and the large polarizability of the charge carrier observed in [Au(CO2 )n ] and [Ag(CO2 )n ]clusters, are both present in [Cu(CO2 )n ] . These observations indicate that Cu is a switching point
between the trends observed in the first row transition metals and in the coinage metals.
It is worth exploring these trends further. The electronegativity of the coinage metals increases moving down the periodic table (1.90 to 1.93 to 2.4 for Cu, Ag and Au, respectively). [49]77This trend is correlated with the maximum amount of excess charge localized on the CO2 ligand of
a formate structure for Au, Ag and Cu (see Figures 3.13 and 3.32). Au, being the most electronegative, allows for the least amount of charge transfer since Au will attract the excess electron morestrongly than either Ag or Cu. This is manifested spectroscopically in the observation of the blue
shifting antisymmetric stretching frequency in [Au(CO2 )n ] clusters for n

10. For [Ag(CO2 )n ] ,there is spectroscopic evidence of Ag participating as a solvent species. This suggests that theelectronegativity of a (CO2 )n cluster is likely on the order of the electronegativity of Ag such thatthe two di↵erent cores ([AgCO2 ]and CO2 based charge carriers) can coexist. Since the elec-

tronegativity of Cu is lower than that of Ag, the increased localization of the excess charge to thebonded CO2 is expected (see Figure 3.32). However, it is interesting to note that (CO2 )n cluster
signatures are not observed in the spectra of [Cu(CO2 )n ] clusters.
Like in both [Co(CO2 )n ]and [Ni(CO2 )n ]clusters, butterfly structures are the dominantcore structure present in [Cu(CO2 )n ] clusters. However, the polarizability of the butterfly core issubstantially higher than those of either Ni or Co complexes (the trace of the polarizability tensor
is calculated to be 82.45, 76.04 and 69.71 Å3 for Cu, Ni and Co, respectively for assigned butterfly
structures). This is manifested in the spectroscopic observations of only minor peak shifts and small
splittings observed for Co and Ni versus the plethora of peaks in Cu. Interestingly, the noticeablee↵ect of solvation on the charge distribution of the core species in Au and Ag is observed for allof the core species in Cu even though some of the core species are quite di↵erent. An additional
explanation for the increased polarizability of the butterfly core in Cu versus Co and Ni concerns
the geometry of the core. In Co, the primary core species has the two CO2 ligands at roughly 90
to each other. In Ni, this angle has been reduced to roughly 45 and in Cu the two CO2 moieties lie
in the same plane. If one considers ⇡ conjugation like in an aromatic system, [102] it is reasonable
to assume that there would likely be more conjugation of the p- and d-orbitals if the butterfly coreis planar than in any other geometry. This would lead to a more polarizable distribution of theexcess charge in the planar core (Cu) rather than a localized and robust distribution in a twisted
butterfly core (Co and Ni). This hypothesis agrees with the experimental observations and the

78
calculations that were performed.
The data in this chapter show interesting trends in the bonding behavior of coinage metals
and first row transitions metals with CO2 , which are in keeping with their positions in the periodic
table. Additionally, the data indicate that Cu is indeed a switching point between the two setsof trends. The coinage metals appear to be more well suited to catalytic reductive activation of
CO2 than Co or Ni, with Ag being the most promising candidate based on these results. The firstrow transition metals present potentially novel chemistry of CO2 and metal anions. It is our hopethat these metal-CO2 studies will illuminate molecular level details that can aid in rational catalyst
design for the development of carbon neutral fuel cycles.

Chapter 4
Infrared Studies of [CoO(CO2 )n ] and [NiO(CO2 )n ] Clusters

This Chapter has been reproduced in part with permission from the following previouslypublished article:
Knurr, B. J., Weber, J. M., “Structures of [CoO(CO2 )n ] and [NiO(CO2 )n ] Clusters Studied by Infrared Spectroscopy.” Journal of Physical Chemistry A, 2015. DOI:10.1021/jp5108608
Copyright 2015 American Chemical Society.

4.1

Introduction
Metal-oxide species are a common occurrence in transition-metal chemistry with many cat-alytic reactions performed on metal-oxide surfaces with reactivity varying from oxide site to oxide
vacancy. [103, 104] Understanding the chemical bonding and interactions at these types of active
sites can aid catalyst design and understanding of surface chemistry. Studies of model metal-oxide
species in vacuo provides detailed structural information by bypassing the difficulties of performing
spectroscopy in the condensed phase environment. A number of groups have studied the structure
and reactivities of various metal-oxide species in the gas phase. The experiments performed onthese species range from reactivity studies [105–117] to photoelectron spectroscopy [111, 118, 119]and vibrational spectroscopy. [117, 120–128]
In Chapter 3, the structure and e↵ects of solvation on [M(CO2 )n ]clusters (M = Au, Ag,

Co, Ni, and Cu) were reported. In this chapter, similar studies are reported for [CoO(CO2 )n ] and[NiO(CO2 )n ] clusters.

80

4.2

Computational Methods
Density functional theory calculations were performed on a number of [CoO(CO2 )n ][NiO(CO2 )n ]andstructures using the TURBOMOLE v. 6.2 suite of programs. [58] For all calcula-tions performed in this chapter, the B3-LYP functional [59, 60] with dispersion correction [90] wasused. For all atoms in these calculations, the def2-TZVPP basis sets were assigned. [61] Harmonicfrequencies were calculated using the AOFORCE program. [66, 67] The calculated antisymmetricfrequencies corresponding to a CO2 bound in an ⌘ 1 motif (described in Chapter 3) were scaled by
0.9380 to account for anharmonicity. All other frequencies were scaled by 0.97537. The manner inwhich these factors were obtained is discussed in Section 3.2.1. Charge analyses were performed
using a natural population analysis. [68]4.3

Spectra and Analysis
Like in the previously discussed [M(CO2 )n ]clusters, the spectra of [CoO(CO2 )n ]and[NiO(CO2 )n ] show spectral signatures in two distinct regions. Peaks in the higher energy region(2250 cm 1 to 2400 cm 1 ) are attributed to solvent type CO2 , which retain the structural natureof a free CO2 molecule. The antisymmetric stretch of these CO2 molecules deviates little fromthat of free CO2 . The second region (1000 cm 1 to 2150 cm 1 ) exhibits spectral features due to
CO2 molecules involved in the charge carrying species and are referred to as “core” CO2 units.
These CO2 units are geometrically distorted from the geometry of a free CO2 molecule and exhibit
di↵erent CO stretching frequencies.

4.3.1

Infrared Spectra of [CoO(CO2 )n ]The spectra of [CoO(CO2 )n ]Clusters

show a single narrow peak in the solvent region and four

pronounced features in the core ion region of the spectra (see Figure 4.1). Additionally, there are
weaker peaks observed in the core ion region similar to what was seen in [Co(CO2 )n ] clusters. As
has been reported in the previously in Chapter 3, the binding energy of a core CO2 is on the order of81intensity [arb. units]

3
4
5
6
7
8
1000

1200

1400

1600

1800

2000

2200

2400

wavenumber [cm-1]
Figure 4.1: Experimental spectra of [CoO(CO2 )n ] for (n = 3 8). The numbers denote thenumber of CO2 molecules present in the cluster. The exploratory experimental data for n = 8 wastruncated after determining that the spectrum was not appreciably evolving as a function of cluster
size. The left and right traces are individually normalized and are not on the same scale.

10000 cm 1 , which is well above the photon energies used in this experiment. Since fragmentation
is first observed at n = 3, the cluster is likely of the form [CoO(CO2 )2 ] ·CO2 so that a solvent
molecule is present (binding energy of solvent CO2 ⇠1600 cm 1 ) and the core species contains two
CO2 units.
Performing DFT calculations on a [CoO(CO2 )2 ] core yields a variety of possible structuralconfigurations (see Figure 4.2). Both singlet and triplet states are possible for cobalt oxide-CO2
anion complexes and give rise to di↵erent structural configurations. The predicted global minimum

82
contains a CO3 and a CO2 moiety, both bound in a bidentate fashion to the Co atom (see Figure
4.2 isomer 1,3 A). The bonding motif in isomer A can be likened to the ⌘ 2 structure of a CO2
molecule bonding to a metal atom where both the CO2 and CO3 exhibit an ⌘ 2 bonding motif. [91]In the singlet configuration, the CO3 and CO2 moieties are nearly perpendicular while in the triplet
configuration they are twisted by about 40 to each other. Additionally, the triplet structures tend
to be predicted to be lower in energy than their singlet counterparts.
The next higher lying core isomer is B where the CO2 unit is instead bound in an ⌘ 1 fashion(see Figure 4.2). In isomer B, the CO2 and CO3 moieties are nearly coplanar. Interestingly, while
the bonding of the CO2 unit is that of an ⌘ 1 configuration, the OCO bond angle is 10 larger thanthe OCO angle of similar structures in [Au(CO2 )n ]and [Ag(CO2 )n ]clusters (155 vs. 145 ,respectively). The structure next higher in energy consists of a CO3 and a CO2 moiety where theCO2 is bound by both of the oxygen atoms to the Co atom (isomer C) rather than by the carbon
atom. Isomer D contains an ⌘ 2 CO2 moiety and an ⌘ 1 bound CO3 where there is only one O-Co
bond. The CO3 in this arrangement is structurally very similar to the formate structures discussed
previously for [Au(CO2 )n ] , [Ag(CO2 )n ] and [Cu(CO2 )n ] clusters. The OCO angle of the nonbonded portion of the CO3 moiety is 136 which is nearly the same, for example, as the OCO angle
calculated for [Ag(CO2 )5 ]at 139 . Based on the Walsh digram for CO2 (see Figure 3.1), thisangle is predicted to decrease as additional charge is polarized onto the bound CO2 moiety until
it reaches approximately 135 . Without performing a charge analysis, the geometry of the CO3
moiety in isomer D suggests that a full excess charge has been taken up by this subunit. Finally,
isomer E shows a lone oxygen atom opposite a C2 O4 moiety bonded to the Co atom.
Figure 4.3 shows a comparison of the calculated isomers in Figure 4.2 to the experimental
spectrum of [CoO(CO2 )3 ] . Only isomers 1,3 A and 1 D recover the features in the experimental
spectrum. Isomers 1,3 A predict the peaks at 1705 cm 1 and 1815 cm 1 with isomer 1 D recovering
the 1655 cm 1 and 1840 cm 1 features. While isomer 3 B does recover the feature at 1655 cm 1 ,the second predicted feature around 1750 cm 1 is not present in the experimental spectrum with
an appropriate intensity compared to the 1655 cm 1 peak. Based on this observation, isomer

83

Figure 4.2: Calculated lowest energy structures for [CoO(CO2 )n ] . The capital letter denotes thecore structure and superscripts 1 and 3 denote singlet and triplet, respectively. Relative energies
are given in meV and are all with respect to isomer 3 A.

84
B, if present in our experiment, only contributes to the experimental spectra as a minor isomer.
Based on the observed relative peak intensities, it is clear that isomer A is the dominant core
structure while there is less population of isomer D. At first glance, it is surprising that isomer D
is a stable structure as it appears that the CO3 could swing to either side making a second Co-O
bond and forming isomer A. However, calculations on such a process revealed at least an 8000
cm 1 barrier to interconversion from isomer D to A. This is likely due to a substantially di↵erent
orbital configuration of isomer A compared to isomer D resulting in the high predicted barrier tointerconversion. Further calculations outside the scope of this work would be required to further
understand this conversion process. None of the other calculated structures in Figure 4.2 recover
the experimental spectrum.
As the degree of solvation increases, the experimentally observed peaks do not shift appreciably. Since isomers 1,3 A and 1 D are the only structures to recover the experimental features, it isnecessary to investigate the e↵ect of solvation on the predicted frequencies of these core structures.
Figure 4.4 shows di↵erent solvation positions for isomers 1,3 A and 1 D. In addition to solvation ofcore structures A and D, a new energetically relevant core isomer is predicted (see isomer F in Figure 4.4). Isomer F results from the addition of the third CO2 molecule to the lone oxygen atom inisomer E creating an ⌘ 2 CO3 moiety opposite the C2 O4 subunit. In 3 F, the two subunits are planar
while in the singlet structure (1 F) the CO3 and C2 O4 are perpendicular. If isomer F is populated
in the experiment, it should not be observable until [CoO(CO2 )4 ]where the first solvent CO2

molecule would be present. The core structure of isomer 3 D is predicted to be dependent on thesolvation environment. If the CO3 moiety is solvated, isomer 3 D is recovered, however if the CO2
subunit is solvated the structure converts to 3 A. As solvation increases, it is unlikely that the CO2
moiety will remain unsolvated, therefore we exclude isomer 3 D from consideration for assignment
to the spectra. This same behavior is not observed for 1 D.
Singlet and triplet structures exhibit a 25 - 30 cm 1 di↵erence in the predicted frequency ofthe lower energy feature of the dominant peaks and a 5 - 10 cm 1 di↵erence for the higher energy
dominant feature. This is on the same order as the predicted di↵erence in signatures for solvationintensity [arb. units]

85

3

A

0 meV

1

A

425 meV3

B

748 meV

3

C1046 meV3

E1209 meV1

1969 meV

D

1000

1200

1400

1600

1800

2000

2200

-1

wavenumber [cm ]Figure 4.3: Comparison of calculated core isomers for [CoO(CO2 )2 ] to the experimental spectrumof [CoO(CO2 )3 ] . Isomer labels and relative energies are given for each structure. See text for
discussion.

86

Figure 4.4: Calculated lowest energy structures for [CoO(CO2 )3 ] . The capital letter denotes
the core structure, the superscript 1 and 3 indicate singlet and triplet configurations, and thenumber/letter combination denote the number and configuration of the solvent molecules. Relative
energies are reported in meV.

87
isomers of the same core structures. In contrast, the splittings between the experimental signatures
are 50 cm 1 and 30 cm 1 for the lower and higher energy features, respectively. Based on thewidth of the experimental peaks and the stable peak position as cluster size increases suggests that
there is likely only one multiplicity present and that theory overestimates solvation e↵ects. While
we cannot definitively assign one multiplicity to the spectra of [CoO(CO2 )n ] , we can distinguishstructures with di↵erent cores. We also believe that theory over estimates the e↵ect of solvationon the core structures. For simplicity, the rest of the discussion will look only at structures with asinglet configuration.
Figure 4.5 shows a comparison of the singlet structures presented in Figure 4.4 to the experimental spectrum of [CoO(CO2 )3 ] . It is clear that isomer F does not recover any of the dominantfeatures. However, it is possible that isomer F does contribute to the spectra of larger clusters(n4) in the region from 1750 cm 1 to 1800 cm 1 . The experimental spectrum is therefore as-signed to isomer A in either a triplet or singlet configuration (1705 cm 1 and 1815 cm 1 features)and isomer 1 D (1655 cm 1 and 1840 cm 1 features). As the number of CO2 molecules present
increases, the peaks assigned to isomer 1 D grow in intensity relative to those assigned to isomers
1,3 A. This suggests that the population of isomer 1 D increases with increasing cluster size relative

to the population of isomers 1,3 A.
Table 4.1: Calculated partial charges for the predicted core structures of [CoO(CO2 )n ] presented
in Figures 4.2 and 4.4. The partial charges are reported in units of e.
Core Structure
3A
1A
3D
1D
3F
1F

Co
+1.10
+0.79
+0.88
+0.68
+1.42
+1.27

CO2
-0.68
-0.61
-0.55
-0.57

CO3
-1.42
-1.18
-1.33
-1.11
-1.12
-1.06

C2 O 4

-1.30
-1.21

The final features to discuss are the minor peaks that appear between 1700 cm 1 and 1800
cm 1 in [CoO(CO2 )4 ]and larger clusters (see Figure 4.1). Since these features first appear in88intensity [arb. units]

1

A(1a)

0 meV

1

A(1b)

86 meVD(1a)1546 meV

1

1671 meV

1

D(1b)

1

1000

1200

1400

1600

1800

2000

F

343 meV

2200

wavenumber [cm-1]
Figure 4.5: Comparison of the calculated singlet structures in Figure 4.4 to the experimental
spectrum for [CoO(CO2 )3 ] . Isomer labels and relative energies are given to the right. See text for
discussion.

89[CoO(CO2 )4 ]clusters, the cluster composition is likely of the form [CoO(CO2 )3 ] ·CO2 . Isomer

F is the most likely candidate and is predicted to exhibit vibrational frequencies within the 1700
cm 1 and 1800 cm 1 region. Exploratory calculations of various solvation positions around isomer
F suggest that the charge distribution of isomer F is more polarizable than that of either A or D
resulting in solvation around the C2 O4 or CO3 moieties giving rise to di↵erent predicted spectra.
These shifts are to be expected since the combination of a C2 O4 and a CO3 subunit is large
compared to all other calculated structures creating a more polarizable charge distribution. Isomer
E also is predicted to exhibit vibrational features in this region and is a possible candidate forassignment of these features.

4.3.2

Infrared Spectra of [NiO(CO2 )n ]Clusters

Figure 4.6 shows the experimental spectra of [NiO(CO2 )n ]clusters (n = 2

6). As was

mentioned at the beginning of this section, there are solvent (2260 cm 1 to 2400 cm 1 ) and core(1000 cm 1 to 2150 cm 1 ) spectral regions. Since the predicted structural motifs present in thisstudy are similar to those predicted for [CoO(CO2 )n ] , the same labeling scheme will be used.
Unlike in [CoO(CO2 )n ]clusters, fragmentation is first observed at n = 2 rather than at

n = 3. Since a solvent peak is present for [NiO(CO2 )2 ]clusters, the form of the species giving

rise to the observed features for n = 2 must be [NiO(CO2 )] ·CO2 . Calculations predict only two
stable structures of this form; isomer G is a CO3 moiety bound in an ⌘ 2 motif to the nickel atom
and isomer H is a lone oxygen atom opposite an ⌘ 2 CO2 unit (see Figure 4.7). These two structures
are then solvated around the CO3 and CO2 moieties, respectively, by the second CO2 molecule.
Solvation around the open side of the Ni atom or the lone oxygen atom result in the formation ofisomer A. While structures G and H are predicted to be stable, they are not predicted to be thelowest energy conformers calculated for [NiO(CO2 )2 ] (see isomer A in Figure 4.7).
The calculated spectra of isomers G and H are compared to the experimental spectrum of[NiO(CO2 )2 ]in Figure 4.8. The combination of isomers G and H recovers the two pronounced

peaks at 1670 cm 1 and 1750 cm 1 in the experimental spectrum, respectively. The signal to noise

90

intensity [arb. units]

2

3

4

5

6
1000

1200

1400

1600

1800

2000

2200

2400

wavenumber [cm-1]
Figure 4.6: Experimental spectra of [NiO(CO2 )n ] . The numbers indicate the number of CO2
molecules present. The left and right traces are individually normalized and are on di↵erent scales.

ratio of this spectrum is quite low compared to the spectra for n

3. As was mentioned previously,

isomers G and H are not the lowest energy structures at this cluster size but are the only calculated
structures that are expected to be experimentally observable for n = 2. It is likely that isomer A
has a much larger population in [NiO(CO2 )2 ] clusters leading to low populations of both G and H
and, consequently, a relatively small number of parent ions containing a solvent CO2 molecule. It
should also be noted that similarly broad features observed in [Ni(CO2 )2 ] clusters were attributed
to multiphoton e↵ects since no solvent feature was observed (see Section 3.3). It is likely that theorigin of the broad feature at 1825 cm 1 is multiphoton dissociation of isomer A (see Figure 4.8)

91

Figure 4.7: Calculated lowest energy structures for [NiO(CO2 )n ] . Structures in the leftmost
column are bare core structures while the structures to the right are solvation isomer(s) of thecore structures. The letters denote the core structure while the letter/number combination indicate the number of solvent molecules and their orientation around the core. Relative energies for [NiO(CO2 )] structures are shown in italics, [NiO(CO2 )2 ] are shown in plain text and[NiO(CO2 )3 ] are shown in bold. No relative energy is presented for F(1) since there is no otherstructure of the form [NiO(CO2 )4 ] to compare it with. All relative energies are reported in meV.

intensity [arb. units]

92

G(1)

H(1)

A

1000

1200

1400

1600

1800

2000

2200

wavenumber [cm-1]
Figure 4.8: Comparison of calculated structures G(1), H(1) and A to the experimental spectrumof [NiO(CO2 )2 ] . See Figure 4.7 for relative energies and text for discussion.

but, due to the low signal to noise ratio of the n = 2 spectrum, the presence of multiphoton e↵ects
is difficult to confirm.
For clusters [NiO(CO2 )3 ]and larger, not only does the signal to noise ratio improve dras-

tically, but many new peaks are observed. This change in the spectrum suggests that multiple
new structures contribute to the experimental spectrum. Figure 4.7 shows all of the relevant core
structures calculated for [NiO(CO2 )n ]clusters with the bare core structures shown in the left

93
column and various solvation isomers in the center and right columns. Like [CoO(CO2 )n ] , isomer A is predicted to be the global minimum structure. Structure D, also found and assigned in[CoO(CO2 )n ]clusters, consists of an ⌘ 2 CO2 moiety and an ⌘ 1 CO3 subunit bound by only oneof the oxygen atoms. Structures E and F are similar as they both contain a C2 O4 moiety opposite
either a lone oxygen atom or a CO3 unit. Structures A, D, and E should all be observable at[NiO(CO2 )3 ] while isomer F should first be observable at [NiO(CO2 )4 ] .
Figure 4.9 shows a comparison of select core isomers with formula [NiO(CO2 )3 ]to theexperimental spectrum of [NiO(CO2 )3 ] . Unlike in [CoO(CO2 )n ] , it is clear that solvation position
makes a di↵erence in the frequencies recovered. Structure A(1a) recovers the two major peaks at
1685 cm 1 and 1880 cm 1 while A(1b) recovers the major peaks at 1685 cm 1 and 1840 cm 1 .
The di↵erence between these two structures can be summarized by whether or not the CO2 moiety
is solvated. The two higher frequency modes represent molecular motion on the CO2 moiety withthe lower frequency of the two modes being recovered when the solvation position behind the CO2
moiety is filled and the higher frequency when this position is vacant. This di↵erence in predictedfrequency as a function of solvation position shows that the charge distribution on the CO2 moiety
is noticeably a↵ected by solvation. Conversely, the charge distribution of the CO3 moiety, encodedin the 1685 cm 1 feature, is not strongly a↵ected by solvation as both A(1a) and A(1b) recover
the 1685 cm 1 peak. The peak at 1840 cm 1 is present for all clusters n

3 indicating that if thesolvation position behind the CO2 moiety is filled, the peak at 1840 cm 1 is recovered.
The features at 1645 cm 1 and 1880 cm 1 are recovered by isomer D(1a) where the CO3
moiety is solvated. The predicted charge distribution of isomers A and D di↵er only on the CO3
moiety and the Ni atom leaving the CO2 unit relatively unchanged. This similarity in partial
charge on the CO2 moiety results in spectral overlap of the 1880 cm 1 feature by isomers A(1a)and D(1a). This overlap is di↵erent from [CoO(CO2 )n ] where the charge distributions of isomers
A and D are completely di↵erent resulting in no overlap of their predicted spectra (see Table 4.1
and Figure 4.4). Another way to interpret the dominant spectral features is that the peaks from1600 cm 1 to 1700 cm 1 encode the connectivity of the CO3 moiety (ie. ⌘ 1 or ⌘ 2 ) while the features94

intensity [arb. units]

A(1a)

A(1b)

D(1a)

E(1)

F

1000

1200

1400

1600

1800

2000

2200

wavenumber [cm-1]
Figure 4.9: Comparison of calculated structures for [NiO(CO2 )3 ] to the experimental spectrumof [NiO(CO2 )3 ] . See Figure 4.7 for relative energies and text for discussion.

95
between 1840 cm 1 and 1880 cm 1 indicate the occupation of the solvation positions around theCO2 moiety. Table 4.2 shows the calculated charge distributions for the reported [NiO(CO2 )n ]core structures.
Table 4.2: Calculated partial charges for the predicted core structures of [NiO(CO2 )n ] presented
in Figure 4.7. The partial charges are reported in units of e.
Core Structure
3G
1H
3A
1D
3E
1F

Ni
+0.56
+0.65
+0.93
+0.72
+1.17
+1.39

CO3
-1.56
-1.41
-1.22

CO2

O

-0.71
-0.52
-0.50

-0.94

-0.76
-1.10

C2 O 4

-1.41
-1.29

The last set of features to address are the weak peaks observed between 1730 cm 1 and 1830
cm 1 . These features can in part be attributed to isomer E(1) which is predicted to exhibit a peakaround 1735 cm 1 (see Figure 4.9). At n = 4, a second feature appears at 1780 cm 1 . With the1745 cm 1 feature being attributed to isomer E(1) and a cluster size of n = 4 at the onset of the1780 cm 1 feature, it is logical to attribute the peak at 1780 cm 1 to isomer F. The comparison in
Figure 4.9 shows that structure F does indeed predict a feature around 1780 cm 1 . Both isomer E
and F persist as minor contributors to the experimental spectra as cluster size increases.

4.4

Conclusions
The structural motifs observed in both [CoO(CO2 )n ] and [NiO(CO2 )n ] clusters are similar.

Structure A is the dominant structure in both sets of spectra with isomer D increasing in prevalenceas the number of CO2 molecules increases. Additionally, structures such as E and F are present as
minor isomers for both CoO and NiO and persist with increasing cluster size.
While the structures and spectra for both clusters are similar, there are small di↵erences.
These di↵erences can be attributed to the charge distributions of the core ions (See Tables 4.1
and 4.2). For [CoO(CO2 )n ]clusters, the charge distribution is dependent on the connectivity of

96
the CO3 moiety while in [NiO(CO2 )n ] the charge distribution changes based on the connectivityof the CO3 moiety and the occupation of specific solvation positions. The overall similarities
and di↵erences of these two cluster species are reminiscent of those between [Co(CO2 )n ][Ni(CO2 )n ] clusters.

andChapter 5
Observation of a Vibrationally Induced Electron Transfer Reaction in a Binary
Molecular Complex

This Chapter has been reprinted with permission from Knurr, B. J., McCoy, A. B., Weber, J.
M., “Vibrationally Induced Charge Transfer in a Bimolecular Model Complex in Vacuo.” Journalof Chemical Physics, 2013. 138: 224301. DOI: 10.1063/1.4808048. Copyright 2014, AIP Publishing
LLC.

5.1

Introduction
Electron transfer is an elementary step in many chemical reactions and biological processes.

[129] The means by which electrons can be transfered during a chemical reaction are many including
detachment and recapture or simply by experiencing a potential di↵erence causing the electron tomove within the system. One unique class of reactions are those where the electron is excited suchthat a transfer event occurs by photodetachment of the electron. In this chapter, a vibrationallyinduced electron transfer reaction in a model system is observed and discussed.
In condensed phase chemistry, the potential energy surface tends to be characterized byrelatively flat landscapes for reactants and products which are separated by a barrier. However,in gas phase reactions, pre-reaction complexes (entrance channel) and post-reaction complexes(exit channel) can be formed due to a lack of solvation e↵ects that would otherwise destabilize
the potential wells supporting these complexes. This is particularly true for reactions involving
ionic species. [130–132] Experiments on mass selected ions in vacuo are well suited to study the

98fundamental properties of entrance and exit channel complexes and the chemical reaction connecting
them because the complexes can be prepared with well defined composition. One interesting classof reactions involves an ion-molecule complex which has an electron excited into the detachmentcontinuum is allowed to evolve. In these systems, the electron wave function is prepared at “halfcollision” and the excited electron can either escape or be recaptured by the complex. [133–141]
In certain systems vibrational excitation can lead to electron transfer on the electronic groundstate potential energy surface. This has been observed, e.g., in CO2 reduction mediated by excitation of a water network in CO2 ·(H2 O)6 clusters, [142] an SN 2 like electron transfer reaction in
Cl ·CH3 Br complexes initiated by the excitation of CH stretching modes, [143] and in hydrated
nitromethane anion. [144]An interesting prototype for this class of electron transfer reactions would be a vibrationallyinduced charge transfer between two molecular partners (rather than atom/molecule [143] ormolecule/cluster [142, 144]). Here, the two molecules could be isolated as a binary complex orsolvated by an inert solvent species (i.e. noble gas atoms). The question is whether or not mode
specific behavior can be observed leading to increased efficiency of the electron transfer process.
Observation of mode specific behavior would indicate that these “special” vibrational modes areconnected efficiently to the transition state of the chemical reaction.
To carry out an experiment of this nature, there are certain requirements for the molecularpartners. The donor must be able to detach an excess electron through vibrational excitation. In
other words, the binding energy of the excess electron must be sufficiently low such that vibrational
excitation can result in electron emission. Nitromethane (CH3 NO2 ) has a well characterized valenceanionic state where the excess electron is bound by ⇠1325 cm 1 . [145] Previous work has shownthat vibrational excitation of CH stretching modes leads to efficient vibrational autodetachment innitromethane as well as in other nitroalkanes. [146–148] These properties of nitromethane make it
an ideal candidate as the electron donor species.
The acceptor must be able to exhibit a clearly detectable signature that electron transfer has
occurred. Additionally, it would be ideal if recapture by the donor molecule is less favorable than99
transfer to the acceptor. Considering these two criteria, iodomethane (methyl iodide or CH3 I) waschosen as an ideal acceptor molecule since it can only dissociatively attach an electron resulting
in the irreversible formation of I

and neutral fragments. The predicted outcome of vibrational

excitation of a complex of nitromethane and iodomethane is shown in Equation 5.1.

CH3 NO2 · CH3 I + h⌫ ! [CH3 NO2 · CH3 I]⇤ ! I + neutral fragments

(R1)

(5.1)

Equation 5.1 (R1) is also shown graphically in Figure 5.1 (see top equation). The dissociative swave capture reaction for CH3 I is strongly exothermic [149] and has been extensively characterized
before. [149, 150]
A second question with regard to these types of reactions is on what time scale the electrontransfer process occurs. Depending on the time scale of the reaction, it may be possible to shut
down the electron transfer reaction by employing a “bath” of solvent Ar atoms. The two competing
processes would then be the electron transfer reaction resulting in the formation of I and relaxationof the vibrational energy to the solvent bath yielding Ar evaporation (see Equation 5.2 and bottom
reaction in Figure 5.1).

CH3 NO2 · CH3 I · Arn + h⌫ ! [CH3 NO2 · CH3 I]⇤ · Arn ! CH3 NO2 · CH3 I + nAr

(R2)

(5.2)

In addition to information on the time scales of the competing reactions, insight can also be garneredon the barrier height of the electron transfer reaction based on the binding energy of Ar atoms andthe number required to quench the reaction. [146, 151, 152]
In this chapter, we show that vibrationally mediated electron transfer does occur in thiscomplex. We also answer questions concerning the structure of the entrance channel complex andgive a lower limit estimate for the barrier height of the reaction. For experimental details pleaserefer to Chapter 2.

100

Figure 5.1: Schematic representations of the potential results of vibrational excitation of the barebinary complex (top) and the Ar solvated complex (bottom). R1 refers to Equation 5.1 while R2
refers to Equation 5.2. See text for additional discussion.

5.2

Computational Methods
Density functional theory calculations were used to determine the structures of reaction

species and the energetics of the reaction. Calculations were performed employing the B3-LYP
functional [59, 60] with dispersion correction [89] as implemented in the TURBOMOLE V. 5.9.1
suite of programs. [58] For all structure calculations the aug-cc-pVTZ-PP basis set [153] was used
for iodine and aug-cc-pVTZ basis sets [65] were used for all other atoms. Potential energy scansof the C-I bond distance were performed using def2-TZVPP basis sets [61] for all atoms. Vibrational frequencies were calculated using the AOFORCE program. [66, 67] The calculated harmonic
vibrational frequencies are scaled to account for the anharmonicities of the CH stretching modes.
The scaling factors were obtained separately for each CH stretching mode in comparison with ex-

101
perimental values for neutral methyl iodide [37] and nitromethane anion. [147] The partial charges
reported were calculated using a natural population analysis. [68]5.3

Results and Discussion
The expected fragment ion to be generated in this reaction is I

(see reaction R(1) at the

top of Figure 5.1) by dissociative attachment of the transfered electron. [149] Previous studies on
I ·CH3 I have been observed to generate I

upon electronic excitation in the equivalent of chargetransfer to solvent. [133–135] Irradiation of [CH3 NO2 ·CH3 I]of Iclusters results in the generation

photo-products (see Figure 5.2 top trace). Five peaks are observed in the experimental

spectrum when monitoring the formation of I

and are all within the region corresponding to

CH stretching frequencies. However, both nitromethane and iodomethane contain CH3 groups
complicating identification of the experimentally observed peaks.
To assign the observed peaks, we carried out experiments on [CH3 NO2 ·CD3 I] clusters. Here,all of the CH stretches belonging to iodomethane are replaced with CD stretches and therefore areshifted to lower frequencies. The spectrum of [CH3 NO2 ·CD3 I] monitoring the generation of I isshown in the bottom trace of Figure 5.2. As can be seen, the two peaks observed at 2950 cm 1
and 3056 cm 1 in the spectrum of the nondeuterated complex are absent in the spectrum of thedeutered complex indicating that those two peaks belong to CH stretches on iodomethane. The
remaining peaks at 2804 cm 1 , 2937 cm 1 and 2977 cm 1 can then all be attributed to CH stretches
on nitromethane (see Table 5.1).
There are a number of configurations in which the two molecules (CH3 NO2 and CH3 I) can bearranged to form stable structures (see Figure 5.3). However, only two of the calculated structures
have the starting molecular species intact (see structures I and II in Figure 5.3). In conformation I,
the methyl group of CH3 I is pointing towards the nitro group in CH3 NO2 , similar to the structureof a SN 2 entrance channel complex. The corresponding SN 2 exit channel is a structure where themethyl group binds to one of the oxygen atoms in the nitro group (see conformation IV in Figure5.3), which is calculated to lie 1.16 eV lower in energy than conformation I. In conformation II, theI fragment yield [arb. units]

102

-

-

CH3NO2 ·CH3I-

CH3NO2 ·CD3I
2600

2700

2800

2900

3000

3100

3200

3300

-1

wavenumber [cm ]Figure 5.2: Photodissociation action spectra of CH3 NO2 ·CH3 I (top trace) and CH3 NO2 ·CD3 I(bottom trace) acquired by monitoring the formation of I photoproducts. The dashed lines indicate the CH stretching modes attributed to the CH3 I molecule within the complex. See text for
further discussion.

103
1

Table 5.1: Experimentally determined and calculated peak positions in cm
for the dominant vibrational features observed in Figure 5.2. The experimental frequencies are reported for
CH3 NO2 ·CH3 I with values for CH3 NO2 ·CD3 I given in parentheses where observed. Calculated
frequencies (see above for discussion) are given for the non-deuterated complex only and are reported in km·mol 1 . The vibrational assignments given are based on structure I in Figure 5.3.

Assignment
Unique CH Stretch CH3 NO2Symmetric CH2 Stretch CH3 NO2
Antisymmetric CH2 Stretch CH3 NO2Symmetric CH3 Stretch CH3 I
Antisymmetric CH3 Stretch CH3 I
Antisymmetric CH3 Stretch CH3 I

Experimental
Wavenumber
2804 (2801)
2937 (2934)
2977 (2978)
2950
3056
Not Observed

Calculated
Wavenumber
2841
2945
2985
2947
3072
3097

Calculated
Intensity (km·mol 1 )
255
58
25
76
73
0.02

iodine atom of CH3 I points towards the nitro group in CH3 NO2 , however this structure is ⇠100 meVhigher in energy than conformation I. Other starting geometries either converged to conformation
I or II or resulted in dissociation of the CH3 I molecule and the subsequent formation of I solvatedby CH3 NO2 and a methyl radical (this family of structures is represented by conformation III). In
principle, conformation III is a possible exit channel complex as it lies ⇠500 meV lower in energy
than the entrance channel complex conformation I.
The calculated IR spectra of conformers III and IV do not match the experimental spectra.
In particular they both lack a dominant peak around 2800 cm 1 , a telltale signature of CH3 NO2(the symmetric methyl stretch). This feature has been observed and well cataloged before in both
CH3 NO2 ·Arn and CH3 NO2 ·(H2 O)n ·Arm clusters. [146,147] In order to confirm if conformation I or
II is dominant or even present in the ion beam, we compare the calculated spectra for conformers I
and II to the experimental spectrum (see Figure 5.4). It is clear in the comparison that conformation
I recovers the experimental spectrum much better than does conformation II. Additionally, withthe excess charge being localized on the nitro group, it is likely that the proximity of the nitro
group to the iodine atom in conformation II will result in a substantially lower barrier to electrontransfer than in conformation I. This lowering of the reaction barrier may be enough to allow

104

Figure 5.3: Calculated lowest energy structures for [CH3 NO2 ,CH3 ,I] complexes. The relative
energy for conformation II is with respect to conformation I. See text for discussion.

the electron transfer to proceed without vibrational excitation if the internal energy of the cluster
is greater than the barrier height. We therefore, based on the predicted spectra and physicalorientation considerations, assign our spectra to conformation I with negligible contribution from
conformation II. Assignments of the vibrational features are given in Table 5.1.
As has been shown previously, Ar predissociation spectroscopy can give an estimate to thebarrier height and information on the relative time scales of reaction to intra- and intermolecularintensity [arb. units]

105

2600

2700

2800

2900

3000

3100
-1

3200

3300

wavenumber [cm ]Figure 5.4: Comparison of the experimental spectrum of CH3 NO2 ·CH3 I·Ar to the simulated spectraof calculated conformers I (middle panel) and II (bottom panel). The spectrum of the argon tagged
ion is used because it is better resolved compared to the spectrum of the bare ion. Note that thepeak at ⇠2950 cm 1 for conformer I contains two close lying transitions (see Table 5.1) which
account for the splitting observed in the experimental spectrum.

106
molecular vibrational relaxation. [146, 151, 152] In the three studies cited, there is always a specific number of Ar atoms where the initially observed process (chemical reaction [151] or electron
emission and vibrational autodetatchment [146,152]) is turned o↵. For CH3 NO2 ·CH3 I·Arn , similar
behavior is observed with a certain number of Ar atoms quenching the electron transfer reaction(see Figure 5.5). For n = 1, the only anionic fragment observed is I

corresponding to R(1) in

Figure 5.1. Spectra are sharper than they are for the bare cluster likely owing to the lower internal temperature of the cluster when Ar solvated. For n = 2, two di↵erent fragment channels areobserved. While a small amount of I

is still observed (red trace), the major pathway is a loss ofAr from the cluster without the formation of I but rather with the CH3 NO2 ·CH3 I cluster intact(blue trace). The latter process corresponds to reaction R(2) as seen in Figure 5.1 and equation
5.2. At n = 3, no I formation is observed indicating that solvation by three Ar atoms is sufficient
to deactivate the reaction. It should be noted that Ar solvation deactivated the electron transfer
reaction for the deuterated cluster completely at n = 2 indicating that the barrier to reaction for
CD3 I is slightly higher than for CH3 I.
The results of the Ar solvation experiments allow for an estimation of the barrier height. The
binding energy of an Ar atom to nitromethane has been extensively measured by photoelectron
spectroscopy and is reported to be ⇠60 meV or 480 cm 1 . [145] We assume that the binding energyof Ar to the CH3 NO2 ·CH3 I complex is roughly the same. Since solvation by two Ar atoms issufficient to mostly suppress the reaction channel, the height of the barrier should be on the orderof the di↵erence between the photon energy and the Ar binding energy. The lowest frequency
peak observed is at ⇠2800 cm 1 . Subtraction of the binding energy of the two Ar atoms gives anestimate of the barrier height at ⇠1800 cm 1 or 230 meV. The typical barrier for an SN 2 reaction
is between ⇠1000 cm 1 and ⇠3000 cm 1 putting our estimate squarely in the typical range for an
SN 2 reaction. [154–157] This estimate, within an error bar of 60 meV (the binding energy of one
Ar atom), is in fact a lower limit of the barrier height as tunneling is possible at higher barrierenergies. A schematic representation of the potential energy surface for the reaction coordinate isshown to scale in Figure 5.6.

107

photofragment action [arb. units]

n=3

n=2

n=1

n=0

2600 2700 2800 2900 3000 3100 3200 3300
-1

wavenumber [cm ]Figure 5.5: Experimental spectra of CH3 NO2 ·CH3 I·Arn (n = 0 3). The di↵erent colors of thespectra indicate a change in the photofragmentation channel. The red curves are monitoring chargetransfer and the subsequent formation of I photoproducts (see R1 in Figure 5.1) while the blue
curves are monitoring Ar evaporation (see R2 in Figure 5.1). Note that both channels are observed
for n = 2.

0

CH3NO2 + CH3I + e

108

-

(a)

-200

rel. energy [meV]

-400
-600

(e)

(b)

_

(c)

I + CH3+ CH3NO2

conf. I

-800
-1000

(d)

-1200
-1400

conf. III_I + CH3NO2CH3

-1600
-1800

conf. IV
reaction coordinate

Figure 5.6: Schematic representation of the potential energy surface along the reaction coordinate ofdissociative electron transfer from CH3 NO2 to CH3 I (energies are to scale). (a) Adiabatic electron
affinity of CH3 NO2 (172 meV). [145] (b) Binding energy of the CH3 NO2 ·CH3 I entrance channel
complex (490 meV, calculated for conformation I). (c) Estimated lower limit of the barrier height
for the electron transfer reaction (230 meV, experimentally determined from Ar solvation). (d)Relative energy of the exit channel complex for structure III (500 meV, solid curve) and structureIV (⇠1150 meV, dashed curve). The dashed line represents the SN 2 product. (e) Exit channel
asymptote for the formation of I and neutral products based on the two possible exit channel
complexes. The solid curve is based on the heat of reaction for dissociative electron attachment
to CH3 I (620 meV). [149] The dashed curve represents the I + CH3 NO2 CH·3 asymptote (⇠1550
meV).

To further characterize the reaction and C-I bond cleavage, we calculated the potential energyof the cluster and the partial charge on the iodine atom as a function of C-I bond distance (seeFigure 5.7). At a C-I bond distance of 2.6 Å the calculated curves show a discontinuity indicating
the likely presence of the transition state. The molecular coordinate showing the most change in thisregion is the out-of-plane angle of the CH3 group of the iodomethane (✓OOP ). ✓OOP changes from

109
15.5 where the C atom is pointing towards the I atom to 5.5 where the C atom is pointing toward
the nitro group with minute changes to a C-I bond distance of 2.6 Å. In addition to lengthening of
the C-I bond, we regard this angle (✓OOP ) as a characteristic geometry parameter in this reaction.
Additionally, the ✓OOP angle is characteristic of an SN 2 reaction giving credence to the idea that
this reaction possesses SN 2 character even though the CH3 NO2 CH·3 radical is an unusual product.
In addition to geometry considerations, we can look at the amount of excess charge localizedon the I atom as a function of C-I distance (bottom panel in Figure 5.7). The amount of chargechanges abruptly from -0.5 e to -0.8 e at the transition state C-I distance of 2.6 Å. As the C-I bonddistance continues to increase, the charge on the iodine approaches

e indicating that the chargetransfer reaction is complete.
The calculated potential energy curves discussed so far are an approximation for the adiabatic
potential. If we fix ✓oop to the values of the entrance and exit geometries while varying the C-I
bond length, we recover good approximations for the diabatic potential curves as well (shown indetail in the inset of the top panel of Figure 5.7). While it appears that we have recovered thepotential energy surface, we caution that the reaction barrier height in this model is only predicted
to be 54 meV. This is much lower than our experiments with Ar solvation suggest. While thequalitative picture of the SN 2 type of molecular motion leading to reaction appears to be correct,
the calculated energetics of the potentials in Figure 5.7 would benefit from higher level calculations
that are outside the scope of this study.
Finally, the question of whether or not mode-specific behavior is observed for this electrontransfer reaction can be addressed. At best, mode specificity is weak for this reaction and several
observations allow this conclusion. First, excitation of vibrational modes on either of the methylgroups results in electron transfer. This observation indicates that the mode into which energy isdeposited is inconsequential for this reaction to proceed. Second, Ar evaporation must be fasterthan the electron transfer process since all observed vibrational frequencies were visible in the Ar
loss spectra and the electron transfer reaction was shut down by addition of three Ar atoms. We
expect intra- and intermolecular vibrational relaxation to prevent strong mode specific behavior

110
relative energy [meV]

relative energy [meV]

400
200
0
-200

200
100
0
-100
2.4

-400

2.6

2.8

C - I distance []

-600
-800
-1000
-1200

-

partial charge on I (e )

0.0
-0.2
-0.4
-0.6
-0.8
-1.0
2.0

2.5

3.0

3.5

4.0

4.5

C - I distance []Figure 5.7: Calculated potential energy curve (top) and charge on the iodine atom (bottom) as
a function of the CI bond distance. The full squares in the top trace are calculated with only
the CI bond distance fixed, letting all other molecular coordinates relax. The open diamonds
represent calculations where the out-of-plane angle of the methyl group of the CH3 I molecule (✓oop )was fixed at 10.0 , corresponding to the geometry at a CI bond distance of ⇠2.575 Å in theunrestricted calculation. The open squares are data points calculated with the same angle (✓oop )fixed at 5.5 , corresponding to the geometry at 2.625 Å CI bond distance in the unrestrictedcalculation. We note that ✓oop changes very rapidly as the transition state is approached. The
resulting approximated diabatic potential energy curves are shown as dashed lines in the inset of
the top panel; the structures shown as insets from left to right represent the entrance channel,transition state, and exit channel geometries, respectively.

111
as it appears that energy likely needs to localize in either the nitro wagging [148] or CI stretchingmodes for the reaction to proceed. Third, since it is unlikely that Ar evaporation is mode-specific,
the close agreement of the relative intensities in the purely reactive (n = 1) and purely evaporative(n = 3) channels again likely indicates that there is no strong mode-specific behavior in the electrontransfer reaction. This observation further suggests that the electron transfer reaction is quite slow
compared to anharmonic coupling processes. Lastly, the CH bond lengths are calculated to change
by only 0.02 Å throughout the entire reaction. These calculations are for geometries corresponding
to the data points in Figure 5.7 where a minimum C-H bond length of 1.07 Å is observed at thetransition state (2.6 Å C-I distance in Figure 5.7). This calculation suggests that the CH stretchingmodes are barely participatory in the electron transfer reaction, consistent with the absence of mode
specificity for the transitions probed.

5.4

Conclusion
In this study we looked at the binary complex of a nitromethane anion and a methyl iodide

molecule as a model system for vibrationally mediated electron transfer. The electron transfer
reaction can be triggered by excitation of any of the CH stretching vibrational modes in the complex,
resulting in the formation of I product ions. The reaction was found to not exhibit mode-specificbehavior. The electron transfer reaction can be partially deactivated by solvation of two Ar atoms
and completely quenched by three Ar atoms. The barrier of the forward reaction is estimated to
be on the order of 230 meV based on the partial quenching behavior of the reaction with two Ar
atoms.

Chapter 6
Infrared Spectra of Naphthalene Water Clusters

This Chapter has been reprinted with permission from Knurr, B. J., Adams, C. L., Weber,J. M., “Structures and Charge Distributions of Cluster Anions Studied by Infrared Photodissociation Spectroscopy.” Journal of Chemical Physics, 2012. 137: 104303. DOI: 10.1063/1.4750371.
Copyright 2014, AIP Publishing LLC.

6.1

Introduction
The solvation of ions is one of the most fundamental aspects of solution chemistry. In particu-

lar, the interaction of water molecules with negative charge has attracted a lot of attention. [158–165]
While hydrated cluster anions are not direct analogs to bulk solutions, they do o↵er a unique opportunity to study ion solvation on a molecular level. Experiments on mass-selected cluster ionsallow for precise control of ion composition and the number of water molecules in the microsolvation
environment. Studies on microsolvated ions can be used to aid in modeling of solvation e↵ects inaddition to lending increased understanding to solution chemistry.
The interactions of water molecules within a cluster containing negatively charged species
depends on the density and the shape of the charge distribution, among other factors. This can beparticularly interesting when considering a di↵use charge distribution. In some cases, a molecularcharge carrier is unstable but becomes stabilized through solvation. The most extreme case of adi↵use charge distribution is that of the solvated electron. The solvated electron is a case where
there is no suitable molecular orbital to accommodate the electron but it can be stabilized through113
interactions with a solvent species, e.g. by species that possess a dipole of sufficient magnitude (µ2.5 D for a rotating dipole [166]). In the case of neat water, an excess electron can be solvatedby two or more water molecules. [158]A second interesting set of di↵use charge carriers are aromatic molecules. Here, the excess
electron can be confined to the aromatic ⇡-system where it will take the general shape of the ⇡system molecular orbital. The adiabatic electron affinities (AEA) of a number of small aromatic
systems have been measured and the general trend is that as the ⇡-system grows, the AEA also
increases. The AEA of benzene (Bz) has been measured to be -1.1 eV [167] while the AEA of
naphthalene (Np) is only -0.2 eV [168–170] and in the case of phenanthrene is thought to be -0.01
eV. [171] Anthracene appears to be the smallest (unsubstituted) aromatic hydrocarbon that canform a stable isolated anion in the gas phase (AEA = +0.53). [172]Experiments by Mikami and co-workers have shown that for a sufficiently large cluster of theform [Bz·(H2 O)n ] , the anion is stable (n6). [159] However, it was unclear whether the excess

charge was localized in the water network or if the ⇡-system of the Bz molecule was involved incontaining the excess electron, even though the data suggested that the valence orbitals of the Bz
were involved.
Naphthalene anion (Np ) has been found to be stable in a number of di↵erent organic solvents
as early as the 1950’s. [173–175] From one perspective, the interaction of water with Np is relevant
for electrochemical applications of Np and other aromatic species. In many of these applications,the presence of water would be an impurity to the reaction and could result in protonation of thearomatic species thus inhibiting the performance of the reaction. [176] This is particularly importantin regards to polycyclicaromatic hydrocarbons (PAH) becoming more common in organic electronics
where exposure to water can be quite damaging to a device. [177–179] From a second perspective,
Np can be viewed as a small model system for graphene or larger PAHs. Graphene and graphite
systems are considered to be hydrophobic. However, the presence of a negative charge can change
the nature of the interaction with water and water networks. In this sense Np can be used as amodel system for studying the interaction of water with a charged carbon-based nanomaterial.

114
Photoelectron spectroscopy studies have shown that one water molecule is sufficient to stabilize an excess electron on the aromatic system of Np. [168–170] Unfortunately, little structural
information could be obtained from these photoelectron studies. Infrared studies have been used togain structural information on isolated gas phase clusters for years. [180–184] The hydrated Np has
been investigated before by Mikami and co-workers where they studied vibrational autodetachment(VAD) after infrared absorption by [Np·(H2 O)2 ] . Based on the infrared spectra, it was determinedthat all of the OH groups were involved in ionic ⇡-H-bonds to the aromatic system of the Np or inwater-water H-bonds. However, no structure was assigned that gave rise to the observed features.
Additionally, no spectra existed on the monohydrate or on higher levels of hydration. One would
expect either solvation of the entire ⇡-system until no more water can be accommodated or theformation of water-water networks that reside on the surface of the Np anion with a competitionof the two processes potentially being at play.
In this chapter, the spectra of [Np·(H2 O)n ] ·Arm (n = 1 6) will be presented. Using density
functional theory, the structural motifs of the water networks and the involvement of the ⇡-systemof the Np will be discussed.

6.2

Computational Methods
Density functional calculations were used to identify structural motifs for [Np·(H2 O)n ] clus-

ters. The B3-LYP functional [59,60] with dispersion correction [89] as implemented in the TURBOMOLE v. 5.9.1 suite of programs was used for all calculations. [58] The def2-TZVP basis sets wereused for all atoms. [64] Harmonic frequencies were calculated using the AOFORCE program. [66,67]
To account for anharmonicity, the calculated harmonic frequencies were scaled by 0.96091. Thisscaling factor was obtained by comparing the calculated and experimentally measured OH stretching frequencies for a free H2 O molecule. For this study, calculated CH stretching frequencies are notreported as the CH stretching region of Np contains a plethora of Fermi resonances. [185–187] It is
likely that the spectra of the anion also exhibit these resonances making harmonic approximation
calculations generally meaningless.

115
It is worth noting that for all identified minimum energy structures, vibrational modes withvery low frequencies ( 50 cm 1 ) were found. These modes can be attributed to the entire water
network sliding, rocking or tilting on the surface of the Np ⇡-system. This indicates that thepotential energy surfaces of these clusters are quite flat and that low energy vibrational modes will
be very anharmonic. As the number of water molecules increases, there are an increasing numberof possible configurations of the water molecules. Additional structures and their relative energies
and predicted spectra for all cluster sizes can be found in Appendix C.

6.3

Results and Discussion

6.3.1

OH Stretching Vibrations

6.3.1.1[Np·(H2 O)]

Figure 6.1 shows the spectrum of [Np·(H2 O)] ·Ar8 acquired by monitoring the loss of 6 Ar
atoms. The region from 2900 cm 1 to 3100 cm 1 contains the CH stretching signatures of the Np
molecule. The broad feature at 3540 cm 1 can be attributed to the water molecule. Interestingly,
no feature is observed at 3700 cm 1 where the frequency of a free OH stretch would be expected.
This observation indicates that both OH groups are interacting with the Np

and that the one

observed OH band is due to the symmetric water stretch. Features due to the antisymmetric water
stretch (expected at higher energy) are predicted to be low in intensity and are likely below thedetection limit of this experiment.
Calculated structures belonging to minima on the potential energy surface and their calculated infrared spectra are shown in Figure 6.1. Conformer 1-1 recovers the experimental spectrum
the best. This structure is also consistent with the predictions of Schiedt et al. [168] and Kawamata
et al. [188] In conformer 1-1, the water molecule is bridging the two carbon rings and binding theexcess electron in the ⇡-system of the Np molecule. Conformer 1-2 has the water molecule aligned
roughly along the length of the Np molecule where the water molecule is only making one ionichydrogen bond with the Np molecule. This results in a free OH group that is only weakly tetheredintensity [arb. units]

116

1-1

1-2

1-3
3000

3200

3400

3600

3800

-1

wavenumber [cm ]Figure 6.1: Experimental spectrum and calculated spectra of [Np·(H2 O)] . Top trace is the experimental spectrum of [Np·(H2 O)] ·Ar8 acquired by monitoring the loss of 6 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The dashed lines are meant
to show ionic ⇡-H-bonds. See text for discussion.

117
to the ⇡-system. The signature of the free OH around 3700 cm 1 is not observed in the experimental spectrum. Conformer 1-3 is similar to conformer 1-1 in that it bridges two carbon atoms but
rather than bridging the two carbon rings, conformer 1-3 bridges one carbon ring perpendicular tothe length of the Np molecule. All of these calculated structures are within 25 meV of each otherand we view them as isoenergetic.
Comparing the predicted harmonic frequencies of the calculated conformers shows that all
conformers generally recover the experimental spectrum. However, as was mentioned above conformer 1-2 exhibits a feature of a free OH group which is not present in the experimental spectrum.
Since the signatures due to the antisymmetric stretch are not recovered in the experimental spectrum we cannot unambiguously assign the experimental spectrum to one conformer. Additionally,
it is likely that there are multiple isomers populated giving rise to the broad nature of the observed
feature at 3540 cm 1 . Alternatively, conformers with the same position of the water moleculewith di↵erent positions of the Ar atoms have been shown to lead to broadening of the observed
vibrational features. It is most likely that there are multiple e↵ects at play that are causing thebroadening of the observed OH stretching feature. Regardless, we can definitively say that thewater molecule is bound to the ⇡-system with both OH groups and likely has a number of di↵erent
positions it can bind to on the surface of the Np .
Aside from the structure of the cluster, the monohydrate o↵ers an opportunity to consider
the stabilization of the excess electron on the Np molecule. It has been firmly established that Np
exists in organic solvents [173, 189] but there has been no observation of Np in aqueous solution.
Neither free Np or water bind an excess electron while water dimer anions can exist as dipole-boundstates. However, since the AEA of Np is much lower than that of water combined with the evidence
from photoelectron spectra by Bowen and co-workers, [170] it is very likely that the excess electron
is localized primarily on the Np molecule.
If the electron binding energy is lower than the transition energies being studied, then excitation can result in VAD. VAD has been observed in a number of systems including [Np·(H2 O)2 ] .
[146–148, 188, 190] To observe VAD, the following condition must be met

118

h⌫

E(Ar) > AEA,

(6.1)

where AEA is the adiabatic electron affinity of the species and E(Ar) is the binding energy of
the Ar atoms evaporated after photon absorption. By describing the system as an evaporative
ensemble, [8] one can approximate

E(Ar) ⇡ m · Eavg ,(6.2)

where Eavg is the average binding energy of one Ar atom.
Figure 6.2 shows a comparison of the action spectra taken for [Np·(H2 O)] ·Arm (m = 4, 5, 8).
Note that the OH stretch is strongly suppressed for m = 4 indicating that the energy remaining after
evaporation of four Ar atoms is enough to undergo VAD. Stated another way, the remaining energy
is greater than the AEA of [Np·(H2 O)] . However, the CH stretches are still observed indicatingthat h⌫ = 3055 cm 1 is insufficient to observe VAD on the µs time scale of the experiment.
Using the condition for VAD given in Equation 6.1 and an AEA of 0.11 eV, [168, 170] the averagebinding energy of an Ar atom can be bracketed between 540 cm 1 and 660 cm 1 by comparing thespectra for m = 4 and 5. This range of calculated binding energies is consistent with a number ofother studies that predict the binding energy of an Ar atom to be on the order of a few hundred
cm 1 . [145, 146] It is also worth noting that the peak position of the OH stretch is not sensitive tothe number of Ar atoms.

6.3.1.2[Np·(H2 O)2 ]The infrared spectrum of [Np·(H2 O)2 ]has been reported previously by Kawamata et al.

[188] In their studies, VAD was used rather than Ar predissociation to acquire the spectra. They
found multiple broad features that were attributed to a combination of higher energy water-waterinteractions and lower energy ionic ⇡-H-bonds.
Figure 6.3 shows the infrared spectrum of [Np·(H2 O)2 ] ·Ar2 monitoring the loss of both Ar

photofragment yield [arb. units]

119

3000

3200

3400

3600

3800

-1

wavenumber [cm ]Figure 6.2: Observation of vibrational autodetachment in [Np·(H2 O)] ·Arm clusters. Top trace:
m = 4, monitoring the loss of all 4 Ar atoms. Center trace: m = 5, monitoring the loss of all 5 Ar
atoms. Bottom trace: m = 8, monitoring the loss of 5 Ar atoms. Note that the loss of 6 Ar atoms(shown in Figure 6.1) yields the same spectrum as the loss of 5 Ar atoms.

120
atoms. The OH stretching region exhibits four well resolved and intense peaks, two around 3600
cm 1 and two around 3515 cm 1 . Again, there is no signature around 3700 cm 1 indicative of a freeOH group. These spectra are consistent with those taken by Kawamata et al. [188] however, thespectra presented here exhibit sharper vibrational features allowing better structural assignment.
This can be attributed to the lower cluster temperatures necessary to achieve Ar solvation.
Selected calculated conformers for [Np·(H2 O)2 ]are shown in Figure 6.3 beneath the ex-perimental spectrum. In the lowest energy conformer (2-1), one water molecule forms two ionic⇡-H-bonds while the other forms one water-water H-bond and one ⇡-ionic H-bond and the watersubcluster is oriented roughly parallel to the long axis of the Np molecule. This structure is reminiscent of the (H2 O)2 anion where the excess electron is in a dipole bound state localized on one
side of the two water molecules. Here, the water dimer is binding into the ⇡-system of the Np
molecule and is acting as a stabilizing presence for the excess electron localized in the ⇡-system.
Conformer 2-2 is isoenergetic with conformer 2-1 but structurally spans only one of the carbon
rings rather than both. Additionally, one of the OH groups in conformer 2-2 is nearly free.
Conformers 2-3 and 2-4 are typical members of the family of structures where there are nowater-water hydrogen bonds formed. Multiple structures were found in this family with both watermolecules on the same side of the Np (2-3) or on di↵erent sides (2-4). This family of structures
is rather high in energy compared to those where there are water-water H-bonds (⇠100 meVhigher). This energetic di↵erence indicates that there is a substantial energetic penalty for notforming water-water interactions and that water-water H-bonds are more energetically stabilizing
compared to ionic ⇡-H-bonds. The structure proposed by Kawamata et al. [188] was not found to
be a stable structure in our calculations.
Comparison of the calculated spectra to the experimental spectrum reveals that only conformer 2-1 accurately recovers the experimental spectrum. Even though conformer 2-2 exhibits anearly identical water subunit, the nature of its interaction with the Np results in a very di↵erent
infrared spectrum specifically in the prediction of a nearly free OH group. Based on our structural
assignment to conformer 2-1, the vibrational modes can also be assigned. The lowest frequency

121

intensity [arb. units]

2-1
2-2
2-3
2-4
3000

3200

3400

3600

3800

wavenumber [cm-1]
Figure 6.3: Experimental spectrum and calculated spectra of [Np·(H2 O)2 ] . Top trace is theexperimental spectrum of [Np·(H2 O)2 ] ·Ar2 acquired by monitoring the loss of 2 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The dashed lines show thehydrogen bonds. See text for discussion.

122
peak corresponds to the symmetric stretching modes of the two water molecules coupled in phase
while the next higher mode corresponds to their out-of-phase combination. These assignments arereminiscent of those of the water hexamer anion where the observed lowest frequency mode has been
assigned to the symmetric stretch of a H-bond acceptor. [164, 165] The next higher (and weakest)
OH stretch is assigned to the antisymmetric stretch of the water with two hydrogen bonds to the ⇡system. The highest energy observed stretching frequency is assigned to the antisymmetric stretchof the H-bond donor water molecule. The weak shoulder feature around 3620 cm 1 is unassigned
but could be due to an isomer where one water molecule is in a slightly di↵erent conformation withrespect to the Npanion. Additionally, it should be noted that the small feature around 3260

cm 1 is likely due to the bend overtone of the water dimer. Since the symmetric stretching modes
are lower in the dihydrate cluster compared to the monohydrate, there can be improved coupling
to the bending overtones. This feature is absent in the monohydrate cluster.

6.3.1.3[Np·(H2 O)3 ]Figure 6.4 show the experimental spectrum of [Np·(H2 O)3 ] ·Ar8 . The broad feature centeredat 3270 cm 1 can be attributed to the bend overtones of water molecules. The OH stretching regionis quite complicated and exhibits far more bands than there are OH groups in the cluster. This
indicates that there are at least two conformers populated to give rise to the number of observed
peaks. The most stable calculated conformer is a homodromic ring (3-1) where the water molecules
form a three membered ring with each water molecule donating and accepting a proton in the formof water-water H-bonds while also engaging in an ionic ⇡-H-bond. A bent chain structure (3-2)
is predicted 67 meV higher in energy. Here two of the water molecules each form a water-water
H-bond and an ionic ⇡-H-bond while the third water molecule has two ionic ⇡-H-bonds. A second
chain conformer (3-3) is predicted to lie 247 meV higher in energy than the ring structure. Bothof the chain conformers are analogous to the water trimer anion but here, the dangling OH groupon the end can bind to the ⇡-system resulting in no free OH groups.
Other minimum energy structures were found that possessed a water dimer and a lone water

123

intensity [arb. units]

3-1 + 3-2

3-1
3-2
3-3
3000

3200

3400

3600

3800

wavenumber [cm-1]
Figure 6.4: Experimental spectrum and calculated spectra of [Np·(H2 O)3 ] . Top trace is theexperimental spectrum of [Np·(H2 O)3 ] ·Ar8 acquired by monitoring the loss of 5 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The second trace from the
top is a combination of conformers 3-1 and 3-2. The dashed lines show the hydrogen bonds. See
text for discussion.

124
as well as three lone water molecules. These structures were all found to be very high in energy
relative to the ring and chain structures and to not recover the experimental spectrum. These
calculations and the energetic di↵erence between the ring (3-1) and the chain structures (3-2 and3-3) further reinforce that it is energetically favorable to form water-water H-bonds rather thanionic ⇡-H-bonds as was seen in the dihydrate case as well. Additionally, this pattern of bonding is
consistent with the observation that the simulated normal modes of the lower frequency modes tend
to correspond with ⇡-H-bonds while the higher energy modes belong to water-water interactions.
This behavior has also been observed before in Cl2 ·(H2 O)3 clusters. [161]The complex nature of the experimental spectrum make a definitive assignment of the structures contributing to the spectrum difficult. However, as was mentioned above, it is clear that there
are at least two conformers populated. Figure 6.4 shows a comparison of the two lowest energy
isomers (3-1 and 3-2) which is consistent with the experimental spectrum. The higher energy andrelatively simple nature of the spectrum of conformer 3-3 suggests that it does not contribute to
the experimental spectrum. Because of the use of Ar solvation, it is likely that multiple isomers
can be “frozen out”. [144, 191–194] Alternatively, the large number of bands could be caused by
combination bands of lower lying vibrational modes. However, the sharp nature of the experimental
features and the large intensity relative to the other peaks suggests that a combination of di↵erent
isomers is more likely. Experiments employing IR-IR hole burning would be useful in determining
the identity of the contributing structures. [195]6.3.1.4[Np·(H2 O)4 ]Figure 6.5 shows the experimental spectrum of [Np·(H2 O)4 ] ·Ar2 monitoring the loss of both
Ar atoms. The experimental spectrum exhibits a strong set of bands between 3450 cm 1 and
3600 cm 1 . Based on the discussion of [Np·(H2 O)3 ] , these features can be attributed to the⇡-H-bonding OH groups in the water network. At lower frequencies, a broad unresolved peak isobserved which can be assigned to the OH groups involved in water-water H-bonds. The pattern ofthe vibrational frequencies observed here is very similar to that observed in I ·(H2 O)4 , [162] which

125
suggests that the same structural motif is present in both cases. Johnson and co-workers [162]tentatively assigned the structure of I ·(H2 O)4 to a homodromic ring, similar to what was seen
previously in other halide water cluster ions [163] in addition to [Np·(H2 O)3 ] .
Of the calculated structures, the lowest energy structure (4-1) is a homodromic ring (seeFigure 6.5). The next higher energy structure (4-2) is the homodromic ring from [Np·(H2 O)3 ]with the fourth water molecule H-bonding to the ring and the ⇡-system of the Np . This conformer
is predicted to be 201 meV higher in energy from conformer 4-1. In other higher energy conformers,
the water network is fragmented (see 4-3 and 4-4 at 241 meV and 351 meV higher in energy than 4-1,
respectively). Only conformers 4-1 and 4-2 are roughly consistent with the experimental spectrum
where the spectrum of 4-1 is too simple while the spectrum of 4-2 exhibits too many peaks. Whileboth structures recover a portion of the experimental spectrum, a definitive spectral assignment
is not possible without additional high level anharmonic calculations. It is again possible that
there are multiple isomers present for the same reasons as are stated above (evaporative coolingand barrier to formation of other structures). Additionally, the presence of higher energy isomers
has precedence in similar experiments. [144, 191–194] The spectrum of [Np·(H2 O)4 ] is tentatively
assigned to conformer 4-1 with contributions from 4-2 and potentially other isomers.

6.3.1.5[Np·(H2 O)5 ]and [Np·(H2 O)6 ]The infrared spectrum of [Np·(H2 O)5 ](see Figure 6.6) is similar to that of [Np·(H2 O)4 ]possessing broad features below 3450 cm 1 and between 3450 cm 1 and 3600 cm 1 . It is likely thatthe broad features observed here are the result of multiple unresolved peaks of similar frequency.
Again, the lower energy features are due to water-water interactions while the higher energy features
are due to ⇡-H-bonds. A number of di↵erent structures were calculated and, like in smaller clusters,
the lowest energy conformer is a homodromic ring (see 5-1 in Figure 6.6). The calculated spectrumof conformer 5-1 is qualitatively consistent with the experimental spectrum however, similar toboth the tri- and tetrahydrates, the water-water interactions are overestimated and the frequenciesof the corresponding signatures are predicted to be too low. The other conformer shown (5-2) is126

intensity [arb. units]

4-1
4-2
4-3
4-4
3000

3200

3400

3600

3800

wavenumber [cm-1]
Figure 6.5: Experimental spectrum and calculated spectra of [Np·(H2 O)4 ] . Top trace is theexperimental spectrum of [Np·(H2 O)4 ] ·Ar2 acquired by monitoring the loss of 2 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The dashed lines show thehydrogen bonds. See text for discussion.

intensity [arb. units]

127

5-1

5-2

3000

3200

3400

3600

3800

wavenumber [cm-1]
Figure 6.6: Experimental spectrum and calculated spectra of [Np·(H2 O)5 ] . Top trace is theexperimental spectrum of [Np·(H2 O)5 ] ·Ar2 acquired by monitoring the loss of 2 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The dashed lines show thehydrogen bonds. See text for discussion.

also roughly consistent with the experimental spectrum and is calculated to be 184 meV above
conformer 5-1. This structure is characteristic of the “cage” family of structures where the water
network begins to favor water-water H-bonds over ionic ⇡-H-bonds. Conformer 5-2 also bears
resemblance to the water pentamer anion where the Np has taken the role of the di↵use electron.
However, neither 5-1 or 5-2 can account for the small features observed above 3600 cm 1 .
The spectrum for [Np·(H2 O)6 ](see Figure 6.7) is similar to that of [Np·(H2 O)5 ] . There128
are again intense broad peaks around 3550 cm 1 , water-water features trailing down to 3150 cm 1and some weak features above 3600 cm 1 . The water hexamer has been studied extensively andhas been shown to exhibit a number of di↵erent structures. [164, 165] While it is unlikely that adefinitive structural assignment can be made, a comparison with a number of di↵erent structures
is warranted.
Conformer 6-1 is similar to the “book” structure that has been observed before in both neutral
and anionic water clusters and recovers the general envelope of the experimental spectrum. [164,165]
Conformation 6-2 is another variation on the “book” structure that is in the vein of a four membered
homodromic ring with a water dimer attached on the side and is isoenergetic with conformer 6-1.
Conformer 6-2 is more consistent with the features above 3600 cm 1 but overestimates the lower
energy features. Conformer 6-3 is a cage structure where two three-membered rings are stacked on
top of each other and is calculated to be only 62 meV above 6-1. The spectrum for conformer 6-3
recovers the general shape of the experimental spectrum but does not predict the features above
or below the bulk of the spectrum from 3350 cm 1 to 3600 cm 1 . The final reported structureis the six membered homodromic ring (6-4) which is calculated to be 68 meV above 6-1. The
calculated spectrum for 6-4 does not recover the experimental spectrum well but it is possible thatthe low energy ring modes could be broadened by ring deformation modes. In summary, noneof the calculated spectra recover the experimental spectrum individually but a combination may.
Higher level anharmonic calculations would be necessary to truly assign the experimental spectrumof [Np·(H2 O)6 ] .
For both the penta- and hexahydrate of Np the experimental spectrum is too congested toallow for a definitive structural assignment. This is likely due to the presence of multiple conformers
and the coupling of various frequencies resulting in broad unresolved features. However, a uniqueinsight for these larger clusters is that it appears as though the water network may be partiallydetached from the Np⇡-system. In particular, the hexahydrate is the first cluster where thelowest energy structure is not the homodromic ring but rather two di↵erent “book” conformations.
This observation again shows that the water-water interaction appears to be favorable compared129

intensity [arb. units]

6-1

6-2

6-3

6-4

3000

3200

3400

3600

3800

wavenumber [cm-1]
Figure 6.7: Experimental spectrum and calculated spectra of [Np·(H2 O)6 ] . Top trace is theexperimental spectrum of [Np·(H2 O)6 ] ·Ar2 acquired by monitoring the loss of 2 Ar atoms. Bottom
traces are calculated conformers and their predicted infrared spectra. The dashed lines show thehydrogen bonds. See text for discussion.

130
to ionic ⇡-H-bonds.

6.3.2

CH Stretching Vibrations

The CH stretches of Np are severely a↵ected by Fermi resonances. [185–187] Similar e↵ects
are expected here making harmonic calculations in this region useless. It is therefore not worth
presenting an depth analysis of calculations in this region, however some useful qualitative insightcan be gained by comparisons of calculated frequencies and trends from one cluster size to another.
It is worth noting that the CH stretching region of [Np·(H2 O)] does not show significant variation
with di↵erent structural orientation of the water molecule indicating that the CH stretches are notstrongly a↵ected by position of the water network and consequently the CH stretches are not a
useful aid in structural assignment (all structures showed an invariance within 2 cm 1 ).
The overall structure of the CH stretching region does not change appreciably with increasinglevel of hydration. However, the CH stretching frequencies shift to the blue by 2 - 10 cm 1 foreach additional water molecule. This can be rationalized by stabilization of the excess electron
into the ⇡-system. At low levels of hydration, the excess electron can be partially localized in *orbitals of C-H bonds resulting in a weakening of the C-H bonds and subsequent red shift of the CH
stretching frequencies. As the level of hydration increases, the excess charge is more concentrated
in the ⇡-system and even the water network itself. This strengthens the C-H bonds resulting in theobserved blue shift of the CH stretching region with increasing solvation. Similar behavior has been
observed before in CH3 NO2 ·(H2 O)n cluster anions. [192, 196] One remaining question is whether
the excess electron will remain localized in the ⇡-system in the limit of bulk solution or if there
is a critical water network size where the excess electron will instead reside in the water network
exclusively. Np has been observed in bulk acetonitrile solution [189] where a solvated electron has
also been detected [197] which lends credence to the idea that the excess electron can stay boundto the ⇡-system even in the limit of bulk solution. However, thus far there has been no detectionof Np in bulk water.

131

6.4

Conclusion
The infrared Ar predissociation spectra of [Np·(H2 O)n ] ·Arm (n = 16) were presented

in this chapter. Analysis of the spectra indicated that the excess electron is mostly localized inthe ⇡-system of the Np molecule with some spill-out into the* CH orbitals. This results in aweakening of the CH bonds of the Np molecule. However, due to the low AEA of Np, the excess
charge is also partially accommodated in the water network. This e↵ect increases with increasing
hydration. The excess charge is stabilized in the ⇡-system through ionic ⇡-H-bonds made by thewater molecules.
The water molecules form networks on the surface of the Np . As hydration increases,
the more favorable water-water interactions dictate polycyclic and cage like structures between
the water molecules. Many of the structures predicated have been observed before in a numberof di↵erent hydrated anion or hydrated electron clusters. It is clear that when three or morewater molecules are present, multiple structures are being populated leading to massive spectralcongestion and an inability to definitively assign structures to the spectra while still allowing for
qualitative understanding of the water network behavior.

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Appendix A
Supplementary Material for Chapter 3.2.2

This appendix contains additional calculated structures for [Au(CO2 )n ] clusters (n = 4, 5).
The number of conformers that can be calculated at larger cluster sizes and the flatness of thepotential energy surface makes identifying larger clusters computationally prohibitive.

148

intensity [arb. units]

0 meV

14 meV

16 meV

48 meV

88 meV

1800

1900

2000
-1

wavenumber [cm ]Figure A.1: Calculated infrared spectra in the region of the partially reduced antisymmetric CO
stretch of the formate core for [Au(CO2 )4 ] clusters. The experimental spectrum for [Au(CO2 )4 ]is shown in the top panel. The calculated spectra are shown in the lower panels with their corresponding structures and relative energies given to the right. Calculated OCO bond angles range
from 139.3 to 142.6 for the bonded CO2 moiety and 175.2 to 178.0 for the solvent CO2 molecules.

149

intensity [arb. units]

0 meV

51 meV

55 meV

69 meV

91 meV
1800

1900

2000
-1

wavenumber [cm ]Figure A.2: Calculated infrared spectra in the region of the partially reduced antisymmetric CO
stretch of the formate core for [Au(CO2 )5 ] clusters. The experimental spectrum for [Au(CO2 )5 ]is shown in the top panel. The calculated spectra are shown in the lower panels with their corresponding structures and relative energies given to the right. Calculated OCO bond angles range
from 139.0 to 140.2 for the bonded CO2 moiety and 175.3 to 178.0 for the solvent CO2 molecules.

Appendix B
Supplementary Material for Chapter 3.3.3

This appendix contains additional calculated structures for [Ni(CO2 )n ]clusters and an

estimation for the barrier height for conversion of isomer E* to isomer D. See Chapter 3.3.3 fordiscussion.

151

intensity [arb. units]

D

F

E

E*

1000

1200

1400

1600

1800

2000

wavenumber [cm-1]
Figure B.1: Calculated spectra of bare core isomers for [Ni(CO2 )2 ]presented here are discussed in Chapter 3.3.3.

clusters. All of the isomers

152

intensity [arb. units]

B

G

H

I

1000

1200

1400

1600

1800

2000

wavenumber [cm-1]
Figure B.2: Additional calculated spectra of bare core isomers for [Ni(CO2 )n ] clusters. Isomers
B and G are discussed in Chapter 3.3.3. Isomers H and I were determined not to contribute to theexperimental spectra and were not included in the discussion.

153

3000

relative energy [meV]

2500

2000

1500

1000

500

0
20

40

60

80

100

120

140

160

180

ONiC bond angle [°]Figure B.3: Estimated barrier height for the conversion of isomer E* to D. Due to the barrierheight, we do not expect this isomer conversion process to occur in our experiment.

Appendix C
Supplementary Material for Chapter 6

This appendix contains additional structures calculated for [Np·(H2 O)n ] clusters. The structure is given on the left with the spectrum on the right and the relative energy for that cluster size
below. See Chapter 6 for discussion.

C.1[C10 H8 ·H2 O]

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 0 meV

3600

3700

3800

155

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

-1

Photon Energy (cm )Relative Energy: 4 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 18 meV

3600

3700

3800

156

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 25 meV

3600

3700

3800

157

C.2

[C10 H8 ·(H2 O)2 ]Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 0 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 9 meV

158Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 92 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 95 meV

159

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 98 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 102 meV

160

C.3

[C10 H8 ·(H2 O)3 ]Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 0 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 67 meV

3600

3700

3800

161
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 218 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 248 meV

3600

3700

3800

162
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 274 meV

3600

3700

3800

163

C.4

[C10 H8 ·(H2 O)4 ]Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 0 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 201 meV

164Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 241 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 270 meV

3600

3700

3800

165
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3700

3800

Photon Energy (cm-1)

Relative Energy: 351 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 370 meV

3600

3800

166
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 433 meV

3600

3700

3800

167

C.5

[C10 H8 ·(H2 O)5 ]Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

Photon Energy (cm-1)

Relative Energy: 0 meV

10

8

2

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 25 meV

3600

3700

3800

3800

168

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3700

3800

Photon Energy (cm-1)

Relative Energy: 81 meV

10

8

2

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 102 meV

3600

169

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 163 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 184 meV

3600

3700

3800

170

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 208 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 225 meV

3800

171

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 243 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 356 meV

3600

3700

3800

172

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 371 meV

3600

3700

3800

173

C.6

[C10 H8 ·(H2 O)6 ]Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 0 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 19 meV

174
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 52 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 57 meV

175
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 62 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 66 meV

176

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 68 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 89 meV

3600

3700

3800

177
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 92 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

Photon Energy (cm-1)

Relative Energy: 95 meV

3600

3700

3800

178
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 111 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

Relative Energy: 113 meV

3000

3100

3200

3300

3400

3500

3600

3700

3800

179

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 161 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 190 meV180
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500

3600

3700

3800

Photon Energy (cm-1)

Relative Energy: 220 meV

Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 325 meV

3600

3700

3800

181
Normalized Intensity (arb. units)

1.0

0.8

0.6

0.4

0.2

0.0
2900

3000

3100

3200

3300

3400

3500
-1

Photon Energy (cm )Relative Energy: 459 meV

3600

3700

3800