Tunable Luminescent Boron Complexes

Home , Double bond rule, Lithium borohydride

A Thesis Submitted to the faculty of the Graduate School of the University of Minnesota By

Christian Lawrence Toonstra

In partial fulfillment of the requirements for the degree of Master of Science

Dr. Paul Kiprof, Advisor

July, 2011

Acknowledgements

I am indebted to Dr. Paul Kiprof for his knowledge and support over the last two years. I am also grateful to Dr. Steven Berry and Melanie Halverson for their help in X-ray crystallography, as well as Dr. Ahmed Heikal for his assistance in measuring quantum yield. I also want to thank Dr. Alan Oyler for his help in obtaining the APCI-MS data. Finally, I greatly appreciate the support from the department, especially Randall Helander.

Abstract

The use of BF2 adducts has a long and important history in the field of auxochromic dyes, most notably in the sundry applications of the BODIPY family of dyes. Previous work in our laboratory on phenyl borinic acid yielded good emissive properties including quantum yield. BF2 was chosen due to the potential increase in the luminescence intensity of the adducts as compared to phenyl borinic acid derivatives. Simple azole based ligands were chosen due to their flexibility in color tuning. The azole N,O-type motif was found to be amenable to formation of adducts with

BF2 as the oxygen readily formed a relatively stable boron ester type bond, while the non-bonding lone pair of electrons on the nitrogen is donated to the boron. Color tuning in this family of ligands is attainable through extension of the π system, or through auxochromic heteroatoms. The characterization of products included NMR, LC-APCI-MS, and X-ray crystallography. The luminescence data was collected using fluorimetry. The emissive nature of the complexes was probed using computational techniques. The TD-DFT data obtained from these computational studies was compared to the absorbance data that was obtained. The current findings as well as short-term future plans will be presented.

Overview

Beyond the appeal of advancements in display clarity and resolution, the development of organic light emitting diodes (OLEDs) represents a further step toward progression in the development of highly-tuned luminescent devices that are likewise energy efficient. This paper presents new research into a family of boron based OLEDs, having interesting properties that make them potentially useful in the development of OLEDs. The scaffold chosen was a tunable azole-based ligand. This thesis is separated into three chapters covering current, relevant research into boron-based OLEDs., rationalization of the photophysical properties of these new complexes, and experimental information. A brief overview of boron and it’s role in OLEDs with respect to the development of the boron compounds presented here is discussed in chapter one. Characterization was completed through a combination of NMR, X-ray crystallography, HR-APCI-MS, UV-vis, Fluorimetry, and Computational TD-DFT, as means to probe both the physical and optical properties of the compounds, this is discussed in the second chapter. The final chapter, along with the appendix, outlines the experimental details of this project.

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TABLE OF CONTENTS

List of Tables v List of Figures vi List of Abbreviations viii Chapter 1 i Background of Boron Chemistry i Application of boron to OLEDs vi Overview of OLEDs ix New Directions for OLEDs Design xxiii

Chapter 2 xxix New directions for luminescent boron adducts xxix Instrumentation xxxi Optical Data (UV-vis and Fluorimetry) xxxi Nodal Plane Theory xxxvi NMR xl HR-APCI-MS xlii X-ray Crystallography xliii Theoretical Chemistry xlv Chapter 3 lxix Instrumentation lxix Materials lxx Synthesis lxxi References lxxviii Appendix lxxxvi

List of Tables

Scheme 1 Ligand synthesis xxvii

Scheme 2 Adduct synthesis xxx

Table 1 Comparison of emission data xxxiii

Table 2 Summarized optical properties of all of the compounds xxxv

Table 3 Comparison of Stoke’s shift values xxxix

Scheme 3 Possible products of the reaction of phenyl boronic acid and 10- Hydroxybenzo[h]quinoline lxiii Table 4 Raw data for quantum yield calculations cxiix

List of Figures

Fig. 1 Boron fragmentation iii

Fig. 2 Examples of medically useful boron derivatives. iv

Fig. 3 ELF isosurface plots of boron-halogen complexes. v

Fig. 4 Examples of boron-based ETLs and HTLs. iix

Fig. 5 Cross Section of a typical OLED x

Fig. 6 Charge “hop” mechanism x

Fig. 7 Examples of the physical basis of OLED operation xii

Fig. 8 Methods of energy transfer xiv Fig. 9 OLED operating mechanism xvi Fig. 10 Synthetic methods for BODIPY xiix Fig. 11 Examples of 2-pyridyl adducts and BORAZAN xx Fig. 12 Examples of quinolato-type adducts xxi Fig. 13 Examples of three-coordinate organoboron compounds xxii Fig. 14 Examples of benzoboroxole adducts xxv Fig. 15 Emission comparison xxxii Fig. 16 Overlay of the emission spectra of the azole adducts xxxiii Fig. 17 Stoke’s shift comparison xxxvi Fig. 18 Nodal plane theory diagram xxxiix Fig. 19 MS diagram xliii

Fig. 20 Crystal structure of BF2(1,2-HNBT) xliii

Fig. 21 Crystal packing unit cell of BF2(1,2-HNBT) xlv

Fig. 22 Molecular orbital plots of BF2(HPBO) and BF2(HPBT) xlvii

Fig. 23 Molecular orbital plots of BF2(1,2-HNBO) and BF2(1,2-HNBT) xlvii

Fig. 24 Molecular orbital plots of BF2(2,3-HNBO) and BF2(2,3-HNBT) xlviii Fig. 25 Calculated band gap energies liii Fig. 26 VMOdes plots liv Fig. 27 Possible substitution points on the azole and phenyl scaffold lx Fig. 28 Calculated excitation spectrum of EDG substituted azole derivatives lxi

Fig. 29 VMOdes and molecular orbital plot of BF2(HPBI) lxv Fig. 30 Calculated excitation spectrum of anthracene derivatives lxviii Fig. 31-82 Raw NMR data lxxxvi Fig. 83-87 Raw HR-MS data cxii Fig. 88-94 Overlay of experimental excitation and emission data cxv Fig. 93-99 Experimental excitation versus calculated excitation spectrum cxviii

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List of Abbreviations

LET LINEAR ENERGY TRANSFER

BNCT BORON NEUTRON CAPTURE THERAPY

ELF ELECTRON LOCALIZATION FUNCTION

OLEDS ORGANIC LIGHT EMITTING DIODES

LCD LIQUID CRYSTAL DISPLAY

ITO INDIUM TIN OXIDE

EML EMISSIVE MATERIALS LAYER

ETL ELECTRON TRANSPORT LAYER

HTL HOLE TRANSPORT LAYER

EWG ELECTRON WITHDRAWING GROUP

EDG ELECTRON DONATING GROUP

SMOLEDS SMALL MOLECULE ORGANIC LIGHT EMITTING DIODE

PHOLEDS PHOSPHORESCENT ORGANIC LIGHT EMITTING DIODE

LUMO LOWEST UNOCCUPIED MOLECULAR ORBITAL

HOMO HIGHEST OCCUPIED MOLECULAR ORBITAL

BODIPY BORON-DI-PYRROMETHENE

BORAZAN LITERATURE TRADEMARK

ALQ3 ALUMINUM-TRISQUINOLATO DDQ 2,3-DICHLORO-5,6-DICYANO-1,4-BENZOQUINONE

PLEDS POLYMERIC LIGHT EMITTING DIODE

IR INFRARED

DCM DICHLOROMETHANE

DMSO DIMETHYL SULFOXIDE

DMF N,N-DIMETHYL FORMAMIDE

THF TETRAHYDROFURANS

APCI ATMOSPHERIC PRESSURE CHEMICAL IONIZATION

HR-MS HIGH-RESOLUTION MASS SPECTROSCOPY

UV-VIS ULTRA VIOLET-VISIBLE

UV ULTRA VIOLET

TD-DFT TIME DEPENDENT- DENSITY FUNCTIONAL THEORY viii

QY QUANTUM YIELD

ESIPT EXCITED STATE INTRAMOLECULAR PROTON TRANSFER

PET PHOTOINDUCED ELECTRON TRANSFER

NMR NUCLEAR MAGNETIC RESONANCE

PPM PARTS PER MILLION

PCM POLARIZABLE CONTINUUM MODEL

VMODES VIRTUAL MOLECULAR ORBITAL DESCRIPTION PROGRAM

FMO FRONTIER MOLECULAR ORBITAL

DRE DEWAR RESONANCE ENERGY

HBQ BENZOHYDROXY QUINOLINE

HQ HYDROXY QUINOLINE

HPBO HYDROXYPHENYL BENZOXAZOLE

HPBT HYDROXYPHENYL BENZOTHIAZOLE

HPBI HYDROXYPHENYL BENZIMIDAZOLE

HNBO HYDROXYNAPHTHYL BENZOXAZOLE

HNBT HYDROXYNAPHTHYL BENZOTHIAZOLE

HNBI HYDROXYNAPHTHYL BENZIMIDAZOLE

DPA DIPHENYL ANTHRACENE

HABO HYDROXYANTHRYL BENZOXAZOLE

HABT HYDROXYANTHRYL BENZOTHIAZOLE

Chapter 1

Boron Chemistry and its Applications

The sundry uses of Boron have made this metalloid a popular topic of study since its isolation in 1808 by Gay-Lussac and Thenard. It was finally recognized as an element in 1924 by Jöns Jakob Berzelius.1 The history of boron dates back as far as the sixteenth century, then called tinkal in Arabia, where it was used to assist melting processes. One of the interesting features of boron compounds is their lack of an octet in many species as well as their Lewis acidity. The first major contribution to the chemistry of the boron-nitrogen bond came in 1926 when Stock and Pohland reacted diborane with ammonia, forming borazine, known colloquially as the “inorganic benzene.”1 Boron is by no means a major constituent in the earth’s crust, accounting for a mere 3 parts per million. This is astonishingly low given the variety of uses for which boron is requisite, especially the glass industry. When obtained as a crude ore, boron is most often found in the form of tourmaline, which is ~10% boron within aluminosilicate.2 The vast majority of boron deposits are in Turkey, while much of the remainder is in California in the United States. Boron has features that are similar to both of its neighbors, carbon and the metalloids, specifically silicon. The unique feature of boron is that, though it maintains some characteristics with its neighbors, it is chemically distinct. For example, group 13 elements, except boron, are all easily ionizable, allowing their cations to play an important role in their chemistry. The energy of ionization is too high to play a role in boron chemistry. This allows researchers to use boron in many applications both organic and inorganic. Boron is known to exist in at least five allotropes, illustrating the similarity to carbon.3

Part of the unique nature of boron lies in its valency prediction. Monovalent boron would have a ground state with only one unpaired electron in the p-orbital. However, the promotion of the s2p1 ground state to the hybridized s1p2 state is energetically small. The promotion requires much less energy compared to the amount of energy that is released by the formation of three covalent bonds, this is the reason why monovalent boron is not able to be isolated, except at high temperatures. The three singly-occupied orbitals in the sp2 state account for the formation of trigonal planar covalent boron complexes. The remaining empty pz orbital is responsible for the common formation of tetrahedral boron complexes, it is also the root of the function of boron as an electron acceptor. 2, 3 The preceding attributes as well as other idiosyncrasies have made boron an attractive element for an extremely diverse range of materials applications. Of the two isotopes of boron, 10B and 11B, 10B has a high neutron absorption cross section. This ability has led to the use of 10B in control rods in nuclear reactors to control neutron flux. When boron captures a neutron, it releases an isotope of lithium and an alpha particle, referred to as high linear energy transfer particles (LET) FIG 1. The alpha particle is lethal to surrounding cells. This feature of boron has been used in the treatment of cancer. The treatment, called boron neutron capture therapy4-6, is administered by inducing a high concentration of 10B into tumors and exposing the patient to low energy neutrons. When the boron captures the low energy

2 neutrons alpha particles are released and destroy the surrounding cancer cells.

Fig. 1 Release of a gamma particle upon neutron collision with boron.

These characteristics have been exploited as a minimally invasive and low toxicity treatment for a variety of types of cancer. This technique rests upon the fact that the constitutive elements of cells have low neutron absorption cross-sections. This allows the damage to tissue to be restricted to only 10B containing cells. Moreover, the distance of recoiling neutrons is a relatively small area (due to high LET), further limiting damage to healthy tissue. The major setback to this innovative technique remains achieving a functional concentration of 10B inside tumor cells. The effective concentration is 109 10B atoms per cell.7 Again the great versatility of boron has helped to partially overcome some of these hindrances. Borons ability to form allotropes, or boron clusters, allows the delivery of boron in a form that is impressively hydrolytically stable as well as metabolically stable FIG 2. Additionally, boron forms strong, hydrolytically stable bonds with various inert elements; also its size is similar to carbon, allowing facile replacement of boron into carbon structures. Usually a targeting agent is coupled to the boron complex to

3 ensure delivery to the appropriate location. These advances have allowed in situ concentrations that are analogous to other cytotoxic anti-cancer drugs.

Fig. 2 Medically useful forms of boron for BNCT. p-Boronophenylalanine (1); Maltose Derivative (2); Boron allotrope (3).

Boron has unique relationships with a variety of different elements, forming bonds that are unlike any other. Binding with halogens forms an important class of boron compounds. With the exception of iodine, boron-halogen complexes form adducts with Lewis bases, forming a dative bond with the non-bonding electrons of the base. The heavier halogen compounds have less π-interaction with the boron, so that they are more aggressive Lewis acids. The binding of halogens to boron are further influenced by the atomic charges.8 The larger the atomic charge the smaller the atomic valence, therefore, with a smaller atomic valence, ionic bonding becomes more significant. Mercier et al. illustrated this by comparing the bond orders of BF3 and BCl3, BF3 has a significantly lower bond order, owing to its higher atomic charge.11 BCl3 has a lower atomic charge, and ionic bonding is less significant, and 4 covalent bonding becomes more important. The higher ionic character of the bond between boron and fluorine results in a less stable bond that is more easily dissociated. The nature of bonding also affects the degree of hard/soft acid base properties, because the overlap of orbitals is negligible between atoms without significant covalent bond character.9 Therefore, heavier halogens form softer acids with boron, whereas fluorine forms more ionic bonds with boron and therefore a harder Lewis acid. The partial atomic charge on boron becomes increasingly positive as the electronegativity of the substituents is amplified FIG 3.10,11 In the case of boron halides, an increase of Lewis acidity is accompanied by a softer character of the Lewis acid.

The empty pz orbital on boron is often exploited for its Lewis acid properties.11 The Lewis base most often chosen to donate its non-bonding electrons is nitrogen.12,13 Numerous complexes have been designed around these conjugate pairs nearly since the beginning of borons history.

Fig. 3 ELF isosurface plots of boron-halogen complexes. Green region indicates bonding basin, while blue indicates non-bonding, lone pair basin, and red, core basin. Values are the basin volumes in arbitrary units.11 Used by authors permission. 5

Doubtless the most well known application of boron is in the field of organic reagents. The breadth of organic chemistry that involves boron is truly impressive, and is beyond the scope of this thesis to cover in detail, however, some the highlights are worth noting to illustrate the significance of boron in many diverse fields of chemistry. One of the most important uses of boron in organic chemistry was serendipitous, because the original study was an inquiry into the nature of boron-carbon bonding. A side product of this study was the discovery that diborane reacted with aldehydes to give the corresponding alcohols following hydrolysis.14 The importance of this finding was largely ignored due to the difficulty in obtaining pure diborane. Interest in diborane peaked during World War II, due to its uses in volatilizing uranium.15 Many developments in boron-based organic reagents came from the work during this time, most notably, the discovery of both sodium borohydride and lithium borohydride.16 These two compounds have become extremely popular due to their contrasting gentle (former) and powerful (latter) reduction capabilities. Due to the dichotomy in the strength of these two reagents, a small degree of selectivity can be achieved. However, numerous stereoselective boron based organic reagents have subsequently become available. Many asymmetric synthesis reagents have since been designed around boron, covering diverse areas.15 Boron again came into prominence by its role in C-C bond forming reactions.16 Since the initial discovery, a wide variety of boronic acids catalysts have been utilized in Suzuki-Miyauri coupling. 17 Mere changes in both the Pd catalyst and the boronic acids can lead to a diverse range of sp2-sp2 and sp2-sp3 coupling.18, 19

Application of Boron to OLEDs Boron’s significance is not limited to synthetic applications; a diverse array of technology exploits some of boron’s unique advantages. The technology most pertinent to this thesis is the application of boron to organic light emitting materials.

Boron has been a popular candidate in OLEDs because of the vacant pz orbitals, 6 which gives it the useful characteristic of Lewis acidity and the formation of stable covalent dative bonds. This interaction allows π-π* transitions through electronic conjugation of the π-electron rich organic ligand and the empty pz orbital on boron.

This vacant pz orbital forms adducts with Lewis bases, and the luminescent properties are easily tuned by changes in the coordination sphere. Boron is also used in the formation of the hole transport layer and the electron transport layer. In the case of hole-blocking layers, these materials exploit boron’s inherent electron acceptor properties and serve to prevent charge spill FIG 4.

Fig. 4 Boron based electron transporting layer amorphous materials (top). Boron based hole transporting layer amorphous materials (bottom).

Overview of OLEDs versus LCD Current interest in OLEDs stems from the need to improve existing technology. The dominant technology today is liquid crystal display (LCD). Twisted nematic liquid crystals are sandwiched between a vertical and a horizontal polarization plate. The vertical plate is behind the indium tin oxide (ITO) anode, and the horizontal plate is between the color filter and ITO cathode. Behind the ITO anode is a white backlight. Operation of the device occurs through both the introduction of incident light from the backlight and application of voltage from the ITO. The liquid nematic crystals first circularly polarize the incident light, and the voltage drive causes the twisted nematic crystals to straighten. With the straightening of the crystals, vertically polarized light is allowed to strike the horizontal plane filter. The vertically polarized pixels that strike the horizontal filter are not visible.20 As a white backlight is used, variation in color is provided by the color filter. The architecture of LCDs provides ample room for improvement. For instance, the quality of the picture is marred by a series of pitfalls that originate in the nature of LCD design. The resolution of the picture is slowed by the time required to straighten the nematic crystals, and the backlight reduces the energy efficiency. Unpolarized light

8 from the white backlight is able to leak through, causing a reduction in darkening efficiency, along with preventing the ability to present true blacks.21 OLEDs overcome both of these shortcomings because they are emissive and not filtering. Additionally, OLEDs have the capacity to construct both thinner and flexible displays. Both of these features are unable to be applied to LCD technology as the liquid crystals alone require 100 nm thickness.

Overview of OLEDs The history of OLEDs reaches back to the early 1950s, when Beranose and colleagues applied an alternating current to various dyes (acridine orange for example) that were fixed to cellophane.22 The authors made the observation that the luminescent intensity increased with increasing voltage, proposing the emission of light to be due to the formation of metastable excited molecules, that emitted light as they returned to the ground state. This early, incorrect hypothesis was replaced when Pope and co-workers discovered a second, delayed electroluminescence peak. Suggesting that the luminescence was not at all due to molecular relaxation, but rather to recombination of charges.23 That same year saw the construction of the first two OLEDs. One of these, constructed by Helfrich and Schneider, was fabricated with a single layer of anthracene and both hole and electron injecting electrodes.24 Progress in the construction of highly efficient OLEDs was hampered by poor conductivity of the organic materials. Therefore, a significant amount of research was devoted to finding organic materials that would overcome the shortcomings of the current components. The sum of all of this research found that OLEDs must consist of a minimum of three layers, the electron transport layer, the emissive materials layer, and the hole transport layer. Any other layers that may be applied serve only to improve luminescent efficiency FIG 5 and FIG 6. It is requisite that the materials be isoenergetic with one another, as any imbalance will lead to a charge bias. A charge bias will require a larger voltage drive in order to operate, decreasing 9 the efficiency. Therefore the LUMO of the semiconducting ETL should be nearly isoenergetic with the EML, likewise, the HOMO of the HTL should be nearly isoenergetic with the EML.

Fig. 5 Cross section of a typical triple layer OLED.86

Fig. 6 Charge “hop” mechanism 10

The most promising organic materials have come from research into polymers. This research began in the early 1960s with the construction of polypyrrole structures. These polymers were found to have extremely low resistivities.25 The results from this study led to an explosion of research into the electronic usefulness of polymers. Hitherto polymers were seen as insulators only, but the new research found applications in organic electronics such as molecular switches26 and organic semiconductors.27 Polymers, in addition to their regular electronic properties, allow for facile modification. A wide variety of changes to the electronic and luminescent properties can be achieved by addition of various electron withdrawing groups (EWGs) and electron donating groups (EDGs) about the polymer scaffold, either in the ring or the vinylene moiety.28,29 Small molecule OLEDs (SMOLEDs) are also quite popular, with an enormous variety of designs. Phosphorescent OLEDs (PHOLEDs) are merely a derivative of either category, but contain phosphorescent material. This allows the OLEDs to capitalize on transitions that are otherwise forbidden.

Physical Basis of Emission in OLEDs The emission of light in OLEDs is analogous to fluorescence in any other emission, however, in OLEDs, it is due to recombination of charges rather than molecular relaxation. It is precisely the charge recombination in OLEDs that allows their function. In order to understand the concept of charge recombination, it is useful to consider the charge transport layer as the “host”, and the emissive layer as the “guest.” The underlying mechanism of OLEDs is referred to as the “charge hop” mechanism, in which the charges “hop” through the π-systems FIG 7. The hopping is instigated by the injection of electrons from the cathode, and the injection of holes (removal of electrons) from the anode. These pairs are referred to as polarons. The drive voltage pushes the charge carriers along toward each other, where they meet

11 in the emissive layer resulting in the release of photonic light. The charges must meet on a single molecule to form a molecular excited state.

Fig. 7 Physical operation of two types of OLEDs. Triple layer (top) and dual layer (bottom).

One significant difference between photoexcitation versus electroluminescence, is that photoexcitation forms excitons consisting solely of singlets, whereas electroluminescence is spin-independent and random. This leads to four possible spin states, three triplets and a lone singlet. As internal conversions in which ∆S≠0 is forbidden30, the internal efficiency is reduced to a mere 25% because it relies exclusively on the singlet state. The luminescent attenuation forms one of the fundamental problems with both small molecule OLEDs (SMOLEDs) and polymer OLEDs. To overcome this limitation , phosphorescent OLEDs (PHOLEDs) have been developed. PHOLEDs can utilize triplet states through intersystem crossing, and increase the quantum yield to unity. A molecular excited state is formed from polarons combining on the same molecule in the EML. If this excited state is not immediately lost to dissociation, but diffuses between particles, it can be considered itself a quasiparticle, these quasiparticles are referred to as excitons. Excitons are able to “hop” among molecules by two main methods of charge transfer, Förster transfer, and Dexter energy (Marcus, or hopping) transfer FIG 8. 34

Fig. 8 Two methods of intermolecular charge transfer: Dipole-dipole coupled (Förster) [top] and Electron exchange (Dexter) [bottom].

Förster energy transfer is a relatively long-range (<100 Å) energy transfer process, and is the result of a dipole-dipole interaction between the excited state host and the ground state guest.31,35 Förster energy transfer is limited to singlet to singlet transitions because the ground state of the guest is always a singlet. Broadly speaking, the location of the HOMO and LUMO orbitals of both the guest and the host are not significant, at least in as much as the excited state of the host must be lower than the excited state of the guest. The emissive properties of the guest dictate the quantum yield of the luminescence.32-34 The other major energy transfer that occurs during electroluminescence is Dexter energy transfer, also referred to as hopping or Marcus transfer. Dexter transfer is a short-range (<10 Å) energy transfer 15 that consists of an electron exchange.34 The HOMO and the LUMO of the guest must be within the HOMO-LUMO gap of the host in order for an electron transfer to occur FIG 9. The rate decreases substantially with increasing distance.31 Dexter energy transfer allows singlet-singlet and triplet-triplet energy transfers36, 37, explaining the efficiency of PHOLEDs.

Fig. 9 Energy orbital diagrams illustrating intermolecular energy transfer.

OLED Motifs The EML of OLEDs often consists of a coordination complex, coordinating an electrophilic, Lewis acid center to a nucleophilic Lewis base ligand. The ligands themselves come in a wide variety of designs, all attempting to chelate the electrophile with the greatest stability possible. There are three basic ligand scaffolds that have been popularized, and many of the patented OLEDs have EMLs that consist of some variation of these designs. The three designs are BODIPY, 2- pyridyl, and quinolato. 16

Perhaps the most famous family of boron-based luminescent complexes, is the Boron-di-pyrromethene (BODIPY) family of dyes. No other boron based dye is so well characterized. BODIPY dyes have been employed in a wide variety of applications including a singlet oxygen generator for cancer therapy.38, 39 energy transfer cassettes,40, 41 and molecular switches.42 BODIPY was first reported in the late 1960s, but was relatively ignored until nearly the 1990s, when their potential application to biological imaging was realized, and BODIPY was found to be a viable, photostable replacement for fluorocein,43 after this initial discovery, there was an explosion of patents into various derivatives of BODIPY, making it the most ubiquitous boron based dye to date. The core of BODIPY is dipyromethene N,N bidentate chelate, which provides excellent binding to boron, however, the appeal of

BODIPY goes beyond the stability.

Fig. 10 Synthetic methods for BODIPY. The alpha, beta, and meso positions correspond to R1, R2 and R3, and finally R4, respectively.

Facile means of substitution allows dramatic changes in BODIPY qualities that can be tuned by substitutions on the scaffold. The scaffold has three main positions from which substitutions can be made, the alpha, beta, and meso positions FIG 10. Until recently, the parent structure of BODIPY, with no substitutions, had eluded researchers due to its instability because none of the pyrrole-based carbons are blocked from electrophilic attack.44 Also the dipyrromethene precursor itself was

18 found to decompose above -30 °C, both factors have proven to be limitations to the synthesis of the parent. However, recently three separate groups reported the synthesis of the bodipy parent complex. The success of the synthesis was based upon a meso-bonded thiomethyl group that serves as an intermediate, binding causes an electron withdrawing effect preserving the structure.45 Another group obtained the desired structure by asymmetric condensation of a pyrrole and an aldehyde.46 Color tuning in BODIPY is unambiguous, and is based upon replacement at any of the three substitution points (alpha, beta, or meso.) The high peak intensities, excellent thermal stability, and high fluorescence quantum yield make BODIPY a very attractive candidate for use in OLEDs50, 51, although there has been relatively little research into this area, especially when compared to the other applications of BODIPY dyes. The setbacks for applying these compounds to OLEDs has been trouble with the films melting upon addition of the voltage and overall weakness in efficiency.47 Another class of OLEDs is based upon the 2-pyridyl scaffold. This bidentate N,N chelate has attracted attention due to similar reasons as BODIPY. The N,N scaffold provides good stability, and the carbon backbone provides many positions for facile substitution FIG 11.48 A related class of N,N chelates is the 2-pyridyl-pyrrolide motif, which has been tuned to achieve a family of dyes that spans the UV spectrum.49 Another 2-pyridyl scaffold analog is the BORAZAN family of dyes, which is an emerging class of N,N- bidentate dyes. Preliminary investigations into the optical properties of these dyes has been propitious.49

Fig. 11 Examples of 2-pyridyl adducts and BORAZAN

By far the most common type of scaffold in OLEDs is the quinolato-type scaffold. This ligand design was popularized by the Alq3-based OLEDs first described in Tang and Van Slyke’s ground-breaking paper.52 Aluminum was found to be less than ideal for chelating to the quinolato moiety, however, and boron has characteristics that make it a natural choice for chelation. Boron was found to provide greater stability than aluminum in quinolato chelates because boron tends to form bonds with ligands that are more covalent in character.53 Aluminum-ligand bonds are more ionic in nature. In the majority of these papers the substituents on boron consist of various aryl groups. These provide a hydrophobic coordination sphere. The boron aryl quinolato complexes are most often used in the formation of the EML in OLEDs FIG 12.54 Multinuclear boron-quinolato polymers have also been reported, leading to greater thermal stability, due to polymerization, and the possibility of functioning in dual roles in OLEDs as EML and ETL due to the chelation of the boron.55

Fig. 12 Examples of quinolato-type adducts. The example on the right has dual binding sites.

Tuning in these complexes can be achieved in a variety of ways. The chelation of the boron to the ligand alone can cause a lowering of the LUMO in the band gap, resulting in a bathochromic shift. Further color tuning, either bathochromic or hypsochromic, can be achieved through the substitution of various EDGs and EWGs. The trends observed are rather obvious, EWGs shift the emission hypsochromically, whereas EDGs and extension of the π system shift the emission bathochromically.56, 57 Further tuning can be achieved by substituting the boron adduct auxiliaries. A limited amount of research has been conducted on these substitutions, and little success has been achieved in shifting the emission.58 However, changes in emission intensity can be attained through such substitutions. An interesting, but less well characterized class of luminescent boron derivatives is the three-coordinate luminescent family of organoboron compounds.

The pπ-π* conjugation of the vacant p-orbital on boron with the π* orbital of the attached carbon π-conjugated moieties is responsible for unique properties that make this class of compounds attractive for absorption and emission in OLEDs FIG 13.59

Fig. 13 Examples of three-coordinate organoboron compounds.

The exceptional π-acceptor properties lend the ability to facilitate remarkable delocalization when coordinated to an organic π-system. These properties are ubiquitous among boron coordination complexes. What makes three-coordinate organoboron species attractive is that boron can function as a σ-donor given its low electronegativity compared to carbon. The pz orbital is readily attacked by nucleophiles cleaving bonds (four-coordinate) or forming four-coordinate species (three-coordinate), therefore, organoboron complexes are constructed of bulky, sterically demanding organic ligands.60 Amid the growing research on three coordinate organoboron compounds, it has been shown that tuning of these complexes can be reached through traditional means, but with greater ease due to facile synthetic properties. Researchers have found that extended π-conjugation is readily attainable through the use of identical π-conjugated moieties. These identical ligands extend the π-conjugation through the vacant p-orbital in the LUMO of boron. This highly delocalized low-lying LUMO causes long wavelength absorbance.61 -63 The interest in applying trivalent organoboron complexes to OLEDs and other molecular electronics is due to the charge transfer ability. This feature of the complexes has been exploited for use in fluoride ion sensors, as fluoride binding to the pz orbital of boron forms a fluoroborate interrupting the π-conjugation with the organic substituents leading to a color change.64 OLEDs have also utilized these

22 same charge transfer abilities in order to produce both emitters and charge transport layers.65, 66

New Directions for Luminescent Organoboron Design

Our group’s interest in luminescent boron complexes started with benzoboroxole. Benzoboroxole is a versatile molecule that has found a large variety of synthetic and medicinal applications. Benzoboroxole was first described in 1957 by Torssell.67 These new compounds were further characterized by Snyder in 1958.68 Benzoboroxole is a polycycle, consisting of a six-membered aromatic ring and a boron heterocycle consisting of a five-membered ring with internal ester, sometimes referred to as a boronolactone.69 Since its initial characterization, there have been many attempts to improve the synthetic procedure in order to improve both yield and ease of synthesis. Synthesis usually begins with ortho-boronobenzyl alcohols, which are unstable and easily dehydrated, leading to the use of a variety of protecting groups in the synthetic method. These methods vary, but there are four general synthetic schemes. The first is the introduction of a hydroxymethyl group followed by intramolecular esterification.70 The next method uses harsher conditions, and starts with ortho-bromobenzyl alcohols treated with butyllithium and a trialkyl borate, leading to dehydration.71 The two final, updated synthetic schemes deviate from the tradition by starting with non-aryl alcohol reagents. The first begins with a series of alkynes, and uses a ruthenium catalyst to form benzoboroxole.72 A 6, 7-substituted benzoboroxole can be formed from a corresponding magnesium heterocycle by treatment with a trialkyl borate, and DDQ.73 Perhaps the most convenient synthesis of benzoboroxole was described by Mereddy’s group. This one-pot synthesis requires only 2-formylphenylboronic acid and sodium borohydride. With the appropriate solvents the reaction gave high yields, with far fewer steps.83 Since these advances in synthesis, the multi-functional 23 uses of benzoboroxole continue to be discovered. Following the research development into the field of organic coupling reactions, benzoboroxoles have been used in Suzuki-Miyaura coupling for regioselectivity, giving ortho-arylsubstituted benzyl alcohols.74, 75 Given benzoboroxole’s excellent water solubility, they have been investigated for use as molecular receptors, requiring no co-solvent to facilitate solubilization. Benzoboroxoles are able to bind glycosides under physiologically relevant conditions, and thus are viewed as promising candidates for oligomeric sensors for selective sugar recognition.76, 77 Benzoboroxoles have also been studied for use as carbohydrate chemosensors, 78 and dyes for the determination of polyols.68 One of the most promising applications is in medicinal research. Benzoboroxoles were discovered to be fungicidal as late as 2005.79 One of the derivatives with the most potential is a species fluorinated at the 5-position that can easily transit through the nail.80, 81 Interestingly, benzoboroxole’s anti-fungal efficacy has been found to be dependent upon the presence of a boron in the five- membered ring. Five-membered rings containing carbon alone are inactive. Additionally, halo-substituents in the benzene ring, especially at position 5, have the highest activity.82 The application of benzoboroxoles in medicine was the original starting point of our group’s research into boron chemistry, with the discovery of some virulent anti-fungicide substituted boron species.83 After a foray into the uses of benzoboroxoles in medicine, our research group focused on luminescent properties of various benzoboroxole adducts. These novel quinolato Alq3 analogs were found to emit at 530 nm. The synthesis of these complexes uses N,O-type ligands to coordinate to the boron, exploiting boron’s Lewis acidity, through a condensation reaction, rapidly forming a boronic ester FIG 14. The heterocycle of benzoboroxole provides a distinctive structure which grants numerous positions for substitution, offering potential tunable species in the future.84 As stated, the benzoboroxole forms a boron-ester type bond with the ligand and the non-bonding lone pair electrons on the nitrogen are donated to the empty pz orbital on the boron.

Fig. 14 Two examples of benzoboroxole adducts. The ligands here are 8- hydroxyquinoline (left) and 10-hydroxybenzo[h]quinoline (right).

However, the stability of the complexes was exiguous as the boron-ester bond is hydrolyzed by acid or water, to the point that the reaction is reversible, unless dry solvents are used.85 Even the relatively gentle environment of silica gel columns has proven too harsh for these compounds. To mitigate these problems Jeffrey Carlson began synthesis with phenyl borinic acid derivatives.86 The aryl based moiety creates a hydrophobic coordination sphere that prevents polar substituents (viz. water or protons) from approaching the sensitive boron-oxygen bond. This improvement greatly increases the stability of these complexes. These phenyl borinic acid complexes are stable enough to be purified using vacuum sublimation. One of the purposes of this research was the synthesis of a family of boron-based luminescent complexes that covered a wide range of the UV-vis spectrum, preferably within the same family of ligands. Unfortunately, the quinoline-based ligands did not allow any facile means of synthetic color tuning, except for addition of EWGs and EDGs. The original ligands that were proposed were polymer-like Schiff bases, and, given the success of PLEDs, this seemed a promising route. However, the lack of rigidity in the Schiff base backbone allowed the loss of energy

25 through rotation rather than loss of a photon. This free rotation resulted in the severe loss of luminescent intensity. Therefore, a more rigid, tunable ligand was sought. Azole based ligands can be easily tuned through what is essentially an auxochromic heteroatom. Further, the ligands that were chosen have two aryl regions, which allows the synthesis of, not only phenyl rings but naphthyl and anthracene as well. This flexibility allows several means in which to tune the luminescent boron complexes. Moreover, the quinoline moiety provides a stable binding site for the boron, forming a six-membered heterocyclic ring. These functional ligands have allowed the synthesis of a family of luminescent boron complexes that spans the UV-vis region. If one follows the unambiguous trend in bathochromic shifts, the addition of anthracene to the basic azole-scaffolding should achieve a red emitting complex, however, synthesis of the anthracene ligand has been stalled by difficulties in purification. Computational studies were pursued in order to predict the properties of the derivatives (see Chapter 2). Alternatively, the addition of EDGs to the scaffold could reduce the rigidity of the ligand, resulting in a lower energy (red) emission (See Chapter 2). The ligand itself consists of a hydroxy substituted phenyl ring and a heteroatom substituted polycyclic azole. The N,O quinolato-type moiety allows stable adduct formation with boron, which serves to extend the π conjugation of the ligand and forms a rigid complex. This rigidity is due to the formation of a dative bond by donation of the non-bonding electrons on nitrogen to the empty pz orbital on boron. This dative bond limits free rotation of the ligand. The adduct formation increases the luminescent quantum yield, while reducing the band gap through contraction of the LUMO energy level.87-89 Synthesis of these ligands is straightforward through an acid catalyzed cyclization, and is dependent upon the desired heteroatom that is used, either 2-aminophenol or 2- aminothiophenol. The position of the naphthyl ring is also easily changed, simply based upon the position of both the naphthoic acid and hydroxy group. When reactions were run using 2-aminophenol, the conditions for reaction were increasingly harsh, requiring potent silicon tetrachloride, and still resulted in lower 26

yields, mandating column chromatography for purification. 2-thiophenol warranted the less potent phosphorus trichloride. Good yields were achieved merely through recrystallization from acetone SCHEME 1.

Scheme 1

The improvements previously discussed are limited to tuning the complexes for emission at various wavelengths, based solely on changes to the ligand scaffold. In order to increase the luminescent intensity of the adducts, the substituents on boron can be changed. The work presented here is based upon the addition of electron withdrawing groups, here fluorine, as substituents on boron. The similar aryl moiety of the substituents between benzoboroxole and phenyl borinic acid did not permit any significant change in luminescent intensity of the final complexes. It was theorized that the addition of EWGs as substituents on boron would increase the luminescent intensity and result in a slight hypsochromic shift. This is the result of the stronger dative bond formed between boron and nitrogen as a result of the electron withdrawing effect of the substituents. The bonding in these complexes with these new substituents is exactly the same, and a multi-dentate adduct forms, where boron is attached through a boronic ester bond and a dative bond with

27 nitrogen. The substituent chosen in this work was fluorine, which withdraws electron density from the boron, causing a stronger electron donation from nitrogen into the empty pz orbital on boron. The increase in bond strength results in increased rigidity of the overall structure, leading to higher energy, as well as improved intensity emissions. While fluorine has provided some improved characteristics, it has presented a number of challenges that must be overcome. The compounds were found to be less soluble due to their increased polar character, but exhibited stronger emission of light in fluorescence studies

Chapter 2: Results and Discussion

One can easily see from the structure that the boron atom in these complexes remains electrophilic, providing ample opportunities for hydrolysis of the boronic ester. The boron-nitrogen bond is merely a dative bond that, even with the increased binding strength due to the electron withdrawing effect of the fluorines, can be easily broken.85 Another problem that presented itself was the ease with which fluoride can be lost so that the boron is able to form an adduct with a donor solvent. Finally, the hydrophobic organic ligand , in conjunction with the hydrophilic fluorides, limits the solubility of these novel complexes, forcing the majority of the reactions to be worked up in dichloromethane (DCM). Even dimethyl sulfoxide (DMSO), and N,N-dimethyl formamide (DMF) proved ineffective in solubilizing these complexes, additionally, these are both donor solvents. Also, given the extreme moisture sensitivity of our adducts all deliquescent solvents, such as DMSO, had to be avoided. The synthetic method used was simple and conditions were maintained in all circumstances. The ligand was dissolved in DCM and allowed to stir under argon, as boron trifluoride diethyl etherate, and triethyl amine were added. A fluoride from BF3 and a proton from the ligand were both lost during the reaction to form hydrogen fluoride. The reaction was slow and was allowed to progress for a week, and in all cases a fluorescent precipitate formed. The precipitate was collected and the resulting solution kept in the freezer to facilitate further precipitation SCHEME 2.90 The formation of trace product was near instantaneous, as monitored by UV lamp. The original work-up included washing the organic layer with water to remove ammonium salts that formed in solution, however, this step was abandoned after the sensitivity to hydrolysis was discovered, and washing with water seriously affected the yield. Some of the issues with instability were discovered after attempting synthesis. The first ligand tested was 8-hydroxyquinoline. After numerous attempts with diverse conditions, the perhaps fugacious complex, was never able to be isolated, and there were no indications of a successful reaction. This 29 was due to the formation of a less favorable five-member heterocycle, that was not stable enough to form a stable adduct. 10-hydroxybenzo[h]quinoline was then chosen as it forms a more stable six membered heterocycle. After reaction the product precipitated out of solution. All azole-based complexes also form six- membered heterocycles and thus maintain moderate stability. Purification has also been a challenge with these complexes due to the moderate stability.

Scheme 2

Given the poor solubilities of these complexes, crystallization has been elusive, and eventually we again relied upon DCM to purify the complexes. Purification through column chromatography was also attempted, however, this method was abandoned as the compounds reacted with the silica gel, probably through loss of a fluoride. The formation of these spurious adducts prevented collection as they could not be flushed from the column. Eventually, the purification method that was 30 adopted was to dissolve the adduct in DCM and stir for two hours with anhydrous sodium carbonate to neutralize traces of the ammonium salt, this was filtered and recrystallized from hexane and DCM. The resulting precipitate was collected and dissolved in DCM, concentrated in vacuo, and placed in the freezer to precipitate. The thermal stability of these complexes has also proven to be unsatisfactory. In the context of HR-APCI-MS, apparently, rather than sublime upon heating to be read as MS data, these complexes decompose. Given the hypsochromic shift compared to the free ligand inherent in these complexes, the realization of a family of luminescent adducts that span the UV spectrum has yet to be realized. The lowest energy emission that has been observed is 521 nm (green). Computational work is currently being pursued to determine the most effective position in which to place EDGs. The addition of EDGs at strategic locations on the ligand scaffold should result in a bathochromic shift. Another approach that was previously mentioned was the construction of a ligand with extended π-scaffold. The extension of this pi electron density could also result in a significant bathochromic shift if the currently observed trend with the addition of a naphthyl ring can be extrapolated to anthracene-based ligands. Many means of characterization were utilized in order to both confirm the final structure as well as probe the luminescent properties. The characterization methods used were 1H, 13C, and 11B-NMR, X-ray crystallography, HR-APCI-MS, UV- visible spectroscopy, quantitative fluorimetry, and TD-DFT computations.

Optical Measurements

When comparing both the absorbance and the emission of the fluorine substituted and phenyl substituted derivatives, the excitation had the same alignment, however there was a hypsochromic shift, when compared to the phenyl substituted species, in emission in the fluorine substituted derivatives, due to the

31 electron withdrawing effects of the fluorine substituents FIG 15.

450

400

350

300

250 BPh2(2,3-HNBO) Ex. Intensity 200 BPh2(2,3-HNBO) Em. (au) 150 BF2(2,3-HNBO) Ex. BF2(2,3-HNBO) Em. 100

0 200 300 400 500 600 -50 Wavelength (nm)

Fig. 15 Comparison of emission and excitation data between phenyl substituted and fluorine substituted species.

The electron withdrawing effect of the fluorines increases electrophilic character of the boron. The non-bonding electrons on nitrogen then form a progressively stronger bond with the boron. This increases the rigidity of the overall complex, resulting in a higher energy emission. This result was expected, as computational work has shown the importance of nitrogen coordination on the emissive properties. Our X-ray crystallography studies have confirmed the stronger coordination in the fluorine substituted derivatives. The fluorine substituted derivatives also qualitatively displayed much higher intensity excitation and emission, as compared to the phenyl substituted derivatives. The augmented intensity could be the result of either enhanced emission of light, or higher

32 absorption of energy. Quantum yield measurements are requisite to determine the cause of this amplified emission.

900

800

700

600

500 BF2(1,2-HNBO) Intensity 400 BF2(2,3-HNBO) (au) BF2(1,2-HNBT) 300 BF2(2,3-HNBT) 200

100

0 350 450 550 650 -100 Wavelength (nm)

Fig. 16 Overlay of the emission spectra of the azole adducts

A series of stepwise shifts in emission was achieved through various means of tuning the complexes; emission ranged from 428 to 524 nm FIG 16. When compared to previously characterized phenyl substituted azole ligands86 TABLE 1, which have a smaller π-system, the bathochromic shift is obvious.

1,2-HNBO 1,2-HNBT 2,3-HNBO 2,3-HNBT

BF2 Derivatives 433 457 494 518 412[1] 438[3]

BPh2 Derivatives 454 489 536 578 476[2] 510[4] 658[5] 655[6]

Table 1 Comparison of the emission maxima between the BF2 and phenyl boronic acid derivatives. (1) and (3-6) correspond to shoulders, while (2) represents a secondary degenerate emission maximum.

The means of tuning the complexes was two-fold: substitution of the auxochromic heteroatom, and extension of the conjugated π-framework. Oxygen and sulfur were both tested as auxochromes. A nitrogen substituted species was also tested for synthesis, however, the ability to transfer a proton between the two nitrogens prevents formation of an adduct. Various means of synthesis were tested, including an attempted synthesis at -78 °C, however, the ligand would first need to be methylated on one nitrogen to lock it into a useful conformation. Further research is required to see the effect of a nitrogen heteroatom, current research is underway. At present, the phenyl substituted, ethylated nitrogen derivatives have been synthesized and characterized. For more details see the end of this chapter. Optical data obtained from the two heteroatoms that were tested revealed a clear bathochromic shift for sulfur substituted ligands. The emission of these complexes is bathochromically shifted compared to the oxygen derivatives which appear at higher energy wavelengths. The reasons for the shifts have been explored through computational means. The other substitution that was made was the extension of the conjugated π-framework, as compared to the phenyl substituted derivatives which have been characterized by another group.90 In both instances the overall bathochromic shift is the result of a decrease in the compound ionization energy. The results of quantum yield measurements have also shown interesting trends. The high intensity emission with moderate quantum yield values may be the result of a high absorption cross section. The disparity in quantum yield, between the 1,2- substituted and the 2,3-substituted derivatives is striking.

Absorptivity Φ Compound Main ε Excitation Emission Calculated Absorbance [1/M*cm] Wavelength Wavelength Wavelengths [nm] [nm] [nm]

HPBI 333 68000000 340 385 0.35 297 380[1] 286 1,2-HNBO 394 370000000 470 433 0.25 375 412[2] 325 311 1,2-HNBT 415 370000000 470 457 0.28 395 438[3] 337 322 2,3-HNBO 395 370000000 470 494 0.04 348 478[4] 332 284 2,3-HNBT 408 370000000 470 518 0.08 366 529[5] 350 288 HBQ 395 37000000 370 482 0.12 289 460[6] 245

Table 2 Summarized optical properties of all compounds. BF2(1,2-HNBO) exhibited a degenerate emission maxima [2]. The rest exhibited shoulders ([1] and [3-6]).

The dramatic difference in quantum yield values is most likely the result of fluorescent quenching due to photoinduced electron transfer (PeT) in the 2,3- substituted ligands versus the 1,2-substituted ligands TABLE 2. It is well documented that the fluorescence of a molecule can be turned on or off through PeT.91, 92, 93 It may be speculated that the charge transfer responsible for this

35 quenching may be due to the shift in electron density from HOMO to LUMO, this has been illustrated through computational studies. In order to understand the nature of the PeT in these compounds, it will be necessary to find the oxidation energies of these derivatives through electrochemistry.

Nodal Plane Theory

The absorbance and emission data that was collected for these compounds revealed a striking trend in Stoke’s shift values. It was observed that the Stoke’s shift of the long-wavelength fluorescence of the 2,3-substituted complexes was much larger as compared to the fluorescence emission of the 1,2-substituted complexes FIG 17.

900 4.50E+02 800 4.00E+02 700 3.50E+02 600 3.00E+02 Intensity 500 2.50E+02 Abs. (au) 400 2.00E+02 Excitation 300 1.50E+02 Emission 200 1.00E+02 UV 100 5.00E+01 0 0.00E+00 200 300 400 500 600 Wavelength (nm)

450 0.2 400 0.18 350 0.16 300 0.14 0.12 250 Intensity 0.1 Excitation 200 Abs. (au) 0.08 150 Emission 0.06 100 0.04 UV 50 0.02 0 0 -50 200 300 400 500 600 -0.02 Wavelength (nm)

450 0.08 400 0.07 350 0.06 300 0.05 250 Excitation Intensity 0.04 200 Abs. (au) 0.03 Emission 150 UV 100 0.02 50 0.01 0 0 -50 200 400 600 -0.01 Wavelength (nm)

250 2.50E-01

200 2.00E-01 Excitation 150 1.50E-01 Intensity Abs. Emission (au) 100 1.00E-01 UV

50 5.00E-02

0 0.00E+00 200 400 600 Wavelength (nm) 37

Fig. 17 Illustration of the change in Stokes shift given the orientation of the naphthyl plane. Featured complexes BF2(1,2-HNBO) [top], BF2(2,3-HNBO) [2nd from top],

BF2(1,2-HNBT) [2nd from bottom], BF2(2,3-HNBT) [bottom].

These results are consistent with nodal plane theory.94,95 Nodal plane theory, as it applies to our ligands, is analogous to excited state intramolecular proton transfer, except that the ESIPT-type process is here described by skeletal distortion of the aromatic ring rather than an intramolecular proton transfer. In an ESIPT, the main change in electronic structure results from the σ-bonding of the hydrogen bonded moiety. In our ligands the change is a result of the S0 -> S1(π) excitation ((π- π*) transition) on the π-bonding scaffold of the aromatic ring moiety.96 The change is more global and encompasses the molecular framework. Nodal plane theory inserts a node in the aromatic rings through the double bonds, which gets rid of lone π-electrons, no electron density can exist within this plane.

Fig. 18 Energy diagrams of the 1,2 and 2,3 isomers.

The nodal plane that runs through the 1,2-substituted compounds is localized on only one of the six-membered rings of naphthalene. Notwithstanding, the nodal plane in the 2,3-substituted compounds passes through both naphthalene rings. The enhanced aromaticity of the nodal plane in the 2,3-substituted derivatives should result in lower energy of the compound overall. Given that the nodal plane passes through a single ring of the 1,2-substituted species, they are not expected to be very stable, whereas the nodal plane passing through two rings could be expected to be much more stable FIG 18. That is exactly what is observed, the energy of the 2,3-substituted derivatives is lower, and results in a lower energy emission, and therefore a larger Stokes shift.

Complex Excitation Emission Stokes Shift

BF2(1,2-HNBO) 375 431 56 394 _ 37

BF2(2,3-HNBO) 332 484 152 348 _ 136

BF2(1,2-HNBT) 395 449 54 413 _ 36

BF2(2,3-HNBT) 350 517 167 366 _ 151

Table 3 Comparison of Stokes Shift values for BF2 substituted derivatives. In the absorption data of each derivative there is a degenerate excitation peak.

The 1,2-subtituted derivatives have a decreased degree of aromaticity, as compared to 2,3-dervatives and the higher energy complexes have correspondingly high Stokes shift values. The increased aromaticity of the 2,3-derivatives may also help to explain the trend of upfield shifts in the 1H-NMR spectrum.

NMR Spectroscopy

Much of the characterization of complexes in this research was based upon a combination of 1H, 13C, and 11B-NMR. NMR was used both to confirm formation of the intended compound as well as determine purity. Dichloromethane (CD2Cl2) and chloroform (CDCl3) were the two deuterated solvents that were used in all of the characterizations. A range of solvents were tested, however, the solvents were limited by poor solubility as well as the sensitivity of the compounds to any type of donor solvents. In most cases, the compounds were only weakly soluble in donor solvents. Given their low hygroscopic nature, both chloroform and dichloromethane proved to be good candidates for characterizing these water sensitive complexes. Two separate regions of aromatic proton residence exist, the aryl and azole moieties. The aryl moiety coincides with the naphthyl rings of the ligand, while the azole moiety correlates to the phenyl ring that is directly attached to the heteroatoms forming the heterocycle. The individual protons have been assigned for these ligands through 2D-NMR in a previous study.86 A careful analysis of proton shifts among the oxygen and sulfur substituted heteroatoms of the azole based ligands reveals a downward shift in proton signals. The shift corresponds with a red-shift in peak fluorescence. This trend is most obvious in the oxygen based derivatives, and less transparent in the sulfur based derivatives. This bathochromic shift seen in the emission data is the result of a reduction in the HOMO-LUMO band gap, due to an increase in the HOMO, or a decrease in the LUMO, this could be the result of an increase in aromaticity. The proton shifts could result from a global increase in aromaticity. The correlation of proton shifts with trends in optical spectroscopy has proven inconclusive for the BF2 derivatives. However, the results of the 1H-NMR have lent support to the other observations about the effects of nodal plane theory on our complexes. The location of the nodal plane, localized on both the phenyl and naphthyl portions of the ligands, precludes significant localization of 40 electron density limited to the naphthyl ring, rather, the location of the nodal plane results in the delocalization of electron density across the entire ligand. This is in contrast to the effect of the nodal plane on the electron density of the 2,3- substituted derivatives, where the nodal plane results in significant delocalization of electron density on the naphthyl ring alone. This obviates the localization of electron density on the phenyl ring in the azole moiety. The validity of these assertions have been confirmed computationally (see molecular orbital plots). The proton shifts of the oxygen species mirror the trend of the emission data, in that higher energy emissions result in less aromatic, and consequently, further downfield shifted signals. The highest energy emission is the phenyl substituted parent compound [BF2(HPBO)] which appears furthest downfield (highest signal is δ = 7.95 ppm).90 Likewise, the furthest signal upfield corresponds to the 2,3- substituted derivative, which also has the lowest energy emission, due to higher aromaticity. The relationship in the sulfur derivatives is less clear. The proton signals for the 1,2-substituted derivative appear further upfield than the 2,3- substituted derivative, which has similarly shifted signals as the phenyl substituted analog. The unexpected results may be due to an increase in the electron density on the phenyl ring in the azole moiety which results in an augmentation in the aromaticity, this increase results in a upfield shift in proton signals as compared to the 2,3-substituted derivatives. More research into this area needs to conducted in order to elucidate the exact trend responsible for these results. The 1H-NMR of the nitrogen-substituted derivatives yielded similar proton spectra. The ligand alone revealed a spectrum similar to the other ligands with the downfield O-H proton, and the overlapping signals. The major difference in the proton spectrum is the addition of the ethyl group which appeared much higher upfield in the alkyl region. Analysis with 13C-NMR revealed thirteen aromatic carbon signals and two alkyl carbon signals. Carbon NMR spectra were difficult to obtain as the poor solubility of the complexes precluded the ability to obtain clear 13C-NMR data. Given the low 41 solubility, all of the complexes required overnight runs. The azole based ligands have seventeen carbons and these signals are all visible with a varying degree of intensity. Boron NMR spectra provided additional support. The peaks that were obtained revealed a slight upfield shift in boron signal from the boron trifluoride etherate standard. The sharp signal is indicative of a four coordinate boron species, in a three coordinate boron, the signal would appear more broad. Though far from conclusive, this data provided further support of the formation of our complexes.

HR-APCI-MS

The relative instability (mostly to moisture) of our complexes has made characterization a challenge. Methods of HR-MS characterization suitable for the complexes were difficult to find. Fortunately, some of the adducts that form are able to give some insight into the chemical behavior of our compounds while confirming the expected structure. Many methods were tested, but most were found to be too harsh for the complexes. When APCI-MS data was gathered, an interesting trend was found. The molecular ion was never found. However, a molecular weight indicative of the loss of a fluorine and the formation of an adduct with the solvent, THF, was observed. These results are not surprising as a loss of fluorine through heating is expected, as is adduct formation with a donor solvent such as THF. The adduct formation with THF could have occurred before or after the MS experiment as the solvent used to dissolve the sample was THF. The other major peak that was observed was the complex with the loss of one of the fluorines without adduct formation FIG 19.

Fig. 19 Explanation of the peaks observed in the MS data. See appendix for experimental data.

X-ray Crystallography X-ray crystallography allowed us to obtain irrefutable evidence of the existence of our compounds. The usefulness was limited to one crystal structure that we were able to obtain FIG 20 and 21. Rich structural information was obtained relevant to photoluminescent behavior. The most significant feature of the crystal structure was the bond length of the boron-nitrogen dative bond. When comparing the length of the dative bond in the boron heterocycle, the length was found to be much closer to a single bond. At 1.56 Å, the bond length was much shorter than the B-N dative bond found in benzoboroxole adducts (1.62 Å), and even shorter than those found in the phenyl borinic acid derivatives (1.63 Å). The significance of these findings become clear when one considers that the length of the bond is a function of its strength.97

C23 O4 C8 C18 C19 C17 C11C12 C9 C6 N5 C10 C15 H231 C24 H171 H81 H181 H191H241 H161 H221 H111C7 B21 C16 C13 S1 C20 H151 C22

H151 H111 H161 S1 C16 C15 C11 H191 C20 C22 C6 H221 C7 C13 C19 C10 C9 C18 N5 C12 C8 H231 C23 H181

O4 B21 C24 C17 F3 H81

H171 H241 F2

Fig. 20 Crystal structure of BF2(1,2-HNBT) refined using SHELX software.

This is no less true in the case of dative bonds though they are only roughly 1/3 of the strength of covalent bonds. The fact that the boron-nitrogen bond was so much shorter than that in the benzoboroxole-8-hydroxyquinoline adduct indicates that bond is much stronger. This enhanced bond strength has significant effects upon the optical spectra of the BF2 derivatives, as discussed in the optical spectroscopy section of the chapter.

Fig. 21 Crystal packing unit cell of BF2(1,2-HNBT).

Theoretical Chemistry

In order to investigate both the nature of the transitions of our complexes, as well as characterize the bonding, computational methods were pursued. The software used throughout was the Gaussian 09 suite of programs.98 TD-DFT was the level of theory used, with B3LYP, and PCM solvent modeling. The fairly ubiquitous B3LYP hybrid functional theory, with the 6-311++G(d,p) basis set was chosen as it has been a popular choice among many small organic dyes.99-101 Given the acceptable results of the chosen basis set, there was no need to test any others(see appendix for a comparison of the experimental UV-vis plot versus the computed UV- vis plot). This basis set was utilized for both the geometry optimization and the excitation energies. 45

When considering the molecular orbital plots that were generated from computation, the trend in excitation becomes clear. The fluorine substituents in the phenyl substituted parent compounds contribute more to the HOMO than in any of the naphthyl substituted complexes, this is further substantiated by VMOdes calculations. Moreover, the electron density of the phenyl substituted complexes is scattered across the entire ligand in both the HOMO and the LUMO FIG 22. This diffuse electron density is consistent with a nodal plane across the phenyl ring of the ligand. A similar case is seen in the 1,2-substituted derivatives, however, there is much less electron density on the fluorine substituents compared to the phenyl derivatives. The electron density is nearly entirely localized on the ligand FIG 23. And the electron density is spread across the entire ring, not just the naphthyl component. This behavior is consistent with what is known from nodal plane theory. The best way to illustrate this is to consider the molecular orbital plots of the 2,3-substituted derivatives. The molecular orbital plots of the 2,3-substituted derivatives shows nearly all density localized on the naphthyl ring in the HOMO FIG 24. If one considers the location of the nodal plane, where it runs the length of the naphthyl rings, it is easy to see that the electron density is mostly delocalized across the naphthyl ring. In the 1,2-substituted series, the nodal plane runs through only one ring in the naphthyl group. This prevents the extensive electron delocalization on the naphthyl ring and spreads it across the entire ligand. The types of bonding among the three analogs, the parent, phenyl substituted derivatives, the 1,2- substituted derivatives, and the 2,3-substituted derivatives are different among all three.

Fig. 22 Molecular Orbital Plots of the parent complexes, BF2(HPBO) [top], and

BF2(HPBT) [bottom]. The HOMO is plotted in the middle column and the LUMO is plotted on the right.

The molecular orbitals in the parent analogs are primarily a global bonding type. While the molecular orbitals in the naphthyl substituted derivatives are primarily a ligand based anti-bonding type, even between the 1,2 and 2,3- substituted derivatives the molecular orbital type is unique.

Fig. 23 Molecular Orbital Plots of BF2(1,2-HNBO) [top], and BF2(1,2-HNBT) [bottom]. The HOMO is plotted in the middle column and the LUMO is plotted on the right. Computational means can also be employed to try to determine the nature of the differences in luminescence based solely on the auxochromic heteroatom that was used in each complex. When considering the optical data, the differences are 47 striking and lucid. In each case, the sulfur substituted species were more bathochromically shifted than their oxygen analogs. The bathochromic shifts observed in the sulfur analogs is due to a lowering in the energy of the LUMO. This reduction in the band gap results in a lower energy emission. When a comparison of the three possible heteroatoms is made, including sulfur, oxygen, and nitrogen, and

Fig. 24 Molecular Orbital Plots of BF2(2,3-HNBO) [top], and BF2(2,3-HNBT) [bottom]. The HOMO is plotted in the middle column and the LUMO is plotted on the right.

specifically focusing on the bonding and aromaticity properties of these three, several fascinating features are explicated. Consider first the nature of carbon bonding among these three heteroatoms. The bonding with carbon consists of both σ and π bonding. And the overlap of the σ and π orbitals of the heteroatoms with the σ and π orbitals of carbon varies with the nature of the heteroatom. The bond strength as well as type of bonding between oxygen and nitrogen is very similar. The bond length of single bonded oxygen, nitrogen and sulfur is 142 pm, 167 pm, and 226 pm.102 As an example, one can consider the bond lengths of thiophene (1.71 Å), furan (1.37 Å), and pyrrole (1.38 Å).103 Given the fact that bond strength is a 48 function of bond length, where shorter bonds correspond to stronger bonding, these results evince the difference in the nature of bonding in these heteroatoms. Part of these differences in bond strength may be rooted in the dissimilitude in bond strength between σ and π bonds. When comparing σ-bond strengths of row 2 and row 3 elements, the σ-bond strength increases toward the right of the periodic table.104 Similarly, this data is in accordance with the double bond rule, which surmises that third row elements form much weaker double bonds than second row atoms.105 The weaker bonds were previously thought to be the result of poor overlap between 2p and 3p orbitals, with further investigation,106 however, this has been disproven. Indeed, it has been shown that there is greater overlap with carbon in third row elements than in the second row. The greater overlap is easily rectified with the decline in polarizability following the trend in electronegativity moving to the right in the table.106 Heavier, less electronegative atoms are more easily polarizable. This indicates that the difference between σ and π-bond strength is larger in the third rather than the second period. In order to speculate on the potential causes of this decrease in bond strength, it is useful to consider periodicity. One of the most accessible reasons that accounts for the attenuated bond strengths of heavier elements, is the lack of electron density between bonding partners, this point is obvious, however, the cause of the lower electron density is more opaque. There is a trend in periodicity that heavier atoms have more diffuse orbitals because the electrons spend more time in the inner lobes that cannot overlap with the orbitals of other atoms, and so the covalent bonds of these atoms are weaker.108 The cause of these diffuse orbitals arises from the fact that only the s-atomic orbitals of the first row elements are occupied, therefore, there is no Pauli repulsion in the unoccupied p-atomic orbitals, and hence little diffusion. The situation is just the opposite in the case of the heavier third row elements; both the s-atomic orbitals and the p-atomic orbitals are occupied.109 The repulsion between these two occupied orbitals results in a greater distance between, and hence, more diffuse orbitals. Similarly, the σ-bonds in the lighter elements are weaker than the π-bonds 49 because of lone pair repulsion. This repulsion is limited to the σ-orbitals. Experimental evidence suggests that the lone-pair repulsion is of reduced importance in higher row atoms. In both second and third row atom bonding there is a degree of hybridization, specifically, isovalent hybridization. It is beneficial for these heteroatoms to undergo isovalent hybridization as there is better atomic orbital overlap, reduced repulsion between bond and lone pair, and finally, hybridization favors larger bond angles.109 The degree of isovalent hybridization is higher in second row atom bonding due to strong overlap of p-valence atomic orbitals. It is advantageous for second row elements to undergo isovalent hybridization because it can reduce the repulsion of the lone pairs of electrons. The effects of σ-bonding in sulfur-carbon compounds are obviously most dramatic in non-aromatic systems. In aromatic sulfur-carbon systems π-bonding is the object of focus. It would be remiss to assume that σ-bonding is the only factor in aromatic sulfur-carbon bonding. In the excited state the π-bonding becomes very important. As previously discussed, the π-orbital overlap is actually greater in sulfur than in oxygen. Furthermore, experimental evidence suggests that sulfur is better able to undergo conjugation with the aromatic π-system of the ligand. This can be seen by a comparison of the VMOdes plots, in all cases, there is greater contribution on sulfur compared to oxygen. Further credence to this assertion is given by contrasting the molecular orbital plots of each of these. Visually, there is more contribution on the sulfur atom than on the oxygen, and this observation can be made in all derivatives. The increased π-orbital overlap may help to facilitate stabilization of the extended π-conjugate system, due, in part, to sulfur’s fundamental ability to improve aromatization when bonded with carbon in a heterocycle, when compared to oxygen. A more detailed discussion of this follows, at the end of this section. Another consideration that must be accounted for in this discussion of bonding is electronegativity effects. Electronegativity is a property that is intrinsic to atoms, and an atom with high electronegativity “pulls” electron density towards itself. 50

When forming bonds with carbon, the nature of the electronegativity of both the carbon and the heteroatom can affect the type of bonding that is favored. If the difference in electronegativity between the carbon and the heteroatom is large then the bonding will be more ionic in nature. Conversely, if the difference in electronegativity is small, then the bonding will be more covalent. The reason for the dramatic differences in bonding is due to the polarizability of the electron density in sulfur. When a comparison is drawn of the electronegativity differences between sulfur and oxygen, there is a striking contrast of 2.5 and 3.5 respectively.110 When juxtaposed with carbon, sulfur shares the same electronegativity, while oxygen is obviously much more electronegative. This disparity affects the bonding of carbon with these two heteroatoms. When carbon bonds with oxygen, the intrinsic electronegativity of oxygen “pulls” the bonding electrons so that there is preferential residence on the oxygen.111, 112 This preferential residence results in bonding that is more ionic in nature, rather than covalent.113 The reduction in covalency correlates to higher energy bonding114, which may explain the effectiveness of the more facile means of synthesis of the sulfur derivatives as compared to the harsher conditions necessary to form the oxygen derivatives. The lower electronegativity of sulfur allows the electron density between carbon and sulfur to be more equally shared. This sharing results in higher bond covalency, and therefore greater bond stability in π and σ bonding.114 In order to decipher nature of bonding in these adducts, other aspects of the heteroatoms must be examined. Perhaps the most significant justification for the observed bathochromic shift in sulfur substituted derivatives, is most appropriately delineated through consideration of the degree of aromaticity in these heterocycles. The increased level of aromaticity in thiophene versus the oxygen and nitrogen analogs results in a decrease in the band gap of the frontier molecular orbital (FMO) of sulfur as compared to the other two heteroatoms.115 The calculated band gap for thiophene is 9.45601 eV, while for pyrrole and furan the values are 10.03549 eV and 10.03978 eV respectively.115 The reduction in the band gap in 51 sulfur equates to a lower energy emission and therefore a bathochromic shift. It is well documented that sulfur-containing heterocycles exhibit greater aromaticity.103, 115- 117 If one continues to consider five-membered heterocycles of these three heteroatoms as a model system, then the degree of aromaticity can be easily shown through comparison of the Dewar resonance energy (DRE) values of each.118 These DRE values can be calculated by finding the difference in energies of atomization between the conjugated molecule of interest and some reference. This method represents the contribution due to the cyclic electron delocalization, essentially adding up the contributions from both the π and σ bonds.118 When these DRE values are compared for thiophene, pyrrole, and furan, 6.5, 5.3, and 4.3116 are obtained, respectively; where the DRE values are measured in kcal/mol. There are multifarious reasons for sulfur’s contribution to heterocyclic aromaticity, the most auxiomatic contributions are due to sulfur’s large size. The larger covalent radius reduces ring strain.103 The longer sulfur-carbon bonds create larger bonding angles that reduce the strain.2 Oxygen is considerably smaller than sulfur and therefore “fits” more closely in the ring, forming shorter bonds. It must also be speculated that the higher electronegativity of oxygen compared to sulfur contributes to a lower degree of delocalization/conjugation of the carbon-oxygen bond by disparate sharing of electron density. Nitrogen should be expected to act intermediately between sulfur and oxygen, as it is larger than oxygen, and more electronegative than sulfur. For a more detailed explanation of nitrogen’s contributions to these adducts, please see the section discussing benzimidazole ligands at the end of chapter 2. All of the reasons listed here for the observed bathochromic shift in sulfur- substituted derivatives are somewhat circumstantial, and some of the reasons for the decrease in the band gap must yet be formulated using more classes of compounds. It is likely that there are other factors that should be considered, and other tools in the field of theoretical chemistry will help to elucidate them.

An analysis of the calculated band gaps was in agreement with the experimental optical data obtained. As the band gap increased, the emission was hypsochromically shifted FIG 26. With a decrease in the band gap the emission was bathochromically shifted. These results coincide with the increase or reduction in the ionization potential. The change in the band gap was a result of either an increase in the LUMO or a decrease in the HOMO. The largest calculated band gap

[BF2(HPBI)] corresponded with the highest energy emission, whereas the smallest band gap [BF2(2,3-HNBT)] correlated with the lowest energy emission.

-1 -2.145 -2.515 -2.494 -2.482 -2.657 -2.635 -2.828 -3

∆= ∆= ∆= ∆= ∆= ∆= ∆= -3.720 -3.536 -3.421 -3.384 -5 -4.323 -4.182 -3.957 -6.018 -6.055 Energy (eV) Energy -6.214 -6.212 -6.469 -6.697 -6.615 -7

-9

-11 HPBO HPBT 1,2-HNBO 1,2-HNBT 2,3-HNBO 2,3-HNBT HPBI

Fig. 25 Calculated molecular orbital diagrams and band gaps (Δ) of all experimental compounds. Refer to Fig. 11 and the appendix for a comparison with the experimental emission maxima.

One of the unique features of our complexes is the electron withdrawing effect due to the fluorine substituents on boron. The effect of these substituents on the optical properties has been discussed. Confirmation of the fluorine’s involvement in adduct bonding was probed using computational means. Through the program VMOdes, we were able to estimate the level of electron density in any region of the complexes. The output files that were used in VMOdes were obtained through Gaussian03120 suite of programs using the B3LYP/6-311++G(d,p) level of theory with (PCM) solvent optimization. The Gaussian output files run on VMOdes were all single-point calculations of optimized structures. The electron density of the complexes remains dominated by the ligand, both in the HOMO and the LUMO, however, when focusing on the fluorines, or more broadly the boron-fluorine region, it is unambiguous that there is a high level of involvement in bonding in the HOMO FIG 26.

Ligand Oxygen Boron Fluorine

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(HPBO)

Ligand Sulfur Boron Fluorine

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(HPBT)

Ligand Oxygen Boron Fluorine

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(1,2-HNBO)

Ligand Sulfur Boron Fluorine

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(1,2-HNBT)

Ligand Oxygen Boron Fluorine

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(2,3-HNBO)

Ligand Sulfur Boron Fluorine

92 90 88 86 84 82 80 78 76 74

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

BF2(2,3-HNBT)

Fig. 26 Comparison of the molecular orbital distributions of the derivatives of interest. Data calculated using VMOdes 03119 at the B3LYP/6-311++G(d,p) with (PCM) solvent optimization. The top charts belongs to the oxygen heteroatom species, while the bottom charts belongs to the sulfur heteroatom species. In all cases the top plots correspond to the parent compounds (HPBO [top] and HPBT [bottom]) the middle is the 1,2-substituted species, 1,2-HNBO [top] and 1,2-HNBT [bottom]. And the bottom plots belong to the 2,3-substituted species, 2,3-HNBO [top] and 2,3-HNBT [bottom].

This involvement is shifted upon excitation as the level of electron density in the LUMO is greatly attenuated as the electron density is further centralized onto the ligand. The significance of these results is delineated when compared to complexes with di-phenyl rings as the substituents on boron. These complexes were a previous project in our lab, and the contribution of the ligand was not near as significant as in these complexes. In comparison to the di-phenyl substituted species our complexes showed several striking and anomalous differences. For instance, the sulfur substituted derivatives showed the expected change in sulfur’s involvement in electron density, whereas the oxygen derivatives showed the reverse, their significance in electron density decreased from HOMO to LUMO. The oxygen heteroatom made a substantial contribution in the HOMO and a very small contribution in the LUMO. This was the reverse of what was observed in the di- phenyl substituted derivatives. Broadly speaking, the sulfur heteroatom maintains its greater involvement in carrying electron density in the LUMO, which may be related to some of sulfur’s unique characteristics as discussed earlier.

Computational analysis of the strategic addition of EDGs

A number of calculation were run in order to determine the effects of addition of EDGs along the periphery of the scaffolding. Given the high degree of 59 electron delocalization across the entire ligand, the π-system functions as a short molecular wire, and so an addition of an electron donating group on either moiety should result in a bathochromic shift. The addition of the EDGs adds electron density which may decrease the level of rigidity and therefore results in lower energy emissions. Obviously certain positions will lead to more dramatic shifts than others FIG 27.

Fig. 27 Possible points for substitution of EWGs and EDGs on the ligand scaffold.

The model system that was chosen here was the phenyl-substituted oxazole derivative. Though the shifts in luminescence were not as dramatic as initially hoped, there was an unambiguous red-shift no matter where the EDG was placed. The results support the assertion that the EDGs could cause a red-shift on either moiety, however, the most dramatic shifts occurred upon addition of EDGs on the phenyl moiety FIG 28.

BF2_HPBO Position 1 Position 2 Position 3 Position 4

190 240 290 340 390 440 Wavelength (nm)

BF2_HPBO Position 5 Position 6 Position 7 Position 8

190 240 290 340 390

Wavelength (nm)

Fig. 28 Overlay of calculated excitation spectrum of BF2(HPBO) substituted with a methoxy group at various points. The spectrum is for the phenyl moiety (top) and azole moiety (bottom).

Phenyl Boronic acid derivatives A related aspect of this project was the synthesis of phenyl boronic acid derivatives as a means to investigate the nature of the boron-nitrogen dative bond. The complexes consist of a quinoline based ligand and phenyl boronic acid. Synthesis with azole based ligands was also attempted, and an adduct was formed as monitored by UV lamp, however, the complexes proved to be too unstable for isolation. Product is formed through a condensation reaction. Adducts formed with 10-hydroxybenzo[h] quinoline were found to be the most stable. Synthesis through azeotropic distillation was chosen, and equimolar amounts of the ligand and phenyl boronic acid were dissolved in toluene, and allowed to reflux for a period of time on a Dean-Stark apparatus. The removal of water through this method drove the reaction to completion. However, the final species isolated is still under speculation. HR-MS data supports the existence of at least two forms, the expected product, and a dimer. NMR data that has been collected has proven inconclusive, and it is speculated that the complex is very sensitive to hydrolysis, leading to the possibility of the existence of a hydrolysis product SCHEME 3. The phenyl boronic acid derivatives yield a large number of peaks in the aromatic region. The main confirmation of a complete reaction is the loss of the signal that appears in the region of ~11-12 ppm. This peak coincides with the –OH proton. After complexation with boron, this signal is lost. However, the evidence that suggests that multiple forms exist has made it difficult to elucidate the exact nature of the species. The phenyl boronic acid derivatives had a distinct fragmentation pattern. The dimers gave some of the clearest MS evidence of the success of their synthesis. The dimer could be easily reconstructed from the weights of the fragmentation pattern, showing a pattern in which the boron-oxygen-boron bond is severed leaving two 62 halves. The other prominent peak corresponded to the loss of a ligand. New methods of synthesis minimally involving Schlenck ware under argon to use of a glovebox may be requisite to isolate and characterize these difficult complexes. Both complexes are novel, however, the single phenyl boron species is more useful for the purpose of studying boron-nitrogen bonding, while the dimer provides better stability. After reaction the purified crystals exhibit green fluorescence, a significant hypsochromic shift from the original orange-red emission.

Scheme 3 Possible products from reaction of phenyl boronic acid and 10- Hydroxyphenylbenzo[h] quinoline.

Benzimidizole ligands

Our family of azole based ligands is contingent upon two unique features, the configuration of the naphthyl rings (1,2 versus 2,3-substituted) and the identity of the auxochromic heteroatom. Four of these are relatively easy to prepare, covering both the oxygen and sulfur heteroatoms. The optical properties related to the changes in luminescence upon changing heteroatoms makes substitution of the heteroatom an interesting feature of our ligands. Given that fact, work on creating another variety of ligand has been pursued, this ligand contains a nitrogen heteroatom. The choice of heteroatoms in these complexes is limited by some of the characteristics that are requisite for the position. For instance, oxygen and sulfur are ideal because they are able to achieve a stable configuration with two bonds. Another atom, phosphorus for example, would require more bonds, and would limit the viability as a ligand. Nitrogen is a feasible alternative as it prefers to be three- coordinate, therefore in the ligand of interest, hydroxynaphthyl benzimidizole, there is a double bond on one nitrogen, while the other nitrogen is bonded to a hydrogen. Our first attempts to bind boron to this ligand most likely failed due to the tautomerization between the two nitrogens. Boron forms a boron-ester bond with the oxygen, however, it may not form a dative bond with the nitrogen due to the tautomerism.121 This made the compound mercurial and it was unstable in solution. Attempts were made to synthesize under milder conditions by lowering the temperature, first to 0˚C, and then to -78.5 ˚C. However, neither method was able to realize any product. In order to ‘lock’ the ligand into a functional configuration one of the nitrogens must be bonded to something that will not undergo tautomerism. To this end, an ethyl group was bonded to the nitrogen bearing the proton. This allows the boron to bind to the oxygen and form a dative bond with the nitrogen. This work is 64 still in its seminal phase, however, given the facile means of substitution, a variety of substituents could be added to the nitrogen to tune the emissive properties.

Currently, an ethyl group has been added, allowing complexation with BF2. The formation of the complex hypsochromically shifted the emission from green (unbound ligand) to blue (adduct). See appendix for NMR confirmation. If one considers the trend in optical properties when comparing the various heteroatoms, substitution of a nitrogen atom should lead to a hypsochromic shift, when compared to sulfur, based solely on electronegativity. Computational means are being pursued to understand, and predict some of the properties. In comparison to the much larger sulfur atom, nitrogen has some distinct advantages. Nitrogen shares the size advantage with oxygen when contrasted with sulfur. In terms of electronegativity, nitrogen is an exact intermediary between sulfur and oxygen, so it would be expected to have increased covalent character compared to oxygen, but decreased covalency character bonding when compared to sulfur.

Ligand Nitrogen Boron Fluorine

83 81 79 77 75 73 71 69 67 65

0% 10% 20% 30% 40% 50% 60% 70% 80% 90% 100%

Fig. 29 Molecular orbital distribution for BF2(HPBI). Data calculated using VMOdes 03103 at the B3LYP/6-311++G(d,p) with (PCM) solvent optimization. Below,

Molecular Orbital Plots of BF2(HPBI). The HOMO is plotted in the middle column and the LUMO is plotted on the right.

The improved attributes of nitrogen are readily seen in the calculation of the molecular orbital distribution Fig. 29. The calculated role of the heteroatom is greater for nitrogen than either oxygen or sulfur.

The increased contribution of nitrogen is consistent with the expectation that nitrogen is the ideal heteroatom, due to its size, and electronegativity it can mitigate some of the shortcomings of the other two heteroatoms. The optical data gave somewhat unexpected results. There is a hypsochromic shift that exceeded even oxygen in energy, reaching a maximum of 333 nm, compared to the literature values of 344 nm and 364 nm for the equivalent oxygen and sulfur analogs.90 Similarly, the emission data or the nitrogen substituted adduct showed a high energy emission value of 385 nm, while the value of the oxygen derivative was only 408 nm, and sulfur 433 nm.90 When the optical data of the ligand is compared to the adduct, a significant increase in both energy and intensity is observed. When compared to the literature data for both the sulfur and oxygen substituted derivatives, a good agreement was observed. The excitation of the adduct was bathochromically shifted in comparison to the ligand. Likewise, the emission of the adduct was hypsochromically shifted from the ligand, most likely due to the increased rigidification as a result of the boron-nitrogen coordination. Both types of shifts were observed for the sulfur and oxygen derivatives as well. Additionally, the intensity of emission of the adduct was much higher than the intensity of emission 66 of the ligand, most likely due to the interruption of a ESIPT process, and perhaps, the increased rigidity again reduces the energy loss through vibration, increasing emission efficiency. The difference in both excitation and emission are small for nitrogen and oxygen, but the difference is significant. The high energy emission, while still maintaining involvement with the rest of the π-conjugate system, could be explained from the lower aromaticity in the model pyrrole, along with the larger band gap, when compared to thiophene. However, the nitrogen derivative would be expected to be intermediate in the energy of emission, as oxygen makes an even smaller contribution to the aromaticity of the ligand than does pyrrole. This is an illustration of the fact that there are still aspects of heteroatom bonding that need to be elucidated.

Computational Analysis of Anthracene Derivatives

Due to synthetic difficulties, experimental data of the optical properties of anthracene substituted azole ligands and their conjugate adducts have not been realized. However, computational means have been pursued in order to predict the expected properties of these species. The major feature of interest was the ability of the extended π-conjugated scaffold to induce red-shifted emission. Without experimental data to compare against, the findings must remain suspect. However, given the close agreement of the computational data and the experimental data in the naphthyl substituted derivatives, the data can be considered with some confidence. The results were in accord with expectations, and a large bathochromic shift in absorbance was observed as compared to the naphthyl substituted derivatives, reaching a maximum for BF2(2,3-HABT) of 554 nm FIG 29. The bathochromic shift that was observed is due, in part, to the stabilizing nature of the extension of the conjugated π-system. The increased conjugation diffuses electron density across a greater area. This is further supported by the agreement with nodal plane theory. The 1,2-substituted derivatives are not as bathochromically shifted as 67 are the 2,3-substituted derivatives. This behavior is the result of greater electron diffusion across the nodal plane. The orientation of the anthracene rings in the 1,2- substituted derivative, forms a nodal plane that precludes contribution of the anthracene in electron density. In the 2,3-substituted derivatives the nodal plane is across the entire anthracene ring and allows greater diffusion of electron density.

BF2(2,3-HABO) BF2(2,3-HABT) BF2(1,2-HABO) BF2(1,2-HABT)

230 330 430 530

Wavelength (nm)

Fig. 30 Overlay of calculated excitation spectrum of the anthracene derivatives.

Chapter 3 Experimental

Instrumentation

NMR

All NMR spectra were collected using a Varian 500 MHz NMR spectrometer. The temperature was maintained at 25°C throughout all data acquisition. During 1H- NMR, the instrument was operated at 500 MHz, while all 13C-NMR data was obtained operating at 126 MHz. All spectra were referenced using either the literature values of the solvent peaks (13C-NMR), or according to literature value of tetramethylsilane (TMS) [1H-NMR]. Please see appendix for all experimental spectra.

HR-MS-APCI

High pressure liquid chromatography (HPLC) was performed using a Thermo Fischer Scientific Finnigan LCQ Advantage Max instrument. This HPLC apparatus was coupled with atmospheric pressure chemical ionization (APCI) mass spectrometry (MS) using an Agilent 1100 series device. The mobile phase solvent used in all cases was tetrahydrofuran (THF).

Optical Spectroscopy

For general data collection, Excitation data was collected using a selection of concentrations ranging from 1.0x10-5 M to 1.0x10-6 M. The instrument used was a JASCO V-670 spectrophotometer with a 1nm band pass. Emission data was collected at the same concentrations using a Varian Eclipse Fluorescence Spectrophotometer. 69

For collection of data for quantum yield measurements, a Varian Cary 50 Series UV- Vis Spectrophotometer was used to collect excitation data. Emission data was collected using a Horiba Jobin Yvon Fluorolog 3 with spectrally corrected PMT and Xenon lamp source.

X-ray

All x-ray data was collected using a Rigaku Saturn Kappa CCD single crystal x-ray diffractometer. The crystal was mounted on a glass pipette, and attached by coating the crystal with epoxy. The raw data was processed using SHELX software package.122, 123

Materials

Both 2,3 and 1,2-Hydroxynaphthoic acid was purchased through TCI America, and used as purchased. The 10-hydroxybenzo[h]quinoline was purchased from Alfa Aesar. The parent ligands, 2-(2-hydroxyphenyl)benzothiazole and 2-(2- hydroxyphenyl)benzoxazole were purchased from Sigma-Aldrich along with the phosphorus trichloride. Some of the solvents, toluene, ether, cyclohexane, and ethyl acetate were purchased from Acros Organics without further drying, while dichloromethane, hexane, tetrahydrofurans, and acetonitrile were taken from a MBraun manual solvent purification system (MB-SPS). Silica gel was purchased from EMD Chemicals. All other chemicals were purchased through Acros Organics. Most images used in this thesis were obtained using ChemDraw Ultra 12.0.124 Any others were used with permission from the author.

Synthesis

Synthesis of 1,2-(Hydroxynaphthyl) benzoxazole (1,2-HNBO)

To a double neck 50 mL round bottom flask was added 14 mL anhydrous chlorobenzene. To this, 11 mmol (1.90 g) silicon tetrachloride was added, followed by 11 mmol (2.07 g) 1,2-hydroxynaphthoic acid. A drying tube was added as the solution was slowly heated, until the formation of HCl gas was visible to litmus. Heating was continued until change in color from yellow to pale green was observed. At color change, 10 mmol (1.10 g) 2-aminophenol was added in small portions over the period of a half hour. The vessel containing 2-aminophenol was rinsed with 3 mL anhydrous chlorobenzene, and the washing was added to the solution to bring the total volume to 17 mL. The solution was allowed to reflux for three hours. After reflux, saturated, aqueous NaHCO3 was added dropwise, quenching the reaction, being careful to not let the solution boil over. NaHCO3 was added until the solution was basic. Vacuum distillation was used to remove water and chlorobenzene. The brown solid was extracted with warm dichloromethane (3 x

100 mL) and filtered. The filtrate was dried over anhydrous MgSO4 and concentrated in vacuo. The crude brown product was purified using silica gel column chromatography (19:1 hexanes:ethyl acetate). The product was further purified through recrystallization from acetone to give fine white needles (27% yield). GC/MS was used for confirmation of the product. See appendix for experimental details.

Synthesis of 2,3-HNBO

The same procedure for the synthesis of 1,2-HNBO was followed for the synthesis of 2,3-HNBO. However 2.07 g (11 mmol) 2,3-hydroxynaphthoic acid was used. This ligand was also purified through silica gel column chromatography (19:1 hexane:ethyl acetate), and recrystallized from acetone to give orange needles (33% yield). See appendix for experimental details. 71

Synthesis of 1,2-(Hydroxynaphthyl) benzothiazole (1,2-HNBT)

1,2-hydroxynaphthoic acid (2.77 g, 10 mmol) was added to double neck 50 mL round bottom flask with boiling toluene (21 mL). To this was added 2- aminothiophenol (1.39 g, 11 mmol). These reactants were allowed to stir at reflux until dissolution was observed. After dissolution, the flask was cooled to 90°C, and

PCl3 (1.37g, 10 mmol) in 5 mL toluene was added dropwise. The reaction was allowed to stir at reflux for two hours. The flask was then cooled to 60°C and the product was precipitated by the addition of cold (4°C) methanol (40 mL). The filtrate was collected through filtration and washed with cold methanol. This product was recrystallized from acetone yellow plate-like crystals (81% yield). See appendix for experimental details.

Synthesis of 2,3-HNBT

The same procedure for the synthesis of 1,2-HNBT was followed for the synthesis of 2,3-HNBT. However 2.77 g (10 mmol) 2,3-hydroxynaphthoic acid was used. The resulting product was also recrystallized from acetone, and similarly gave yellow plate-like crystals (84% yield). See appendix for experimental details.

Synthesis of tosylated 2-(2-Hydroxyphenyl)-1H-benzimidazole (HPBI)

HPBI (1.05 g, 5mmol) and triethylamine (3.5 mL, 25 mmol) were dissolved in

10 mL of CH2Cl2. The mixture was cooled to 0°C, and 4-methylbenzenesulfonyl chloride [Tosyl Chloride] (1.03 g, 5 mmol) dissolved in 10 mL CH2Cl2 was added dropwise with stirring. The mixture was allowed to stir for 1 hour. The reaction was quenched by the addition of 50 mL water, and was subsequently extracted with 72

CH2Cl2 (3 x 30 mL). The organic layer was collected and dried over MgSO4. The solvent was removed in vacuo, and the resulting yellow oil was purified by flash column chromatography using hexane/ethyl acetate (6:4). The solvent was removed to give a white powder (1.71 g, 90%).121 The product and purity were confirmed through GC/MS analysis. Refer to appendix for experimental details.

Synthesis of ethylated HPBI

In an argon atmosphere, Ts-HPBI (760 mg, 2 mmol) and tetrabutylammonium bromide [TBAB] (64.5 mg, 0.2 mmol) were dissolved in 10 mL DMSO. In a three- neck round bottom flask. Aqueous K2CO3 (2M, 3.3 mL) was added dropwise and then stirred for 30 minutes. Ethyl bromide (0.16 mL, 2.2 mmol) was added dropwise. A balloon attached to a syringe was added to the reflux column, and the flask was heated to 70°C and stirred at this temperature for 12 hours. The flask was cooled to room temperature and the solution was poured into 100 mL water, and extracted with dichloromethane (3 x 30 mL). The organic layer was dried over

MgSO4. The solvent was removed in vacuo, and the residue was a pale yellow oil. This oil contained the crude ethylated Ts-HPBI. The crude product was dissolved in 20 mL methanol, and 5 mL aqueous NaOH (2M, 5 mL) was added. The mixture was refluxed for 2 hours and then cooled to room temperature and neutralized by the addition of dilute aqueous HCL (0.1 M). The mixture was extracted by dichloromethane (3 x 30 mL). The organic layer was dried over MgSO4. The solvent was removed and the resulting yellow oil was purified by flash column chromatography using hexane/ethyl acetate (8:1).121 The product was a white, blue fluorescent powder (75% Yield). The confirmation of purity was through GC/MS.

Synthesis of 1,2-(Hydroxynaphthyl) benzimidazole (1,2-HNBI)

To 20 mL of 85% H3PO4, was added P2O5 (31.11 g, 0.22 mol) portionwise. The mixture was heated to 110 ˚C for 2 hours with stirring, to form homogeneous polyphosphoric acid (PPA). 1,2-Hydroxynaphthoic acid (1.882 g, 10 mmol) was added, along with o-phenylenediamine, and the temperature of the mixture was increased to 200˚C. The reaction was stirred for four hours at this temperature. The reaction was quenched by pouring the solution in to ice water under stirring. The pH of the mixture was adjusted to pH 7 by addition of aqueous NaOH (2 M). The precipitate was collected as a deep purple powder. Testing with TLC and GC/MS revealed no starting materials. The crude mixture exhibited blue fluorescence under UV lamp. No product was isolated.

Synthesis of BF2(1,2-HNBO)

1,2-HNBO (720 mg, 2.33 mmol) was added to a three-neck round bottom flask, fitted with a valve, and an airtight septum. To this was added 33 mL DCM, and an atmosphere of nitrogen was introduced. Boron trifluoride diethyl etherate (7.7 mmol, 1.1 g) and triethylamine (13.87 mmol , 1.4g) were added via syringe. The reaction produced hydrogen fluoride gas, which was visible in the flask. The solution was allowed to stir for a week, during which time the insoluble product precipitated. The precipitate was collected and the filtrate placed at 4°C for further precipitation. The crude product, according to NMR, consisted of both the product and the deprotonated ammonium base. The product was dissolved in CH2Cl2 and treated with Na2CO3, in order to regenerate the ammonium base. The solution was allowed to stir for two hours. The Na2CO3 was filtered and any residual product was extracted with CH2Cl2. Hexane was added to the CH2Cl2 solubilized product and the

CH2Cl2 was removed in vacuo. Purified product precipitated out as white crystals, and was collected. The hexane containing the regenerated base was washed with

CH2Cl2 until the hexane layer had no visible fluorescence to UV lamp, the CH2Cl2 74 layer was collected and the hexane layer was discarded. The resulting solution was concentrated in vacuo. This solution was filtered via pastuer pipette, and placed at 4°C in order to crystallize. The resulting product crystallized as white needles. (480 mg, 66.6% yield, M.P. 260°C). The product and purity were confirmed through 1H- NMR. Refer to the appendix for experimental details.

Synthesis of BF2(1,2-HNBT)

Using the same procedure as for BF2(1,2-HNBO), BF2(1,2-HNBT) was synthesized from 1,2-HNBT (646 mg, 2.33 mmol) and Boron trifluoride diethyl etherate (7.7 mmol, 1.1 g), in DCM (33mL) and triethylamine (1.4 g, 13.87 mmol). The crude product was purified by recrystallization from first hexane, and then DCM, to give greenish-yellow needle-like crystals (414.84 mg, 54.8% yield, M.P. 263.5°C). X-ray quality crystals were obtained through dissolving the resulting crude mixture in boiling hexane and filtering via Pasteur pipette into a heated vial. The vial was sealed and left to evaporate, leaving yellow block crystals.

Synthesis of BF2(2,3-HNBO)

Using the same procedure as for BF2(1,2-HNBO), BF2(2,3-HNBO) was synthesized from 2,3-HNBO (720 mg, 2.33 mmol) and Boron trifluoride diethyl etherate (7.7 mmol, 1.1 g), in DCM (33mL) and triethylamine (1.4 g, 13.87 mmol). The crude product was purified by recrystallization from first hexane, and then DCM, to give pale green needle-like crystals (436 mg, 60.5% yield, M.P. 285.7°C ). Both product and purity were confirmed through 1H-NMR.

Synthesis of BF2(2,3-HNBT)

Using the same procedure as for BF2(1,2-HNBO), BF2(2,3-HNBT) was synthesized from 2,3-HNBT (646 mg, 2.33 mmol) and Boron trifluoride diethyl etherate (7.7 mmol, 1.1 g), in DCM (33mL) and triethylamine (1.4 g, 13.87 mmol). The crude product was purified by recrystallization from first hexane, and then DCM, to give yellow needle-like crystals (334 mg, 44.1% yield, M.P. 293°C). Both product and purity were confirmed through 1H-NMR.

Synthesis of BF2(HBQ)

Using the same procedure as for BF2(1,2-HNBO), BF2(HBQ) was synthesized from HBQ (566 mg, 2.33 mmol) and Boron trifluoride diethyl etherate (7.7 mmol, 1.1 g), in DCM (33mL) and triethylamine (1.4 g, 13.87 mmol). The crude product was purified by recrystallization from first hexane, and then DCM, to give yellow needle-like crystals. (226 mg, 40%, M.P. 194°C). Both the product and the purity were confirmed by 1H-NMR.

Synthesis of BF2(HPBI)

Using the same procedure as for BF2(HBQ) , BF2(HPBI) was synthesized from ethylated HPBI (1.37 mmol, 326 mg) and Boron trifluoride diethyl etherate (4.53 mmol, 643 mg), in DCM (33mL) and triethylamine (8.16 mmol, 826 mg). The crude product was purified by recrystallization from first hexane, and then DCM, to give white needle-like crystals (225 mg, 57.3% yield, M.P. 226°C). The product was confirmed through 1H-NMR.

Synthesis of BPh(HBQ)2

To a round bottom flask containing toluene (60 mL), was added 10- hydroxybenzo[h]quinoline (390 mg, 2 mmol) and phenyl boronic acid (122 mg, 1 76 mmol). This was fitted with a Dean-Stark apparatus, and refluxed for 12 hours. The solvent was removed in vacuo, leaving a yellow powder. This was dissolved in DCM and an equal part ether was added. The solution was cooled to 4°C for 48 hours, during which time yellow crystals precipitated out of solution. 1H-NMR analysis revealed that the starting materials were used up, however, the final nature of the product was unclear.

References

1) Niedenzu, K. , Dawson, J.W. Boron-Nitrogen Compounds, 1st Ed.; Springer- Verlag: Berlin, 1965. 2) Muetterties, E.L. The Chemistry of Boron and its Compounds, 1st Ed.; John Wiley and Sons, Inc.: New York, 1967 3) Cotton, F.A., Wilkinson, G., Murillo, C.A., Bochmann, M. Advanced Inorganic Chemistry, 6th Ed.; Wiley Interscience: New York, 1999. 4) Soloway, A.H.; Tjarks, W.,; Barnum, B.A.,; Rong, F.,; Barth, R.F.; Codogni, I.M.; and Wilson, J.G. The Chemistry of Neutron Capture Therapy. Chem. Rev. 1998, 98, 1515 5) Morris, J. Chemistry in Britain. 1991, 27, 331. 6) Hawthorne, M.F. Angew, Chem. Int. Ed. Engl. 1993, 32, 950. 7) Barth, R.F., Coderre, J.A., Graca, M., Vicente, H. Boron Neutron Capture Therapy of Cancer: Current Status and Future Prospects. Clin. Cancer. Res. 2005, 11, 3987. 8) Angyan, J.G. J. Molec. Struc. (Theochem.) 2000, 501, 379. 9) Frenking, G.; Fau, S.; Marchand, C.M.; Grutzmacher, H. J. Am. Chem. Soc. 1997, 119, 6648. 10) Shriver and Atkins Inorganic Chemistry 4th Ed.; Oxford University Press 11) Mercier, H.P.A.; Moran, M.; Schrobilgen, G.J.; Suontamo, R.J. Journal of Fluorine Chemistry. 2004, 125, 1563. 12) Alkorta, I.; Elguero, J.; Del Bene, J.E.; Mo, O.; Yanez, M. Chem. Eur. J. 2010, 16, 11897. 13) Del Bene, J.E.; Alkorta, I.; Elguero, J.; Mo, O.; Yanez, M. J. Phys. Chem. A. 2010, 114, 12775. 14) Brown, H.C.; Schlesinger, H.I.; Burg, A.B. J. Am. Chem. Soc. 1939, 61, 673. 15) Schlesinger, H.I.; Brown, H.C. J. Am. Chem. Soc. 1953, 75, 219.

16) Schlesinger, H.I.; Brown, H.C.; Finholt, A.E.; Gilbreath, J.R.; Hoekstra, H.R.; Hyde, E.K. J. Am. Chem. Soc. 1953, 75, 215. 17) Carey, F.A.; Sundberg, R.J. Advanced Organic Chemistry: Part B: Reactions and Synthesis 5th Ed. Springer, 2007. 18) Miyaura, N.; Suzuki, A. Chem. Rev. 1995, 95, 2457. 19) Yin, L.; Liebsher, J. Chem. Rev. 2007, 107, 133. 20) Boer, W.D. Active Matrix Liquid Crystal Displays; Newnes: Oxford, 2005. 21) Physical Properties of Liquid Crystals: Nematics; Dunmur, D.; Fukuda, A.; Luckhurst, G., Eds. Inspec/lee: London, 2001. 22) Bernanose, A.; Comte, M.; Vouaux, P. J. Chim. Phys. 1953, 50, 64. 23) Sano, M.; Pope, M.; Kallmann, H. J. Chem. Phys. 1965, 43, 2920. 24) Helfrich, W.; Schneider, W. Phys. Rev. Lett. 1965, 14, 229. 25) Bolto, B.A.; McNeill, R.; Weiss, D.E. Australian Journal of Chemistry, 1963, 16, 1090. 26) McGinness, J.; Corry, P.; Proctor, P. Science, 1974, 183, 853. 27) De Surville, R.; Jozefowicz, M.; Yu, L.T.; Pepichon, J.; Buvet, R. Electrochimica Acta. 1968, 13, 1451. 28) Greenham, N.C.; Moratti, S.C.; Bradley, D.D.C.; Friend, R.H.; Holmes, A.B. Nature. 1993, 365, 628. 29) Marsella, M.J.; Fu, D.K.; Swager, T.M. Adv. Mater. 1995, 7, 145. 30) Ball, D.W. Physical Chemistry, 1st Ed.; Brooks/Cole; Pacific Grove, CA., 2003. 31) Fret and Film Techniques, Gadella, T.W.J., Ed.; Elsevier: Oxford, 2009. 32) Bulovic, V.; Deshpande, R.; Thompson, M.E.; Forrest, S.R. Chem. Phys. Lett. 1999, 308, 317. 33) Organic Light Emitting Materials and Devices; Li, Z., Meng, H., Eds.; CRC Press, Taylor & Francis Group: Boca Raton, 2007. 34) Organic Electroluminescence, Kafafi, Z., Ed.; CRC Press, Taylor & Francis Group: Boca Raton, 2005. 35) Förster, T. Discussions Faraday Soc. 1959, 27, 7. 79

36) Dexter, D.L. J. Chem. Phys. 1953, 21, 836. 37) Closs, G.L.; Johnson, M.D.; Miller, J.R.; Piotrowiak, P. J. Am. Chem. Soc. 1989, 111, 3751. 38) Atilgan, S.; Ekmekci, Z.; Lale Dogan, A.; Guc, D.; Akkaya, E.U. Chem. Commun. 2006, 4398. 39) Yogo, T.; Urano, Y.; Ishitsuka, Y.; Maniwa, F.; Nagano, T. J. Am. Chem. Soc. 2005, 127, 12162. 40) Burghart, A; Thoresen, L.H.; Chen, J.; Burgess, K.; Bergström, F.; Johansson, L.B. Chem. Commun. 2000, 2203. 41) Ulrich, G.; Ziessel, R.; Harriman, A. Angew. Chem. Int. Ed. 2008, 47, 1184. 42) Kollmannsberger, M.; Rurack, K.; Resch-Genger, U.; Daub, J. J. Phys.Chem. A. 1998, 102, 10211. 43) Monsma, F.J.; Barton, A.C.; Kang, H.C.; Brassard, D.L.; Haughland, R.P.; Sibley, D.R. J. Neurochem. 1989, 52, 1641. 44) Loudet, A.; Burgess, K. Chemical Reviews. 2007, 107, 4891. 45) Arroyo, I.J.; Hu, R.; Merino, G.; Zhong Tang, B.; Peña-Cabrera, E. J. Org. Chem. 2009, 74, 5719. 46) Schmitt, A.; Hinkeldey, B.; Wild, M.; Jung, G. J. Fluoresc. 2009, 19, 755. 47) Chen, H.-Y.; Chi, Y.; Liu, C.-S.; Yu, J.-K.; Cheng, Y.-M.; Chen, K.-S.; Chou, P.- T.; Peng, S.-M.; Lee, G.-H.; Carty, A.J.; Yeh, S.-J.; Chen, C.-T. Adv. Funct. Mater. 2005, 15, 567. 48) Chen, J.-L.; Lin, C.-H.; Chen, J.-H.; Chi, Y.; Chiu, Y.-C.; Chou, P.T.; Lai, C.-H.; Lee, G.-H.; Carty, A.J. Inorg. Chem. 2008, 47, 5154. 49) Liddle, B.J.; Silva, R.M.; Morin, T.J.; Macedo, F.P.; Shukla, R.; Lindeman, S.V.; Gardinier, J.R. J. Org. Chem. 2007, 72, 5637. 50) Bonardi, L.; Kanaan, H.; Camerel, F.; Jolinat, P.; Retailleau, P.; Ziessel, R. Adv. Funct. Mater. 2008, 18, 401. 51) Hepp, A.; Ulrich, G.; Schmechel, R.; von Seggern, H.; Ziessel, R. Synthetic Metals. 2004, 146, 11.

52) Tang, C.W.; VanSlyke, S.A. Appl. Phys. Lett. 1987, 51, 913. 53) Wu, Q.; Esteghamatian, M.; Nan-Xing, H.; Popovic, Z.; Enright, G.; Breeze, S.R.; Wang, S. Angew. Chem. Int. Ed. 1999, 109, 2055. 54) Li, H.; Jäkle, F. Macromolecules. 2009, 42, 3448. 55) Hou, Q.; Zhang, H.; Wang, Y.; Jiang, S. Journal of Luminescence. 2007, 126, 447. 56) Qin, Y.; Kiburu, I.; Shah, S.; Jäkle, F. Org. Lett. 2006, 8, 5227. 57) Kappaun, S.; Rentenberger, S.; Pogantsch, A.; Zojer, E.; Mereiter, K.; Trimmel, G.; Saf, R.; Möller, K.C.; Stelzer, F.; Slugovc, C. Chem. Mater. 2006, 18, 3539. 58) Wu, Q.; Esteghamatian, M.; Nan-Xing, H.; Popovic, Z.; Enright, G.; Tao, Y.; D’Iorio, M.; Wang, S. Chem. Mater. 2000, 12, 79. 59) Cui, Y.; Wang, S. J. Org. Chem. 2006, 71, 6485. 60) Yamaguchi, S.; Shirasaka, T.; Tamao, K. Organic Letters. 2000, 26, 4129. 61) Entwistle, C.D.; Marder, T.B. Chem. Mater. 2004, 16, 4574. 62) Miller, D.S.; Leffler, J.E. J. Phys. Chem. 1970, 74, 2571. 63) Yuan, Z.; Collings, J.C.; Taylor, N.J.; Marder, T.B.; Jardin, C.; Halet, J. Journal of Solid State Chemistry, 2000, 154, 5. 64) Jia, W.; Song, D.; Wang, S. J. Org. Chem. 2003, 68, 701. 65) Yamaguchi, S.; Akiyama, S.; Kohei, T. J. Am. Chem. Soc. 2001, 123, 11372. 66) Noda, T.; Ogawa, H.; Shirota, Y. Adv. Mater. 1999, 11, 283. 67) Torssell, K. Ark. Kemi. 1957, 10, 507. 68) Snyder, H.; Lennarz, W. J. Am. Chem. Soc. 1958, 80, 835. 69) Adamczyk-Wozniak, A.; Cyranski, M.K.; Zubrowska, A.; Sporzynski, A. Journal of Organometallic Chemistry. 2009, 694, 3533. 70) Hawkins, R.T.; Lennarz, W.J.; Snyder, H.R. J. Am. Chem. Soc. 1960, 80, 3053. 71) Zhdankin, V.V.; Persichini, P.J.; Zhang, L.; Fix, S.; Kiprof, P.; Tetrahedron Lett. 1999, 40, 6705. 72) Yamamoto, Y.; Ishii, J.; Nishiyama, K.; Itoh, J. J. Am. Chem. Soc. 2005, 127, 9625.

73) Tessier, P.E.; Nguyen, N.; Clay, M.D.; Fallis, A.G. Org. Lett. 2005, 7, 767. 74) Tan, Y.L.; White, A.J.P.; Widdowson, D.A.; Wilhelm, R.; Williams, D.J.; J. Chem. Soc., Perkin Trans. 1. 2001, 3269. 75) Nicolaou, K.C.; Li, H.; Boddy, C.N.C.; Li, H.; Koumbis, A.E.; Hughes, R.; Natarajan, S.; Jain, N.F.; Ramanjulu, J.M.; Bräse, S.; Solomon, M.E. Chem. Eur. J. 1999, 5, 2602. 76) Dowlut, M.; Hall, D.G. J. Am. Chem. Soc. 2006, 128, 4226. 77) Bérubé, M.; Dowlut, M.; Hall, D.G. J. Org. Chem. 2008, 73, 6471. 78) Ricardo, A.; Frye, F.; Carrigan, M.A.; Tipton, J.D.; Powell, D.H.; Benner, S.A. J. Org. Chem. 2006, 71, 9503. 79) Baker, S.J.; Zhang, Y.-K.; Akama, T.; Lau, A.; Zhou, H.; Hernandez, V.; Mao, W.; Alley, M.R.K.; Sanders, V.; Plattner, J.J. J. Med. Chem. 2006, 49, 4447. 80) Hui, X.; Baker, S.J.; Wester, R.C.; Barbadillo, S.; Cashmore, A.K.; Sanders, V.; Hold, K.M.; Akama, T.; Zhang, Y-K.; Plattner, J.J.; Maibach, H.I. J. Pharm. Sci. 2007, 96, 2622. 81) Baker, S.J.; Hui, X.; Maibach, H.I.; Ann. Rep. Med. Chem. 2005, 40, 323. 82) Baker, S.J.; Zhang, Y.-K.; Akama, T.; Wheeler, C.; Plattner, J.J.; Rosser, R.M.; Reid, R.P.; Nixon, N.S. J. Labelled Compd. Rad. 2007, 50, 245. 83) Gunasekara, D.S.; Kiprof, P.; Zhdankin, V.V.; Reddy, M.V. Abstracts, 37th Great Lakes Regional Meeting of the American Chemical Society. 2006. 84) Beck, R.M. Fluorescent Materials Derived from Benzoboroxole: MS Thesis UMD, 2004, pp. 10-19. 85) Brown, H.; Bhatt, N.; Somayaji, V. Organometallics 1983, 2, 1311. 86) Carlson, J.C. Strategic Synthetic Color Tuning of a Family of Simple Azole Based Luminescent Organoboron Compounds: MS Thesis UMD, 2009. 87) Wakamiya, A.; Taniguchi, T.; Yamaguchi, S. Angew. Chem. Int. Ed. 2006, 45, 3170. 88) Monkman, A.P.; Pålsson, L.-O.; Higgins, R.W.T.; Wang, C.; Bryce, M.; Batsanov, A.S.; Howard, J.A.K. J. Am. Chem. Soc. 2002, 124, 6049.

89) Yamaguchi, S.; Shirasaka, T.; Akiyama, S.; Tamao, K. J. Am. Chem. Soc. 2002, 124, 8816. 90) Kwak, M.J.; Kim, Y. Bull. Korean Chem. Soc. 2009, 12, 2865. 91) Matsumoto, T.; Urano, Y.; Shoda, T.; Kojima, H.; Nagano, T. Org. Lett. 2007, 9, 3375. 92) Guo, H.; Jing, Y.; Yuan, X.; Ji, S.; Zhao, J.; Lib, X.; Kanc, Y. Org. Biomol. Chem. 2011, 9, 3844. 93) Bozdemir, O.A.; Guliyev, R.; Buyukcakir, O.; Selcuk, S.; Kolemen, S.; Gulseren, G.; Nalbantoglu, T.; Boyaci, H.; Akkaya, E.U. J. Am. Chem. Soc. 2010, 132, 8029. 94) Nagaoka, S.; Nagashima, U.; Ohta, N.; Fujita, M.; Takemura, T. J. Phys. Chem. 1988, 92, 166. 95) Nagaoka, S.; Nagashima, U. J. Phys. Chem. 1990, 94, 1425. 96) Nagaoka, S.; Kusunoki, J.; Fujibuchi, T.; Hatakenaka, S.; Mukai, K.; Nagashima, U. Journal of Photochemistry and Photobiology A: Chemistry. 1999, 122, 151. 97) Haaland, A. Angew. Chem. Int. Ed. Engl. 1989, 28, 992. 98) Gaussian 09, Revision A.1, M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A. Robb, J. R. Cheeseman, G. Scalmani, V. Barone, B. Mennucci, G. A. Petersson, H. Nakatsuji, M. Caricato, X. Li, H. P. Hratchian, A. F. Izmaylov, J. Bloino, G. Zheng, J. L. Sonnenberg, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O. Kitao, H. Nakai, T. Vreven, J. A. Montgomery, Jr., J. E. Peralta, F. Ogliaro, M. Bearpark, J. J. Heyd, E. Brothers, K. N. Kudin, V. N. Staroverov, R. Kobayashi, J. Normand, K. Raghavachari, A. Rendell, J. C. Burant, S. S. Iyengar, J. Tomasi, M. Cossi, N. Rega, J. M. Millam, M. Klene, J. E. Knox, J. B. Cross, V. Bakken, C. Adamo, J. Jaramillo, R. Gomperts, R. E. Stratmann, O. Yazyev, A. J. Austin, R. Cammi, C. Pomelli, J. W. Ochterski, R. L. Martin, K. Morokuma, V. G. Zakrzewski, G. A. Voth, P. Salvador, J. J. Dannenberg, S. Dapprich, A. D. Daniels, Ö. Farkas, J. B. Foresman, J. V. Ortiz, J. Cioslowski, and D. J. Fox, Gaussian, Inc., Wallingford CT, 2009.

99) Jacquemin, D.; Perpète, E.A.; Scuseria, G.E.; Ciofini, I.; Adamo, C.J. Chem. Theory Comput. 2008, 4, 123. 100) Jacquemin, D.; Perpéte, E.A.; Ciofini, I. Adamo, C.J. Acc. Chem. Res. 2009, 42, 326. 101) Yanai, T.; Tew, D.P.; Handy, N.C. Chem. Phys. Lett. 2004, 393, 51. 102) Ebbing, D.D.; Gammon, S.D. In General Chemistry; Houghton-Mifflin: Boston, 2002, 7th Ed. p. 382. 103) Joule, J.A.; Mills, K. In Heterocyclic Chemistry; Blackwell Science Press: Oxford, 2000, 4th Ed. p 6. 104) Schmidt, M.W.; Truong, P.N.; Gordon, M.S. J. Am. Chem. Soc. 1987, 109, 5217. 105) Pitzer, K.S. J. Am. Chem. Soc. 1948, 70, 2140. 106) Mulliken, R.S. J. Am. Chem. Soc. 1950, 72, 4493. 107) Hinchliffe, A.; Soscún, H.J. J. Molec. Struc. (Theochem). 1995, 331, 109. 108) Wulfsberg, G. Inorganic Chemistry; University Science Books: Sausalito, CA, 2000, 1st Ed. p 93. 109) Kutzelnigg, W. Angew. Chem. Int. Ed. Eng. 1984, 23, 272. 110) Zumdahl, S.S. Chemistry, 1st ed.; Heath: Lexington, MA, 1986. 111) Gámez, J. A.; Serrano-Andrés, L.; Yáñez, M. Phys. Chem. Chem. Phys. 2010, 12, 1042. 112) Guo, X.; He, J.; Liu, Z.; Tian, Y. Phys. Rev. B. 2006, 73, 104115. 113) Lupinetti, A.J.; Fau, S.; Frenking, G.; Strauss, S.H. J. Phys. Chem. A. 1997, 101, 9551. 114) Thomas, T.D.; Saethre, L.J.; Børve, K.J.; Bozek, J.D.; Huttula, M.; Kukk, E. J. Phys. Chem. A. 2004, 108, 4983. 115) Jursic, B.S. J. Molec. Struc. (Theochem). 1999, 468, 171. 116) Balaban, A.T.; Oniciu, D.C.; Katritzky, A.R. Chem. Rev. 2004, 104, 2777. 117) Cook, M.J.; Katritzky, A.R.; Linda, P. Adv. Heterocyclic Chem. 1974, 17, 255. 118) Simkin, Y.A.; Minkin, V.I.; Glukhovtsev, M.N. Adv. Heterocyclic Chem. 1993,

56, 303. 119) VMOdes Program, Revision A 7.2 V. N. Nemykin, P. Basu; University of Minnesota Duluth and Duquesne University; 2001, 2003, 2005. 120) Gaussian 03, Revision C.02, Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Montgomery, Jr., J. A.; Vreven, T.; Kudin, K. N.; Burant, J. C.; Millam, J. M.; Iyengar, S. S.; Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani, G.; Rega, N.; Petersson, G. A.; Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.; Li, X.; Knox, J. E.; Hratchian, H. P.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Zakrzewski, V. G.; Dapprich, S.; Daniels, A. D.; Strain, M. C.; Farkas, O.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul, A. G.; Clifford, S.; Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.; Wong, M. W.; Gonzalez, C.; and Pople, J. A.; Gaussian, Inc., Wallingford CT, 2004. 121) Xu, H.; Xu, Z.-F.; Yue, P.-F.; Wang, B.; Jia, L.-W.; Li, G.-M.; Sun, W.-B.; Zhang, J.-W. J. Phys. Chem. C. 2008, 112, 15517. 122) Sheldrick, G.M. Acta Cryst. 2010, D66, 479. 123) Sheldrick, G.M. Acta Cryst. 2008, A26, 112. 124) ChemDraw Ultra 12.0

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aro_Kiprof_1_25_2011_35-02_1 1/25/2011 7:08:31 AM C.T.(Kiprof): 35-02; 2 uL 100% THF; 0.5 mL/min; split 1:1 to APCI

RT: 0.00 - 4.63 0.71 NL: 100 8.27E7 TIC M S 80 aro_Kiprof_1 _25_2011_35- 60 02_1

20 Relative Abundance 0.10 0.16 0.36 0.46 0.58 0.77 0.82 1.01 1.13 1.33 1.41 1.49 1.65 1.80 1.88 2.11 2.19 2.33 2.42 2.64 2.72 2.90 3.09 3.18 3.24 3.41 3.54 3.73 3.83 3.92 4.11 4.18 4.34 4.51 4.62 0 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0 2.2 2.4 2.6 2.8 3.0 3.2 3.4 3.6 3.8 4.0 4.2 4.4 4.6 Time (min)

aro_Kiprof_1_25_2011_35-02_1 #151-158 RT: 0.69-0.72 AV: 8 NL: 2.01E7 T: + c APCI corona Full ms [50.00-600.00] 361.53 100

60 290.44 40 362.64 20 288.39 291.43 576.15 Relative Abundance 132.96 144.93 171.85 206.41 219.02 248.22 262.15 314.97 358.45 363.62 384.49 413.90 448.15 466.70 498.92 531.36 542.77 585.64 0 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 480 500 520 540 560 580 600 m/z

112

Fig. 83 Chromatogram and MS data from HR-APCI-MS of BF2(1,2-HNBO). The mobile phase solvent was THF.

aro_Kiprof_1_25_2011_42-01_1 1/25/2011 7:29:32 AM C.T.(Kiprof): 42-01; 2 uL 100% THF; 0.5 mL/min; split 1:1 to APCI

RT: 0.00 - 0.84 0.69 NL: 100 0.70 1.81E8 0.69 TIC M S 80 aro_Kiprof_1 _25_2011_42- 60 01_1

20 Relative Abundance 0.03 0.04 0.08 0.12 0.15 0.16 0.18 0.20 0.24 0.25 0.29 0.32 0.33 0.36 0.38 0.40 0.44 0.47 0.49 0.52 0.54 0.57 0.59 0.62 0.64 0.78 0.82 0 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 Time (min)

aro_Kiprof_1_25_2011_42-01_1 #151-158 RT: 0.68-0.71 AV: 8 NL: 4.41E7 T: + c APCI corona Full ms [50.00-600.00] 377.50 100

80 306.41 60

20 307.31 378.58

Relative Abundance 72.90 87.28 109.05 142.47 165.10 176.78 206.47 220.85 247.10 278.22 292.21 309.54 332.27 375.19 380.65 404.82 430.08 451.95 475.80 499.29 518.02 546.72 563.90 585.33 0 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 480 500 520 540 560 580 600 m/z

Fig. 84 Chromatogram and MS data from HR-APCI-MS of BF2(1,2-HNBT). The mobile phase solvent was THF.

113

aro_Kiprof_1_25_2011_43-01_1 1/25/2011 7:03:44 AM C.T.(Kiprof): 43-01; 2 uL 100% THF; 0.5 mL/min; split 1:1 to APCI

RT: 0.00 - 0.88 0.68 0.69 NL: 100 1.55E8 TIC M S 80 aro_Kiprof_1 0.70 _25_2011_43- 60 01_1

20 0.74

Relative Abundance 0.76 0.03 0.05 0.07 0.09 0.11 0.15 0.17 0.18 0.21 0.24 0.27 0.30 0.33 0.35 0.40 0.41 0.43 0.45 0.49 0.51 0.54 0.56 0.58 0.61 0.62 0.80 0.83 0.84 0 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 0.85 Time (min)

aro_Kiprof_1_25_2011_43-01_1 #161-167 RT: 0.67-0.70 AV: 7 NL: 2.45E7 T: + c APCI corona Full ms [50.00-500.00] 361.50 100 262.26

80 290.36 60

40 348.42 362.47 263.26 20 291.39 413.92

Relative Abundance 58.80 74.43 92.87 102.15 110.95 133.11 145.05 169.00 198.11 217.40 234.22 248.18 311.42 347.27 363.56 403.70 427.24 444.48 459.59 468.29 496.07 0 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 480 500 m/z

Fig. 85 Chromatogram and MS data from HR-APCI-MS of BF2(2,3-HNBO). The mobile phase solvent was THF.

aro_Kiprof_1_25_2011_44-01_1 1/25/2011 7:31:51 AM C.T.(Kiprof): 44-01; 2 uL 100% THF; 0.5 mL/min; split 1:1 to APCI

RT: 0.00 - 0.88 0.70 NL: 100 2.17E8 TIC M S 80 aro_Kiprof_1 _25_2011_44- 60 01_1

20 Relative Abundance 0.02 0.05 0.09 0.11 0.13 0.16 0.18 0.19 0.24 0.27 0.31 0.33 0.35 0.37 0.40 0.44 0.47 0.48 0.51 0.54 0.56 0.60 0.64 0.65 0.78 0.81 0.83 0.85 0 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 0.85 Time (min)

aro_Kiprof_1_25_2011_44-01_1 #154-157 RT: 0.70-0.71 AV: 4 NL: 4.92E7 T: + c APCI corona Full ms [50.00-600.00] 377.48 100 306.45 80

40 278.22 376.41 307.43 20 279.20 378.57

Relative Abundance 102.38 115.38 145.10 191.11 219.63 250.47 309.27 347.80 375.02 406.61 416.53 454.44 464.28 494.72 522.12 553.04 563.36 582.54 0 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 480 500 520 540 560 580 600 m/z

Fig. 86 Chromatogram and MS data from HR-APCI-MS of BF2(2,3-HNBT). The mobile phase solvent was THF.

114

aro_Kiprof_1_25_2011_10-03_1 1/25/2011 7:21:46 AM C.T.(Kiprof): 10-03; 2 uL 100% THF; 0.5 mL/min; split 1:1 to APCI

RT: 0.00 - 1.00 0.68 NL: 100 0.69 1.66E8 TIC M S 80 aro_Kiprof_1 _25_2011_10- 60 03_1 0.70 40

20 Relative Abundance 0.01 0.06 0.11 0.12 0.15 0.19 0.22 0.24 0.28 0.29 0.34 0.38 0.40 0.42 0.47 0.49 0.53 0.57 0.58 0.62 0.75 0.79 0.84 0.87 0.89 0.94 0.95 0.98 0 0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 0.85 0.90 0.95 Time (min)

aro_Kiprof_1_25_2011_10-03_1 #148-153 RT: 0.67-0.69 AV: 6 NL: 5.03E7 T: + c APCI corona Full ms [50.00-600.00] 295.62 100

60 224.41 40 296.65 20 222.45 225.34 445.03

Relative Abundance 88.01 101.94 130.07 144.81 163.08 196.44 250.28 291.92 297.71 318.17 349.20 361.19 390.18 418.83 478.17 488.50 515.03 536.20 555.15 572.11 585.08 0 60 80 100 120 140 160 180 200 220 240 260 280 300 320 340 360 380 400 420 440 460 480 500 520 540 560 580 600 m/z Fig. 87 Chromatogram and MS data from HR-APCI-MS of BF2(HBQ). The mobile phase solvent was THF.

180000 2.50E-02 160000 2.00E-02 140000 120000 1.50E-02 100000 1.00E-02 80000 Emission

Intensity (au) Intensity 60000 5.00E-03 Absorbance (au)Absorbance Absorbance 40000 0.00E+00 20000 0 -5.00E-03 0 200 400 600 800 1000 Wavelength (nm)

Fig. 88 Overlay of excitation and emission data for BF2(1,2-HNBO).

115

180000 0.04 160000 0.035 140000 0.03 120000 0.025 100000 0.02 80000 0.015 Emission

60000 0.01 Intensity (au) Intensity

40000 0.005 (au)Absorbance Absorbance 20000 0 0 -0.005 0 200 400 600 800 1000 Wavelength (nm)

Fig. 89 Overlay of excitation and emission data for BF2(1,2-HNBT).

16000 0.06 14000 0.05 12000 0.04 10000 0.03 8000 0.02 Emission

6000 Intensity (au) Intensity

4000 0.01 (au)Absorbance Absorbance 2000 0 0 -0.01 0 200 400 600 800 1000 Wavelength (nm)

Fig. 90 Overlay of excitation and emission data for BF2(2,3-HNBO).

116

35000 0.025

30000 0.02 25000 0.015 20000 0.01 15000 Emission

Intensity (au) Intensity 0.005 10000 (au)Absorbance Absorbance

5000 0

0 -0.005 0 200 400 600 800 1000 Wavelength (nm)

Fig. 91 Overlay of excitation and emission data for BF2(2,3-HNBT).

250000 0.07 0.06 200000 0.05

150000 0.04 0.03 Emission

100000 0.02 Intensity (au) Intensity

Absorbance (au)Absorbance Absorbance 0.01 50000 0 0 -0.01 0 200 400 600 800 1000 Wavelength (nm)

Fig. 92 Overlay of excitation and emission data for BF2(HPBI).

117

Compound Excitation Range Integrated Wavelength Integrated Fluorescence [nm] Intensity

DPA(340 nm) 340 350-700 28132450

BF2(HPBI) 340 350-700 14367220

BF2(1,2-HNBO) 370 380-700 10225960

BF2(1,2-HNBT) 370 380-700 11519865

BF2(2,3-HNBO) 370 380-700 1562345

BF2(2,3-HNBT) 370 380-700 3082295

BF2(HBQ) 370 380-700 12169970

DPA(370 nm) 370 380-700 37322305

Table 4 Raw data for quantum yield calculations.

120000 2.50E-01

100000 2.00E-01

80000 1.50E-01 60000 1.00E-01 Theoretical

40000 Experimental Intensity (au) Intensity

5.00E-02 (au)Absorbance 20000

0 0.00E+00 195 245 295 345 395 -20000 -5.00E-02 Wavelength (nm)

Fig. 93 Overlay of UV-vis spectra of BF2(1,2-HNBO) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height. 118

100000 0.08

0.07 80000 0.06

60000 0.05

0.04 40000 Theoretical 0.03

Experimental Intensity (au) Intensity

20000 0.02 (au)Absorbance

0.01 0 200 250 300 350 400 450 0 -20000 -0.01 Wavelength (nm)

Fig. 94 Overlay of UV-vis spectra of BF2(1,2-HNBT) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

140000 0.2 0.18 120000 0.16 100000 0.14

80000 0.12 0.1 60000 Theoretical 0.08 Experimental

Intensity (au) Intensity 40000

0.06 (au)Absorbance

20000 0.04 0.02 0 0 190 240 290 340 390 440 490 -20000 -0.02 Wavelength (nm)

119

Fig. 95 Overlay of UV-vis spectra of BF2(2,3-HNBO) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

140000 2.50E-01

120000 2.00E-01 100000 1.50E-01 80000

60000 1.00E-01 Theoretical

Experimental Absorbance Intensity (au) Intensity 40000 5.00E-02 20000 0.00E+00 0 190 290 390 490 -20000 -5.00E-02 Wavelength (nm)

Fig. 96 Overlay of UV-vis spectra of BF2(2,3-HNBT) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

120

80000 0.07

70000 0.06 60000 0.05 50000 0.04 40000 0.03 Theoretical 30000 Experimental

Intensity (au) Intensity 0.02

20000 (au)Absorbance 0.01 10000

0 0 190 240 290 340 390 -10000 -0.01 Wavelength (nm)

Fig. 97 Overlay of UV-vis spectra of BF2(HPBI) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

60000 0.16

0.14 50000 0.12 40000 0.1 30000 0.08 Theoretical Experimental

20000 0.06 Intensity(au)

0.04 Absorbance(au) 10000 0.02 0 0 190 240 290 340 390 -10000 -0.02

Fig. 98 Overlay of UV-vis spectra of BF2(HPBT) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

121

70000 0.16

60000 0.14 0.12 50000 0.1 40000 0.08 30000 Theoretical 0.06 Experimental

Intensity (au) Intensity 20000 0.04 (au)Absorbance 10000 0.02

0 0 190 240 290 340 390 -10000 -0.02 Wavelength (nm)

Fig. 99 Overlay of UV-vis spectra of BF2(HPBO) in DCM with Gaussian 09 generated spectra calculated by B3LYP/6-311G(d,p)++PCM, TD-DFT methods. Gaussian distributions were generated using 0.1 eV half width at half height.

122