University of California, Irvine

The School of Physical Sciences

Department of Chemistry

Borovik Research Laboratory

Effects of Intramolecular Hydrogen Bonds in Synthetic Iron-Thiolate Complexes

A thesis presented

by

Hannah Boon

to

The Department of Chemistry

in partial fulfillment of the requirements

for honors in the subject of

Chemistry

June 2015

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Abstract Hydrogen bonds (H-bonds) mediate the stability and reactivity of active sites in

many proteins. These non-covalent bonds interact with substrates to position them for

specific binding to a metal ion and also influence the physical properties around a metal

ion. To study the effect of H-bonds in the secondary coordination sphere, synthetic iron-

thiolate complexes were prepared with a varying number of H-bonds under the premise

that as the number of H-bonds increased, the oxidation potential would become more

positive. Contrary to the hypothesis, the complex containing no H-bonds to the terminal

thiolate ligand had a more positive redox potential (-1.24 V vs [FeCp2]0/+) while the

complex with three H-bonds had a more negative redox potential (-1.55 V vs [FeCp2]0/+).

However, as adjusting the number of H-bonding donors also alters the primary

coordination sphere about the iron ion, it is suggested that the change in the primary

coordination sphere has a greater influence on physical properties rather than the

secondary coordination sphere in these ironII/III-thiolate complexes. Electron paramagnetic

resonance spectroscopy also suggested that the iron(III)-thiolate complex has rhombic

symmetry with g-values of 4.26 and 8.97.

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Introduction

I. Structure-Function Relationship of Active Site(s) in Biological Systems

The active sites in metalloproteins have unique structural properties that aid in

regulating a wide range of functions.1 The physical properties of the active sites in

metalloproteins can be characterized by the primary and secondary coordination spheres

surrounding the metal ion. The primary coordination sphere is defined by the ligands

directly coordinated to the metal center, and the secondary coordination sphere involves

the non-covalent, such as hydrogen bonding (H-bonding), interactions surrounding the

active site.2 The metal ion is an electron pair acceptor and acts as a Lewis acid;

conversely, the ligands are electron pair donors and act as a Lewis base.3 The

coordination spheres dictate the reactivity and other properties of the metal ion such as

the stereochemistry, electronic structure and Lewis acidity. Additionally, external

reagent binding (substrates) and selectivity can be investigated through consideration of

the secondary coordination sphere.

Examples of the effects of primary and secondary coordination spheres can be

observed within the active site of cytochrome P450 enzymes (P450s), a naturally

occurring heme metalloprotein that has the ability to activate C—H bonds (Scheme 1).

The active site of this protein contains a heme-iron center with an axial thiolate ligand

that serves to activate the iron center to bind dioxygen.4 The porphyrin ring and the

thiolate ligand define the primary coordination sphere to bind heme cofactor to the

protein. The role of hydrogen bonds in the secondary coordination sphere of the active

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site are to facilitate substrate binding and handle proton transfers key to performing the

catalytic cycle.5

Scheme 1: The catalytic cycle of P450 with ferryl species compound I and compound II labeled Inset: The active site of heme metalloprotein, P450, in the ferric state in the presence of camphor substrate and threonine participating in H-bonding. (PDB: 1HHO)

II. Synthetic C3-Symmetric Ligand Framework

Taking inspiration from biological systems like the P450s, the Borovik Group has

designed and synthesized C3-symmetric ligands that build a specific architecture

including noncovalent interactions to mediate metal ion reactivity. More specifically,

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urea-based ligands have been developed because of their inherent ability to form H-bond

networks as well as having characteristically high N—H bond dissociation energies

(>100 kcal/mol in DMSO)6 to withstand the high oxidation potentials required to produce

highly reactive species relevant to biology (Figure 1). The ligand precursor, tris[(Nʹ′-tert-

butylureayl)-N-ethylene]amine (H6buea), is deprotonated three times to yield [H3buea]3-,

where three anionic nitrogen atoms form an equatorial plane along with an apical neutral

amine nitrogen to provide a tetradentate binding pocket around a metal ion. Upon binding

of an external ligand, X, a favorable 6-membered ring is formed with the H-bonding N-H

groups of the urea functionality. Additionally, bulky tert-butyl groups interlock on the

top of the cavity to protect the highly reactive terminal ligands.1

Figure 1: Design of a tripodal urea-based ligand

III. Investigation of Physical Properties in Synthetic Metal Complexes Inspired by the structure/function relationship of the active sites of

metalloproteins like P450s, the purpose of the research discussed in this thesis is to

examine the role of H-bonding to an external ligand and gauge the strength of this

interaction by examining the physical properties of the metal ion, most notably by a

change in the oxidation potential. Synthetic iron(II) and iron(III) complexes were

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synthesized with tripodal ligand systems that interact with a bound thiolate ligand

through a varying number of H-bonds. One objective was to determine if varying the

number of H-bonds in the system has an effect on the physical properties of the complex.

The following four complexes provide a series of second coordination sphere frameworks

with a range of zero to three hydrogen bond interactions to the thiolate (Figure 2). The

urea-based ligands provide the H-bond while the amide-based ligands do not. Two of the

four complexes, [FeIIH3buea(SCH3)]2- and [FeII0cyp(SCH3)]2-, have been characterized

spectroscopically and electrochemically.

Figure 2: The four iron(II) complexes from left to right are: [FeIIH3buea(SCH3)]2-, [FeIIH22cyp(SCH3)]2-, [FeIIH1cyp(SCH3)]2- and [FeII0cyp(SCH3)]2-

Previous studies have been done with iron bound to terminal hydroxo ligands with

the same set of tripodal chelating systems but contain isopropyl groups instead of

cyclopentyl groups on the amidate arm(s): [FeIIIH3buea(OH)]-, [FeIIIH22iPr(OH)]-,

[FeIIIH1iPr(OH)]- and [FeIII0iPr(OH)]-.(1, 7-11) One goal of the research was to exchange the

hydroxo ligand for a thiolate ligand and explore the differences in the physical properties

of the complexes. Since sulfur atoms are known to be a weaker H-bond acceptor in

comparison to oxygen atoms, a greater affect in reactivity could possibly be identified

with various H-bond networks in the thiolate complexes. The iron-thiolate complexes

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reported here represent preliminary spectroscopically and electrochemically

characterizations.

Results and Discussion

I. Preparation the FeII- and FeIII-thiolate complexes

The two extremes of the Fe(II)-thiolate series were synthesized and characterized.

[FeIIH3buea(SCH3)]2- (1) and [FeII0cyp(SCH3)]2- (2) were prepared using the synthetic route

illustrated in Scheme 2. The starting ligand (H3buea and H30cyp) was deprotonated with 3

equivalents of KH in DMA under an argon atmosphere. The trianionic ligand was treated

with FeII(OAc)2, then NaSCH3, and finally metathesized with Me4NOAc which resulted

in the formation of the FeII-thiolate salt. [FeIIIH3buea(SCH3)]- (3) was prepared similarly

to compound 1 with the exception of FeIIICl3 instead of FeII(OAc)2. It is worth noting that

3 was more difficult to isolate than 1 due to iron(III) being less polarizable (a hard acid)

and therefore having a greater affinity for hydroxo species (a hard base). Conversely,

iron(II) is less electropositive and has a lower charge-to-radius ratio compared to

iron(III), therefore, iron(II) has a greater affinity for soft bases such as SCH!!.3

The isolation of 1, 2 and 3 set up the preliminary work for the series of related

tripodal ligands that have similar primary coordination spheres but different secondary

coordination sphere environments. Compounds 1 and 3 contain up to three H-bond donor

groups that are formed between the urea NH groups from the ligand and the thiolate

group that is bound to the FeII/FeIII center. Conversely, compound 2 does not have any H-

bond interactions to the thiolate group. These complexes were studied to determine if

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there was a relationship between different H-bond networks and physical properties, such

as redox potential.

Scheme 2. Preparation of A. (Me4N)2[FeIIH3buea(SCH3)] and B. (Me4N)2[FeII0cyp(SCH3)] complexes

II. Spectroscopic Properties

The [FeIIH3buea(SCH3)]2- salt was characterized by infrared (FTIR) spectroscopy

revealing three sharp peaks at high energy, 3179, 3227 and 3331 cm-1, using Nujol mull,

a result consistent with N—H vibrations from the urea-based ligand system (Figure 3).

There were no IR peaks observed in the same region for the [FeII0cyp(SCH3)]2- salt

confirming that were no hydrogen bond interactions to the axial thiolate group and

supporting the notion that the FeII-thiolate complexes would be less prone to adventitious

water binding resulting in the previously reported FeII-hydroxo complexes. Both complex

1 and 2 displayed vibrational peaks at 1571 and 1573 cm-1, respectively, consistent with

the (CO) group found in the urea and amide-based ligand systems.

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Figure 3. Overlaying vibrational spectra for [FeIIH3buea(SCH3)]2-, shown in black, and [FeII0cyp(SCH3)]2-, shown in red

Table 1. Summary of Selected Spectroscopic and Electrochemical Results for the FeIISCH3 Complexes

[FeIIH3buea(SCH3)]2- [FeII0cyp(SCH3)]2-

ν (NH) (cm-1) 3179, 3227, 3331 ν (CO) (cm-1) 1571 1573 E1/2 (V)a -1.55 -1.24 Scan Rate (mV) 500 100

aMeasured as DMA solutions using [FeCp2]0/+ as the reference.

Compound 3 was characterized using electron paramagnetic resonance (EPR)

spectroscopy which is a branch of magnetic resonance spectroscopy that depends on the

absorption of electromagnetic radiation to elicit an electronic transition.12 For a simple

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case with one unpaired electron, the measure of the energy difference, ΔE, between the

two possible spin states ms = +1/2 and ms = -1/2 of an electron when an external magnetic

field is applied gives insight into the identity and structure of the sample. The energy

differences studied in EPR spectroscopy are due to the interaction in paramagnetic

systems, a system with unpaired electrons. The electron spin that opposes the magnetic

field (ms = +1/2) will increase in energy and conversely, the electron spin that aligns with

the magnetic field (ms = -1/2) will decrease in energy. The resulting difference in E can

be calculated by

𝛥𝐸 = 𝑔𝛽𝐵! = ℎ𝜈 (Eq. 1)

An important parameter is the g-value(s), a measure of electronic interaction between the

unpaired electron and applied magnetic field, which splits the energy levels of the ground

state by an amount of ΔE (Eq. 1). Since the magnitude of the magnetic field, B0, can be

measured and β is the Bohr-magneton constant, g can be calculated to determine the

value of ΔE by irradiating the sample with microwaves. The g-value can then be

calculated from ν(GHz) and B0 (gauss) using

𝑔 =   !!!!!

(Eq. 2)

where h is Planck’s constant (6.626x10-34 J�s) andβis a constant (9.274x10-28 J�G-1).12, 13

Perpendicular mode EPR spectra were obtained at 77 K for compounds 1 and 2,

and found to be silent, consistent with a high spin FeII center (S = 2). An EPR spectrum

was obtained for [FeIIIH3buea(SCH3)]- in perpendicular mode as a DMA solution at 77 K

and displayed a sharp peak at g = 4.26 and a low intensity peak at g = 8.97 (Figure 4).

These features are indicative of a high-spin, S = 5/2 FeIII center.12-14 The characteristics

of these peaks suggest that the complex is rhombic, gx ≠ gy ≠ gz. The g-value of particular

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interest is the g = 8.97 because it suggests that the majority of the FeIII ions present in the

solution are a single, rhombic species. A single species will exhibit an intense g-value

signal around 4.3 with a derivative line shape while a signal at g ≈ 9 will be much less

intense but relatively sharp and well defined. If there were multiple, rhombic species, the

intense g = 4.3 signal will be present, but the g = 9 feature would likely be either

completely unobservable or extremely broad.

Figure 4. X-band EPR Spectrum of [FeIIIH3buea(SCH3)]- recorded as a DMA solution at 77 K III. Electrochemical Properties Previous studies for the FeIIIOH complexes with varying hydrogen bond networks

had similar physical properties because both the number of ureido donors and the number

of hydrogen bond interactions influenced the redox potential. As the number of H-bonds

increase, the ligand field around the metal center is weakened and should show a more

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positive oxidation potential. However, the number of H-bonds correlates to the number

of ureido donors, which counteracts the H-bond effect and shifts the oxidation potential

negatively. However, the electrochemical properties of complexes 1 and 2 illustrate a

different trend (Figure 5A and 5B). The approximate E1/2 for the oxidation of both

complexes was obtained by taking the average of the cathodic and anodic potentials. The

redox potential for the first reversible one-electron process for 1 was E1/2 (DMA) = -1.55

V versus [FeCp2]0/+ and assigned to the FeII/IIISCH3 couple. The second reversible process

was at E1/2 (DMA) = -1.91 V versus [FeCp2]0/+ and assigned to the FeII/III(OH) couple,

matching the previously reported value and present due to adventitious water in the

solvent.9 The redox potential for 2 was E1/2 (DMA) = -1.24 V versus [FeCp2]0/+ and

assigned to the FeII/IIISCH3 couple. There also appears to be an additional oxidative peak,

represented by the more negative potential, E1/2 (DMA) = -1.64 V, which could be the

Fe(II)-aquo specie.

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Figure 5. (A) Cyclic voltammogram of [FeIIH3buea(SCH3)]2- and (B) cyclic voltammogram of [FeII0cyp(SCH3)]2-.

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The initial hypothesis was that varying H-bond networks would influence the

physical properties of the complexes. As the number of H-bonds increase, the redox

potential should become more positive. Therefore, complex 1 with the most H-bonding

interactions should have the more positive oxidation potential in comparison to 2, which

has no H-bonds. However, CV experiments indicated the reverse trend, in which

compound 1 had a more negative E1/2 in comparison to 2.

By comparison, previous studies on the [FeIIIH3buea(OH)]- and [FeIII0iPr(OH)]-

complexes displayed a similarly perplexing trend in the reduction potentials between one

another. With decreasing H-bonding, the redox potential is expected to decrease, as there

should be greater electronic donation to the metal center. However, in the exchanging of

the urea-nitrogen donors (pKa ≈ 26.9) with amide-nitrogen donors (pKa ≈ 25.5) would

result in a slight decrease of the electron density at the metal center, which would yield

an increase in the reduction potential.15 Since the urea-based ligands are slightly more

basic compared to the amide-based ligands, the iron center becomes easier to oxidize and

therefore the oxidation potential will be more negative. The conclusion was that in

varying the number of H-bonds, the altering of the primary coordination sphere

counteracts the H-bonding effect. When considering the thiolate complexes, the

influence of H-bonding is expected to be lessened due to the poor ability of sulfur to

accept hydrogen bonds compared to oxygen. Therefore, the altering of the primary

coordination sphere must have a greater influence on redox potentials of the

FeIISCH3/FeIIISCH3 complexes compared to the FeIIIOH complexes previously reported.

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Experimental Section

I. General Procedures All manipulations, unless otherwise stated, were performed under an argon

atmosphere in a Vac-atmospheres dry box. All chemicals were purchased from

commercial sources. Dimethylacetamide, DMA, was obtained from a purification

solvent system and further purified by stirring with barium oxide for three days, refluxing

for one hour under nitrogen and distilling under reduced pressure.16 The distilled solvent

was stored over molecular sieves in a dry box. Sodium thiomethoxide was purified by

dissolving the salt in ethanol and adding ether in a 1:10 ratio, respectively. The

precipitate was cooled in an ice bath, collected on a medium fritted funnel and washed

with cool ether. The resulting white power was dried under vacuum and stored in a dry

box under an argon atmosphere.16

II. Physical Methods

1H NMR were recorded on a GN500 spectrometer. Fourier transform infrared

(FTIR) spectra were collected on a Varian 800 Scimitar Series FTIR spectrometer.

Perpendicular-mode X-band electron paramagnetic resonance (EPR) spectra were

collected using a Bruker EMX spectrometer at 77K using liquid nitrogen. Cyclic

voltammetry experiments were conducted using a CH1600C electrochemical analyzer.

III. Cyclic Voltammetry Methods For a typical cyclic voltammetry (CV) experiment, an electrolytic solution was

prepared using tetrabutylammonium hexafluorophosphate (TBAP, 0.1162 g, 100mM, 3

mL DMA). Approximately 0.005 g of sample was added to the electrolyte solution and

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dissolved. A glassy carbon electrode was used for the working electrode, a platinum wire

was used for the counter electrode and a silver wire was used for the reference electrode.

IV. Synthetic Methods

1. bis(tetramethylammonium){Tris[(N’-tert-butylurealato)-N-ethyl]aminato(thiomethoxide)ferrate(II)} (NMe4)2[FeIIH3buea(SCH3)] H6buea (0.150 g, 0.338 mmol) was dissolved in 8 mL of DMA and deprotonated

with KH (0.043 g, 1.07 mmol) under an Ar atmosphere. After H2 evolution ceased,

Fe(OAc)2 (0.060 g, 0.346 mmol) was added as a solid and the solution was stirred for 1

hour. This solution was treated with 1 equivalent of NaSCH3 (0.025 g, 0.352 mmol) and

the resulting pale yellow solution was stirred for 1.5 hours. Cation metathesis was

carried out by treating this solution with Me4NOAc (0.090 g, 0.676 mmol) and allowed to

stir for 1.5 hours. The resulting solution was filtered to collect byproduct salts KOAc and

NaOAc. The filtrate was concentrated in vacuo and solid product precipitated with

diethyl ether to yield 0.165 g of (NMe4)2[FeIIH3buea(SCH3)] (79% yield) as a pale yellow

solid. FTIR (Nujol, cm-1): ν(NH) 3179, 3227, 3331 (urea); ν(CO) 1571 (urea). UV/vis

(DMA): no λmax. EPR (X-Band, DMA, 77 K): silent. E1/2 (FeII/IIISCH3) = -1.55 V (vs

[FeCp2]0/+, scan rate: 500 mV/s, DMA/TBAP).

2. bis(tetraethylammonium){Tris(N’’-cyclopentylcarbamoylmethyl)-aminato(thiomethoxide)ferrate(II)} (NMe4)2[FeII0cyp(SCH3)]

H30cyp (0.025 g, 0.069 mmol) was dissolved in 8 mL of DMA and deprotonated

with KH (0.043 g, 1.07 mmol) under an Ar atmosphere. After H2 evolution ceased,

Fe(OAc)2 (0.012 g, 0.069 mmol) was added as a solid and the solution was stirred for 1

hour. This solution was treated with 1 equivalent of NaSCH3 (0.005 g, 0.076 mmol) and

the resulting pale yellow solution was stirred for 45 minutes. Cation metathesis was

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carried out by treating this solution with Me4NOAc (0.020 g, 0.150 mmol) and allowed to

stir for 1 hour. The resulting solution was filtered to collect byproduct salts KOAc and

NaOAc. The filtrate was concentrated in vacuo and solid product precipitated with

diethyl ether to yield 0.032 g of (NMe4)2[FeII0cyp(SCH3)] (81% yield) as a pale yellow

solid. FTIR (Nujol, cm-1): ν(CO) 1573 (amide). UV/vis (DMA): no λmax. EPR (X-Band,

DMA, 77 K): silent. E1/2 (FeII/IIISCH3) = -1.24 V (vs [FeCp2]0/+, scan rate: 100 mV/s,

DMA/TBAP).

3. Potassium{Tris[(N’-tert-butylurealato)-N-ethyl]aminato(thiomethoxide)ferrate(III)} K[FeIIIH3buea(SCH3)]

H6buea (0.215 g, 0.485 mmol) was dissolved in 10 mL of DMA and deprotonated

with KH (0.058 g, 1.45 mmol) under an Ar atmosphere. After H2 evolution ceased,

FeIIICl3 (0.079 g, 0.485 mmol) was added as a solid and the solution was stirred for 2

hours. This solution was treated with 1 equivalent of NaSCH3 (0.034 g, 0.485 mmol) and

the dark red solution was stirred for 1 hour. The resulting solution was filtered to collect

byproduct salts KOAc and NaOAc. The filtrate was concentrated to 6mL of DMA in

vacuo. Half an aliquot, 3 mL, was metathesized with Et4NOAc (0.093 g, 0.0491 mmol)

and stirred for 1 hour. The resulting solution was filtered and the filtrate was

concentrated in vacuo and solid product precipitated with diethyl ether to yield 0.081 g of

K[FeIIIH3buea(SCH3)] (67% yield, with respect to ~1/2 aliquot theoretical yield) as a red-

orange solid. UV/vis (DMA): λmax/nm = 402. A solution of [FeIIIH3buea(SCH3)]- was

transferred to an EPR tube and sealed with a rubber septum. The tube was brought out of

the dry box and placed in a 77 K liquid nitrogen bath and allowed to equilibrate for 2

minutes. EPR (X-Band, DMA, 77 K): g = 8.97, 4.26.

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Conclusion

The effects of hydrogen bonds in the secondary coordination sphere in metal

complexes vary greatly from system to system. The FeII- and FeIII-thiolate complexes

prepared have similar N4 primary coordination spheres with the terminal thiolate ligand

positioned within the cavity. H-bonding effects are more prominent in metal complexes

with terminal hydroxo ligands in comparison to terminal thiolate ligands due to the sulfur

atom acting as a weaker H-bond acceptor. Ultimately, it was determined that the physical

properties of the FeII- and FeIII-thiolate complexes were dominated by the primary

coordination sphere environment with varying ureido/amidate tripodal ligands exhibiting

a greater influence on the oxidation potential than altering H-bonding interactions. These

results illustrate the importance of considering both the primary and secondary

coordination spheres in determining the properties of a metal complex.

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Acknowledgments

Thank you to Professor Andy Borovik for the opportunity to gain experience in a

graduate research lab and for mentoring me the past two years. Thank you to all the

graduate students in the Borovik lab for assisting me with the spectroscopic and

electrochemical techniques.

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REFERENCES

(1) Borovik, A. S. Bioinspired Hydrogen Bond Motifs in Ligand Design: The Role of noncovalent Interactions in Metal Ion Mediated Activation of Dioxygen, Acc. Chem. Res. 2005, 38, 54-61. (2) (a) Werner, A. Ann Chem. 1912, 386, 1.(b) Werner, A. Ber. Dtsch. Chem. Ges. 1912, 45, 130. (3) Pearson, R. G., Hard and Soft Acids and Bases, J. Am. Chem. Soc. 1963, 85, 3533-3539 (4) Green, M. T. C-H Bond Activation in Heme Proteins: The Role of Thiolate Ligation in Cytochrome P450, Chem. Bio. 2009, 13, 84-88 (5) Schlichting, I., Berendzen, J., Chu, K., Stock, A. M., Maves, S. A., Benson, D. E., Sweet, R. M., Ringe, D., Petsko, G.A., Sligar, S. G. The Catalytic Pathway of Cytochrome P450cam at Atomic Resolution, Science 2000, 287, 1615-1622. (6) Holm, R. H.; Kennepohl, P.; Solomon, E. I. Structural and Functional Aspects of Metal Sites in Biology. Structural and Functional Aspects of Metal Sites in Biology, Chem. Rev. 1996, 96, 2239-2314. ��� (7) MacBeth, C. E.; Golombek, A. P.; Young, V. G., Jr.; Yang, C.; Kuczera, K.; Hendrich, M. P.; Borovik, A. S. O2 Activation by Nonheme Iron Complexes: A Monomeric Fe(III)-Oxo ComplexDerived from O2, Science 2000, 289, 938-941. ���

(8) MacBeth, C. E.; Gupta, R.; Mitchell-Koch, K. R.; Young, V. G., Jr.; ���Lushington, G. H.; Thompson, W. H.; Hendrich, M. P.; Borovik, A. S. Utilization of Hydrogen Bonds to Stabilize M-O(H) Units: Synthesis and Properties of Monomeric Iron and Manganese Complexes with Terminal Oxo and Hydroxo Ligands, J. Am. Chem. Soc. 2004, 126, 2556-2567. ���

(9) MacBeth, C. E.; Hammes, B. S.; Young, V. G., Jr.; Borovik, A. S. Hydrogen- Bonding Cavities about Metal Ions: Synthesis, Structure, and Physical Properties f or a series of Monomeric M-OH Complexes Dervied from Water, Inorg. Chem. 2001, 40, 4733- 4741.

(10) Gupta, R.; MacBeth, C. E.; Young, V. G., Jr.; Borovik, A. S. Isolation of Monomeric MnIII/II-OH and MnIII-O Complexes from Water: Evaluation of O-H Bond Dissociation Energies, J. Am. Chem. Soc. 2002, 124, 1136-1137. (11) Mukherjee, J.; Lucas, R. L.; Zart, M. K.; Powell, D. R.; Way, V. W.; Borovik, A. S. Synthesis, Structure, and Physical Properties for a series of Monomeric Iron(III) Hydroxo Complexes with varying Hyrogen-Bond Networks, Inorg. Chem. 2008, 47, 5780-5786.

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(12) Drago, R.S.; Physical Methods for Chemists, 2nd Ed.; Saunders, 1992, Ch 9, 13 (13) Fatai, T. A., Electron Paramagnetic Resonance Spectroscopic Studies of iron and Copper Proteins, J. Spectrosc. 2003, 17, 53-63 (14) Weil, J. A., Bolton, J. R. Electron Paramagnetic Resonance Spectroscopy: elementary Theory and Application, 2nd Ed.; 2007, Wiley-Interscience.

(15) Bordwell, F. G., Guo-Zhen, J., Effects of Structural changes on Aciditites and Homolytic Bond Dissocation Energies of the H—N Bonds in Amidines, Carbozamides, and Theocarboxamides, J. Am. Chem. Soc. 1991, 113, 8398-8401.

(16) Armarego, W. L. F., Chai, C. L. L. Purification of Laboratory Chemicals, 5th Ed.; 2003, Elsevier-Science