Understanding the Mechanisms of Transition Metal Catalysed …688638/s4204895_final_the… ·...

Understanding the Mechanisms of Transition Metal Catalysed

Redox Reactions

Timothy Justin Zerk

BSc(Hons)/BEd

A thesis submitted for the degree of Doctor of Philosophy at

The University of Queensland in 2017

School of Chemistry & Molecular Biosciences

Abstract

Transition metal complexes catalyse a number of important synthetic chemical reactions. Two such

reactions which rely on copper and ruthenium complexes respectively are atom transfer radical

polymerisation (ATRP) and the Ley-Griffith oxidation of alcohols.

In ATRP a copper(I) complex bearing a chelate ligand ‘L’ homolytically cleaves the carbon–halogen

bond of an organo halide initiator (R–X) to produce an alkyl radical and the corresponding copper(II)-

halido complex (CuIL + RX → R• + CuIILX). The reverse reaction is coined ‘deactivation’ and

provides control to the process by keeping the concentration of the radical low. It is experimentally

difficult to determine the rate of this reaction because it is so fast. Herein, a new method for measuring

the rate of deactivation is developed using cyclic voltammetry coupled to simulations. The

mechanism of deactivation is also unknown and this aspect is investigated by a kinetic study of halide

substitution reactions on three ATRP-relevant CuIILX complexes. Together, the electrochemical and

kinetic results unveil the influence of the chelate ligand, halide and solvent on the rate and mechanism

of deactivation.

Radicals formed in ATRP also have the potential to react with CuIL by a radical transfer reaction

generating the organometallic complex CuIILR. While this has been postulated to occur for a number

of copper-based ATRP systems, the organometallic complex has never been observed. Here this

species is identified for the first time using a combination of electrochemistry and spectroscopy. The

rate and equilibrium constants from the central atom- and radical-transfer reactions are also measured

using cyclic voltammetry and reveal which solvent/initiator/catalyst combinations direct the system

towards the radical- or atom-transfer products (CuIILR or CuIILX respectively).

The Ley-Griffith reaction is catalysed by tetrapropylammonium perruthenate (TPAP) and also

requires a stoichiometric quantity of N-methylmorpholine N-oxide (NMO). Despite the popularity of

this method, the catalytic cycle has not been elucidated. Here a suite of electrochemical and

spectroscopic methods are applied to the reaction and unveil each of the steps in the cycle. Kinetic

studies show that the transition state during oxidation is comprised of a single perruthenate anion and

a single alcohol molecule. The products of this reaction are a highly unstable RuV species along with

water and the corresponding aldehyde/ketone. The N-oxide then re-oxidises RuV to perruthenate

before it irreversibly disproportionates to ruthenium dioxide and (RuVI). EPR spectroscopy also

shows that NMO forms an outer-sphere associated complex with perruthenate which may be

important in facilitating this reaction. Finally, synergistic EPR and UV-vis spectroscopy demonstrate

that even with a large excess of NMO, a small amount of ruthenium dioxide still forms during the

Ley-Griffith oxidation and acts as an accelerant for the reaction – i.e. the reaction is auto-catalytic.

Declaration by author

This thesis is composed of my original work, and contains no material previously published or

written by another person except where due reference has been made in the text. I have clearly

stated the contribution by others to jointly-authored works that I have included in my thesis.

I have clearly stated the contribution of others to my thesis as a whole, including statistical

assistance, survey design, data analysis, significant technical procedures, professional editorial

advice, and any other original research work used or reported in my thesis. The content of my thesis

is the result of work I have carried out since the commencement of my research higher degree

candidature and does not include a substantial part of work that has been submitted to qualify for

the award of any other degree or diploma in any university or other tertiary institution. I have

clearly stated which parts of my thesis, if any, have been submitted to qualify for another award.

I acknowledge that an electronic copy of my thesis must be lodged with the University Library and,

subject to the policy and procedures of The University of Queensland, the thesis be made available

for research and study in accordance with the Copyright Act 1968 unless a period of embargo has

been approved by the Dean of the Graduate School.

I acknowledge that copyright of all material contained in my thesis resides with the copyright

holder(s) of that material. Where appropriate I have obtained copyright permission from the

copyright holder to reproduce material in this thesis.

Publications during candidature

Peer-reviewed papers:

Zerk, T. J.; Bernhardt, P. V., Organo-Copper(II) Complexes as Products of Radical Atom

Transfer. Inorg. Chem. 2017, 56 (10), 5784-5792

Gavrilov, M.; Zerk, T. J.; Bernhardt, P. V.; Percec, V.; Monteiro, M. J., SET-LRP of

NIPAM in water via in situ reduction of Cu(II) to Cu(0) with NaBH4. Polym. Chem. 2016,

7 (4), 933-939

Zerk, T. J.; Moore, P. W.; Williams, C. M.; Bernhardt, P. V., N-Oxides rescue Ru(V) in

catalytic Griffith-Ley (TPAP) alcohol oxidations. Chem. Commun. 2016, 52 (67), 10301-

10304

Zerk, T. J.; Martinez, M.; Bernhardt, P. V., A Kinetico-Mechanistic Study on CuII

Deactivators Employed in Atom Transfer Radical Polymerization. Inorg. Chem. 2016, 55

(19), 9848-9857

Zerk, T. J.; Bernhardt, P. V., New Method for Exploring Deactivation Kinetics in Copper-

Catalyzed Atom-Transfer-Radical Reactions. Inorg. Chem. 2014, 53 (21), 11351-11353

Publications included in this thesis

Zerk, T. J.; Bernhardt, P. V., New Method for Exploring Deactivation Kinetics in Copper-Catalyzed Atom-

Transfer-Radical Reactions. Inorg. Chem. 2014, 53 (21), 11351-11353

– Incorporated as Chapter 2

Contributor Statement of contribution

Timothy Zerk Designed and conducted experiments (100%)

Wrote and edited the paper (85%)

Paul Bernhardt Wrote and edited the paper (15%)

Zerk, T. J.; Martinez, M.; Bernhardt, P. V., A Kinetico-Mechanistic Study on CuII Deactivators Employed in

Atom Transfer Radical Polymerization. Inorg. Chem. 2016, 55 (19), 9848-9857





Manel Martinez Designed and conducted experiments (5%)



Zerk, T. J.; Bernhardt, P. V., Organo-Copper(II) Complexes as Products of Radical Atom Transfer. Inorg.

Chem. 2017, 56 (10), 5784-579






Zerk, T. J.; Moore, P. W.; Williams, C. M.; Bernhardt, P. V., N-Oxides rescue Ru(V) in catalytic Griffith-

Ley (TPAP) alcohol oxidations. Chem. Commun. 2016, 52 (67), 10301-10304

‒ Incorporated as Chapter 5




Peter Moore Designed and conducted experiments (5%)


Craig Williams Wrote and edited the paper (5%)

Paul Bernhardt Designed and conducted experiments (5%)


Contributions by others to the thesis

Paul Bernhardt:

Paul has contributed primarily to the conceptual design of the project, the interpretation of the data

and to editing various manuscripts. While his contribution is largely intangible throughout this thesis,

without his input the document would certainly lack much of the cogent experimental focus and

interpretation which it presents.

Peter Moore:

Peter has been an important collaborator in the work which focuses on ruthenium-catalysed alcohol

oxidation reactions. He has produced and characterised many of the reagents which are analysed in

Chapter 5 and Chapter 6 and through various discussions has contributed to the conceptual design of

a number of the experiments contained therein. He has been an invaluable partner in conducting

various experiments (such as the 1H NMR studies) which required two sets of hands and has been

involved in writing up the experimental details of these.

Joshua Harbort:

Joshua has been involved in performing and interpreting the Electron Paramagnetic Resonance (EPR)

measurements referred to in Chapter 6. He produced the 2D EPR Figure in Chapter 6 and was

consulted in compiling the discussion of these results.

Jeffrey Harmer:

Jeffery has also been involved in collecting and interpreting the EPR data as Joshua’s principal

supervisor.

Statement of parts of the thesis submitted to qualify for the award of another degree

None

Acknowledgements

“I have always been blessed both in my official and my unofficial teachers”

C.S. Lewis

At a recent conference dinner I was seated next to a man who, during the course of our conversation,

mentioned that the most important part of any thesis is the Acknowledgements because “it’s personal

while the rest is just science”. Truly, there is more to life than science. Science is a beautiful, intricate

and satisfying pursuit but it is people who do science, who listen to it, who use it and who make it

possible. I am indebted to a few special people for the science contained herein.

Paul, you are a champion, a bastion of proper chemistry. Mercifully, you are not an egoist or a political

animal; instead you are brilliant and genuinely interested in chemistry. Thank you for your patience,

direction and hospitality as well as your rumbling witticism. More importantly, thank you for being

my great mentor and friend.

Lawrie, the above likewise apply to you although you do not ‘rumble’ so much. Thank you for the

kindness you have shown me and for the encouragement you provide when I need it most. Most of

all, thank you for being an example in integrity, grit and wisdom. Paul and I are both beneficiaries of

your company. One evening at the conference already mentioned I was genuflecting on these last

years and was overwhelmed with gratitude for you both. I ask you to recall those words from my

wedding speech if ever you doubt your legacy or significance. You don’t need a Nobel Prize to be

great, you just need to take a young man under your wing and invest heavily in him.

Following closely on the heels of these two must come Manel. Manel you are fantastic (the perfect

adjective). What Lawrie began in first year you have fanned into flame, an inclination towards

reaction mechanisms through kinetics. Thank you for investing your vibrancy into a reserved stranger

– my time in Barcelona was the best part of my PhD.

Michael, thank you for all the advice, encouragement and genuine interest you have shown in my

work. Many a snare has been avoided because of your well-informed guidance. You are a model of

dedication to the pursuit of knowledge and every time I part your company I walk away with a

renewed energy.

Craig, collaborating with you has been an unexpected yet total pleasure. You are different from Paul

in ways that have honed edges which would otherwise have remained blunt. I am so grateful for the

enthusiasm which you have demonstrated and for the encouragement to aim high. Thank you for

introducing me to the literally black art of ruthenium oxidation chemistry.

To the two musketeers, Myles and Jess, thank you for your longsuffering patience, your

companionship, your encouragement and for your friendship. It is a delightful thing to discover that

there are other people in this world who find chemistry both fun and funny. I would also like to thank

my good friend and close collaborator Peter, for your unflappable, sanguine steadiness which kept

me going through the slough of despond in a tough year investigating ruthenium oxidation chemistry.

It was very satisfying to reach the end of this project with you.

Mum and Dad, I am beginning to understand how fortunate I am to have you as my parents. Your

ongoing support, in tangible and (more importantly) intangible ways, is beyond value. Thank you for

your selflessness.

To Gail, I have much to say but I would rather say it in person. For now, I am content to summarise

it by affirming that two are indeed better than one.

There is one final person to whom I owe a great deal and to whom I dedicate this work, David Denner.

Dave, I have been blessed in my official and unofficial teachers, you were the first to occupy both

positions and I owe you a debt I can never repay. So much of the above applies to you. It was you

who taught me the greatest lesson of all – the value of people. Thank you for the hours and hours you

have poured into my life – they have not been in vain. To borrow again from C.S. Lewis “I found…

that the ripest are kindest to the raw and the most studious have the most time to spare”. Thank you

Dave, for everything.

Keywords

Atom transfer radical polymerisation (ATRP), Ley-Griffith oxidation, organometallic mediated

radical polymerisation (OMRP), ruthenium, copper, catalyst, redox reaction, deactivation, kinetics,

mechanism.

Australian and New Zealand Standard Research Classifications (ANZSRC)

ANZSRC code: 030601 Catalysis and Mechanisms of Reactions, 40%

ANZSRC code: 030207 Transition Metal Chemistry, 30%

ANZSRC code: 030604 Electrochemistry, 30%

Fields of Research (FoR) Classification

FoR code: 0306 Physical Chemistry (incl. Structural), 70%

FoR code: 0302 Inorganic Chemistry, 30%

i

TABLE OF CONTENTS

...................................................................................................................... 1

Introduction ...................................................................................................................................... 1

Cyclic voltammetry including heterogeneous and homogeneous electron transfer ........................ 7

Voltammetry of ATRP catalysts .................................................................................................... 16

.................................................................................................................... 21

Introduction .................................................................................................................................... 21

Results & Discussion ..................................................................................................................... 22

Cyclic Voltammetry ................................................................................................................... 22

Fitting kact ................................................................................................................................... 25

Fitting kdeact ................................................................................................................................ 30

Conclusion ..................................................................................................................................... 34

Experimental .................................................................................................................................. 35

.................................................................................................................... 37

Introduction .................................................................................................................................... 37


A model system .......................................................................................................................... 40

Anation kinetics ......................................................................................................................... 41

Solvent exchange kinetics .......................................................................................................... 48

Halide exchange kinetics ........................................................................................................... 49

Effect of the chelate ................................................................................................................... 53

Effect of the halide ..................................................................................................................... 55

Effect of the solvent ................................................................................................................... 56

Conclusion ..................................................................................................................................... 57

Experimental .................................................................................................................................. 58

Synthesis .................................................................................................................................... 58

Kinetics ...................................................................................................................................... 59

.................................................................................................................... 60

Introduction .................................................................................................................................... 60


Electrochemistry of [CuII(tpa)Br]+ ............................................................................................. 66

Spectroelectrochemistry of [Cu(tpa)Br]+ ................................................................................... 67

ii

Simulating the voltammetry of [CuII(tpa)Br]+ ........................................................................... 71

Hydrolytic decomposition of [CuI(tpa)R] .................................................................................. 75

Factors controlling ATRP versus OMRP................................................................................... 76

Electrochemistry of [Cu(Me6tren)Br]+ ....................................................................................... 78

Spectroelectrochemistry of [Cu(Me6tren)Br]+ ........................................................................... 79

Simulating the voltammetry of [CuII(Me6tren)X]+ .................................................................... 81

Conclusion ..................................................................................................................................... 83

Experimental .................................................................................................................................. 84

Synthesis .................................................................................................................................... 84

Physical Methods ....................................................................................................................... 85

Fitting Process ............................................................................................................................ 86

Effect of H2O on voltammetry ................................................................................................... 90

.................................................................................................................... 91

Introduction .................................................................................................................................... 91


The role of NMO........................................................................................................................ 95

NMO and RuO4 electrochemistry ............................................................................................ 106

Conclusion ................................................................................................................................... 109

Experimental ................................................................................................................................ 110

Synthesis .................................................................................................................................. 110

Physical methods...................................................................................................................... 111

.................................................................................................................. 112

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

Results & Discussion ................................................................................................................... 115

The mechanism of oxidation – bi-phasic kinetics .................................................................... 115

The mechanism of oxidation – 1H NMR ................................................................................. 120

The mechanism of oxidation – EPR ........................................................................................ 123

The mechanism of oxidation – RuO2.2H2O catalysis .............................................................. 127

Conclusion ................................................................................................................................... 133

Experimental ................................................................................................................................ 134

Synthesis .................................................................................................................................. 135

Kinetics .................................................................................................................................... 137

Effect of RuO22H2O................................................................................................................. 137

References ................................................................................................................ 138

iii

Appendices ............................................................................................................... 153

Appendix 2.1 ................................................................................................................................ 153

Appendix 3.1 ................................................................................................................................ 154

Appendix 3.2 ................................................................................................................................ 155

Appendix 3.3 ................................................................................................................................ 156

Appendix 3.4 ................................................................................................................................ 157

Appendix 3.5 ................................................................................................................................ 158

Appendix 4.1 ................................................................................................................................ 167

Appendix 4.2 ................................................................................................................................ 167

Appendix 5.1 ................................................................................................................................ 168

Appendix 5.2 ................................................................................................................................ 168

Appendix 6.1 ................................................................................................................................ 169

iv

LIST OF FIGURES

Figure 1-1 Simulated cyclic voltammogram for the reversible, one-electron reduction of O. ......................... 8

Figure 1-2 Schematic representation of the reversible redox reaction involving O and R at a planar electrode.

.................................................................................................................................................................. 8

Figure 1-3 Time-dependent concentration gradient profile for O at a three fixed potentials. Solid line – no

applied potential; dashed line – E < E0', dotted line – E << E0'. .............................................................. 10

Figure 1-4 Simulated CV waveforms for the ECcat mechanism as a function of [A]. Red – 0.0 mM, Green –

1.0 mM, Yellow – 10.0 mM, Blue – 50.0 mM A. Rate constant k = 1.0 × 103 M-1 s-1. .......................... 11

Figure 1-5 Representation of a reversible electron transfer at the electrode coupled to an irreversible,

catalytic chemical reaction in solution. ................................................................................................... 11

Figure 1-6 Kinetic zone diagram and simulated CV waveforms as a function of the dimensionless

parameters λ and γ. ................................................................................................................................. 12

Figure 1-7 Simulated CV waveforms for the ECcat mechanism as a function of the rate constant k. ............. 14

Figure 1-8 Representation of a reversible electron transfer at the electrode coupled to a reversible, catalytic

chemical reaction in solution. ................................................................................................................. 15

Figure 1-9 Experimental and simulated cyclic voltammograms of 1.0 mM [CuII(Me6tren)Br]Br at four

different concentrations of EBriB. ......................................................................................................... 18

Figure 2-1 Structurally characterised copper(II) complexes of PMDETA. .................................................... 22

Figure 2-2 Cyclic voltammetry of 1.0 mM [CuII(PMDETA)Br(OSMe2)]Br in with added EBriB and

TEMPO. .................................................................................................................................................. 23

Figure 2-3 [RBr] dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess

TEMPO. .................................................................................................................................................. 26

Figure 2-4 Sweep rate dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess

TEMPO and 5.0 mM EBriB. .................................................................................................................. 26

Figure 2-5 Sweep rate dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess

TEMPO and 12.0 mM BnBr. .................................................................................................................. 27

Figure 2-6 The CV of a blank solution of DMSO + electrolyte. .................................................................... 27

Figure 2-7 Sensitivity of the digital trace to kact. ............................................................................................ 28

Figure 2-8 [RBr] dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with no TEMPO. ..... 30

Figure 2-9 Sweep rate dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with 5.0

mM EBriB and no TEMPO. ................................................................................................................... 31

Figure 2-10 Sweep rate-dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with

12.0 mM BnBr and no TEMPO. ............................................................................................................. 31

Figure 2-11 Sensitivity of the digital trace to kdeact. ........................................................................................ 32

Figure 3-1 Relevant structures determined for [Cu(L)X]+ deactivating complexes with the chelating ligands

Me6tren, Et6tren or tpa. ........................................................................................................................... 38

Figure 3-2 A) Time-resolved spectral changes for the reaction of 2.0 × 10-4 M [CuII(Me6tren)(NCMe)]2+ with

Br– (0.0025 M) at 298 K and I = 0.1 M LiClO4; B) Time resolved spectral changes for the reaction of

3.8 × 103 M [CuII(Me6tren)(NCMe)]2+ with DMF (1.29 M) at 298 K and I = 0.1 M LiClO4. ................ 40

Figure 3-3 Left) Plot of kobs versus [Br-] for the [Cu(Me6tren)(NCMe)]2+ + Br- reaction at different

temperatures and at I = 0.1 M LiClO4; Right) kon as a function of temperature and pressure. ............... 42

Figure 3-4 Left) Plot of kobs versus [Br-] for the [Cu(Me6tren)(OSMe2)]2+ + Br- reaction at 298 K; Right) Plot

of kobs versus [Br-] for the [Cu(Me6tren)(DMF)]2+ + Br- reaction at 288 K. ........................................... 42

v

Figure 3-5 Representation of different ligand substitution mechanisms. ....................................................... 44

Figure 3-6 Representation of the reaction coordinate during an anation reaction in which the rate

determining step proceeds by partial dissociation of the coordinated solvent (Id).................................. 46

Figure 3-7 Plot of kobs versus [Br-] for the [Cu(Me6tren)(DMSO)]2+ + Br- reaction in DMSO/MMA. .......... 47

Figure 3-8 Plots of kobs vs. [Y-] for the halido ligand exchange reaction [Cu(Me6tren)X]+ + Y- in MeCN, I =

0.1 M LiClO4. Left) X = Br, Y = Cl and Right) X = Cl, Y = Br. ........................................................... 49

Figure 3-9 Plots of kobs vs. [Y-] for the halido ligand exchange reaction [Cu(Me6tren)X]+ + Y- in DMF, I =

0.1 M LiClO4. Left) X = Br, Y = Cl and Right) X = Cl, Y = Br. ........................................................... 50

Figure 3-10 Plots of kobs vs. [Y-] for the reaction of Left) [CuII(Et6tren)Cl]+ + Br- at different temperatures or

Right) [CuII(tpa)Cl]+ + Br- at 288 K. ...................................................................................................... 50

Figure 3-11 Comparison of the CuII‒Cl coordinate bonds (Å) across the homologous series [CuII(tren)Cl]+

(BPh4- salt); [CuII(Me6tren)Cl]+ (ClO4

- salt) and [CuII(tpa)Cl]+ (Cl- salt). .............................................. 54

Figure 3-12 Comparison of the CuII‒N coordinate bonds (Å) across the homologous series

[CuII(Me3tren)(NCMe)]2+ (ClO4- salt); [CuII(Me6tren)(NCMe)]2+ (BPh4

- salt); and [CuII(tpa)(NCMe)]2+

(ClO4- salt). ............................................................................................................................................. 54

Figure 4-1 Cobalt catalysts for OMRP bearing tetramesitylporphyrin (1) acetylaceconate (2) or dioxime (3)

chelating ligands. .................................................................................................................................... 61

Figure 4-2 Olefinic monomers........................................................................................................................ 62

Figure 4-3 Chelate ligands of highly active copper catalysts. .................................................................... 63

Figure 4-4 Proposed formation of a organocopper(II) species which leads to increased R-R terminated

products via a Catalytic Radical Termination (CRT) pathway. .............................................................. 64

Figure 4-5 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br + [RBr] in DMSO (A & B) or MeCN (C & D). 66

Figure 4-6 A) Spectra measured every five seconds during electrolysis of 6.0 mM [CuII(tpa)Br]Br + 48 mM

in MeCN. Potential was held at -850 mV vs. Fc+/0. B) Spectra of 6 mM [CuII(tpa)Br/R/CN]+ in MeCN.

................................................................................................................................................................ 68

Figure 4-7 Experimental (top) and simulated (bottom) X-band (9.3708 GHz) EPR spectra at 130 K of 1 mM

[CuII(tpa)Br]+. ......................................................................................................................................... 70

Figure 4-8 Experimental (top) and simulated (bottom) X-band (9.3708 GHz) EPR spectra at 130 K of 1 mM

[CuII(tpa)(CH2CN)]+ formed after bulk electrolysis of 1 mM [CuII(tpa)Br]+ in the presence of

bromoacetonitrile. ................................................................................................................................... 70

Figure 4-9 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br in DMSO. .......................................................... 71

Figure 4-10 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br + [RBr] with added TEMPO (0.2 M).. ........... 72

Figure 4-11 Red) Experimental (solid) and simulated (broken) voltammetry of 1.0 mM [CuII(tpa)Br]Br in

MeCN (0.2 M TEMPO) with 4.0 mM EBriB. ........................................................................................ 73

Figure 4-12 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br in MeCN (0.1 M (Et4N)(ClO4)) with 3.0 mM

bromoacetonitrile and various added [H2O]. .......................................................................................... 75

Figure 4-13 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)Br]Br (A) or 1.0 mM [CuII(Me6tren)Cl]Cl (B) in

MeCN + [RX]. ........................................................................................................................................ 78

Figure 4-14 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)Br]Br in DMSO + bromoacetonitrile. ............... 79

Figure 4-15 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)X]X + [RX] with added TEMPO (0.2 M). A) X =

Br-, Solv. = MeCN; B) X = Cl-, Solv. = MeCN. C). X = Br-, Solv. = DMSO. ....................................... 79

Figure 4-16 A) Spectra measured every 20 seconds during electrolysis of 6 mM [CuII(Me6tren)Br]Br + 48

mM bromoacetonitrile in MeCN. Potential was held at -1000 mV vs. Fc+/0. B) Spectra of 6 mM

[CuII(Me6tren)X]+ in MeCN (X- = Br-, CN- or R = NCCH2-). ................................................................ 80

vi

Figure 4-17 A) Spectra measured every 20 seconds during electrolysis of 6 mM [CuII(Me6tren)Cl]+ + 48 mM

chloroacetonitrile in MeCN. Potential was held at -900 mV vs. Fc+/0. B) Spectra of 6 mM

[CuII(Me6tren)Br]+ + 48 mM bromoacetonitrile in DMSO measured every 20 seconds during reduction

at -1000 mV vs. Fc+/0. ............................................................................................................................. 80

Figure 4-18 Spectra recorded during the addition of 0.1 mM aliquots of (Et4N)Br to a 1.0 mM solution of A)

[CuII(Me6tren)(OSMe2)](ClO4)2 in DMSO and B) [CuII(tpa)(OSMe2)](ClO4)2

in DMSO. ................... 86

Figure 4-19 Experimental (A) and simulated (B) voltammetry during the titration of Br- into 1.0 mM

[CuII(tpa)(NCMe)](ClO4)2 in acetonitrile. .............................................................................................. 89


[CuII(tpa)(OSMe2)](ClO4)2 in DMSO. .................................................................................................... 89

Figure 4-21 Cyclic voltammetry of the carefully dried 1.0 mM [CuII(tpa)Br]Br in MeCN +

[bromoacetonitrile]. ................................................................................................................................ 90

Figure 5-1 Popular reactions for alcohol oxidation. ....................................................................................... 91

Figure 5-2 Tetrahedral ruthenium tetroxide/perruthenate and trigonal bipyramidal ruthenate. ..................... 93

Figure 5-3 Cyclic voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN. ............................................................ 96

Figure 5-4 Spectrum of n-Pr4N[RuO4] measured in MeCN. .......................................................................... 96

Figure 5-5 Experimental (A) and simulated (B) voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN in the

vicinity of the RuVII/VI couple. ................................................................................................................. 97

Figure 5-6 UV-vis spectra of 1.0 mM [RuO4]- in MeCN during electrochemical reduction electrolysis at -

2300 mV in an anaerobic glovebox (O2 < 10 ppm). ............................................................................... 98

Figure 5-7 Time resolved speciation profiles (Left) and de-convoluted spectra (Right) from the

spectroelectrochemical reduction of [RuO4]-. ......................................................................................... 99

Figure 5-8 Structure of Bis-2-hydroxy-2-ethylbutyrato(oxo)-ruthenate(V). ................................................ 100

Figure 5-9 Experimental (A) and simulated (B) voltammetry of 0.8 mM n-Pr4N[RuO4] in MeCN. I = 0.1 M

(Bu4N)(BF4). ......................................................................................................................................... 102

Figure 5-10 Experimental (A) and simulated (B) CVs of 0.8 mM [RuO4]- in MeCN with NMO

concentrations of 0 mM (red), 10 mM (green), 20 mM (yellow) and 30 mM (blue). .......................... 102

Figure 5-11 Cyclic voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN with pyridine N-oxide concentrations

of 0.0 mM (red), 10.0 mM (green), 20.0 mM (yellow) and 30 mM (blue)........................................... 103

Figure 5-12 Various N-oxides relevant to this work and corresponding N+–O- bond lengths determined from

the corresponding crystal structures...................................................................................................... 104

Figure 5-13 Experimental (A) and simulated (B) CVs of 0.8 mM [RuO4]- in MeCN with TMNO

concentrations of 0 mM (red), 10 mM (green), 20 mM (yellow) and 30 mM (blue). .......................... 104

Figure 5-14 Structural valance forms of pyridine N-oxide. .......................................................................... 105

Figure 5-15 Cyclic voltammetry of 1 mM n-Pr4N[RuO4] in MeCN with 0.0 mM (red), 10 mM (green), 20

mM (yellow) and 30 mM (blue) NMO - A or pyridine N-oxide - B. ................................................... 106

Figure 5-16 n-Pr4N[RuO4] synthesis, post-reaction. Flask A contains n-Pr4N[RuO4] in solution. .............. 111

Figure 6-1 Time-resolved spectra following the oxidation of 12.5 mM diphenylmethanol by 0.25 mM n-

Pr4N[RuO4] and 67 mM NMO in MeCN (303 K). ............................................................................... 115

Figure 6-2 A) Spectrum of 1.00 × 10-2 M benzophenone in MeCN. ʎmax ~ 336 nm, ɛ = 119.7 M-1 cm-1. B)

Time resolved spectra from Figure 6-1 with the ‘constant’ spectrum of n-Pr4N[RuO4] subtracted. .... 116

Figure 6-3 Maximum rate of oxidation during the induction and catalytic periods is determined by the slope

of the steepest tangent within each region. ........................................................................................... 117

Figure 6-4 [RuO4]--dependent kinetics.. ....................................................................................................... 117

Figure 6-5 [Diphenylmethanol]-dependent kinetics. .................................................................................... 118

vii

Figure 6-6 [NMO]-dependent kinetics. ........................................................................................................ 118

Figure 6-7 [H2O]0-dependent kinetics. ......................................................................................................... 120

Figure 6-8 1H NMR spectra (500 MHz) of NMO (brown), 1-octanol (purple) and 1:1 NMO:1-octanol

(green) in d3-acetonitrile. ...................................................................................................................... 121

Figure 6-9 X-band (νav = 9.7041 GHz) CW EPR spectra measured for n-Pr4N[RuO4] in acetonitrile, with

differing additives. A) n-Pr4N[RuO4], B) n-Pr4N[RuO4] + NMO, C) n-Pr4N[RuO4] + NMO +

diphenylmethanol.................................................................................................................................. 124

Figure 6-10 X-band HYSCORE spectrum at 4K recorded in deuterated acetonitrile of A) n-Pr4N[RuO4] with

no additives (B0 = 352.9 mT, τ = 120 ns, ν = 9.561267 GHz), B) n-Pr4N[RuO4] + NMO (B0 = 358.65

mT, τ = 108 ns, ν = 9.733610 GHz). ..................................................................................................... 125

Figure 6-11 Hydrogen-bonded NMO/[RuO4]- adduct. ................................................................................. 126

Figure 6-12 Left) UV-vis and concentration profile for benzophenone (inset) during a reaction between 0.25

mM n-Pr4N[RuO4], 6.0 mM diphenylmethanol and 60 mM NMO in MeCN (T = 303 K). Right) frozen

X-band spectra (νav = 9.7041 GHz) measured at the intervals indicated in the inset (T = 6K). ............ 126

Figure 6-13 X-band (νav = 9.6766 GHz) CW EPR spectra measured at 6 K showing the decay of

perruthenate EPR signal over time after addition of substrate alcohol in the absence of co-oxidant

NMO. .................................................................................................................................................... 127


Pr4N[RuO4] and 60 mM NMO in MeCN (303 K). A) – no added RuO2.2H2O, B) 16 µL of RuO2.2H2O

stock solution added at t0. ..................................................................................................................... 128


Pr4N[RuO4] (commercial) and 60 mM NMO in MeCN (303 K). ......................................................... 132

viii

LIST OF ABBREVIATIONS

ATRP Atom transfer radical polymerization

BnBr Benzyl bromide

CCT Catalytic chain transfer

CRT Catalytic radical termination

CSD Cambridge structural database

CV Cyclic voltammetry

D Diffusion coefficient (cm2 s-1)

DCM Dichloromethane

∆H0298K Bond dissociation enthalpy in the gas state

DMF Dimethylformamide

DMSO Dimethylsulfoxide

DPn Degree of polymerisation

E Potential applied at the working electrode

E0’ Formal redox potential

EN T

Normalised Dimroth-Reichardt parameter

EBriB Ethyl 2-bromoisobutyrate

EPR Electron paramagnetic resonance (spectroscopy)

EtOH Ethanol

Et6tren Tris[2-diethylamino(ethyl)]amine

Fc Ferrocene

EXAFS Extended x-ray absorption fine structure analysis

FRP Free radical polymerisation

I Ionic strength

ip Peak current

IUPAC International union of pure and applied chemistry

KIE Kinetic isotope effect

k0 Standard heterogeneous rate constant (cm s-1)

LAM Less active monomer

M Olefinic monomer

ix

MBrP Methyl 2-bromopropionate

MeOH Methanol

Me6tren Tris[2-dimethylamino(ethyl)]amine

MeCN Acetonitrile

Mn Number average molecular weight

Mw Weight average molecular weight

NIR Near infrared (spectroscopy)

NMM N-methylmorpholine

NMO N-methylmorpholine N-oxide

NMR Nuclear magnetic resonance (spectroscopy)

OMRP Organometallic mediated radical polymerisation

(DT) OMRP Degenerative transfer OMRP

(RT) OMRP Reversible termination OMRP

Pn• Propagating polymer radical

PnX Halogen-capped polymer chain

PMDETA N,N',N',N'',N''-pentamethyl-diethylenetriamine

PRE Persistent radical effect

RX Alkyl halide initiator

RDRP Reversible deactivation radical polymerisation

SP-PLP Single pulse-pulsed-laser polymerisation

TBAP Tetrabutylammonium perruthenate

TEMPO 4-amino-2,2,6,6-tetramethylpiperidine-1-oxyl

TMNO Trimethylamine N-oxide

tpa Tris-[2-pyridyl(methyl)]amine

tpa* Tris-[4-methoxy-3,5-dimethyl(pyridin-2-yl)methyl]amine

TPAP Tetrapropylammonium perruthenate

tren Tris[2-amino(ethyl)]amine

UV-vis Ultraviolet & visible (spectroscopy)

1

Redox reactions catalysed by transition metal complexes

Introduction

The rate determining step in a variety of synthetic organic reactions involves breaking an inert

chemical bond. In the absence of an external thermodynamic driving force these reactions proceed at

negligible rates. However, they can sometimes be accelerated by a transition metal complex. If the

metal complex is consumed during the reaction a stoichiometric quantity is necessary. A more

desirable scenario is one in which the added complex is regenerated throughout by a second chemical

reaction. In this case, only a catalytic quantity is required which is advantageous in terms of cost,

experimental workup and for various environmental reasons.

Two classes of reaction which rely on such transition metal complexes are atom transfer radical

polymerisation (ATRP)1-2 and the Ley-Griffith oxidation of alcohols3 (Scheme 1-1 and Scheme 1-3).

The overarching aim of this thesis is to provide a mechanistic understanding of the central reactions

which underpin these two methods. While there are many differences between the two, both

ultimately rely on a homogeneous redox reaction involving a transition metal complex, referred to

informally as a ‘catalyst’. The function of the catalyst in both cases is to effect the rupture of an inert

chemical bond. This reaction consumes the starting complex but a second reaction regenerates it

throughout, leading to the use of the term ‘catalyst’. It can be argued that these are not catalysts in

the formal sense as defined by IUPACa however the term is applied ubiquitously throughout the

relevant literature and is therefore adopted (informally) here.

ATRP

Various transition metals, including those in groups 4 (Ti)4, 6 (Mo)5, 8 (Fe)6-7 and 10 (Ni & Pd)8-9

have been used in ATRP, but the vast majority of studies utilise Cu10-16 and the highest efficiencies

have been achieved by very simple copper(I) complexes chelated by multi-dentate, N-donor ligands

‘L’.17-19 The catalyst is therefore CuIL and the reaction it catalyses is the homolysis of a carbon-

halogen bond through a halogen atom transfer reaction. Adding CuIL to a solution containing an alkyl

halide initiator ‘RX’ forms an alkyl radical (R•) and the corresponding copper(II) complex CuIILX

a IUPAC defines a catalyst as “A substance that increases the rate of a reaction without modifying the overall standard

Gibbs energy change in the reaction. The catalyst is both a reactant and product of the reaction.”

https://goldbook.iupac.org/html/R/R05163.html

https://goldbook.iupac.org/html/P/P04861.html

2

(Scheme 1-1). This process is coined ‘activation’ and is characterised by the rate constant kact. During

activation the coordination number and the oxidation state of the central metal ion increase by one.

Scheme 1-1 Central ATRP reactions. R-R is the radical-radical coupled product; RH and R= are the radical-radical

disproportionation products.

The key to the success of ATRP is reversibility of the halogen atom transfer. The reverse reaction

between R• and CuIILX regenerates the catalyst and re-caps the alkyl radical; this process is called

‘deactivation’ and is characterised by the rate constant kdeact. Typically, kdeact is several orders of

magnitude larger than kact so the position of the central equilibrium is biased strongly towards the left

(KATRP = kact/kdeact = 10-7 – 10-3).20 This ensures that the concentration of radicals is kept low

throughout the reaction. Radicals produced by activation initiate (or continue) chain propagation (kp)

adding olefinic monomers sequentially to a growing chain before they are rapidly re-capped by

deactivation. In this way, propagating polymers spend most of the reaction in the dormant halogen-

capped form. As a result, each chain is extended in a controlled fashion and, by minimising the

concentration of radicals, unwanted termination products (kt) are avoided. For these reasons ATRP

produces polymers with well-defined molecular weights and molecular weight distributions21 and in

some instances ultra-high molecular weight products can be synthesised.22-23

During the initial stage of an ATRP reaction the concentrations of CuIILX and R• are low which

favours termination (rate = 2kt[R•]2) over deactivation (rate = kdeact[R

•][CuIILX]); kt and kdeact being

of comparable magnitudes.24-25 While this forms the undesirable termination products, it also

produces a buildup of CuIILX which redirects the position of the central equilibrium further to the

left by a process known as the persistent radical effect (PRE).26-28 This provides additional control to

the polymerisation and is essential for producing well-defined products. The transition metal complex

therefore confers several advantages to the reaction.

Overall, ATRP is one of a number of techniques for reversibly forming an alkyl radical from a capped

precursor for the purpose of initiating/continuing a polymerisation (Scheme 1-2). Each of these

techniques falls under the general banner of ‘reversible deactivation radical polymerisation’

(RDRP)29 and this term will be referred to several times throughout this thesis.

3

Scheme 1-2 Popular methods for RDRP include RAFT – Reversible Addition-Fragmentation Chain Transfer, NMP –

Nitroxide Mediated Polymerisation, ITP – Iodine Transfer Polymerisation, ATRP – Atom Transfer Radical

Polymerisation and DT/RT OMRP – Degenerative Transfer/Reversible Termination Organometallic Mediated Radical

Polymerisation.

Ley-Griffith Oxidation

The Ley-Griffith protocol utilises a transition metal catalyst to cleave the α-C–H bond of alcohols,

forming the oxidised aldehyde or ketone products. The reaction is performed in organic solvent

(typically dichloromethane or acetonitrile) using a catalytic quantity of tetrapropylammonium

perruthenate (TPAP) and a stoichiometric amount of the co-oxidant N-methylmorpholine N-oxide

(NMO – Scheme 1-3).30 Dry, powdered molecular sieves are added to the solution to remove water

generated by the reaction and dry reagents and solvent are also generally utilised. Apart from the

necessity for dry conditions,31 the reaction is straightforward and oxidises a variety of primary and

secondary alcohols to their products in high yields in less than one day.32 A particular advantage of

the technique is that it selectively oxidises primary alcohols to aldehydes without over-oxidation to

the corresponding acid. Moreover, this is achieved using reagents which are reasonably benign. Very

little is known about the specific redox reactions involved in this process despite its widespread

adoption by synthetic chemists.33-34 A more comprehensive description of the development and

application of this technique is given in the introduction to Chapter 5.

4

Scheme 1-3 The Ley-Griffith alcohol oxidation utilising TPAP and NMO

Conceived in 1995 and 1987 respectively, ATRP and the Ley-Griffith oxidation have since garnered

widespread attention in both industrial and academic settings because they utilise fairly simple metal

reagents to efficiently conduct their targeted reactions with high selectivity and at reasonable rates.

There are thousands of reports in the literature related to the application or assessment of ATRP and

hundreds of papers which utilise the Ley-Griffith protocol. Yet despite the prevalence of these

methods, various mechanistic details remain obscured.

For example, very little is known about the deactivation reaction in ATRP. It is widely acknowledged

that the solvent, the catalyst and the identity of the halogen capping agent each affect the position of

the central equilibrium (KATRP).20, 35-37 However, this effect has mostly been rationalised in terms of

changes to the activation rate (kact).38-48 Activation rates are much easier to measure than deactivation

rates because they are slower (kact ~ 10-3 – 103 M-1 s-1, kdeact ~ 104 – 108 M-1 s-1).20 Accordingly, the

activation rate constant has been measured with different catalysts,39, 41, 46, 49-50 solvents43, 48, 51 and

initiators38, 45, 47, 52-53 and the effects of each on the kinetics are now well established.

Measuring the deactivation rate constant is more difficult as the reaction can approach diffusion-

controlled limits (107 – 108 M-1 s-1). Up until the commencement of this work, two methods were

utilised to determine kdeact. The most common method relied on separately measured values of KATRP

and kact to determine kdeact empirically from the relation KATRP = kact/kdeact.20, 40

A second, experimental method has also been applied in a limited number cases using a competitive

clock reaction.40 Here a radical R•, which mimics the propagating polymer chain, is generated by

thermal dissociation of a stable nitroxide precursor (Scheme 1-4).54-56 In the presence of both the

deactivating complex CuIILX and TEMPO, R• is competitively consumed to generate the dormant

organo-halide (RX) by deactivation and the radical-TEMPO adduct by radical-radical termination

(kcomb2). The ratio of the products at the end of the reaction is related to the relative rates of kdeact and

kcomb2. Where kcomb2 is known, kdeact can be calculated from Equation 1.1. This equation holds where

kcomb1 and kdis2 are small – this being the case for the nitroxide radical precursors utilised.57

5

Scheme 1-4 The competitive clock reaction for determining kdeact. With kcomb1 and kcomb2 known, kdeact can be calculated

from Eqn. 1.1.

d[R-X]

d[R-TEMPO] =

kdeact

kcomb2

[CuII

LX][R•]

[TEMPO][R•]

kdeact = kcomb2

[TEMPO]

[CuIILX]

[R-X]

d[R-TEMPO] (1.1)

In 2012 a new method was published for measuring kdeact using single-pulse‒pulsed laser

polymerisation coupled to electron paramagnetic resonance spectroscopy (SP‒PLP‒EPR).58 Here, a

solution of an olefinic monomer was irradiated with a single laser pulse to photo-initiate the formation

of the radical. The time resolved decay of the radical concentration was followed by EPR in the

absence and presence of CuIILX. In the absence of CuIILX, the signal diminished as radical-radical

termination (rate = 2kt[R•]2) occurred and kt was isolated by fitting a second order decay function to

the radical concentration profile. In the presence of the deactivator, the radical was consumed by

concurrent termination and deactivation. Therefore, Equation 1.2 was used along with the determined

value of kt to calculate kdeact. Since its inception the SP-PLP-EPR method has been utilised to measure

deactivation rates with both copper58 and iron catalysts.59

d[R•]

dt = -2kt[R

•]2 – kdeact[CuII

LX][R•] (1.2)

6

Each of these methods suffers from a number of practical drawbacks. The empirical method for

determining kdeact is only applicable to systems for which kact can be measured experimentally; that

is, systems with lower activation rate constants. The competitive clock reaction relies on reagents

which are reasonably difficult to synthesise and there are only a few published values of kcomb2 which

further limits the application of this technique. The EPR method is the most promising of the three

however the copper(II) deactivator complex also absorbs within the region of the radical signal and

this is not taken into account when fitting the time-resolved decay profile of the radical signal. The

cumulative result of these limitations is that the number of accurately (independently) determined

values of kdeact in the literature is far outweighed by the corresponding values of kact.

This is problematic in light of one of the major challenges facing ATRP – the development of catalysts

which have very high activation rates for the polymerisation of particularly inert monomers.60-62 The

rate of polymerisation (from Scheme 1-1) is given by Equation 1.3 where KATRP appears in the

numerator. More active catalysts therefore lead to faster polymerisations. However, the molecular

weight distribution of the polymer products (Mw/Mn) is only kept narrow (small Mw/Mn) if

deactivation is also fast (Equation 1.4 – DPn is the degree of polymerisation and p is the monomer

conversion).63 Therefore an increase in kact must be accompanied by an increase in kdeact if control

over the propagating radical concentration is to be maintained. Understanding the various influences

of the catalyst, the solvent and the monomer/initiator on the deactivation rate is essential for the

rational development of these systems but a method for rapidly and accurately determining kdeact is

not yet available.

KATRP=[Cu

IILX][Pn

• ]

[CuIL][PnX]

= kact

kdeact

d[M]

dt = kp[M][Pn

• ] = kpKATRP[PnX][CuIL][M]

[CuIILX]

(1.3)

Mw

Mn

= 1 + 1

DPn

+ (kp[PnX]

kdeact[CuIILX]

) (2

p – 1) (1.4)

7

Scheme 1-5 Central OMRP reactions

Another aspect which requires further attention is the competition between atom transfer and radical

transfer. In addition to termination, propagation or deactivation, radicals formed under ATRP

conditions may also react with CuIL to generate an organometallic complex (CuIILR) in which R is

formally a carbanion ligand. A separate branch of RDRP reactions, known as ‘organometallic

mediated radical polymerisations’ (OMRP), rely on such complexes to control the release of

propagating radicals through homolytic scission of their metal-carbon bonds (Scheme 1-5).64-67 Like

ATRP, the central equilibrium of OMRP involves a transition metal complex which undergoes an

increase in coordination number and oxidation state; however, the positions of these equilibria are

diametrically opposed (KATRP << 1; KOMRP >> 1). While the goal in both processes is the controlled,

reversible release of radicals, the capping agent is different.

The possibility of concurrent atom transfer and radical transfer has recently been realised for some of

the most active Cu-based ATRP systems68-71 and is invoked in order to explain unusual experimental

phenomena.72 However, the organometallic complex has never been directly observed in these

systems. Furthermore, there is virtually no understanding of what chemical or physical factors

expedite radical versus atom transfer where both are possible. Thus the fate of radicals during ATRP

is particularly important for providing control to the system, however there is still much that is

unknown about these reactions.

Cyclic voltammetry including heterogeneous and homogeneous electron transfer

Both the atom transfer and the proposed radical transfer pathways in ATRP involve a change in the

oxidation state of the central metal ion and a coupled chemical reaction. In this regard, cyclic

voltammetry (CV) is a uniquely suitable technique for studying these reactions in order to elucidate

various mechanistic details. A single CV experiment can reveal a host of qualitative and quantitative

information about redox and chemical reactions in solution. While several analytical techniques are

employed throughout this project, a heavy emphasis is placed on cyclic voltammetry coupled to

simulations. Accordingly, this section provides a brief introduction to CV experiments which involve

coupled electron transfer and chemical reactions.

8

Figure 1-1 Simulated cyclic voltammogram for the reversible, one-electron reduction of O. Arrow indicates the starting

position and direction of the sweep. Sweep rate = 100 mV s-1; [O] = 1.0 mM. Diffusion coefficients of O and R (DO = DR)

= 1.0 × 10-5 cm2 s-1. α = 0.5.

Figure 1-2 Schematic representation of the reversible redox reaction involving O and R at a planar electrode. Fick’s

second law determines the rate of diffusion (flux) of O and R towards and away from the electrode.

The cyclic voltammogram of a reversible, single electron transfer reaction involving the soluble

species ‘O’ and ‘R’ is shown in Figure 1-1. At the beginning of the sweep only the oxidised species

‘O’ is present in solution and the applied potential (E) is positive of the formal redox potential of O

(E0'). At this point, only a small non-faradaic current flows as the electric double layer is established.

If the kinetics of the electron transfer reaction (k0) are fast then the applied potential defines the ratio

of the concentrations of O and R at the electrode surface ([O]0/[R]0) by the Nernst equation (Eqn.

1.5). As E is swept in the negative direction and approaches E0', species ‘O’ is reduced within a thin

layer of the solution (thickness of μm) adjacent to the electrode surface by a heterogeneous electron

transfer reaction; as a result, faradaic current begins to flow.

One might expect that at more extreme negative potentials, the current should continue to increase as

predicted by the Butler-Volmer equation (Eqn. 1.6). However, at these potentials, the concentration

of O at the electrode surface is depleted so the current is instead diminished. Thus, the current peaks

E0'

-ve +ve

Potential

𝑖p

𝑥 = 0 𝛿

∂[R](𝑥,t)

∂t= 𝐷O (

∂2[R](𝑥,t)

∂𝑥2)

∂[O](𝑥,t)

∂t= 𝐷O (

∂2[O](𝑥,t)

∂𝑥2)

9

at a value of ‘𝑖p’ and then diminishes as the sweep continues beyond E0'. The reason the current does

not return to zero is because O is resupplied to the electrode by mass transport from the bulk solution

(Figure 1-2). By using a high concentration of electrolyte and an unstirred solution, diffusion is the

sole, relevant form of mass transport.

E = E 0'+RT

Fln

[𝑂]0

[𝑅]0

(1.5)

𝑖 = -FAk0 [[O]te-αf (E – E0') – [R]te

(1 – α)f (E – E0')] (1.6)b

E = Ei – ν t (1.7)

During the forward (cathodic) sweep, the concentration of O has a minimum value [𝑂]0 at the surface

of the electrode (𝑥 = 0) and a maximum value of [𝑂]𝛿 in the bulk solution so a concentration gradient

is established. The zone between the electrode and the point where [O] becomes constant is defined

as the diffusion layer (thickness of δ cm) and the concentration gradient is approximately linear across

this range (Figure 1-3). It is the slope of this gradient which ultimately defines the rate of diffusion

(or ‘flux’) of O towards the electrode by Fick’s second law (Figure 1-2 inset equation).

For an unstirred solution the diffusion layer thickness depends on the elapsed time ‘t’ (δ ~ (Dt)1/2

where 𝐷 is the diffusion coefficient of species O in cm2 s-1) and the applied potential (through the

Nernst equation).73 These two variables are related by Equation 1.7 which describes the applied

potential as a function of time and sweep rate (v). This deceptively simple equation subsumes a wealth

of information about the diffusion layer thickness.

b 𝐸0′ is the formal redox potential of O. The coefficient ‘R’ in eqn. 1.5 is the universal gas constant. CO(t) and CR(t) are

the concentrations of O & R at the electrode surface at time ‘t’, k0 and α describe the rate and ‘symmetry’ of the

heterogeneous electron transfer reaction. ‘A’ is the surface are of the electrode f = F/RT. ν is the scan rate (mV s-1).

Note: The morphology of the electrode affects the diffusional process and throughout this thesis, only the case of a planar

electrode is considered.

10

Figure 1-3 Time-dependent concentration gradient profile for O at a three fixed potentials. Solid line – no applied

potential; dashed line – E < E0', dotted line – E << E0'.

At very negative potentials (longer timescale, larger δ and ‘complete’ removal of O at the electrode

surface - Figure 1-3 dotted profile) a steady state current is achieved where the rate of electron transfer

is limited by the diffusion rate; this results in the plateau region observed in Figure 1-1. Here the

current is exactly matched by the flux of O.

At the switching potential, the scan direction is reversed and essentially the same process in reverse

(i.e. for R) describes the oxidative current. The reversibility of the profile in Figure 1-1 demonstrates

that the electron transfer kinetics (described by the rate constant k0) are fast (electrochemical

reversibility) and that the reduced and oxidised forms of the reagent are stable on the timescale of the

experiment (chemical reversibility).

[O]

𝑥

0

t2 < t

2 < t

3

[O]

11

Figure 1-4 Simulated CV waveforms for the ECcat mechanism as a function of [A]. Red – 0.0 mM, Green – 1.0 mM,

Yellow – 10.0 mM, Blue – 50.0 mM A. Rate constant k = 1.0 × 103 M-1 s-1. All other parameters are identical to Figure

1-1.

Figure 1-5 Representation of a reversible electron transfer at the electrode coupled to an irreversible, catalytic chemical

reaction in solution.

When a homogeneous electron transfer reaction is coupled to heterogeneous electron transfer the

situation is more complex. For example, if species ‘A’ is added to the solution – this species being

electrochemically inert in the scanned potential window – a homogeneous, second order chemical

reaction with R may yield the product ‘B’ and concomitantly regenerate O (Figure 1-5). This

mechanism is formally an electron transfer (‘E’ step) coupled to a catalytic chemical reaction (‘C’

step) so the overall process is termed an ECcat mechanism. When the voltammetry is remeasured with

the addition of A, the cathodic current is amplified because O is now replenished by diffusion and

the homogeneous electron transfer reaction (Figure 1-4). The rate limiting step in this process is

catalytic regeneration of O which proceeds at a rate = k[A][R] (Equation 1.8). The current, defined

by the concentration of O at various distance/time (𝑥, 𝑡) intervals (i.e. the flux), is now governed by

-ve +ve

Increasing [A]

Potential

∂[A](𝑥,t)

∂t= DA (

∂2[A](𝑥,t)

∂𝑥2)

𝑥 = 0 𝛿

12

the applied potential as well as the rate at which A (usually in excess) is consumed by the catalytic

reaction.

Rate = k[A][R] = kobs[R] (1.8)

kobs = k[A]

Saveant and co-workers have rigorously mapped out the various CV waveforms for an ECcat

mechanism which change according to the two dimensionless parameters λ and γ, taken as

coordinates of the “kinetic zone diagram” shown in Figure 1-6.74,c This figure is particularly helpful

because it concisely summarises how an ECcat voltammogram is defined by the magnitude of the

various physical and chemical parameters which constitute it, including kobs. Several of these zones

are particularly relevant to the work discussed herein.

Figure 1-6 Kinetic zone diagram and simulated CV waveforms as a function of the dimensionless parameters λ and γ.

No catalysis

This region amounts to the reversible behaviour already described. The chemical reaction proceeds

at a negligible rate due to slow catalysis (bottom section) or too little substrate (top left) and bears no

influence on the voltammetry of O.

c Note: where the concentrations of O, R and A etc. are given without a subscript, these refer to the concentrations in bulk

solution before a potential is applied to the electrode.

0

log(

λ)

log(γ)

NO

CA

TA

LY

SIS

K

KS

𝑣

𝑘

[A]

[R]

0 -1 -2 1 2 3

0

1

2

3

-1

-2

KT

λ = (RT

F) (

𝑘[O]

𝑣)

γ = [A]

[O]

KD

KG

13

Pure kinetics conditions – zone KS

This behaviour occurs when the substrate (A) is rapidly reduced by R within the diffusion layer. A

specific set of conditions are required to observe such behaviour: the heterogeneous electron transfer

process must be fast; ‘A’ must be in a pseudo first order excess ([A] >> [O]); the homogeneous

electron transfer (chemical) step must be rate limiting; the diffusion coefficients of O and R must be

close to identical (DO/DR = 1), and A must not exhibit redox activity at the electrode within the

potential window which is measured.

The resulting waveform is characteristically S-shaped and tends towards a limiting current at low

potentials. In this region, the rate of electrolysis of O is exactly compensated for by its replenishment

through diffusion and the catalytic reaction. If the chemical reaction is too fast (large γ), substrate

depletion will occur (Figure 1-7). This may be experimentally countered by using a very high

concentration of A (large λ) or by increasing the sweep rate to minimise the timescale for reaction.

Pure catalytic conditions are particularly desirable because the rate constant k can be extracted

mathematically from the limiting current. The mathematical theory describing the CV response of an

ECcat mechanism was developed by Delahay and Stiehl,75 Nicholson and Shain,76-77 and Saveant and

co-workers74, 78-81 who demonstrated that when pure kinetic conditions are fulfilled, the current is

described by:

𝑖 =-𝑛FA[O]√𝐷𝑂kobs

1 + exp [𝑛FRT

(E – E0'

)] (1.9)

This equation produces the characteristic, symmetrical, S-shaped waveform with the anodic sweep

exactly tracing the cathodic sweep. At sufficiently negative potentials the second term in the

denominator disappears leading to

𝑖lim = -𝑛FA[O]√DOkobs (1.10)

from which kobs can be extracted directly if the diffusion coefficient and electrode area are known.

This expression also indicates that the current will tend towards the same limiting value at low

potentials independent of the sweep rate.

14

Substrate diffusion – zones K and KG

If turnover of the catalyst leads to the depletion of [A]0 then diffusion from the bulk limits the current

at low potentials and the waveform changes. Under these conditions it is not possible to generate a

closed-form E-i equation (like Eqn. 1.9) to describe the voltammetry. Instead, numerical methods

must be employed. An explanation of this approach is beyond the scope of this review and is covered

in detail elsewhere.76 Simply put, a dimensionless E-i expression must be generated and solved

numerically for discrete potential steps, the results of which are then plotted to yield the waveform

(e.g. Figure 1-6). Electrochemistry software packages such as Digisim82 perform this calculation

computationally and can generate a simulated voltammogram based upon an input mechanism along

with the values which are defined for the associated physical and chemical constants (vide infra).

Qualitatively, the effect of substrate depletion during catalysis is similar to that already discussed for

a reversible CV where an expanding diffusion layer thickness leads to a peak followed by a drop in

current at high overpotential. Here the contribution of diffusion becomes more and more significant

with larger values of γ and the waveform is increasingly distorted away from the classical S-shaped

profile of the pure kinetics region (Figure 1-7).

Figure 1-7 Simulated CV waveforms for the ECcat mechanism as a function of the rate constant k. Red – k = 1.0 × 102 M-

1 s-1, Green – k = 1.0 × 103 M-1 s-1, Yellow – k = 1.0 × 104 M-1 s-1, Blue – k = 1.0 × 105 M-1 s-1. [A] = 50.0 mM. All other

parameters are identical to Figure 1-1.

The compass rose in the bottom left of Figure 1-6 indicates that this behaviour will be observed for

particularly active catalysts (large k) or when an insufficient excess of A is employed. It is also

possible to migrate across the substrate depletion boundary to the pure kinetic zone by decreasing the

scan rate. This allows sufficient time for the substrate to be replenished within the reaction layer

following the chemical reaction.

-ve +ve

Increasing k

Potential

15

If A is consumed within the reaction layer during the forward sweep and is not replenished rapidly

by diffusion, then a peak current may be observed during the anodic sweep (zone KG). This situation

can also be encountered when an insufficient excess of A is added to the cell.

Total catalysis – zone KT

When the catalytic reaction is exceptionally fast, only a small amount of R is required for the turnover

of A. At the onset of O reduction, a small amount of R is produced and the substrate A is immediately

consumed within the diffusion layer to regenerate O. The voltammogram displays a catalytic peak at

E < E0' and then a second wave appears at E0' as the remainder of O is reduced.

Figure 1-8 Representation of a reversible electron transfer at the electrode coupled to a reversible, catalytic chemical

reaction in solution. The product B is removed by a rapid, irreversible second order reaction.

An alternate form of the ECcat mechanism is presented in Figure 1-8 in which heterogeneous electron

transfer is coupled to a reversible chemical reaction as well as a second, irreversible reaction. The

flux of A towards the electrode now depends on the rate at which it is consumed by the catalytic

reaction (k1[R][A]) as well as the rate of the back reaction (k-1[O][B]) and the second order

consumption of B (2k2[B]2).

Again a numerical approach is required in order to describe the E-i response; the full derivation of

the dimensionless expressions which lead to these waveforms is found elsewhere.80 It has been shown

that when k2 is large, a steady state approximation can be applied to [B] throughout the sweep and

the following three parameters characterise the voltammetry: 𝛾, 𝜆1 and 𝜆/𝜆2 (see below).

𝑥 = 0 𝛿

16

γ = [A]

[O] (𝟏. 𝟏𝟏) λ1 = (

RT

F) (

k1[A]

ν) (𝟏. 𝟏𝟑)

λ = (RT

F) (

k2

ν) (𝟏. 𝟏𝟐) λ2= (

RT

F) (

k-1[O]

ν) (𝟏. 𝟏𝟒)

Two limiting scenarios can be envisaged based on the ratios of k2 and k-1. If k2 << k-1 then kinetic

control of the overall chemical reaction is provided by the self-termination of B with the reversible

chemical reaction acting as a pre-equilibrium. The two parameters which the system depends on are

𝜆𝜆1/𝜆2 and 𝛾.

Alternatively, if k2 >> k-1 then kinetic control of the overall chemical reaction is conferred by the

forward reaction between A and R. Under these conditions the third parameter (𝜆/𝜆2) becomes

irrelevant and the voltammetry can be described by 𝜆1 and 𝛾 in a manner analogous to that shown in

Figure 1-6. Thus, changing the sweep rate, the rate constant k1 and the concentrations of A and O

elicit the same effects on the voltammetry as previously discussed for the case where the catalytic

reaction was irreversible. This behaviour is particularly relevant to the voltammetry of ATRP

catalysts which are the subject of the following Chapter as well as Chapter 4.

Voltammetry of ATRP catalysts

The study of ATRP complexes using electrochemical methods is a recent development in this

burgeoning field. The first use of CV in this vein was to determine the CuII/I redox potentials of a

series of catalysts with differing reactivity.39, 83 Copper(I) complexes can be oxidised by molecular

oxygen and are often unstable, undergoing rapid, bimolecular disproportionation to Cu0 and CuII.

Accordingly, the voltammetry was measured using stable solutions of copper(II) deactivator

complexes (CuIILX) in degassed acetonitrile (a suitable solvent for ATRP). Quasi-reversible

reduction to CuILX was observed in most cases and an inverse correlation between the redox potential

and activity of the catalyst was found; i.e. complexes which facilitated fast activation were formed at

more negative potentials.

The discrepancy between these experiments and the mechanism of ATRP was noted in 2011 by

Genarro et al. who showed that CuILX, which is formed by reducing CuIILX, is not the species which

activates the alkyl halide.84 These authors utilised chronoamperometry to monitor the current

response of a rotating disk electrode immersed in a solution containing only CuIL. In the absence of

any other reagents, a diffusion limited current was observed when CuIL was oxidised to CuIIL (solvent

bound in place of X) and the current corresponded to the Levich equation (Eqn. 1.15).

17

𝑖 = 0.62FAD 2/3ν-1/6ω1/2[CuIL]

* (1.15)

Levich Equation: 𝐷 is the diffusion coefficient of CuIL, 𝜔 is the angular velocity of the rotating disk electrode (s-1) and

[CuIL]* is the concentration of CuIL in the bulk solution.

When an alkyl halide was added to the cell in a pseudo first order excess, CuIL was consumed in the

bulk solution by the activation reaction at a rate = kobs[CuIL]* (kobs = kact[RX]) so the current from

oxidative electrolysis also diminished over time. The back reaction (deactivation) was prevented by

adding a large excess of the radical trap 2,2,6,6-tetramethylpiperidine 1-oxyl (TEMPO). From the

current-decay rate, the value of kobs was extracted and plotted for various concentrations of RX to

yield kact.

When the experiment was repeated using a solution of CuILX instead of CuIL, activation proceeded

at a negligible rate leading the authors to conclude that the ‘active’ catalyst is CuIL. While this paper84

demonstrated the power of electrochemistry in accessing mechanistic and kinetic information about

ATRP catalysts, the chronoamperometric method suffered from a number of limitations. For instance,

alkyl halides which were highly active (kact > 10 M-1 s-1) could not be used in a pseudo first order

excess because they consumed the copper(I) catalyst almost immediately so the current decay was

too fast to follow. Using lower concentrations of these initiators complicated the kinetic analysis.

Furthermore, preparing solutions of copper(I) compounds is not ideal for the reasons described

earlier.

Later in the same year, an alternate method for determining kact was reported by Bernhardt et al. using

cyclic voltammetry on stable solutions of CuIILX complexes.51 The chelate ligand was the same N-

donor polyamine used by Gennaro, tris-[2-dimethylamino(ethyl)]amine (Me6tren – Scheme 1-6);

acetonitrile and dimethylsulfoxide were used as solvents and bromide was the auxiliary halide ligand

‘X’. In the absence of other additives, [CuII(Me6tren)Br]+ was reversibly reduced to [CuI(Me6tren)Br]

by sweeping the potential in the negative direction (Figure 1-9).

When the alkyl halide initiator ethyl 2-bromoisobutyrate (EBriB - Scheme 1-6) was added to the cell

the waveform became distinctly catalytic. The radicals formed by homolysis of the carbon-halogen

bond of EBriB mimic propagating methyl methacrylate radicals making this one of the most popular

initiators. Increasing the concentration of EBriB led to an amplified cathodic current and a diminished

anodic peak on the return sweep with the wave becoming increasingly asymmetric at higher

concentrations of EBriB. As described earlier, this behaviour (shown in Figure 1-9) is typical for a

catalytic system in which substrate depletion occurs within the diffusion layer. The diffusion limited

behaviour was conserved even with higher concentrations of EBriB which precluded the use of the

steady-state equations (Eqn. 1.9 and Eqn. 1.10) to determine the catalytic rate constant. Therefore,

18

the authors determined kact by simulating the voltammetry across a range of sweep rates and

concentrations of EBriB using Digisim.

Figure 1-9 Experimental and simulated cyclic voltammograms of 1.0 mM [CuII(Me6tren)Br]Br at four different

concentrations of EBriB. Sweep rate = 50 mV s-1, I = 0.1 M (Et4N)(ClO4).

Scheme 1-6 Catalytic mechanism for initiator activation used to simulated the voltammetry shown in Figure 1-9.

simulation

E (mV vs Fc+/0

)

-1200 -1000 -800 -600 -400 -200 0

no EBriB

0.6 mM EBriB

1 mM EBriB

2.5 mM EBriB

10 mM EBriB

20 A

experimental

E (mV vs Fc+/0

)

-1200 -1000 -800 -600 -400 -200 0

no EBriB

0.6 mM EBriB

1 mM EBriB

2.5 mM EBriB

10 mM EBriB

20 A

19

The reaction mechanism which was used in the simulations is shown in Scheme 1-6 and satisfies the

observations of both Matyjasewski83 and Gennaro;84 that CuIILX is reversibly reduced to CuILX but

the active catalyst is CuIL. By starting with the solid complex [CuII(Me6tren)Br]Br, the second

equivalent of Br-, along with the large binding constant KII,Br, ensures that [Cu(Me6tren)Br]+ is the

only species present at the start of the sweep. Scanning the potential in the negative direction forms

[CuI(Me6tren)Br] and the reversibility of the EBr redox couple (in the absence of EBriB) reveals that

[CuI(Me6tren)Br] is stable on the timescale of the experiment. However, solvation of the halido

ligand (KI,Br) is still expected and the position of this equilibrium should lie further towards the solvato

complex than the corresponding equilibrium with CuII (KII,Br) due to the reduced electrostatic affinity

of Br- for Cu+ versus Cu2+. Thus, a small amount of the catalyst still forms.

In the presence of initiator, [CuII(Me6tren)Br]+ is regenerated by activation and the voltammetry

changes. The radicals released during activation undergo rapid self-termination due to their high

concentration and the large termination rate constant (kt ~ 109 M-1 s-1) so the steady state

approximation for [CuIL]+ is valid. The reverse deactivation reaction is also possible so the

mechanism essentially mirrors the one shown in Figure 1-8.

Because of the complexity of the mechanism, kact was determined by simulating the experimental

voltammetry using Digisim. The program performs a least squares regression analysis in which the

difference between the experimental and simulated E-i curves is minimised by allowing the unknown

parameters in the mechanism to refine.

As with any parameterization program, there is a danger of finding a false minimum during the

regression analysis if too many variables are allowed to vary at once. Most of the chemical and

physical constants in Scheme 1-6 were determined independently by additional experiments (EBr, Esol,

KII,Br etc.) or taken from the literature (kt) before attempting the simulation. In fact, the only

parameters which were allowed to resolve during the iterative fitting process were the rate constants

kId,Br and kact. The resulting simulated voltammograms accurately reproduced the experimental

behaviour across the full range of sweep rates and concentrations of EBriB and generated values of

kact which were within the expected range for this combination of highly-active catalyst and initiator.

This electrochemical method, which monitors the transient consumption of [CuIL] within the reaction

layer as opposed to within the bulk, was not limited to slowly activating systems. The values of kact

which the authors reported were two orders of magnitude larger than the uppermost limit of the values

which were determined by chronoamperometry (kact – 104 M-1 s

-1). So far this is the only experimental

method capable of determining kact for such reactive systems. While this represents a significant

advancement in the study of ATRP mechanisms, the ability to measure kdeact is more difficult again

20

and much less is known about the kinetics of this reaction. Expanding the capability of this

electrochemical method to determine kdeact is the subject of the following Chapter.

21

An Electrochemical Method for Determining kdeact

Introduction

In the electrochemical method developed by Bernhardt et al.85 the deactivation rate was fixed at 107

M-1 s-1 while kact was allowed to vary until a satisfactory fit between the simulated and experimental

CVs was obtained. The activation and deactivation rate constants are inextricably linked through the

ATRP equilibrium constant KATRP. Therefore fixing kdeact will necessarily have an impact on the

simulated value of kact. A better experiment would be one in which kact and kdeact are determined

independently without any assumptions.

Scheme 2-1 Simplified mechanism for electrochemically-initiated atom transfer. Loss of X- from CuILX is excluded for

clarity.

It has been shown elsewhere that the activation reaction in ATRP can be made unidirectional by

adding a large excess of TEMPO.50 Under these conditions, the deactivation reaction is quenched as

the olefinic radical is rapidly and irreversibly trapped by the stable nitroxide radical. In the

electrochemical experiment, deactivation consumes CuIILBr at the electrode surface where radicals

are produced and suppresses the catalytic cathodic current (Scheme 2-1). By adding TEMPO to the

electrochemical cell the effect of deactivation on the catalytic current can therefore be abolished and

kact can be determined without any assumptions regarding kdeact. Remeasuring the catalytic

voltammetry under identical conditions but in the absence of TEMPO allows kdeact to be isolated from

the difference in the catalytic currents.

22

Results & Discussion

Cyclic Voltammetry

The stable precursor complex employed in this study is [CuII(PMDETA)Br2] (PMDETA is the linear,

tridentate ligand N,N,N′,N′′,N′′-pentamethyl-diethylenetriamine – Figure 2-1). Copper complexes

chelated by PMDETA are amongst the most widely used catalysts for ATRP86-92 as they facilitate fast

and well-controlled polymerisations. Bromide was used as the halide and DMSO the solvent as these

are both suitable for ATRP.

Crystal structures reveal that copper(II) complexes of PMDETA adopt a square pyramidal geometry

with the chelate occupying three of the four equatorial binding sites (Figure 2-1). Halides bind

strongly to the remaining equatorial site (X) and can also be bound by a more distant interaction at

the axial site (Y).93-95 The axial bonds of these compounds are always slightly elongated due to the

(pseudo/secondary) Jahn-Teller effect.96-97

The complex [CuII(PMDETA)Br2] was synthesised and a 1.0 mM solution prepared in DMSO which

contained 0.1 M (Et4N)(ClO4) as the supporting electrolyte. In solution, the weakly coordinated halide

in the axial position does not remain bound but rapidly exchanges with a solvent ligand;98-99 a number

of crystal structures illustrate this binding pattern.93-95, 100-101 The species formed upon dissolution of

[CuII(PMDETA)Br2] in DMSO is therefore [CuII(PMDETA)Br(OSMe2)]+.

Figure 2-1 Structurally characterised copper(II) complexes of PMDETA. Complexes have the generic formula

[CuII(PMDETA)(X)(Y)]n.

The voltammetry of a 1.0 mM solution of [CuII(PMDETA)Br(OSMe2)]Br in DMSO is shown in

Figure 2-2 and relevant heterogeneous and homogeneous reactions are given in Scheme 2-2. At the

start of the experiment the applied potential is positive of the formal redox potential of the complex.

Sweeping the potential, initially in the negative direction, reveals a quasi-reversible, CuII/I redox

couple (Figure 2-2 - red trace).

23

Figure 2-2 Cyclic voltammetry of 1.0 mM [CuII(PMDETA)Br(OSMe2)]Br in with added EBriB and TEMPO. Sweep rate

= 50 mV s-1. I = 0.1 M (Et4N)(ClO4).

Scheme 2-2 Proposed mechanism for electrochemically induced atom transfer starting from a stable solution of

[CuII(PMDETA)Br2] in DMSO. KI-Br = kIa,Br/kId,Br; KII-Br = kIIa,Br/kIId,Br.

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

+ 3.0 mM TEMPO

+ 5.0 mM TEMPO

5 A

E / mV vs. Fc+/0

24

The maximum number of ligands which can formally coordinate to copper(I) is four.102 Careful

analysis of the structures of copper(I) complexes which are assigned with higher coordination

numbers reveals that at least one of the ‘bonds’ is significantly elongated. Solid-state crystal structures

of CuI/PMDETA complexes are four coordinate adopting slightly distorted, tetrahedral geometry for

example in [CuI(PMDETA)MeCN](ClO4)103 or in the series of η2-coordinated olefin complexes

[CuI(PMDETA)(Y)](BH4) (Y = styrene, octene, methyl acrylate).104 Therefore the electrochemical

reduction of [CuII(PMDETA)Br(OSMe2)]+ must be accompanied by the loss of one ligand. Here this

requirement is fulfilled by dissociation of the weakly-bound solvent ligand upon reduction to give

[CuI(PMDETA)Br] (Scheme 2-2). This proposal is consistent with EXAFS measurements on

solutions of CuI/(PMDETA) in the presence of bromide.98

Electrochemical reduction must also be coupled to solvation of the remaining bromide ligand in order

to generate the active catalyst [CuI(PMDETA)(OSMe2)]+.84, 99 These two steps, namely a reduction

in coordination number and halide count, must be fulfilled en route to generating the active catalyst

otherwise activation would result in the continual accumulation of halido ligands (which is not

observed experimentally) and the process would not be catalytic.84

On the reverse sweep, [CuII(PMDETA)Br(Solv)]+ is regenerated by oxidising [CuI(PMDETA)Br]

(Scheme 2-2 - right hand side). In the absence of an initiator, the [CuII/I(PMDETA)Br]+/0 redox couple

is reversible which indicates that both forms of CuI/(PMDETA) are stable on the timescale of the

experiment. The solvent complex [CuII(PMDETA)(OSMe2)2]2+ is also included in the mechanism of

Scheme 2-2 (in grey) but the two equivalents of bromide coming from the solid starting material

ensure that the major species at the start of the sweep is [CuII(PMDETA)Br(Solv)]+ (vide infra).

When EBriB is added to the cell the waveform becomes asymmetric as the cathodic current is

amplified and the anodic peak on the return sweep is diminished. Together these observations are

consistent with the catalytic mechanism proposed in Scheme 2-2 and the behaviour described earlier

for the Cu/Me6tren system with the same initiator. The waveform is clearly diffusion-limited at low

potentials meaning EBriB is consumed within the diffusion layer during the cathodic sweep. When

TEMPO is also added to the cell, the catalytic current is further amplified as the activation reaction

is made irreversible and the oxidative current on the reverse sweep is further diminished as the

equilibrium KI,Br is shifted to the left. Five equivalents of TEMPO versus the added alkyl halide is

sufficient to entirely quench the deactivation reaction.

Having established that TEMPO makes the activation reaction unidirectional, the voltammetry with

an excess of TEMPO was measured with increasing concentrations of EBriB at different sweep rates.

Similar experiments were performed with methyl 2-bromopropionate (MBrP) and benzyl bromide

25

(BnBr) initiators as these mimic the chain ends of propagating methyl acrylate and styrene polymers

respectively (Figure 2-3). The data were then simulated using the program Digisim according to the

mechanism of Scheme 2-2.82

Fitting kact

In order to minimise the number of variables which float during the iterative fitting of kact, the majority

of the thermodynamic, kinetic and physical constants from Scheme 2-2 were independently

determined or taken from the literature and fixed during the simulations.

The redox potential EBr as well as the diffusion coefficients (D) and heterogeneous electron transfer

rate constants (k0) were determined by simulating the sweep-rate dependent voltammetry of

[CuII(PMDETA)Br(OSMe2)]+ in the absence of initiator. To obtain the equivalent parameters for the

solvato-complexes, a 1.0 mM solution of [CuII(PMDETA)(OSMe2)]2+ was prepared by dissolution

of [CuII(PMDETA)(EtOH)(H2O)](ClO4) in DMSO (0.1 M (Et4N)(ClO4)) and assessed in an identical

manner. The relevant constants are collected in Table 2.2.

The binding constant KII,Br has been determined previously, from spectrophotometric titrations of

(Et4N)Br into [CuII(PMDETA)(OSMe2)2]2+, as 3.63 × 103 M-1.99 Therefore under the conditions

employed here (i.e. with two equivalents of bromide) ~81% of the total [CuII] is bound by a single

bromido ligand. The remainder exists as the solvato-complex [Cu(PMDETA)(OSMe2)2]2+ and is

reduced directly to the active catalyst at potential ESol. This reduction is not distinctly observed in the

voltammetry because ESol and EBr are separated by only 20 mV (Table 2.2).

With EBr, ESolv and KII,Br measured, KI-Br is fixed by the Nernst equation (Eqn. 2.1 – See Appendix

2.1 for a derivation).

ESolv – EBr = -59.2

nln

KII,Br

KI,Br

(2.1)

The rate constant kIIa,Br was set (103 M-1 s-1) in accordance with previous work and changing this value

had no effect on the fits.99 The deactivation rate constant was set close to zero (making the rate

constant equal to zero is not possible within the constraints and operation of the simulation software),

however, its value was irrelevant under these conditions (i.e. even values at the diffusion limit of 109-

1010 M-1 s-1 had no effect on the fit due to the quenching of R• by TEMPO). The radical termination

rate constants kt and kR-T were set as 1.0 × 109 M-1 s-1.70, 105

26

With the majority of the parameters from Scheme 2-2 determined, only kId,Br and kact were allowed to

float while fitting the concentration and sweep-rate dependent voltammetry. In each case, a single set

of rate constants were determined which accurately reproduced the experimental data. The

experimental and simulated CVs are shown in Figure 2-3, Figure 2-4 and Figure 2-5. The rate

constants which produced these profiles are collected in Table 2.1.

Figure 2-3 [RBr] dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess TEMPO. Solid

lines – experimental data, broken lines – simulated data. A) RBr – EBriB, sweep rate = 50 mV s-1. B) RBr – MBrP, sweep

rate = 20 mV s-1. C) RBr – BnBr, sweep rate = 50 mV s-1. I = 0.1 M (Et4N)(ClO4). Note that the nonspecific reduction at

low potentials in 2.4 B was not modelled.

Figure 2-4 Sweep rate dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess TEMPO and 5.0

mM EBriB. Solid lines – experimental data, broken lines – simulated data. I = 0.1 M (Et4N)(ClO4).

-1200 -1000 -800 -600 -400 -200

20 mVs-1

50 mVs-1

100 mVs-1

300 mVs-1

5 A

E / mV vs. Fc+/0

A B C

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5.0 mM

5 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

0.0 mM

3.0 mM

6.0 mM

9.0 mM

5 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

0.0 mM

3.0 mM

6.0 mM

9.0 mM

2 A

E / mV vs. Fc+/0

27

Figure 2-5 Sweep rate dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with excess TEMPO and 12.0

mM BnBr. Solid lines – experimental data, broken lines – simulated data. I = 0.1 M (Et4N)(ClO4).

Figure 2-6 The CV of a blank solution of DMSO + electrolyte is shown in red. To the solution is added 1.0 mM TEMPO

(green) and a further 4.0 mM TEMPO (blue). The orange curves are the experimental and simulated voltammetry of 1.0

mM [CuII(PMDETA)Br2]+ with 9.0 mM MBrP and 40 mM TEMPO. Sweep rates are all 20 mV s-1. I = 0.1 M

(Et4N)(ClO4).

It may be noted from Figure 2-3 that the simulated plots do not appear to fit the ‘downturn’ in the

voltammetry at very low potentials (For example plot B, ~ -1100 mV). However, a blank experiment

with increasing concentrations of TEMPO reveals that non-specific, heterogeneous reduction of

TEMPO at the electrode is responsible for this extra cathodic current (Figure 2-6). This additional

reaction was not modelled in Digisim resulting in a small discrepancy between the simulated and

experimental voltammetry. This small discrepancy does not affect the value of kact which comes from

the fit within the catalytic region near EBr.

-1200 -1000 -800 -600 -400 -200

20 mVs-1

50 mVs-1

100 mVs-1

300 mVs-1

5 A

E / mV vs. Fc+/0

-1400 -1200 -1000 -800 -600 -400

2 A

E / mV vs. Fc+/0

28

It is also important to highlight that this method is highly sensitive to the value of kact. The rate

constants reported in Table 2.1 are unique and other values of kact do not successfully reproduce the

experimental voltammetry (See Figure 2-7). This observation, coupled with the good agreement

between the experimental and digital traces across a range of concentrations of RBr and a range of

sweep rates, lends credence to the rate constants reported in Table 2.1.

The determined values of kact vary across two orders of magnitude with the expected order of EBriB

(tertiary bromide) > MBrP (secondary bromide) > BnBr (primary bromide). Radicals centred on a

tertiary carbon atom are stabilised compared to those centred on a primary carbon atom providing a

greater driving force for activation with EBriB.106 The same trend has been reported elsewhere with

[CuI(PMDETA)Br] in acetonitrile using a different methodology.47

Figure 2-7 Sensitivity of the digital trace to kact. Experimental (solid orange) and simulated (broken orange) voltammetry

of 1.0 mM [CuII(PMDETA)Br(OSMe2)]+ in DMSO with excess TEMPO and 5.0 mM EBriB. Sweep rate = 50 mV s-1. I

= 0.1 M (Et4N)(ClO4). The broken orange line is for kact as reported in Table 2.1 (i.e. 2.4 × 103 M-1 s-1). The broken red,

green and blue lines illustrate how the voltammetry changes when kact is multiplied by 0.5, 2, or 10 respectively.

-1200 -1000 -800 -600 -400 -200

5 A

E / mV vs. Fc+/0

29

Table 2-1 Summary of the activation and deactivation rate constants from Scheme 2-2 with EBriB, MBrP and BnBr

initiators (RBr). T = 298 K.

EBriB MBrP BnBr

KATRP 1.3 × 10-3 3.3 × 10-5 5.2 × 10-5

kact (M-1 s-1) 2.4 × 103 2.5 × 102 4.5 × 101

kdeact (M-1 s-1) 1.8 × 106 7.6 × 106 8.6 × 105

Table 2-2 Summary of the key chemical and electrochemical constants from Scheme 2-2 with [CuII(PMDETA)Br2] in

DMSO. T = 298 K.

Chemical constants Electrochemical constants

KII,Br (M-1) 3.6 × 103 ESolv (mV vs Fc+/0) -611

kIIa,Br (M-1 s-1) 1.0 × 103 k0 (cm s-1) 5.0 × 10-3

KI,Br (M-1) 1.7 × 103 α 0.5

kIa,Br (M-1 s-1) 2.0 × 104 EBr (mV vs Fc+/0) -630

k0 (cm s-1) 5.0 × 10-3

kRT (M-1 s-1) 1 × 108 α 0.5

k-RT (s-1) 1 × 10-4

kt (M-1 s-1) 1 × 109

k-t (s-1) 1 × 10-2

30

Fitting kdeact

Having determined kId,Br and kact, the voltammetry was remeasured and simulated in the absence of

TEMPO. Identical sweep rates and concentrations of initiator were employed. The catalytic current

at EBr is attenuated by the only unknown reaction rate constant from Scheme 2-2, kdeact.

Thus, in simulating the experimental data the deactivation rate was allowed to float (all other

parameters having been determined and fixed) in order to reproduce the diminished catalytic currents.

Each time this parameter was typically set to an unreasonably small number – to give an over

amplified catalytic current – and then allowed to increase until a fit was obtained. Once kdeact was

determined, the simulations were perturbed by readjusting this value to be unreasonably large and the

iterative fitting process reengaged (i.e. the fit was approached from the opposite direction). The

software minimised kdeact until the same value was achieved proving the robustness of these numbers.

The concentration and sweep rate dependent fits are illustrated in Figure 2-8, Figure 2-9, and Figure

2-10. The optimised values of kdeact which generated these fits are included in Table 2.1. As with kact,

the simulations are sensitive to the magnitude of kdeact (Figure 2-11).

Figure 2-8 [RBr] dependent voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with no TEMPO. Solid lines –

experimental data, broken lines – simulated data. A) RBr – EBriB, sweep rate = 50 mV s-1. B) RBr – MBrP, sweep rate

= 20 mV s-1. C) RBr – BnBr, sweep rate = 50 mV s-1. I = 0.1 M (Et4N)(ClO4).

A B C

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5.0 mM

5 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5.0 mM

2 A

E / mV vs. Fc+/0

31

Figure 2-9 Sweep rate dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with 5.0 mM EBriB

and no TEMPO. Solid lines – experimental data, broken lines – simulated data. I = 0.1 M (Et4N)(ClO4).

Figure 2-10 Sweep rate-dependent catalytic voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with 12.0 mM BnBr

and no TEMPO. Solid lines – experimental data, broken lines – simulated data. I = 0.1 M (Et4N)(ClO4).

-1200 -1000 -800 -600 -400 -200

20 mVs-1

50 mVs-1

100 mVs-1

300 mVs-1

5 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

20 mVs-1

50 mVs-1

100 mVs-1

300 mVs-1

5 A

E / mV vs. Fc+/0

32

Figure 2-11 Sensitivity of the digital trace to kdeact. Experimental (solid orange) and simulated (broken orange)

voltammetry of 1.0 mM [CuII(PMDETA)Br2] in DMSO with no TEMPO and 5.0 mM EBriB. Sweep rate = 50 mV s-1. I

= 0.1 M (Et4N)(ClO4). The broken blue line yields kdeact as reported in Table 2.1 (i.e. 7.6 × 106 M-1 s-1). Broken red, green

and blue lines illustrate how the voltammetry changes when kdaact is multiplied by 0.5, 2, or 10 respectively.

The deactivation rate constants reported in Table 2.1 have not been measured previously in DMSO.

In acetonitrile kdeact has been determined for MBrP and EBriB as 4.3 × 107 and 1.9 × 107 M-1 s-1

respectively using the empirical method (from the measured value of kact and the relationship KATRP

= kact/kdeact).20 These values are roughly one order of magnitude larger than the values measured here

in DMSO. Whilst the relationship between kdeact and the solvent is not well understood, two papers

suggest that deactivation is slower in more polar solvents.40, 48

Various empirical parameters have been developed in order to rank solvents according to their

polarity.107-111 One of the most successful of these is the normalised Dimroth-Reichardt parameter

‘ENT ’ which varies between 1.000 (water) and 0.000 (tetramethylsilane) for the extremes of polar and

nonpolar solvents respectively.112-113 Comparing the rate constants measured here in DMSO

(EN T = 0.444) with those measured in acetonitrile (EN

T = 0.460) would suggest, in contrast to the

previous work, that deactivation is faster in more polar solvents. It is important not to draw too many

conclusions from these results as one of the papers cited in the paragraph above only compares two

solvents (acetonitrile and ethyl acetate), and the other determines kdeact empirically. Clearly, further

study is needed with an expanded range of solvents in order to clarify its effect on kdeact.

Unlike activation, there is no immediately obvious correlation between the magnitude of kdeact and

the properties of the EBriB, MBrP and BnBr radicals; the observed order is MBrP > EBriB > BnBr.

Styryl radicals are more nucleophilic than their acrylate counterparts and so, on the basis of electron-

donating effects only, would be expected to abstract the halogen atom more rapidly.114 Earlier studies

examining the kinetics of atom transfer reactions involving metal halide salts and vinyl radicals also

-1200 -1000 -800 -600 -400 -200

5 A

E / mV vs. Fc+/0

33

indicate that the order of reactivity should be styryl radicals > methyl acrylate/methyl methacrylate

radicals.115-116 It must be noted however that the mechanism of halogen atom transfer was not

conserved across the series.114 Once again, a meaningful discussion of the structure-reactivity

relationships should be suspended until a broader array of initiators has been surveyed. These should

include the chlorido-analogues of the alkyl halides already investigated here; i.e. ethyl-2-

chloroisobutyrate, methyl-2-chloropropionate and benzyl chloride. Another particularly interesting

target is ethyl-2phenylbromopropionate which is the most reactive alkyl halide initiator reported to

date.20

34

Conclusion

Electrochemistry has been used to determine the kinetics of both atom transfer reactions relevant to

the central ATRP equilibrium. Cyclic voltammetry of a resting solution of

[CuII(PMDETA)Br(OSMe2)]+ in the presence of an appropriate alkyl halide initiator and TEMPO

leads to a catalytic current from which the activation rate constant kact can be determined without any

assumptions regarding the reverse, deactivation reaction (kdeact). With kact isolated, the same

voltammetry in the absence of TEMPO reveals an attenuated catalytic current due to the consumption

of [CuII(PMDETA)Br(OSMe2)]+ by the deactivation reaction. The rate constant kdeact is determined

by simulating the difference between the currents in the presence and absence of TEMPO. In a single

experiment, both the forward and reverse atom transfer rate constants are isolated experimentally

representing a significant advancement in the kinetic analysis of the central ATRP equilibrium.

Here it is shown that the activation rate is sensitive to the identity of the initiator and is well correlated

with the stability of alkyl radical product. The kinetics of deactivation are less sensitive to the identity

of the initiator than activation and no meaningful structure-reactivity correlations can be drawn for

kdeact with the limited number of initiators which are considered here.

35

Experimental

Safety note: perchlorate salts are potentially explosive. Although no problems were experienced here

these should never be heated in the solid state or scraped from sintered glass frits. The electrolyte

(Et4N)(ClO4) was selected as the reagent of choice because it is easily prepared from inexpensive

starting materials and is not hygroscopic as prepared.

All reagents were obtained commercially (including PMDETA ligand, Aldrich 99%) and used

without further purification.

Synthesis

Salts

(Et4N)(ClO4) was prepared by adding dropwise 68.31 g of 70% w/w HClO4 solution (0.476 mol) to

100 g of (Et4N)Br in 150 mL of H2O. The addition of HClO4 was stopped at several points and the

precipitate that formed was collected by filtration and washed with cold water. After the filtrate was

collected the titration with HClO4 resumed. After the addition was complete, the collected solid

product was recrystallised three times from hot H2O and finally from 70/30 isopropanol/acetonitrile

before being dried under high-vacuum. The salts (Et4N)Br and (Et4N)Cl were recrystallised from

70/30 isopropanol/acetonitrile and dried overnight under vacuum before use.

CuII PMDETA complexes

[CuII(PMDETA)Br2] was synthesised by the addition of 0.1774 g (1.0 mmol) of PMDETA to 0.223

g (1.0 mmol) of CuBr2 suspended in CH2Cl2 (20 mL). The solid was precipitated by slow addition of

diethyl ether. The product was collected by filtration and washed with cold diethyl ether to remove

any residual ligand. Anal. Calcd for C9H23CuN3Br2: C, 27.25; H, 5.80; N, 10.60. Found: C, 26.95; H,

5.89; N, 10.51.

[CuII(PMDETA)(EtOH)(OH2)](ClO4)2 was synthesised by the drop-wise addition of 10 mL of

ethanol containing 0.0887 g (0.49mmol) of PMDETA to a solution of CuII(ClO4)2.6H2O (0.182 g,

0.49 mmol) in hot ethanol (10 mL). A deep blue colour ensued. The solution was stirred at ~60°C for

10 minutes before being allowed to cool to room temperature. Diethyl ether (20 mL) was added

slowly and the suspension placed in the refrigerator overnight. The resulting blue solid was collected

by filtration and washed with cold ethanol before being dried under vacuum. Anal. Calcd. for C-

11H30O10N3CuCl2: C, 26.5; H, 6.06; O, 32.1; N, 8.42. Found: C, 26.6; H, 5.93; O, 32.0; N, 8.19. X-

ray quality crystals were obtained by slow diffusion of diethyl ether into an ethanolic solution of the

solid and the crystal structure was reported previously.99

36

The solvent complex [CuII(PMDETA)(OSMe2)2]2+ was prepared in situ by dissolving

[CuII(PMDETA)(EtOH)(OH2)](ClO4)2 in DMSO.

Physical Methods

Cyclic Voltammetry

Cyclic voltammetry was performed on a BAS100B/W potentiostat employing a glassy carbon

working electrode, platinum auxiliary electrode and a non-aqueous Ag/Ag+ (c.a. 0.01 M AgNO3)

reference electrode in dimethylsulfoxide (DMSO). Measurements were made at 298 K. DMSO was

dried by the recommended method which involved vacuum distillation over calcium hydride onto 3

Å molecular sieves – discarding the first 20% of collected liquid.117 Ferrocene was used as an external

standard and all potentials are cited versus Fc+/0. The supporting electrolyte was 0.1 M (Et4N)(ClO4)

and all solutions were purged with argon before measurement. All electrochemical solutions were 5.0

mL total volume.

Electrochemical Simulation

All simulations were carried out with DigiSim version 3.0.82 The specific kinetic and thermodynamic

parameters are summarised in Table 2.1 and Table 2.2 while Scheme 2-2 defines each of these values

in terms of the accepted mechanism for Cu-catalysed atom transfer reactions. Other generic

parameters are: diffusion coefficients DRBr (EBriB, MBrP and BnBr) 1 × 10-5 cm2 s-1, DCuL (4 × 10-6

cm2 cm-1, all forms of Cu), DBr 1 × 10-5 cm2 s-1; electrode surface area 0.053 cm2; heterogeneous

electron transfer rate constant k0 5 × 10-3 cm s-1; double layer capacitance 3 × 10-6 F; temperature 298

K; transfer coefficient α 0.5.

37

A kinetico-mechanistic study on CuII deactivators

Introduction

From the existing literature it is clear that the ATRP deactivation reaction is influenced by several

factors. For instance it has been noted that the solvent partially (or completely) solvates the auxiliary

halide ligand on CuIILX leading to differences in the rate and efficiency of deactivation (Scheme

3-1).40, 118-121 The identity of the auxiliary halide ligand itself is important so that deactivation is

slower with chloride than with bromide.40, 122-123

Scheme 3-1 General representation of organic radical (R•) deactivation by [CuIILX]+. Charges omitted for clarity.

The steric bulk of the chelating ligand also appears to have a significant impact on the deactivating

properties of [CuIILX]+. This was demonstrated by experiments which compared polymerisations

using Cu/Me6tren and Cu/Et6tren catalysts (Figure 3-1 – Et6tren is tris[2-diethylamino(ethyl)]amine).

Where [CuII(Et6tren)Br]+ was employed as the deactivator the polymerisations were slow and

exhibited poor control over the molecular weight distributions compared with [CuII(Me6tren)Br]+

under the same conditions;124 inefficient deactivation was suggested as the cause of this problem.

With tris[2-pyridyl(methyl)]amine (tpa – Figure 3-1) the deactivating properties were found to be

different again20, 36, 124 and could be further tuned by incorporating electron donating groups on the

pyridine rings.61, 125

38

Figure 3-1 Relevant structures determined for [Cu(L)X]+ deactivating complexes with the chelating ligands Me6tren,

Et6tren or tpa.

Crystal structures of CuII complexes of tpa and Me6tren systematically reveal five-coordinate, trigonal

bipyramidal geometry with the three terminal amine donors of the ligand occupying the equatorial

coordination sites and the tertiary amine coordinated via a slightly shorter bond at one of the axial

positions (Figure 3-1). The remaining axial site is occupied by a monodentate ligand (‘Y’ in Figure

3-1)126 such as MeCN127, H2O,127-129 HCO2-,130 Br-,131 Cl- 132 or CN-.133 The bond length at this

position is shorter than the trans Cu-Naxial bond for C-, N- or O-donors such as MeCN, CN-, CF3SO3-

or H2O but lengthens for the Cl- and Br- ligands due to their increased covalent radii (See Figure 3-11

and Figure 3-12).132 The only time this binding pattern is violated is when a sterically bulky co-ligand

such as PPh3- is introduced which cases one of the arms of the chelate to dissociate.132, 134 While there

are no crystal structures for CuII/Et6tren, EPR measurements reveal that these complexes have trigonal

bipyramidal geometry in solution.135-136 This geometry has also been confirmed for CuII/Me6tren and

CuII/tpa complexes in solution.135, 137

It is clear that the different deactivating properties of [CuII(tpa)X]+, [CuII(Me6tren)X]+ and

[CuII(Et6tren)X]+ are not due to a difference in geometry. Instead, the unique steric and electronic

influences of each of these ligands must affect the strength of the CuII‒X bond. In fact, each of the

aforementioned variables (halide, chelate and solvent) is likely to affect this bond. This is an

important but seemingly neglected observation. Deactivation is an atom transfer reaction comprising

concomitant bond breaking and electron transfer (Scheme 3-1). Any change in the reaction conditions

which alters the stability of the CuII-X bond is almost certainly going to be correlated with differences

in the atom transfer kinetics (and propagate through to differences in the polymer products).

To better understand the qualitative and (albeit limited) quantitative experimental observations

regarding deactivation, a series of kinetic experiments were undertaken following halide substitution

reactions on copper(II) complexes of Me6tren, Et6tren and tpa. Three types of reactions were

examined including: 1) anation, 2) solvent exchange and 3) halide exchange (Scheme 3-2). These

39

were performed in a variety of ATRP-relevant, organic solvents, with or without added monomer,

using both bromide and chloride. The first two reactions (Scheme 3-2 A and B) are useful in

establishing the mechanism for substitution processes occurring at the axial site of such trigonal bi-

pyramidal copper(II) complexes and the halide exchange process (reaction C) is particularly relevant

to deactivation as a halide is initially coordinated.

Scheme 3-2 A) Anation, B) solvent exchange and C) halide exchange reactions studied.

40


A model system

Cu/Me6tren is one of the most active and widely utilised catalysts for ATRP.19 Fortuitously, the

chelating behaviour and steric bulk of Me6tren also slow down exchange reactions at the remaining

coordination site on [CuII(Me6tren)Y]n to the point where they are observable by stopped-flow

spectroscopic techniques (vide infra).138 Another benefit of the Me6tren chelate is that it restricts the

number of co-ligands to one (i.e. ‘Y’) which simplifies the kinetic analysis. This makes copper(II)

complexes of Me6tren ideal candidates for the intended kinetics experiments. These prerequisite

conditions are also met by copper(II) complexes of Et6tren and tpa as will be discussed.

Anation and solvent exchange reactions on CuII/Me6tren were carried out in dimethylsulfoxide

(DMSO), dimethylformamide (DMF), acetonitrile (MeCN), ethanol (EtOH) and methanol (MeOH)

(common ATRP solvents) and followed using stopped-flow spectrophotometry to establish the

mechanism of ligand exchange. Figure 3.3 shows the typical spectral changes obtained for the solvato

complex anation (A) and the solvent interchange reactions (B) which are complete in less than 0.5 s.

Figure 3-2 A) Time-resolved spectral changes for the reaction of 2.0 × 10-4 M [CuII(Me6tren)(NCMe)]2+ with Br– (0.0025

M) at 298 K and I = 0.1 M LiClO4; B) Time resolved spectral changes for the reaction of 3.8 × 103 M

[CuII(Me6tren)(NCMe)]2+ with DMF (1.29 M) at 298 K and I = 0.1 M LiClO4. Insets show the absorbance changes at the

wavelengths indicated by the arrows.

400 450 500 550 600 650 7000.0

0.2

0.4

0.6

0.0 0.1 0.2

0.01

0.02

0.03

0.04

Ab

so

rba

nce

@4

71

nm

t /s

Ab

so

rba

nce /

a.u

.

/ nm

250 300 350 400 450 5000.0

0.2

0.4

0.6

0.8

1.0

0.0 0.1 0.2 0.3 0.4

0.3

0.4

0.5

Ab

so

rba

nc

e@

35

0 n

m

t /s

t /s

Ab

so

rba

nce /

a.u

.

/ nmλ / nm λ / nm

B A

41

Anation kinetics

For the anation reactions, the ligand-to-metal charge transfer maximum in the range 300-325 nm was

followed as it underwent a significant bathochromic shift upon complexation by halide. All anation

reactions were measured under pseudo first order conditions with the concentration of the halide

being at least 10 times that of the copper complex. The pseudo first order rate constants kobs were

determined by global analysis of the time-resolved changes in the 300-325 nm region (inset of Figure

3-2 A) using the program SPECFIT.139 The values of kobs were plotted as a function of the

concentration of [X-] as well as temperature.

Reactions of [CuII(Me6tren)(Solv)]2+ with chloride or bromide in DMSO, MeCN, DMF, EtOH and

MeOH displayed typical [X-]-limiting substitution dependence of kobs on [X-] (Figure 3-3 and Figure

3-4). This behaviour agrees with the rate law and mechanism indicated in Scheme 3-3 and Equation

3.1, where an outer-sphere, precursor complex accumulates prior to the rate limiting halide

coordination (kon) and the reverse solvolysis reaction (koff) is negligibled.140-144 The rate law obtained

in Equation 3.1 is derived in Appendix 3.1. At high concentrations of [X-] the second term in the

denominator dominates (KOS[X-] > > 1) so kobs tends towards kon in the manner typified by Figure 3-3.

d The assumption that koff is negligible can be confirmed as follows: For the reaction shown below, the halide binding

constant (KII,Br) has been measured in acetonitrile as 2.77 × 106 M-1.51

Comparing this reaction with Scheme 3-3, k1 = KOS.kon and k-1 = 1/KOS.koff. Therefore,

2.77 × 106 M−1 = k1

k-1

= KOSkon

1/(KOSkoff)

koff = KOS

2 kon

2.77 × 106 M-1

Substituting the relevant values of KOS and kon Table 3.1 gives

koff = 0.32 s-1

This confirms that bromide dissociates from [CuII(Me6tren)Br]+ at a negligible rate and validates the approximation of

the rate law used to fit the kinetic data.

42

kobs = KOSkon[X-]

1+KOS[X-] (3.1)

Scheme 3-3 Mechanism and observed rate law for anation of [CuII(Me6tren)(Solv)]2+ by Br- or Cl-.

Figure 3-3 Left) Plot of kobs versus [Br-] for the [Cu(Me6tren)(NCMe)]2+ + Br- reaction at different temperatures and at I

= 0.1 M LiClO4; Right) kon as a function of temperature and pressure.

Figure 3-4 Left) Plot of kobs versus [Br-] for the [Cu(Me6tren)(OSMe2)]2+ + Br- reaction at 298 K; Right) Plot of kobs

versus [Br-] for the [Cu(Me6tren)(DMF)]2+ + Br- reaction at 288 K. I = 0.1 M LiClO4 for both.

0.000 0.005 0.010 0.0150

50

100

150

ko

bs / s

-1

[Br-] / M

0.000 0.005 0.010 0.015 0.0200

10

20

30

40

ko

bs / s

-1

[Br-] / M

0.000 0.006 0.012 0.0180

20

40

60

80

100 288 K

293 K

298 K

303 K

308 K

ko

bs /s

-1

[Br-] / M

0.0033 0.0034 0.0035

-2.5

-2.0

-1.5

-1.0

0 300 600 900 1200 15002.4

2.6

2.8

T-1/K

-1

lnk

on

ln(k

on/T

)

P /atm

298 K

43

In Scheme 3-3 the rate limiting step is kon. In order to ascertain the mechanism by which this step

proceeds (Associative (A), Dissociative (D) or Interchange (I) – Figure 3-5) the thermal activation

parameters ∆S‡ and ∆H‡ were determined where possible. This was done by fitting the temperature-

dependent values of kon to the linear form of the Eyring equation (Eqn 3.2, Figure 3-3). The

dependence of lnkon on pressure was measured to give the activation volumes (∆V‡) where possible

from Equation 3.3145 and the relevant kinetic and thermodynamic parameters are collected in Table

3.1. The full range of measured rate constants for each of the reactions discussed in this Chapter are

provided in Appendix 3.5.

lnkon

T =

∆H‡

R.1

T + (

∆S‡

R + ln

kb

h) (3.2)

lnkon = -∆V‡

RT.P + lnk

o (3.3)

Table 3-1 Kinetic and activation parameters for the anation reaction [CuII(Me6tren)(Solv)]2+ + X- (X = Br or Cl) in

different solvents. I = 0.1 M LiClO4.

Solv X– 298KOS

(M-1)

298kon

(s-1)

ΔH‡

(kJ mol-1)

ΔS‡

(J K-1mol-1)

ΔV‡

(cm3 mol-1)

MeCN Br- 1.3 × 102 53 66 ± 1 7.0 ± 3 6 ± 1(289 K) a

Cl- 1.4 × 102 54 67 ±5 11 ± 16 -

DMF Br- 2.7 × 102 b 1.7 × 102 b - - -

MeOH Br- 75 36 69 ± 1 16 ± 4 -

Cl- 61 48 72 ± 2 28 ± 7 -

EtOH Br- 64 47 72 ± 3 25 ± 9 -

Cl- 69 66 72 ± 4 31 ± 13 -

100% DMSO Br- 1.1 × 102 1.1 × 102 -

25% MMA c Br- 3.3 × 102 1.1 × 102 -

50% MMA c Br- 5.3 × 102 1.1 × 102 -

25% styrene c Br- 2.0 × 102 1.0 × 102 -

50% styrene c Br- 3.5 × 102 1.1 × 102 - a determined at [Br–] = 0.045 M where kobs ≈ kon, see Figure 3.4; b measured at 288 K; c solutions containing an olefinic

monomer (styrene or methyl methacrylate) are quoted as a percentage volume in DMSO.

44

Before progressing with an analysis of the data in Table 3.1 it is helpful to consider what is already

known about exchange reactions on copper(II) complexes. The kinetic and mechanistic details of

ligand exchange reactions on copper(II) compounds has a rich experimental history as these have

been extensively employed as model compounds for a broad array of biological metalloenzymes.146

Discarding the titanic amount of work focused on oxygen uptake and transformation,147-150 a much

smaller subset of the literature focuses on monodentate ligand exchange reactions which are relevant

to this work. From the existing literature, the following observations emerge and serve as a helpful

framework for interpreting the results in Table 3.1.

Solvent exchange reactions on CuII complexes lacking any co-ligands, such as [Cu(Solv)6]2+, are

extremely fast (e.g. kH2O ~ 109 s-1, kMeCN ~ 107 s-1)151-152 and in DMF153, H2O153 and MeOH154 the

reaction proceeds by a dissociative interchange (Id) mechanism (presumably) at the axial coordination

site due to the inherent elongation of this bond as a consequence of the Jahn-Teller effect.155-156 Id

behaviour indicates that partial dissociation of the initially-coordinated ligand accompanies the

formation of the transition state during the rate determining step. Support for this mechanism is

provided by a combination of definitively positive activation entropies (∆S‡), small, positive

activation volumes (∆V‡) and a dependence of the rate determining step on the identity of the outgoing

ligand.

Figure 3-5 Representation of different ligand substitution mechanisms.

45

Whilst some ambiguity surrounds the exact coordination number and geometry of solvated CuII

complexes,153, 157-158 the addition of a tripodal, tetradentate ligand such as tris(2-aminoethyl)amine

(tren) or its methylated analogues (Me3tren and Me6tren) simplifies the situation. These ligands

favour trigonal bipyramidal geometry for CuII (vide supra) which removes the orbital degeneracy of

the d9 ground state present in an octahedral ligand field. The trigonal bipyramidal d9 ground state has

a nondegenerate (dz2) ground state so no Jahn-Teller distortion is operative and no coordinate bonds

are weakened;159-160 accordingly, the substitution kinetics of the remaining co-ligand are slower. For

example the rate constant for acetonitrile exchange (kMeCN) decreases from ~ 107 s-1 to 5.1 × 103 s-1

on going from [CuII(MeCN)6]2+ to [CuII(tren)(MeCN)]2+.152 The exchange mechanism is also

different, proceeding via an Ia pathway.161-162 Ia behaviour is evident from kinetics which depend on

the identity of the incoming ligand and have negative activation entropies and volumes of activation.

Increasing the steric bulk of tren by mono- or di-methylation of each of the three terminal amines

(Me3tren and Me6tren respectively) further alters the exchange behaviour. For Me3tren the mechanism

remains the same (Ia), however the rate constant is much decreased; for H2O → pyridine exchange,

the rate constant is three orders of magnitude less with CuII/Me3tren than CuII/tren.163 Further

methylation, producing Me6tren, causes a definite shift towards an Id mechanism as shown for DMF

and diethylformamide interchange164-165 or azide- and thiocyanate-water substitution.138

With this background in mind, along with the data collected in Table 3.1 it is evident that the anation

reactions involving [CuII(Me6tren)(Solv)]2+ proceed via the expected Id mechanism. To begin with,

kon is insensitive to the identity of the incoming halide. However, kon is sensitive to the identity of the

outgoing solvent, being fastest in DMF which is the most sterically bulky ligand of this series.

The entropies of activation in Table 3.1 are small but definitively positive which is expected for partial

dissociation of the initially bound solvent en route to the transition state (Figure 3.8). The enthalpies

of activation are large and a small, positive volume of activation is observed. Again these observations

are consistent with those reported previously for solvent exchange on [Cu(Me6tren)(DMF)]2+ which

proceeds through an Id mechanism.165 The two latter parameters have to be considered along with the

fact that the entering halide anions are already poorly solvated in the outer-sphere encounter complex

due to charge compensation.

46

Figure 3-6 Representation of the reaction coordinate during an anation reaction in which the rate determining step

proceeds by partial dissociation of the coordinated solvent (Id).

Elsewhere an alternative exchange mechanism has been reported in which partial dissociation of one

of the arms of Me6tren precedes rapid coordination of the incoming ligand. This behaviour has been

proposed for [CuI(Me6tren)Br] and [CuII(Me6tren)(OH2)]2+ but only occurs where the concentration

of the incoming ligand is very high (~200× [CuII]).132, 134, 163, 166 This mechanism leads to a rather

complicated profile for the dependence of kon on the incoming ligand and produces a final UV-Vis

spectrum typical of a tetragonally-elongated octahedral CuII complex. For the reactions measured

here, the characteristic visible-NIR spectral feature of a trigonal bipyramidal

[CuII(Me6tren)(ligand)]n+ complex is always preserved and the kinetics follow the straightforward

rate law of Eqn. 3.1; thus there is no evidence here for partial dissociation of the Me6tren ligand at

any stage. Each of the explored anation reactions in DMSO, DMF, MeCN, EtOH or MeOH proceeds

via an Id interchange mechanism with partial dissociation of the solvent occurring during the rate

determining step.

The effect of added monomer on the anation kinetics in DMSO was also examined by the addition of

styrene or methyl methacrylate (Table 3.1). Increasing the percentage volume of monomer had no

effect on the rate determining step (kon) but was correlated with an increase in the outer-sphere

association constant KOS. Figure 3-7 plots the kinetic profiles with different concentrations of methyl

methacrylate and illustrates that while kon is conserved, the curvature from KOS is markedly steeper

with a higher concentration of monomer. Such behaviour is typical for kinetics with significant outer-

sphere complex formation in which the polarity of the medium is tuned. The monomer has no direct

effect on the copper(II) complex but does lower the dielectric constant of the medium and facilitates

greater association between the 2+ cationic complex and the incoming anion. This behaviour was

observed for both methyl methacrylate and styrene. As a point of clarification, the ionic strength was

47

always maintained at a constant 0.1 M by LiClO4 despite the increasing percentage of monomer. This

is imperative to avoid complicating ‘salt effects’ on the rate constant kon.

0.00 0.05 0.100

10

20

30

40

50

60

70

80

90

100

110

kobs /

s-1

[Br-] / M

Figure 3-7 Plot of kobs versus [Br-] for the [Cu(Me6tren)(DMSO)]2+ + Br- reaction in DMSO/MMA. Black squares – 50

% (vol.) MMA, Red circles – 25 % (vol.) MMA, Blue triangles - 0 % MMA. Fitted functions (using Eqn. 3.1) are

extrapolated to high [Br-] where kobs kon. I = 0.1 M LiClO4.

A recent study examined the influence of monomers on the ATRP equilibrium constant (KATRP) for a

series of polymerisations catalysed by Cu/Me6tren and Cu/tpa.36 As the reaction mixture was adjusted

from pure DMSO to bulk methyl acrylate, KATRP decreased by four orders of magnitude. The data

from Table 3.1 indicate that increasing the monomer concentration is more likely to affect the

activation reaction (kact) involving copper(I) rather than the deactivation reaction (kdeact) involving

copper(II). This proposal is corroborated by crystallography which demonstrates that monomers have

a reasonable affinity for copper(I) compounds, binding side-on as η2-ligands through the π-orbitals

of the olefin moiety.167-171 The proposal that monomer effects copper(I) as opposed to copper(II) is

also consistent with an earlier study which found that the deactivation rate was significantly less

sensitive than the activation rate to the identity of the monomer.172

Having assayed the solvent- and halide-dependent kinetics of anation, the effect of the chelate was

also considered by attempting the same reactions with complexes bearing the Et6tren and tpa ligands.

However, these reactions were completed on the mixing timescale so only the final complex

[Cu(L)X]+ was observed. Reducing i) the temperature (down to 253 K) or ii) the concentration of [X-

] to stoichiometric amounts did not resolve this problem.

48

Solvent exchange kinetics

The solvent exchange kinetics (Scheme 3-2 B) were also screened however these reactions were much

more difficult to follow experimentally. Substitution of MeCN for DMF or DMSO could not be

studied directly by co-injection of an MeCN solution of [Cu(Me6tren)(NCMe)]2+ with a solution of

neat DMF or DMSO in the second syringe as the reaction was too fast. Careful concentration

screening of DMF/MeCN or DMSO/MeCN solvent mixtures in the second syringe, led to measurable,

time-resolved spectral changes (see Figure 3-2 B).

The spectral changes for the solvent interchange reaction are much smaller because the reaction does

not go to completion; i.e. the final spectrum in Figure 3-2 B is a mixture of [Cu(Me6tren)(NCMe)]2+

and [Cu(Me6tren)(DMF)]2+. The reaction was thus performed with higher concentrations of copper

complex (3.75 × 10–3 M) and followed by monitoring the d-d absorption bands in the visible region.

For the MeCN DMF interchange, a final DMF concentration of 1.29 M in MeCN was necessary

for the reaction to be observable. Higher DMF concentrations led to reactions that were too fast, while

lower concentrations resulted in spectral changes that were too small and unreliable due to the

equilibrium favouring the reactant [CuII(Me6tren)(NCMe)]2+. Similarly, for the MeCN DMSO

interchange only a DMSO concentration of 7.05 M produced reliable results for the process. In all

cases the reverse solvent exchange reactions were too fast to measure under the conditions of the

study. The kinetic data collected in Table 3.2 correspond to values acquired under these specific

conditions and indicate a relatively slow pseudo first order reaction. The sensitivity of these reactions

precluded measurement of the thermal activation parameters.

49

Halide exchange kinetics

Especially relevant to deactivation are the halide exchange reactions denoted in Scheme 3-2 C. These

were conducted by co-injecting a solution of [CuII(Me6tren)X]+ (where X = Br- or Cl-) with a solution

containing the alternate halide Y-. Regardless of the identity of X, the pseudo first order rate constants

kobs for these exchange reactions showed a linear dependence on the [Y-] and a non-zero intercept

was often observed (Figure 3-8 and Figure 3-9). This behaviour was conserved in both MeCN and

DMF.

Because these reactions were significantly slower than anation it was also possible to measure the

kinetics with Et6tren and tpa under a limited number of circumstances. The same linear correlations

were again observed (Figure 3-10). The lack of significant curvature in all of the pseudo first order

plots (in contrast to Figure 3-3 and Figure 3-4) indicates that there is no observable build-up of the

outer-sphere complex (KOS is small) in the general reaction sequence of Scheme 3-2 C. This decrease

in KOS is consistent with the lower reactant charges involved i.e. +/- for halide exchange versus 2+/-

for halide anation.

Figure 3-8 Plots of kobs vs. [Y-] for the halido ligand exchange reaction [Cu(Me6tren)X]+ + Y- in MeCN, I = 0.1 M LiClO4.

Left) X = Br, Y = Cl and Right) X = Cl, Y = Br.

0.000 0.005 0.010 0.0150

1

2

3

4

5

6

7

8 288 K

293 K

298 K

303 K

308 K

ko

bs / s

-1

[Cl-] / M

0.000 0.005 0.010 0.015 0.0200.0

0.5

1.0

1.5

2.0 288 K

293 K

298 K

303 K

308 K

ko

bs / s

-1

[Br-] / M

50

Figure 3-9 Plots of kobs vs. [Y-] for the halido ligand exchange reaction [Cu(Me6tren)X]+ + Y- in DMF, I = 0.1 M LiClO4.

Left) X = Br, Y = Cl and Right) X = Cl, Y = Br.

Figure 3-10 Plots of kobs vs. [Y-] for the reaction of Left) [CuII(Et6tren)Cl]+ + Br- at different temperatures or Right)

[CuII(tpa)Cl]+ + Br- at 288 K. Solvent – MeCN, I = 0.1 M LiClO4. The reaction of [Cu(tpa)Cl]+ was too fast to measure

above 288 K.

0.000 0.005 0.010 0.015 0.0200.00

0.75

1.50

2.25

3.00 288 K

293 K

298 K

303 K

308 K

ko

bs / s

-1

[Br-] / M

0.000 0.005 0.010 0.015 0.0200

50

100

150

288K

293K

298K

303K

308K

ko

bs / s

-1

[Cl-] / M

0.00 0.01 0.02 0.030

50

100

150

200

250

300 288K

293K

298K

303K

308K

ko

bs / s

-1

[Br-] / M

0.000 0.003 0.006 0.0090

100

200

300

400

500

ko

bs / s

-1

[Br-] / M

51

Scheme 3-4 represents the expected mechanism for halide substitution given the consistent

observation that substitution reactions on CuII/Me6tren proceed via a dissociatively-activated

pathway. Here the concentration of [Cu(Me6tren)(Solv)]2+ is expected to be small; it is not observed

experimentally and halido complexes are formed with a stoichiometric equivalent of Br- or Cl-, see

Experimental/Materials section and other references.51 Applying the steady-state approximation

produces the observed rate law in Equation 3.4 A (See Appendix 3.2 for a derivation). According to

this rate law kobs should tend towards k1 at high concentrations of Y-. This is not consistent with the

experimental observation that kobs has a linear dependence on [Y-] even at high concentrations.

kobs(s.s.) = k1k2[Y-] + k-1k-2[X-]

k-1[X-] + k2[Y-] (3.4 A)

kobs(p.eq.)= (k1

k-1

) k2[Y-] + (k-2

k2

) k-1[X-] (3.4 B)

Scheme 3-4 Plausible mechanism and rate laws using the steady state (s.s) or pre-equilibrium (p.eq) approximations for

halide exchange reactions on [Cu(Me6tren)X]+.

The alternate rate law in Equation 3.4 B is determined using the preequilibrium approximation (again,

valid given that k1 and k-2 are very low, - Appendix 3.3). At low concentrations of X-, such as those

present in the experiments here ([X-] = 2 × [CuII]), this rate law should not produce the definite

intercept observed. Furthermore, the value for the K–2 (= k–2/k2) equilibrium constant is expected to

be low which again diminishes any intercept.

Thus, despite the large solvent concentration and the dissociatively-activated nature of the processes

(vide infra),143, 173 the direct X → Y substitution reaction is the only mechanism agreeing with the

data (Scheme 3-5). The low donor strength of the solvent, relative to the halide, and an early transition

state with little dissociation of the exiting ligand (generally not very positive activation entropies),

can explain this observation. For this mechanism kobs = kXY[Y-] + kYX[X-] (Appendix 3.4).

Therefore, the slope of the plot of kobs vs. [Y-] is kXY and the intercept is kXY[X-]. Here the small

outer-sphere association constants KOS are de facto included in the values of kXY and kYX. The

kinetic parameters which were determined by fitting the kobs vs. [Y-] data to Equation 3.5 C are

52

collected in Table 3.2. The temperature dependence of kXY was fit according to the Eyring equation

and the resulting values for ∆H‡ and ∆S‡ are also included in Table 3.2.

It must be emphasised that the values of kY-X should be taken with extreme caution given the large

inherent uncertainty involved in their determination (errors as large as 70-80 % of the reported value,

see Figure 3-9 and Figure 3-10). Despite the complexity in interpreting the mechanism for halide

exchange on these complexes, a number of key observations can be made from the data.

kobs = kX-Y[Y-] + kY-X[X-] (3.5)

kX-Y = konKOS(Y)

kY-X = koffKOS(X)

Scheme 3-5 Halide-exchange reaction mechanism and rate law for reaction of Y- with [Cu(Me6tren)X]+.

53

Table 3-2 Key kinetic and thermodynamic constants for halide and solvent interchange on [CuII(L)X]+ and [CuIIL(Solv)]2+

(X = Br- or Cl-, Solv = DMF, MeCN, DMSO).

Solvent Br- → Cl- 288kX→Y

(M–1 s-1)

288kY→X

(M-1 s-1)

ΔHon‡

(kJ mol-1)

ΔSon‡

(J K–1 mol-1 )

MeCN [Cu(Me6tren)Br]+ + Cl- 97 75 58 ± 2 -6 ± 7

DMF [Cu(Me6tren)Br]+ + Cl- 1.7 × 103 6.0 × 103 65 ± 3 40 ± 8

MeCN [Cu(Et6tren)Br]+ + Cl- 4.3 × 105 2.3 × 104 Too fast

Cl- → Br-

MeCN [Cu(Me6tren)Cl]+ + Br- 13 1.0 × 102 71 ± 1 15 ± 2

DMF [Cu(Me6tren)Cl]+ + Br- 8.7 7.5 × 102 92 ± 6 92 ± 20

MeCN [Cu(Et6tren)Cl]+ + Br- 1.1 × 103 1.9 × 104 74 ±3 73 ± 2

MeCN [Cu(tpa)Cl]+ + Br- 3.5 × 105 3.8 × 105 Too fast

Solv → Solv'

DMF [Cu(Me6tren)(NCMe)]2+ +

DMF 36 a

DMSO [Cu(Me6tren)(NCMe)]2+ +

DMSO 42 b

a in s–1, at 288K; b in s–1 at 298K.

Effect of the chelate

The large values of ∆H‡ and generally positive values of ∆S‡ in Table 3.2 indicate that the exchange

reactions are dissociatively activated (Id) for both Me6tren and Et6tren. Previous work has established

that exchange reactions on complexes of the less sterically hindered tren, Me3tren and tpa ligands

(Figure 3-11 & Figure 3-12) proceed by an Ia mechanism161-162, 174 so the effect of steric crowing at

the auxiliary coordination site appears to be significant.

It is interesting to compare the crystal structures of copper(II) complexes of tren, Me6tren and tpa

which have been published with chloride as a common auxiliary ligand (Figure 3-11). The Cu-Cl

bond length is virtually insensitive to the nature of the chelate. This is also true when comparing

structures of copper(II) complexes of Me6tren, Me3tren and tpa which have acetonitrile as the

common auxiliary ligand (Figure 3-12). It was stated earlier that exchange reactions proceed via an

Ia mechanism for the less-hindered tren, Me3tren and tpa ligands but an Id mechanism operates for the

more hindered Me6tren. From Figures 3.12 and 3.13 it is evident that this change in mechanism is not

due to destabilisation of the auxiliary ligand by steric influences from the chelate. Instead, steric

restrictions imparted on the incoming ligand must explain the observed changeover. Thus the

dissociative nature of the exchange with both Me6tren and Et6tren is easily resolved; these both inhibit

the close approach of Y-.

54

Figure 3-11 Comparison of the CuII‒Cl coordinate bonds (Å) across the homologous series [CuII(tren)Cl]+ (BPh4- salt);175

[CuII(Me6tren)Cl]+ (ClO4- salt)176 and [CuII(tpa)Cl]+ (Cl- salt).177 Structures obtained from the Cambridge Structural

Database and rendered with Mercury (vers. 3.5.1). The N-donors are shown in blue colour, Cl- donors in green colour.

Figure 3-12 Comparison of the CuII‒N coordinate bonds (Å) across the homologous series [CuII(Me3tren)(NCMe)]2+

(ClO4- salt);178 [CuII(Me6tren)(NCMe)]2+ (BPh4

- salt);128 and [CuII(tpa)(NCMe)]2+ (ClO4- salt).179 Structures obtained from

the Cambridge Structural Database and rendered with Mercury (vers. 3.5.1). The N-donors are shown in blue color.

While the mechanism of exchange is the same for both Et6tren and Me6tren, the rate of exchange is

different. The rate of the reaction [CuII(L)Cl]+ + Br-, where ‘L’ is Et6tren, is ~ 3 orders of magnitude

faster than the corresponding rate when ‘L’ is Me6tren. The reverse reaction between [Cu(Et6tren)Br]+

and Cl- was too fast to be measured. Overall these observations are consistent with increased steric

crowding around the axial coordination site and the dissociative nature of the exchange. It is worth

noting the significant difference in the entropy for activation between the two reactions with Et6tren

or Me6tren. Apparently the driving force for faster exchange with Et6tren is not enthalpic in origin

but rather related to the entropic rearrangement of the coordination sphere on approach to the

transition state.

Halide exchange reactions on the tpa complex were very fast, even at low temperatures, which

precluded any measurement of the thermal activation parameters. The rapidity of these reactions is

interesting in its own right. If a dissociative mechanism was operative, it would be reasonable to

expect, given the above correlations, that exchange would be slower as there is less steric repulsion

at the auxiliary position from tpa than Me6tren or Et6tren. However the exchange reactions were much

2.233

2.149

2.062

2.237

2.076

2.048

2.253

2.099

2.080

1.960

2.007

2.045

1.968

1.995

2.152

1.982

2.080

2.030

55

faster even than those for Et6tren which suggests instead that the mechanism is associatively activated

– this being consistent with previous work.174

In all, these observations make good sense of the observed differences in the deactivation behaviour

of [CuII(Me6tren)X]+, [CuII(Et6tren)X]+ and [CuII(tpa)X]+. Whilst two new methods have recently

been developed for the direct measurement of fast deactivation rate constants,58, 180 very few

experimentally determined rates exist in the literature; this is especially true for complexes which

activate initiators and halogen-capped polymers rapidly. However, as highlighted in the previous

chapter, it is possible to measure kact (or estimate it)39-40, 46, and, using known values of KATRP for a

given solvent, initiator or temperature, determine kdeact from the simple relation KATRP = kact/kdeact.20

This last reference provides a direct comparison of kdeact for the complexes [CuII(Me6tren)Br]+ and

[CuII(Et6tren)Br]+ (kdeact = 1.5 × 106 M-1 s-1 and 4.7 × 107 M-1 s-1 respectively). These rates are

consistent with the observation here that, while operating via a similar mechanism, dissociatively-

activated halide exchange is faster with the bulkier Et6tren ligand and the driving force for this

reactivity is entropic in origin. It is important to distinguish this remark from the synthetic observation

that deactivation is less efficient with Et6tren which relates to the kact/kdeact ratio.124

The value of kdeact reported for [CuII(tpa)Br]+ is 3.3 × 106 M-1 s-1 20 which is intermediate between

those for the Me6tren and Et6tren complexes. The kinetic data measured here suggests that

deactivation should be faster for this ligand but it must be remembered that the redox potential of the

CuII/I complex also influences kdeact. The CuI/II oxidation potential is much more positive for Cu/tpa

than for Cu/Me6tren meaning that tpa stabilises copper(I) to a greater extent.39, 83 The greater

electronic stabilisation of [CuI(tpa)]+ compared to [CuI(Me6tren)]+ explains why deactivation is

slower despite faster ligand exchange kinetics.

Effect of the halide

The enthalpies of activation for halide exchange reactions on [CuII(Me6tren)X]+ suggest that breaking

the Cu–X bond is more energetically demanding for [CuII(L)Cl]+ than for [CuII(L)Br]+. This is

consistent with a dissociative mechanism and the longer Cu–Br (2.392 Å)131 versus Cu–Cl (2.259

Å)132 bond. Accordingly, the rate of halide exchange is much faster where the outgoing ligand is

bromide rather than chloride. This may explain why in various polymer syntheses, the deactivation

reaction is slower with chloride than bromide.40, 59 As a whole, the data indicate that the deactivation

reaction to produce R–X, and [CuI(L)]+ species will be more favoured for X = Br- than when X = Cl-

.

56

Effect of the solvent

The role that solvent plays in tuning kdeact represents a crucial point. The data from Table 3.2 indicate

that the exchange reactions are influenced by the solvent despite the fact that it is not directly involved

in the mechanism (Scheme 3-5). Given that a solvato-complex is not formed during the exchange,

the observed dependence must be related to stabilisation of the transition state along the reaction

coordinate.

Halide substitution shows a more favourable entropic driving force in more polar solvents (ET N =

0.7775 for DMF, ET N = 0.460 for MeCN)113 which is consistent with the implicit inclusion of KOS in

the kX-Y term (see above). That is, more effective desolvation of the ions (copper and halide) by the

more polar DMF yields a larger entropic term. This driving force is likely the reason for the increased

rate of deactivation in more polar solvents.40 However it must be noted that the enthalpic term is also

slightly larger in DMF which should have the opposite effect on the rate. Because only a limited

number of systems are compared here, absolute statements correlating parameters such as polarity

with the deactivation rate will be avoided pending further work.48

57

Conclusion

The substitution reactions on [CuII(Me6tren)X]+ and [CuII(Me6tren)(Solv)]2+ have been studied across

a range of polar organic solvents and halides and were found to undergo substitution via a

dissociatively activated mechanism. This behaviour is consistent with observations for similar ligand

substitutions in aqueous and organic conditions. When solvent is initially coordinated to the copper

centre, the size of this outgoing ligand determines the rate of substitution; the rate is not sensitive to

the nature of the incoming nucleophile. When a halide is originally bound to the copper centre, the

rate of substitution is faster when the outgoing ligand is the larger, and weaker, Br- nucleophile. The

more compact Cl- nucleophile produces slower reaction rates. The full behaviour is consistent with

slower rates of deactivation in ATRP when using CuII–Cl deactivators instead of the CuII–Br

analogues.

While solvents which have the capacity to bind copper(II) influence the rates of these atom transfer

reactions, introducing an olefinic monomer such as styrene or methyl methacrylate does not. The only

effect of the monomer is in altering the outer-sphere association constant KOS.

In regards to halide exchange reactions, the results here demonstrate that this is dissociatively

activated and does not produce an intermediate solvato-complex, but the rate is still sensitive to the

nature of the solvent. In the more polar DMF, the entopic driving force is greater which indicates

more effective solvation of the copper and halide ions on approach to the transition state.

Finally, for a dissociatively activated atom transfer reaction, as observed for complexes of Me6tren

and Et6tren, increasing steric bulk around the site of exchange increases the kinetics of exchange. The

additional driving force has its origins in entropic rearrangement of the first and second coordination

spheres during the reaction. Again the increased kinetics for halide exchange with Et6tren can explain

the faster deactivation kinetics reported for this ligand over its relative Me6tren.

The dependence of the rates of anation and halide exchange are therefore sensitive to the solvent,

halide and chelating ligand ‘L’ and these sensitivities logically predict the observed variations of kdeact

as each variable is modified. It would be predicted that deactivation kinetics are fastest for copper(II)-

bromido complexes bearing ligands which crowd the coordination site of the CuIIL–X bond if

deactivation is dissociatively activated. The effect of the solvent is complex and for the moment

remains unpredictable pending an exploration of additional solvents. In this regard water is a

particularly interesting candidate because it is so effective at solvating halides and can lead to

inefficient deactivation in aqueous ATRP.37, 120-121 Exploring the kinetics and stability constants for

halide binding and substitution reactions in water should provide some interesting insights.

58

Experimental

All general chemicals used were commercially available. Solvents, of at least HPLC grade, were used

without further purification. (Bu4P)Br and (Bu4N)Cl were gently melted under high vacuum over the

course of 4h; cooling under vacuum afforded the dry salts.

Synthesis

Free Ligands

Tris-[2-dimethylamino(ethyl)]amine (Me6tren) was prepared using the method previously

described.181 1H NMR (CDCl3): δ 2.55 (m, 6H, NCH2CH2NMe2), δ 2.32 (m, 6H, CH2NMe2), δ 2.16

(s, 18H, NMe).

Tris-[2-diethylamino(ethyl)]amine (Et6tren) was prepared as previously described from tris[2-

amino(ethyl)]amine (tren).182 1H NMR (CD3CN): δ 2.49 (q, 12H, NCH2Me2), δ 2.38 (m, 12H,

CH2CH2NCH2Me2), δ 0.89 (t, 18H, NMe).

Tris-[2-pyridyl(methyl)]amine (tpa) was purchased from Sigma Aldrich and used without further

purification.

CuII Me6tren complexes

[CuII(Me6tren)(OH2)](ClO4)2183 and [CuII(Me6tren)Br]Br129 were prepared by well-established

methods and isolated as crystalline solids. [Cu(Me6tren)(OH2)](ClO4)2: Anal. Calc’d for

C12H32Cl2CuN4O9: C, 28.21; H, 6.31; N, 10.97. Found: C, 28.32; H, 6.34; N, 10.86.

[Cu(Me6tren)Br]Br Anal. Calc’d for C12H30Br2CuN4: C, 31.76; H, 3.66; N, 12.35. Found: C, 31.90;

H, 6.72; N, 12.35.

[CuII(Me6tren)(Solv)]2+ (Solv = DMSO, MeCN, DMF, MeOH or EtOH) complexes were generated

in situ by dissolution of [CuII(Me6tren)(OH2)](ClO4)2 in neat solvent. [CuII(Me6tren)Cl]+ was also

generated in situ by adding two equivalents of (Bu4N)Cl to [CuII(Me6tren)(OH2)](ClO4)2 in the

relevant solvent. In all cases no UV-vis spectral changes occur after dissolution or halide addition

indicating that displacement of the aqua ligand with the solvent or halide occurs on the mixing

timescale. This was likewise true for the Et6tren and tpa complexes below.

CuII Et6tren complexes

[CuII(Et6tren)(Solv)]2+ complexes were formed in situ by adding 1.02 equivalents of Et6tren to

Cu(ClO4)2.6H2O in the relevant solvent. Bromido and chlorido complexes were generated by the

further addition of 2.0 equivalents of (Bu4P)Br or (Bu4N)Cl respectively to [Cu(Et6tren)(Solv)]2+. The

59

UV-vis spectra observed for these complexes were very similar to those of the equivalent Me6tren

complexes.

CuII tpa complexes

The complex [CuII(tpa)(OH2)](ClO4)2 was prepared as previously described.184-185 The crystalline

solid was pure from elemental analysis and its redox and spectral behaviour were consistent with

previous reports. Anal. Calc’d for C18H20Cl2CuN4O9: C, 37.87; H, 3.53; N, 9.82. Found: C, 37.78; H,

3.48; N, 9.70.

[CuII(tpa)Cl]+ was formed in situ by adding 2.0 equivalents of (Bu4N)Cl to [CuII(tpa)(OH2)](ClO4)2

in the relevant non-aqueous solvent.

Kinetics

Standard kinetic measurements within the range 15 - 35 °C were performed using a stopped-flow

mixing unit from Applied Photophysics. For experiments run at variable pressure a previously

described pressurized stopped-flow mixing unit setup was used.186-187 All setups were connected with

fibre optics to a J&M TIDAS instrument, as described,187 allowing for the measurement of time

resolved spectra.

Substitution experiments were carried out under pseudo-first-order conditions ([X-]/[CuII] ≥ 10) in all

solvents. The concentration of the copper(II) complex was kept constant at 2.0 × 10-4 M. In all cases

the full spectrum (250 - 850 nm) was collected and analysis was carried out using the programs

SPECFIT139 or REACTLAB Kinetics.188 The observed rate constants were derived from global

analysis of the time-dependent spectral data. The greatest changes were seen in the range 250 - 450

nm. The weaker d-d electronic transitions (ca. 600 - 850 nm) were used to analyze the data only for

solvent-solvent substitution reactions. In these reactions spectral changes were small and high copper

complex concentrations were needed (3.8 × 10–3 M) to enable a significant change in the visible

absorption to be measured. In all cases the time resolved spectral changes agree with the observation

of a single first-order process; no secondary or parallel processes were detected. Typically, errors in

the observed first order rate constant (kobs) were less than 10%. The kobs values were plotted as a

function of the concentration of the reagent in excess and modelled with a rate law appropriate to that

mechanism i.e. Eqn. 3.1 accurately fits the data in Figure 3-3. KOS and kon are calculated from the fit.

The desired halide solutions were prepared with (Bu4P)Br or (Et4N)Cl; ionic strength was kept

constant at 0.1 M for all experiments using LiClO4. Appendix 3.5 collects all the values of kobs as a

function of the different concentration variables used in this study.

60

Organo-copper(II) complexes as products of radical transfer

Introduction

In ATRP, propagating radicals are capped by a halogen atom supplied by the transition metal

complex.106 However, transition metal complexes with a vacant coordination site can also react

directly with radicals to produce an organometallic species. This reaction has been exploited as an

alternative equilibrium for reversibly deactivating propagating radicals under the banner of

‘organometallic mediated radical polymerisation’ (OMRP - Scheme 4-1).

Scheme 4-1 Propagating radicals react with a halogen (ATRP) or transition metal complex (OMRP). ‘M’ is a transition

metal ion with an oxidation state of ‘n’.

OMRP has been successfully conducted using a variety of transition metals including vanadium,189-

190 chromium,191 molybdenum5 and iron.69, 192-193 The pioneering work of Wayland et al.64 utilised a

tetramesitylporphyrinato complex of cobalt as the capping agent and cobalt remains the most popular

transition metal for effecting these transformations.65-66, 194-197 Acetylacetonate and dioxime chelating

ligands are useful aternatives to porphyrins (Figure 4-1).192, 198-201

A key point of interest is the mechanistic flexibility of these organometallic complexes which can

participate in reversible termination (RT) or degenerative transfer (DT) mechanisms, or both (Scheme

4-1).66, 202-203 Here the identity of the ligand/s and the concentration of free radicals determine which

pathway is followed. In either case, the organic products are not metal-free, requiring a workup with

61

TEMPO before further post-reaction modifications are possible (Scheme 4-2).204 High catalyst

loadings are also necessary, particularly for RT OMRP as these form the stoichiometric chain ends.

Figure 4-1 Cobalt catalysts for OMRP bearing tetramesitylporphyrin (1) acetylaceconate (2) or dioxime (3) chelating

ligands.

Scheme 4-2 Synthetic workup with TEMPO to remove the metal capping agent.

There are two methods for initiating an OMRP reaction. The first involves the thermolysis or

photolysis of a stable radical precursor, such as azoisobutylnitrile (AIBN), in the presence of MnL

and monomer which generates Mn+1LR.194, 205 Spontaneous, homolytic scission of the Mn+1L–R bond

(ka,OMRP) releases the alkyl radical which activates polymerisation and the equilibrium shown in

Scheme 4-1 is established. The second method draws upon a limited number of ‘stable’ (but labile)

organometallic complexes (Mn+1LR) which are prepared and added to the reaction mixture.64 Heating

the solution results in homolysis of the Mn+1L –R bond and initiates the reaction.

OMRP is particularly interesting because of its similarity with ATRP. Both processes involve a

transition metal complex which undergoes a change in coordination number and oxidation state as it

reacts to reversibly liberate alkyl radicals. In ATRP, the complex in its lower oxidation state reacts

with an alkyl halide (kact) to release a radical which initiates (or continues) activation; in OMRP

effectively the reverse reaction occurs where the radical is deactivated (kd,OMRP) by coordination as a

carbanion (R-). Therefore, while both processes employ transition metal catalysts, the positions of the

central equilibria are diametrically opposed; KATRP << 1; KOMRP >> 1 (Scheme 4-3). This difference

forms the basis for the selective application (and operation) of ATRP or OMRP.

62

Scheme 4-3 Central atom transfer and radical transfer reactions for ATRP and OMRP

Figure 4-2 Olefinic monomers

Monomers differ in several ways but the most important parameter is the reactivity of the formed

radical (R•). Less-active monomers (referred to in the literature as ‘LAMs’) are those which form

unstable radicals; these include simple olefins (e.g. ethylene, propylene), vinyl acetate and vinyl

chlorides (Figure 4-2). The dormant form of a LAM (olefinic or capped) is particularly inert and the

propagating radical is especially reactive.

OMRP is better suited to these monomers than ATRP because it utilises a much higher concentration

of the deactivator (in this case MnL) so termination is still avoided. Furthermore, the lability of the

Mn+1L–R bond ensures that an appreciable concentration of radicals is maintained by the reverse,

activation reaction and polymerisation proceeds at a reasonable rate. Conversely, when a LAM radical

is instead capped by a halogen atom (ATRP), the stability of the carbon-halogen bond makes

reactivation difficult. An additional boon for OMRP is that the metal complex can easily be

synthetically modified in order to tune the strength of the metal-carbon bond. Thus for more reactive

radicals the Mn+1–R bond can be weakened to maintain the propagation rate, and for less reactive

monomers it can be strengthened to ensure a low concentration of radicals.

63

In ATRP the energy required to break the dormant alkyl-halogen bond is provided by the catalyst

(MnL). An additional driving force is provided by the thermodynamic stability of Mn+1LX which is

the product of activation. In theory then, ATRP catalysts can also be tuned to cater for LAMs. To

date, ATRP has mostly been used to polymerise more reactive monomers such as acrylates,

methacrylates and acrylonitriles which do not require such active catalysts. Designing a catalyst

which has the potential to activate LAMs represents one of the great challenges for this field and in

this regard, only moderate success has been achieved.206

It should be noted before progressing that of the two techniques, ATRP has been much more widely

adopted as it is suitable for a larger array of monomers, is tolerant to a variety of functional groups

and impurities, uses only catalytic quantities of metal and has the added benefit of producing halogen-

terminated chains which are conducive to post-reaction modifications.

The activity of an ATRP catalyst is generally correlated with KATRP (kact/kdeact). Currently the most

active catalysts are those of copper bearing Me6tren, tpa and substituted tpa, as well as macrocyclic

chelating ligands (Figure 4-3).20, 61, 206-208 Using a modified cyclam catalyst, a recent report presented

the first stand-alone ATRP polymerisation of vinyl acetate with good control over the polymer

molecular weight distribution.209 However, similar success has not yet been achieved for other LAMs

such as vinyl chloride.

Figure 4-3 Chelate ligands of highly active copper catalysts.

64

The deployment of highly active copper catalysts has also been correlated with unusual experimental

phenomena such as polymerisations which are slower than expected.61, 72 Synergistic atom transfer

(ATRP) and radical transfer (OMRP) has been posited as the source of the discrepancy. While the

potential overlap between ATRP and OMRP appears obvious on paper, experimental evidence for

their parallel operation is a much more recent discovery. Several reports provide evidence to this

effect from different perspectives.

For example, the uncontrolled, free-radical polymerisation (FRP) of n-butyl acrylate initiated by

AIBN became much more controlled when the ATRP catalyst [CuI(tpa*)]+ (tpa* = tris((4-methoxy-

3,5-dimethylpyridin-2-yl)methyl)amine) was present. This suggests the suppression of radical

termination by formation of the species [CuII(tpa*)R]+.71

A further study which utilised [CuI(tpa)]+ and [CuI(Me6tren)]+ also indicated that these compounds

may react with propagating methyl acrylate radicals.207 By comparing the theoretical (calculated)

versus experimental yields of terminated polymer chains, the authors found significantly higher-than-

expected concentrations of the terminated products. Thus, in addition to conventional termination (kt)

(including disproportionation) an additional pathway was required. The reaction of methyl acrylate

radicals with [CuI(L)]+ to form [CuII(L)R]+ was proposed. This complex can react with a second

radical in a step known as catalytic radical termination (CRT) to form the additional terminated

products (Figure 4-4).

EPR measurements following the reaction of [CuI(tpa)]+ generated in the presence of butyl acrylate

radicals also indicated the formation of a unique copper(II) species.210 Under these conditions, a new

signal, distinct from any known copper(II)-halido complex of tpa, was observed.

Figure 4-4 Proposed formation of a organocopper(II) species which leads to increased R-R terminated products via a

Catalytic Radical Termination (CRT) pathway.

65

Whilst some metals, such as iron, have been successfully utilised in both OMRP and ATRP,69, 193, 211-

213 copper complexes, which are the most active and widely used ATRP catalysts,17, 61, 125, 214 have

not been successfully applied to OMRP. Relatively little is known in general about cupric

organometallic complexes with amino ligands. A survey of the Cambridge Structural Database (CSD)

indicates that alkyl ligands can coordinate polyamine complexes of copper(II), but very few examples

are extant.215-217 By contrast, there are literally thousands of CuI organometallic complexes in the

CSD, the majority bearing N-heterocyclic carbene ligands. So while concurrent Cu-based ATRP and

OMRP is theoretically possible, unequivocal evidence of the organometallic complex CuIIL-R is

lacking. Identifying this species is the focus of this Chapter.

Electrochemistry is particularly suitable for this purpose. Both ATRP and OMRP rely on CuIL

undergoing a reversible change in coordination number and oxidation state. Chapter 2 illustrated how

CuIL is generated electrochemically from a resting solution of CuIILBr and how activation in the

presence of RX (re)generates CuIILBr and a radical R•. By reducing a stable solution of CuIILBr in

the presence of RX, each of the prerequisite components for both the ATRP and OMRP equilibria

are present. If radicals produced near the electrode by kact combine with CuIL instead of terminating

(kt) or deactivating (kdeact), CuIILR should form within the diffusion layer and be evidenced by the

appearance of a novel redox couple.

To this end a series of electrochemical experiments were carried out on resting solutions of CuIIL-Br

with Me6tren or tpa chelating ligands in DMSO and MeCN. Ethyl α-bromoisobutyrate (EBriB) and

bromoacetonitrile were used as the alkyl halide initiators.

66


Electrochemistry of [CuII(tpa)Br]+

The CV of a 1.0 mM solution of [CuII(tpa)Br]Br in DMSO (0.1 M (Et4N)(ClO4)) is shown in Figure

4-5. In the absence of any initiator (RX), quasi-reversible reduction to [CuI(tpa)Br] is observed. When

EBriB is added the cathodic current is amplified, consistent with the ECcat mechanism described in

Chapters 1 & 251, 99, 180 and a second wave also appears at lower potentials. When bromoacetonitrile

is used as the initiator the lower potential wave is much more obvious. The control experiment in the

absence of copper does not produce either wave; a very small amount of nonspecific reduction of the

initiator is observed at ~ -1200 mV vs. Fc+/0 (Appendix 4.1). When the solvent is changed to MeCN

analogous behaviour is observed albeit at different potentials for bromoacetonitrile (Figure 4.4D); no

second wave is observed for EBriB.

Figure 4-5 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br + [RBr] in DMSO (A & B) or MeCN (C & D). The relevant

initiators are shown above each data voltammogram. Sweep rate = 200 mV s-1. Solid curves are the experimental data

and broken lines the simulated data. I = 0.1 M (Et4N)(ClO4).

-1200 -1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

5.0 mM

20 A

E / mV vs. Fc+/0

A B

C

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

D

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5.0 mM

20 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

ER EBr

67

Scheme 4-4 Mechanism of electrochemically-induced, concomitant halogen atom transfer (ATRP) and radical transfer

(OMRP) from the ATRP deactivator CuIILBr.

The appearance of a lower potential wave is consistent with competing reversible halogen atom

transfer and radical transfer pathways as illustrated in Scheme 4-4. As the potential is swept in the

negative direction towards EBr, the activation reaction produces the radical (R•) which reacts with

[CuI(tpa)]+ also in the diffusion layer to generate the organometallic species [CuII(tpa)R]+. The

auxiliary ligand R- thus generated is formally a strongly σ-donating carbanion and consequently

[Cu(tpa)R]+ is reduced in a second electron transfer reaction at a lower potential (ER) than the bromido

complex. Reduction of the radical transfer product [CuII(tpa)R]+ is coupled to rapid dissociation of

R- (k-R) which regenerates the active catalyst [CuI(tpa)]+ so the waveform is also catalytic at this

potential. The loss of R− from [CuI(tpa)]+ is driven by the strong preference of CuI for a coordination

number of four which is satisfied by the tetradentate tpa ligand. The fate of the strongly basic

carbanion R− is not known, but it is most likely protonated by traces of water. The specific role of

water is addressed later.

Spectroelectrochemistry of [Cu(tpa)Br]+

In order to support the proposal that the species being reduced at lower potential is [CuII(tpa)R]+,

spectroelectrochemistry of 6.0 mM [CuII(tpa)Br]Br in MeCN (0.1 M (Et4N)(ClO4)) with 48 mM

bromoacetonitrile was performed while holding the potential slightly below EBr (Figure 4-6 A). At

this potential [CuI(tpa)]+ is continuously regenerated by electrolysis and the •CH2CN radical is formed

by activation (kact). Both [CuI(tpa)]+ and •CH2CN are generated within the electrode diffusion layer

and react to form the organometallic complex [CuII(tpa)(CH2CN)]+. Because the applied electrolysis

potential is higher than the [CuII/I(tpa)(CH2CN)]+/0 couple, the organocopper(II) complex accumulates

at the working electrode.

68

Figure 4-6 A) Spectra measured every five seconds during electrolysis of 6.0 mM [CuII(tpa)Br]Br + 48 mM in MeCN.

Potential was held at -850 mV vs. Fc+/0. B) Spectra of 6 mM [CuII(tpa)Br/R/CN]+ in MeCN. The red trace is the final

spectrum from A. I = 0.1 M (Et4N)(ClO4).

Accordingly, the spectrum of [CuII(tpa)Br]+ displayed a hypsochromic shift during electrolysis with

retention of a clear isosbestic point during the transformation. The final spectrum has a peak at 750

nm and a higher energy shoulder. The fine details of the higher energy transition are obscured because

of uncorrected light scattering from the honeycomb electrode. Nevertheless, the final spectrum

indicates the formation of a distinctly different species, assigned here to the C-bound

[CuII(tpa)(CH2CN)]+ complex. That the electronic transitions for this species occur at higher energy

is consistent with the position of a strongly σ-donating carbanion in the spectrochemical series. This

proposal was supported by a separate experiment where a stoichiometric amount of sodium cyanide

(from a stock solution in 2:8 v/v H2O/MeCN) was added to 6.0 mM [CuII(tpa)Br]Br in MeCN to

generate the spectrum shown in Figure 4-6 B. Clearly the spectra of [CuII(tpa)(CN)]+ and putative

[CuII(tpa)(CH2CN)]+ are similar and distinct from the bromido analog. Note that the spectrum of N-

bound [CuII(tpa)(NCCH3)]2+ displays a vis−NIR spectrum with a maximum at 850 nm,218 thus ruling

out this new species bearing an N-bound co-ligand.

The new species observed at low potential is not the hydrido-complex [CuII(tpa)H]+. The formation

of hydrido complexes has been proposed during the [Cu(tpa*)]+-catalysed polymerisation of

acrylate71 and is grounded in the cobalt-mediated OMRP of methacrylates.202, 219 The hydrido

complex, along with an olefin, is formally generated by β-hydride elimination from Mn+1LR or by

direct hydrogen atom abstraction from a radical (Scheme 4-5). Hydrido complexes are formed by this

second reaction, known as catalytic chain transfer (CCT), in various cobalt-mediated

polymerisations220-221 as well as iron-based OMRP222 however CCT has not yet been reported with

copper.

400 600 800 1000

0.0

0.5

1.0

Ab

so

rban

ce

Wavelength / nm

400 600 800 1000

0.0

0.5

1.0

[CuII(TPMA)Br]

+

[CuII(TPMA)R]

+

[CuII(TPMA)CN]

+

Ab

so

rban

ce

Wavelength / nm

A B

69

Scheme 4-5 Radical and hydrogen transfer reactions which compete with the central OMRP equilibrium. CRT – catalytic

radical termination, CCT – catalytic chain transfer. Charges omitted for clarity.

There are various pieces of evidence which suggest that a hydrido-copper(II) complex is not formed

here. Firstly, the C-bound NCCH2- anion (originating from bromoacetonitrile) cannot undergo β-

hydride elimination or direct hydrogen abstraction. For experiments involving EBriB this is still a

possibility. DFT calculations indicate that an organometallic species is much more energetically

accessible than a hydrido complex from a reaction between Cu/tpa and a methyl acrylate radical

(Scheme 4-5, CCT over OMRP)223 which explains why CCT has not been observed with copper.

The most direct support for the organometallic complex versus the hydrido complex comes from the

voltammetry in Figure 4-5. The redox potential of the low potential complex in the DMSO/EBriB

system is very similar to that of [CuII(tpa)CH2CN]+ (Table 4.1) which suggests that it too is most

likely the carbanion-bound complex.

The possibility of coordination by a second halido ligand was also excluded. Adding excess free

bromide (as (Et4N)Br) in place of bromoacetonitrile had no impact on the spectrum of [CuII(tpa)Br]+.

The spectrum of [CuBr2] in MeCN (0.1 M (Et4N)(ClO4)) is also distinctly different (Appendix 4.2).

Final confirmation that ligand exchange had occurred was provided by measuring the EPR

spectroscopy of [CuII(tpa)Br]+ (before) and [CuII(tpa)(CH2CN)]+ (after electrolysis). The frozen

solution spectra in MeCN are shown in Figure 4-7 and Figure 4-8. The features of an axially

symmetric CuII complex with trigonal bipyramidal geometry (gx,y > gz) are retained in both cases but

electrolysis is accompanied by a significant shift in the spin Hamiltonian parameters which indicates

ligand substitution. The EPR spectroscopy of CuII/tpa complexes of this kind has been discussed

elsewhere224 and the spectra reported here are consistent with these.

70

Figure 4-7 Experimental (top) and simulated225 (bottom) X-band (9.3708 GHz) EPR spectra at 130 K of 1 mM

[CuII(tpa)Br]+. Solvent - MeCN:toluene 1:1. Spin Hamiltonian parameters: gx,y = 2.183 (Ax,y = 105 G), gz = 2.005 (Az =

70 G).

Figure 4-8 Experimental (top) and simulated225 (bottom) X-band (9.3708 GHz) EPR spectra at 130 K of 1 mM

[CuII(tpa)(CH2CN)]+ formed after bulk electrolysis of 1 mM [CuII(tpa)Br]+ in the presence of bromoacetonitrile. Solvent

- MeCN:toluene 1:1. Spin Hamiltonian parameters: gx,y = 2.178 (Ax,y = 100 G), gz = 1.948 (Az = 98 G).

2D Graph 1

H / Gauss

2800 3000 3200 3400 3600

Col 1 v Col 2

Col 3 v Col 4

exp.

sim.

2800 3000 3200 3400 3600

H / Gauss

exp

sim

H / Gauss

2800 3000 3200 3400 3600

exp.

sim.

2800 3000 3200 3400 3600

H / Gauss

exp

sim

71

Simulating the voltammetry of [CuII(tpa)Br]+

Using the mechanism shown in Scheme 4-4 the experimental voltammetry was successfully

reproduced with the program Digism.82 A number of alternate mechanisms were also examined (vide

infra) but none reproduced the experimental behaviour. The simulated voltammograms are shown in

Figure 4-5 as broken lines.

The fitting process was undertaken in a similar manner to that described in Chapter 2 by first

determining the majority of the physical and chemical constants such as KII,Br, kIId,Br, E0, KI,Br, and

kId,Br. The routine methods for determining these parameters are reiterated in the Experimental section

of this Chapter.

With these parameters fixed, the catalytic voltammetry was measured with a large excess of TEMPO

in order to isolate the activation rate (kact). This was done for both EBriB and bromoacetonitrile.

Adding TEMPO provides further support for the mechanism of Scheme 4-4 as the second, low

potential wave disappears (Figure 4-9). The radical capture of R• by TEMPO therefore prevents the

formation of the organometallic complex. This behaviour was reproduced by each of the relevant

initiator/solvent combinations (Figure 4-9).

Figure 4-9 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br in DMSO. Black – no bromoacetonitrile or TEMPO; Red –

added 1.0 mM bromoacetonitrile; Blue – added 0.2 M TEMPO. Sweep rate = 100 mV s-1. I = 0.1 M (Et4N)(ClO4).

-1400 -1200 -1000 -800 -600 -400

10 A

E / mV vs. Fc+/0

72

Figure 4-10 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br + [RBr] with added TEMPO (0.2 M). Solvent = DMSO (A

& B) or MeCN (C & D). The relevant initiators are shown above voltammogram. Sweep rate = 200 mV s-1 for all. Solid

lines are experimental data and broken lines simulated data. I = 0.1 M (Et4N)(ClO4).

The voltammetry with added TEMPO was simulated according to the mechanism shown in the right-

hand cycle of Scheme 4-4. From this mechanism kact is the only unknown parameter because

deactivation is quenched and the other physical and chemical constants are known (see

Experimental). This was the sole variable which was allowed to float during the fitting process and

the values determined for each initiator/solvent combination are collected in Table 4.1. In each case

the simulations accurately reproduce the experimental voltammetry across a range of concentrations

of RBr and sweep rates.

The fits are also very sensitive to the magnitude of kact. Figure 4-11 illustrates how changing the

simulated value of kact from the value reported in Table 4.1 (broken red) to a value comparable with

those calculated in the literature20 for this same catalyst/solvent/initiator combination (broken blue)

does not reproduce the experimental data.

-1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

-1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

-1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

-1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

A B

C D

73

Having determined each of the parameters from the right hand cycle of Scheme 4-4, excluding kdeact,

the catalytic voltammetry was measured without TEMPO and simulated. The only parameters which

were allowed to refine during the iterative fitting were kdeact, kd,OMRP and KOMRP and the dissociation

rate of R- from [CuI(tpa)R] (kId,R). The results for the Cu/tpa system are collected in Table 4.1.

Figure 4-11 Red) Experimental (solid) and simulated (broken) voltammetry of 1.0 mM [CuII(tpa)Br]Br in MeCN (0.2 M

TEMPO) with 4.0 mM EBriB. When kact is perturbed away from the value determined here (Red – kact = 3.6 × 104 M-1 s-

1) to a smaller value (Green - kact = 3.6 × 103 M-1 s-1 or Blue - kact 3.6 × 102 M-1 s-1) the catalytic current is not accurately

reproduced. Sweep rate = 200 mV s-1.

-1000 -800 -600 -400

20 A

E / mV vs. Fc+/0

74

Table 4-1 Thermodynamic and kinetic parameters for the [CuII(tpa)Br]Br system with bromoacetonitrile or EBriB as

initiators and DMSO or MeCN as solvent. Potentials are cited as mV vs. Fc+/0.

MeCN DMSO

EBriB bromoacetonitrile EBriB bromoacetonitrile

KATRP 1.7 × 10-3 3.3 × 10-1 5.2 × 10-3 1.8 × 10-2

kact (M-1 s-1) 3.6 × 104 4.0 × 104 2.4 × 104 1.8 × 104

kdeact (M-1 s-1) 2.1 × 107 1.2 × 105 4.6 × 106 9.9 × 105

KOMRP (M-1) - 2.2 × 108 1.1 × 104 3.9 × 107

kd,OMRP (M-1 s-1) - 3.6 × 107 1.8 × 107 1.8 × 107

ka,OMRP (s-1) - 0.16 1.6 × 103 0.44

KI,R (M-1) - 1.5 × 104 1.9 × 103 2.4 × 104

kId,R (s-1) - 150 2.2 × 103 12

KII,Br (M-1) 3.4 × 107 3.4 × 107 1.1 × 104 1.1 × 104

kIIa,Br (M-1 s-1) 2.4 × 107 2.4 × 107 1.2 × 107 1.2 × 107

KI,Br (M-1) 5.4 × 103 5.4 × 103 410 410

kIa,Br (M-1 s-1) 5.7 × 105 5.7 × 105 7.6 × 105 7.6 × 105

ESol (mV vs. Fc+/0) -425 -425 -610 -610

EBr (mV vs. Fc+/0) -650 -650 -695 -695

ER (mV vs. Fc+/0) - -960 -925 -950

75

Hydrolytic decomposition of [CuI(tpa)R]

The possible involvement of water in the reaction mechanism was considered in a separate

experiment which is worth noting before further discussion. Here a 10 mL solution of MeCN was

prepared which contained 1.0 mM (Et4N)Br and 3.0 mM bromoacetonitrile in addition to the

electrolyte (0.1 M (Et4N)(ClO4)). Each of these reagents was thoroughly dried before use (see

Experimental) and the water content in this solution was determined by NMR as less than 1.0 mM.

To this solution was added 1.0 mM [CuII(tpa)(H2O)](ClO4)2 which raised the water content by 1.0

mM as the aqua ligand is replaced by MeCN (18 M) or Br- giving a maximum concentration of about

2.0 mM H2O. The voltammetry of this solution displays the usual features (Figure 4-12).

Adding a further 10, 30 or 50 mM H2O had no significant effect on the catalytic voltammetry at EBr

or ER (Figure 4-12). The small shift in the peak potentials observed during the addition of water is

due to uncorrected changes in the water content of the solvent relative to the water-free solvent in the

reference electrode. If water is required to liberate [CuI(tpa)]+ at ER by protonating the carbanion

([CuI(tpa)R] + H2O [CuI(tpa)]+ + RH + OH-) then the catalytic current at this potential should be

amplified by the addition of water. The insensitivity of the current to H2O indicates that the rate-

limiting step in regenerating [CuI(tpa)]+ is the unimolecular reaction [CuI(tpa)R] [CuI(tpa)]+ + R-.

Similar anion dissociation occurs at EBr to (re)generate [CuI(tpa)]+ from [CuI(tpa)Br].

Figure 4-12 Cyclic voltammetry of 1.0 mM [CuII(tpa)Br]Br in MeCN (0.1 M (Et4N)(ClO4)) with 3.0 mM

bromoacetonitrile and various added [H2O]. Sweep rate = 100 mV s-1.

As mentioned earlier, CuI complexes of tpa (and CuI complexes in general) strongly prefer four-

coordinate tetrahedral coordination geometries. A survey of the Cambridge Structural Database

shows that all 20 CuI/tpa complexes (including its substituted analogues) are either genuinely four

coordinate or, if a fifth donor atom is present, the coordinate bond to one donor atom is exceptionally

long (> 2.4 Å) in comparison with the remaining four CuI coordinate bonds (2.0 − 2.1 Å). The

-1200 -1000 -800 -600 -400 -200

2 mM

12 mM

32 mM

52 mM

20 A

E / mV vs. Fc+/0

76

elongated bond in pseudo-five-coordinate CuI/tpa complexes is typically the Cu−N bond to the

tertiary amine. Therefore, there is a thermodynamic driving force for dissociation of the monodentate

coordinated alkyl ligand.

Factors controlling ATRP versus OMRP

The data in Table 4.1 can be compared with calculated and measured rates for the central atom transfer

reactions published elsewhere.20, 226 For particularly active catalysts such as Cu/Me6tren and Cu/tpa

the activation rate constant (kact) has been calculated in acetonitrile for both EBriB and

bromoacetonitrile.20 Recently however, electrochemical measurements by Matyjaszewski and

Gennaro et al. generated values which were two orders of magnitude larger than those published

earlier.226 For Cu/tpa, kact was found to be slightly greater with bromoacetonitrile than EBriB; both

were reported as ~ 104 M-1 s-1. These results are comparable to those determined here in Table 4.1.

While the order is reversed in DMSO, the difference is not significant.

Both of the aforementioned studies (refs 20 and 226) utilise a measured value of KATRP along with

kact to obtain kdeact (KATRP = kact/kdeact). KATRP was determined by monitoring the formation of the

persistent radical ([CuIILX]+ in this case) during the reaction between R• and [CuIL]+ and fitting the

concentration profile to the modified Fischer equation.35 In light of the results which are presented

here, both [CuIILX] and [CuIILR]+ will be formed under these conditions where R• is the

bromoacetonitrile or EBriB radical. This was not considered in the derivation of the modified Fischer

equation so the published values of KATRP and the corresponding values of kdeact serve only as a rough

comparison.

The results shown in Table 4.1 confirm that KATRP is larger for bromoacetonitrile than EBriB.

Interestingly, the solvent appears to have little impact on the magnitude of kact for these particularly

highly activating initiator/catalyst systems but as Chapter 3 showed, the influence of solvent on atom

transfer reactions involving trigonal bi-pyramidal copper(II) complexes is much more complicated

than originally anticipated.227 No clear trend emerges here for the values of kdeact except for the

observation that deactivation is faster than activation.

Of particular interest are the rate and equilibrium constants for the OMRP reaction. It is known that

alkyl228 and aliphatic radicals rapidly coordinate copper(I) in aqueous solutions both in the absence

(106 M-1 s-1)229 and presence (108 – 109 M-1 s-1)230 of a chelating ligand. In particular, the rate constants

for the reaction between •CH2COO- or •CH2CH2COO- and [CuIL2]+ (‘L2’ is the linear tetraamine

2,4,8,11-tetramethyl-2,5,8,11-tetraazadodecane) have been reported as 2.7 × 107 M-1 s-1 and 6.3 × 106

M-1 s-1 respectively.230 Here the simulations reveal that acetonitrile radicals coordinate [CuI(tpa)]+ in

organic solvents at approximately the same rate (kd,OMRP ~ 107 M-1 s-1).

77

The identity of the radical affects the value of KOMRP (formation of the organometallic complex) with

EBriB < bromoacetonitrile and the source of this difference is the rate constant ka,OMRP (homolysis of

the CuII‒R bond); ka,OMRP is larger with the more sterically encumbered EBriB radical. The values of

ka,OMRP which are reported here are also similar to those published by Meyerstein et al. for homolysis

of the [CuI(L2)CH2R]+ Cu–C bond.230 It is also interesting to note that the position of the OMRP

equilibrium is always shifted towards the organometallic complex (kd,OMRP >> ka,OMRP) consistent with

the expectations for RT OMRP.

The most important conclusion to draw from Table 4.1 is that [CuI(tpa)R]+ is observed for

combinations of highly active catalysts and initiators which also have large values for KOMRP. The

most distinct observation of the organometallic complex occurs when bromoacetonitrile is the

initiator. For this initiator kact is fast (104 M-1 s-1) leading to a large concentration of R• in the vicinity

of the electrode and KOMRP is also large (107 – 109) which favours the formation of the CuII

organometallic complex. The necessary synergy between these parameters is highlighted by the

EBriB/DMSO system where kact is large but KOMRP is small. Here the organometallic adduct is only

observed to a minimal extent due to steric repulsion at the site of coordination. These conclusions are

consistent with the previous work already mentioned using the most highly activating catalyst

[CuI(tpa*)]+.61 With this catalyst activation is very fast leading to a high concentration of radicals

available to coordinate the CuI centre.

The fate of the radical also depends on the ratio of kdeact/kd,OMRP and this represents a significant point.

In MeCN, deactivation is very fast (107 M-1 s-1) so the radical is reconverted to the dormant, halogen-

capped form before reacting with [CuI(tpa)]+. In DMSO the deactivation reaction is slower and a

small amount of [CuII(tpa)R]+ is observed.

As is highlighted by Scheme 4-5, as well as a number of papers which have already appeared in the

literature, there are further reactions, in addition to the OMRP equilibrium, in which [CuII(tpa)R]+

may participate. The mechanism assumed in Scheme 4-4 does not include CRT to regenerate

[CuI(tpa)]+ + R-R or β-hydride elimination to produce the hydrido-complex [Cu(tpa)H]+. While the

reasons for excluding the latter reaction have already been discussed, CRT is still a plausible reaction

under these conditions. However, including this step in the reaction mechanism made no difference

to the quality of fit (i.e. the rate could be set as zero with no change to the fit). This does not preclude

its possible involvement but merely emphasises that [CuI(L)]+ is regenerated at the electrode from

[CuII(tpa)R]+ by reduction and subsequent loss of R-. The forcing conditions from the applied

potential mean that [CuII(tpa)R]+ does not survive for long enough to participate in CRT or related

reactions.

78

Electrochemistry of [Cu(Me6tren)Br]+

Having examined the effect of the initiator and solvent on the formation of the organometallic

complex, the influence of the chelating ligand was also considered. As Chapter 2 demonstrated, the

steric bulk of the terminal functional groups on tripodal chelating ligands such as tpa, Me6tren and

Et6tren has a significant effect on ligand substitution reactions at the remaining axial coordination

site.227 Therefore the voltammetry and spectroelectrochemical behaviour of the related, but more

sterically hindered, saturated tetra-amine complex [CuII(Me6tren)X]X was also investigated in the

presence of bromoacetonitrile and chloroacetonitrile. Measurements with EBriB were not made as

the voltammetry has already been reported with this combination of catalyst/initiator and displays no

evidence of the organometallic complex.51

The CVs of 1.0 mM [CuII(Me6tren)X]X (in MeCN) in the presence of bromoacetonitrile or

chloroacetonitrile are shown in Figure 4-13. Bromide was the initial auxiliary ligand when

bromoacetonitrile was added and chloride when chloroacetonitrile was added. While

[CuII(Me6tren)Cl]+ is reduced at a lower potential than [CuII(Me6tren)Br]+, in both cases the second

wave appears at a similarly cathodically shifted potential (Table 4.2) providing further support for

the proposal that the common auxiliary ligand of this new CuII complex is R-. In DMSO, the CVs

likewise show that the organometallic species is formed (Figure 4-14). While organocopper

complexes have been suggested to form with tpa and tpa*, this data represents the first suggestion

that the OMRP equilibrium may also be important with the popular Cu/Me6tren catalyst.

Figure 4-13 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)Br]Br (A) or 1.0 mM [CuII(Me6tren)Cl]Cl (B) in MeCN +

[RX]. Relevant initiators are shown above each voltammogram. Sweep rate = 200 mV s-1 for all. Solid curves are the

experimental data and broken lines the simulated data. I = 0.1 M (Et4N)(ClO4).

-1200 -1000 -800 -600 -400

0.0 mM

2.0 mM

4.0 mM

6.0 mM

8.0 mM

20 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400

0.0 mM

2.0 mM

6.0 mM

12.0 mM

20.0 mM

20 A

E / mV vs. Fc+/0

A B

79

Figure 4-14 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)Br]Br in DMSO + bromoacetonitrile. Sweep rate = 500 mV s-

1. Solid curves are the experimental data and broken lines the simulated data. I = 0.1 M (Et4N)(ClO4).

Figure 4-15 Cyclic voltammetry of 1.0 mM [CuII(Me6tren)X]X + [RX] with added TEMPO (0.2 M). A) X = Br-, Solv. =

MeCN; B) X = Cl-, Solv. = MeCN. C). X = Br-, Solv. = DMSO. Sweep rate = 200 mV s-1 for all. Solid lines are

experimental data and broken lines simulated data. I = 0.1 M (Et4N)(ClO4).

The addition of TEMPO again quenched the radical transfer reaction (Figure 4-15) meaning kact could

be obtained by simulating the data according to the right hand cycle of Scheme 4-4.

Spectroelectrochemistry of [Cu(Me6tren)Br]+

The spectrum of [CuII(Me6tren)Br]Br in MeCN with added bromoacetonitrile underwent a similar

hyposchromic shift to [CuII(tpa)Br]Br during electrolysis at -1000 mV vs. Fc+/0 (Figure 4-16 A). The

spectra of the initial [CuIILBr]+ and final [CuIILR]+ complexes were similar for L = Me6tren, tpa;

however, the higher energy shoulder of the organometallic complex was less intense with Me6tren.

-1400 -1200 -1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400

0.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400 -200

0.0 mM

2.0 mM

6.0 mM

12.0 mM

20.0 mM

20 A

E / mV vs. Fc+/0

-1200 -1000 -800 -600 -400

0.0 mM

1.0 mM

2.0 mM

3.0 mM

4.0 mM

20 A

E / mV vs. Fc+/0

A B C

80

Changing the solvent to DMSO did not impact the peak positions of the final spectrum nor did

changing to the [CuII(Me6tren)Cl]Cl/chloroacetonitrile system in MeCN (Figure 4-17). This again

supports the proposal that the common auxiliary ligand is R- as opposed to a second halide or solvent

ligand. Adding KCN to [Cu(Me6tren)Br]Br shifted the main peak to a higher energy, comparable

with that of putative C-bound [Cu(Me6tren)(CH2CN)]+ (Figure 4-16 B).

Figure 4-16 A) Spectra measured every 20 seconds during electrolysis of 6 mM [CuII(Me6tren)Br]Br + 48 mM

bromoacetonitrile in MeCN. Potential was held at -1000 mV vs. Fc+/0. B) Spectra of 6 mM [CuII(Me6tren)X]+ in MeCN

(X- = Br-, CN- or R = NCCH2-). Red spectrum is the final spectrum from A. I = 0.1 M (Et4N)(ClO4).

Figure 4-17 A) Spectra measured every 20 seconds during electrolysis of 6 mM [CuII(Me6tren)Cl]+ + 48 mM

chloroacetonitrile in MeCN. Potential was held at -900 mV vs. Fc+/0. B) Spectra of 6 mM [CuII(Me6tren)Br]+ + 48 mM

bromoacetonitrile in DMSO measured every 20 seconds during reduction at -1000 mV vs. Fc+/0. I = 0.1 M (Et4N)(ClO4)

for both.

400 600 800 10000.0

0.5

1.0

1.5

2.0

[CuII(Me

6tren)Br]

+

[CuII(Me

6tren)R]

+

[CuII(Me

6tren)CN]

+

Ab

so

rban

ce

Wavelength / nm

400 600 800 10000.0

0.5

1.0

1.5

2.0

Ab

so

rban

ce

Wavelength / nm

A B

400 600 800 10000.0

0.5

1.0

Ab

so

rban

ce

Wavelength / nm

400 600 800 10000.0

0.5

1.0

1.5

Ab

so

rban

ce

Wavelength / nm

A B

81

Simulating the voltammetry of [CuII(Me6tren)X]+

The mechanism from Scheme 4-4 was used to simulate the catalytic voltammetry of

[CuII(Me6tren)X]X in both MeCN and DMSO and Table 4.2 collects the relevant parameters for the

simulations. It is important to note that the ATRP-relevant parameters reported here are again mostly

consistent with established trends in activation and deactivation rates and efficiencies.20, 40, 46-47 For

example, kact is faster for bromoacetonitrile than chloroacetonitrile and activation rates with

Cu/Me6tren were also generally faster than those in the analogous experiments with Cu/tpa.

Deactivation was also slower for the chloro- versus the bromo-complex which is expected from the

kinetics data presented in Chapter 3.

Table 4-2 Thermodynamic and kinetic parameters for the [CuII(Me6tren)X]X system in MeCN or DMSO. For

bromoacetonitrile, X = Br-, for chloroacetonitrile, X = Cl-.

[CuII(Me6tren)X]X

in MeCN

[CuII(Me6tren)Br]Br

in DMSO

bromoacetonitrile chloroacetonitrile bromoacetonitrile

KATRP 8.7 × 10-3 2.0 × 10-2 2.1 × 10-2

kact (M-1 s-1) 1.0 × 105 5.9 × 102 3.6 × 105

kdeact (M-1 s-1) 1.2 × 107 2.9 × 104 1.7 × 107

KOMRP (M-1) 3.0 × 108 3.0 × 108 1.8 × 107

kd,OMRP (M-1 s-1) 3.0 × 107 3.0 × 107 1.6 × 107

ka,OMRP (s-1) 0.1 0.1 8.8 × 10-1

KI,R (M-1) 5.6 × 104 5.6 × 104 10

kId,R (s-1) 28 28 16

KII,Br (M-1) 1.2 × 107 7.5 × 108 7.6 × 103

kIIa,Br (M-1 s-1) 6890 7560 1.2 × 104

KI,Br (M-1) 1.6 × 103 3.6 × 103 337

kIa,Br (M-1 s-1) 1.9 × 106 2.6 × 104 3.3 × 106

ESol (mV vs. Fc+/0) -475 -475 -730

EBr (mV vs. Fc+/0) -705 -825 -810

ER (mV vs. Fc+/0) -975 -975 -1080

Crucially, the experimental CVs with bromoacetonitrile and chloroacetonitrile (Figure 4-15) were

accurately simulated using identical values of KOMRP, kd,OMRP, KI,R and kId,R as radicals produced by

homolytic cleavage of the R‒Br or R‒Cl bonds are identical. The experimental differences in the CVs

82

are therefore resolved by thermodynamic and kinetic differences in the ATRP cycle rather than the

OMRP cycle. This provides further support for the mechanism shown in Scheme 4-4.

As before, KOMRP and kact determine the extent to which the organometallic complex is formed.

Bromoacetonitrile is more rapidly activated by [CuI(Me6tren)]+ in MeCN than chloroacetonitrile (kact

= 1.0 × 105 M-1 s-1 vs. 5.9 × 102 M-1 s-1) so [CuI(Me6tren)(CH2CN)]+ is more distinctly observed with

bromoacetonitrile. The effect of the chelating ligand on the OMRP equilibrium is interesting and

perhaps initially surprising. Again, steric interference between the incoming radical and the terminal

functional groups of the chelating ligand has minimal effect on kd,OMRP. In addition, KOMRP values for

the Me6tren complexes are similar to those for the corresponding tpa complexes. This indicates that

[CuIILR]+ is equally (or more) stable when L = Me6tren which is surprising given the steric bulk of

this ligand. Both of these unexpected phenomena are resolved by recalling that Me6tren lowers the

CuII/I redox potential by comparison with tpa.83 The lower redox potential of [CuII(Me6tren)R]+

indicates that copper(II) is stabilized to a greater extent by this ligand and therefore homolysis of the

Cu−C bond of [CuII(Me6tren)R]+ requires a larger driving force.

83

Conclusion

Cyclic voltammetry has revealed that alkyl radicals produced under ATRP conditions do react with

active copper(I) catalysts to form organometallic copper(II) complexes. Radicals produced from the

activation of bromoacetonitrile and chloroacetonitrile, and to a lesser extent EBriB, react with

[CuI(tpa)]+ and [CuI(Me6tren)]+ in DMSO and acetonitrile to produce the C-bound species

[CuII(L)R]+. This species has been observed directly via spectroelectrochemistry and EPR

spectroscopy. In all cases the experimental behaviour is consistent with a complex which has the

carbanion bound in the auxiliary coordination site.

The catalytic cyclic voltammetry of [CuII(tpa)Br]Br, [CuII(Me6tren)Br]Br and [CuII(Me6tren)Cl]Cl in

the presence of these initiators was successfully simulated using a mechanism in which parallel

halogen atom transfer and radical transfer reactions are coupled to the redox steps. For combinations

of highly active initiators, solvents and ligands, [CuII(L)R]+ is clearly observed if KOMRP is also large.

Furthermore, the rate of radical reaction with copper(I) (kd,OMRP) must be competitive with the

deactivation rate (kdeact) in order to observe the organometallic adduct.

The use of cyclic voltammetry coupled to simulations allows the simultaneous determination of the

key kinetic and thermodynamic parameters for the concerted atom transfer and radical transfer

reactions in a single experiment. Further extension of the protocol to a larger range of solvent-

initiator-catalyst combinations would provide an unprecedented understanding of the key variables

involved in modulating the interplay between these two important equilibria.

84

Experimental

Safety note: perchlorate salts are potentially explosive. Although no problems were experienced with

the perchlorate salts used here they should never be heated in the solid state or scraped from sintered

glass frits.

All solvents were obtained as HPLC-grade from Aldrich or Merck and used without further

purification. Initiators, tpa and Me6tren were likewise obtained from Aldrich and used without further

purification. (Et4N)(ClO4), (Et4N)Br and (Et4N)Cl were prepared as described in the Experimental

section of Chapter 2.

Synthesis

[CuII(Me6tren)(OH2)](ClO4)2

Prepared as described previously (Chapter 3, Experimental).

[Cu(tpa)(OH2)](ClO4)2

[CuII(tpa)(OH2)](ClO4)2 was synthesised according to the method first reported by Karlin et al.231 To

a solution of Cu(ClO4)2.6H2O (0.98 g, 2.65 mmol) in acetone (25 mL) was added dropwise a solution

of tpa (0.77 g, 2.65 mmol) in acetone (25 mL). The mixture was heated to 35 °C and stirred for 2h as

the colour changed from light blue to dark blue/aqua. The solution was removed from heat and cooled

to room temperature before carefully layering with ~75 mL of diethyl ether and cooling at -20 °C

overnight in a spark proof freezer. Blue crystals developed overnight along with a green

microcrystalline solid. Previous studies in the Bernhardt group have shown these green and blue

species to be crystallographically identical compounds in spite of their apparently different colours.

The solid was collected by filtration and washed with ether before drying under vacuum for several

hours. CHN elemental analysis of the solid confirms the above formula which also matches its

published crystal structure.232 Anal. Calcd for C18H20CuN4O9Cl2: C, 37.87; H, 3.53; N, 9.82. Found:

C, 37.80; H, 3.48; N, 9.79.

[CuII(L)(Solv)]2+ and [CuII(L)X]+

The solvato complexes [CuIIL(Solv)]2+ (Solv = MeCN or DMSO) were generated in situ by

dissolution of [Cu(L)(OH2)](ClO4)2 (L = tpa or Me6tren) in neat solvent. Halido (X = Cl- or Br-)

complexes [Cu(tpa)X]+ and [Cu(Me6tren)X]+ were generated in situ by adding 2.0 equivalents of

(Et4N)X (from 0.3 M stock solutions) to solutions of the solvato complexes. In all cases the UV-vis

spectral changes upon dissolution or halide addition are rapid and equilibrium is established within

the mixing timescale.

85

Physical Methods

UV-vis spectroscopy

UV-vis spectra were acquired with an Agilent 8453 UV-Visible spectrophotometer equipped with a

multi-cell holder.

Spectroelectrochemistry

Absorbance measurements were made using an Ocean Optics USB2000 fibre optic

spectrophotometer coupled to a DT-MINI-2-GS miniature deuterium/tungsten/halogen UV-Vis-NIR

light source. A Pine Instruments honeycomb spectroelectrochemical cell kit was used employing a

gold working electrode and a separated gold auxiliary electrode. The reference electrode was non-

aqueous Ag/Ag+ (AgNO3 ca. 0.01 M) dissolved in the solvent of interest - MeCN or DMSO), which

was calibrated with an external ferrocene/ferrocenium couple. A BAS100B/W potentiostat was used

in constant potential electrolysis mode to set the applied potential. All measurements were recorded

under a nitrogen atmosphere (< 20 ppm O2) in a Belle Technology glovebox. The cuvette (1.7 mm

path length) contained a 1.0 mL solution of 6.0 mM copper.

Cyclic voltammetry


working electrode, platinum auxiliary electrode and the same non-aqueous reference electrode.

Ferrocene was used as an internal standard and all potentials are cited versus Fc+/0. The supporting

electrolyte was 0.1 M (Et4N)(ClO4) and all solutions were purged with nitrogen before measurement.

Before each new scan the electrode was polished using an alumina nanoparticle paste, washed and

carefully dried. The voltammetry was also measured with a gold working electrode to check the

suitability of using a gold honeycomb electrode for the spectroelectrochemistry; changing the

electrode had no significant impact on the electrochemistry. All voltammetric simulations were

carried out using the Digisim software package.82 Unknown thermodynamic and kinetic parameters

were determined by simulating the experimental voltammetry across a range of sweep rates (50 –

1000 mV s-1) and concentrations of initiator. The fitting process was carried out by the following

series of steps.

86

Fitting Process

The majority of the kinetic, thermodynamic and physical constants in Scheme 4-4 were determined

for both the Cu/tpa and Cu/Me6tren systems before simulating the catalytic voltammetry. These

parameters were measured or taken from the literature as follows.

Halide binding constant – KII,X

The binding constants were determined by titrating (Et4N)X from a concentrated stock solution into

a UV-vis cell containing 1.0 mM [Cu(L)Solv]2+ and the spectrum was recorded each time after

mixing. Figure 4-18 shows the bathochromic shift which occurs as bromide displaces the initially

coordinated solvent during these titrations.

Figure 4-18 Spectra recorded during the addition of 0.1 mM aliquots of (Et4N)Br to a 1.0 mM solution of A)

[CuII(Me6tren)(OSMe2)](ClO4)2 in DMSO and B) [CuII(tpa)(OSMe2)](ClO4)2

in DMSO. All titrations included 0.1 M

(Et4N)(ClO4).

In a typical titration, 0.1 mM aliquots of halide (from 0.3 M stock solutions of (Et4N)Br or (Et4N)Cl

in the relevant solvent) were sequentially added to 2.0 mL of a 1.0 mM solution of [CuII(L)(Solv)]2+

in the solvent of choice with 0.1 M (Et4N)(ClO4) as supporting electrolyte (to mirror conditions in

CV experiments). Bathochromic shifts were always observed as the halide replaced the stronger field

solvent ligand (Figure 4-18). Even when the stoichiometric endpoint had been reached, several further

aliquots of halide were added to ensure that the final halido complex had been formed. Each titration

for a given ligand/solvent combination was performed in triplicate. The halide binding constants KII,X

were obtained by global analysis (350 – 1100 nm) of the spectra using the program REACTLAB

EQUILIBRIA.188 The average value of the halide complex formation constant from the three

experiments is reported in each case (Table 4.1 and Table 4.2) and the variation between individual

measurements was less than 10 %.

600 800 10000.0

0.2

0.4

0.6

0.8

1.0

Ab

so

rban

ce

Wavelength / nm

600 800 10000.0

0.1

0.2

0.3

0.4

0.5A

bso

rban

ce

Wavelength / nm

A B

87

In DMSO the halide binding constants are sufficiently low that they can be obtained directly from the

UV-vis titrations i.e. more than a stoichiometric equivalent of halide is required to complete the

complex formation. In MeCN the halide binding constants are much higher and only a lower-bound

of 105 M-1 can be estimated from global analysis i.e. the reaction is 100% complete with the addition

of just one equivalent of halide. In these cases the value of KII,X was determined by fitting the

electrochemical halide titration experiments (see following). Table 4.3 (below) collects the relevant

values for this parameter.

Table 4-3 Halide binding constants (M-1) to CuII/L

CuII/Me6tren CuII/tpa

DMSO MeCN DMSO MeCN

KII,Br (M-1) 7.6 × 103 1.2 × 107 1.1 × 104 3.4 × 107

KII,Cl (M-1) - 7.5 × 108 - -

Copper(II) substitution kinetics – kII,X

Anation rate constants for the reaction [CuII(Me6tren)Solv]2+ + X- → [CuII(Me6tren)X]+ + Solv were

measured in Chapter 3 for X = Br- and Cl- and the total reaction was shown to consist of two steps –

outer-sphere association (KOS) followed by substitution (kon). The rate constant for the net reaction,

kIIa,X in Scheme 4-4, is equal to KOS.kon. The values of kIIa,X for the CuII/Me6tren in Table 4.2 were

calculated this way using the relevant values of KOS and kon from Table 3.1.

Chapter 3 also showed that anation reactions with CuII/tpa are too fast to measure so this method for

determining kIIa,X is not applicable to the CuII/tpa system. However, the halide substitution reactions

([CuII(L)Cl]+ + Br → [CuII(L)Br]+ + Cl-), characterised by the rate constant kCl→Br, were measureable

for both tpa and Me6tren.227 Therefore as an estimate of kIIa,Br for the tpa complexes, the measured

value of kCl→Br for tpa was divided by the measured value of kCl→Br for Me6tren to give a scaling

coefficient for the unmeasured rate constant kIIa,X (Eqn. 4.1).

kIIa,X (tpa) = kCl→Br(tpa)

kCl→Br(Me6tren) × kIIa,X(Me6tren) (4.1)

The values of kIIa,X had a negligible effect on the fit of the voltammetry and can be found in Table 4.1

88

Redox potentials – E0Solv & E0

X

The CVs of 10 mL solutions of 1.0 mM [CuII(L)(Solv)]2+ were measured in the absence and presence

of 2.0 mM (Et4N)X and simulated to give ESolv and EX respectively along with the relevant diffusion

coefficients and heterogeneous electron transfer rate coefficients (k0). The data are collected in Table

4.4. Figure 4-19 and Figure 4-20 illustrate some exemplar CVs. In all cases α = 0.5.

Table 4-4 Redox potentials and heterogeneous electron transfer constants for the redox reactions in Scheme 4-4.

CuII/Me6tren CuII/tpa

DMSO MeCN DMSO MeCN

ESolv (mV vs Fc+/0) -730 -475 -610 -425

k0 (cm s-1) 0.01 0.01 0.006 0.01

EBr (mV vs Fc+/0) -810 -705 -695 -650

k0 (cm s-1) 0.01 0.03 0.01 0.02

ECl (mV vs Fc+/0) - -825 - -

k0 (cm s-1) - 0.01 - -

ER – bromoacetonitrile

(mV vs Fc+/0) -1080 -975*

-950 -960

k0 (cm s-1) 0.01 0.005 0.003 0.02

ER – EBriB (mV vs Fc+/0) - - -925 -

k0 (cm s-1) - - 0.03 -

* ER was also = -1045 mV vs. Fc+/0 where chloroacetonitrile was used as the initiator.

Copper(I) binding constant - KI-X

KI,X is determined from the Nernst equation given in Chapter 2 (Equation 2.1) using the measured

values of ESolv, EX and KII-X.

Copper(I) substitution kinetics - kIa,X

Titrating (Et4N)X (X = Br- or Cl-) into an electrochemical cell containing 1.0 mM [CuII(L)(Solv)]2+

was accompanied by the appearance of the [CuII(L)X]2+/+ redox couple at lower potentials (Figure

4-19 and Figure 4-20) The data were simulated using the simple mechanism shown in Scheme 4-6 in

order to extract the unknown rate constant kIa,X for each solvent/halide combination; each of the other

parameters was already measured. In acetonitrile, KII,X was also allowed to float during the fitting

process. The simulated voltammograms are shown in Figure 4-19 B and Figure 4-20 B and the

relevant parameters are collected in Table 4.1 and Table 4.2.

89

Scheme 4-6 Mechanism used to simulate electrochemical halide titrations.


[CuII(tpa)(NCMe)](ClO4)2 in acetonitrile. Sweep rate = 200 mV s-1. I = 0.1 M (Et4N)(ClO4).


[CuII(tpa)(OSMe2)](ClO4)2 in DMSO. Sweep rate = 500 mV s-1. I = 0.1 M (Et4N)(ClO4).

-1000 -800 -600 -400 -200

20 A

E / mV vs. Fc+/0

0.0 mM

0.2 mM

0.4 mM

0.6 mM

0.8 mM

1.0 mM

2.0 mM

-1000 -800 -600 -400 -200

20 A

E / mV vs. Fc+/0

A B

-1000 -800 -600 -400 -200

0.0 mM

0.4 mM

8.0 mM

1.0 mM

2.0 mM

10 A

E / mV vs. Fc+/0

-1000 -800 -600 -400 -200

10 A

E / mV vs. Fc+/0

A B

90

Effect of H2O on voltammetry

In order to assess the potential influence of water on the electrochemistry, an experiment was

conducted using freshly distilled MeCN which was stored under argon with 10% (v/v) molecular

sieves overnight in a dry glovebox before use. Each of the salts (metal and electrolyte) was freshly

recrystallised and dried under high-vacuum for several hours before use. The electrochemical titration

of bromoacetonitrile into a dry solution of 1.0 mM [CuII(tpa)Br]+ in the glovebox using each of these

reagents still exhibited catalytic behaviour at both EBr and ER (Figure 4-21); i.e. hydrolysis is not the

driving force for catalysis. The concentration of water in the MeCN/electrolyte solution for this

experiment was quantified by NMR. 300 µL of the solution was diluted with 300 µL of fresh d3-

acetonitrile. Integrating the water peak against the known concentration of the electrolyte

(Et4N)(ClO4) revealed the [H2O]net. The concentration of H2O in the deuterated solvent was measured

independently for a sample containing only d3-acetonitrile + (Et4N)(ClO4) and this was subtracted

from [H2O]net. The [H2O]cell was < 1 mM.

Figure 4-21 Cyclic voltammetry of the carefully dried 1.0 mM [CuII(tpa)Br]Br in MeCN + [bromoacetonitrile]. Sweep

rate = 100 mV s-1. I = 0.1 M (Et4N)(ClO4).

-1200 -1000 -800 -600 -400 -200

0.0 mM

1.0 mM

3.0 mM

5.0 mM

20 A

E / mV vs. Fc+/0

91

The Mechanism of the Ley-Griffith Oxidation Part 1: The Role of NMO

Introduction

The controlled oxidation of a primary alcohol to an aldehyde is a highly utilised chemical reaction

deployed by synthetic chemists, both in academic and industrial settings,233 to access versatile

chemical building blocks, synthetic intermediates, and/or final targets and products. Three protocols

dominate the synthetic landscape in this regard; these are 1) the Swern oxidation,234 2) use of

hypervalent iodides such as the Dess-Martin periodinane (DMP)235-236 or 2-iodoxybenzoic acid

(IBX),237 and 3) the Ley-Griffith reaction (Figure 5-1).3 Each of these methods is both versatile yet

selective and crucially, each facilitates the transformation under mild conditions without over-

oxidation to the acid.33-34, 238

Figure 5-1 Popular reactions for alcohol oxidation.

The Ley-Griffith oxidation,3, 30, 239-242 which is catalysed by tetra-n-propylammonium perruthenate

(n-Pr4N[RuO4] or TPAP), is synonymous in this regard and has been adopted for the synthesis of an

extensive array of fine chemicals, including pharmaceuticals and natural product targets.243

Surprisingly, however, unlike most other popular oxidation protocols, the mechanism remains

92

unknown having been largely unexamined since the inception of the technique in 1987. In a review

published in 1992, Griffith calls attention to the fact that “the mechanism of oxidation by TPAP is not

clear… and this aspect needs investigation”.239 A more recent volume published by the same author

in 2013 indicates that little progress has been achieved in this regard.243

The original high-valent ruthenium oxidising agent is the extremely reactive ruthenium tetroxide,

RuO4.244-245 The RuO4 complex is stabilised in the hypervalent +8 oxidation state (d0) by four,

strongly σ-donating oxido ligands. It is produced in situ by oxidising tetravalent ruthenium dioxide

hydrate (RuO2.nH2O) or tri-valent ruthenium trichloride with sodium periodate (NaIO4) or bromate

(NaBrO3) in a bi-phasic, aqueous/organic-solvent mixture.246-247 The formed ruthenium tetroxide

partitions into the organic phase (typically carbon tetrachloride)248 while periodate (or bromate) and

their reduced forms remain in the aqueous phase.

The high oxidation state of the ruthenium ion along with its position in the periodic table means that

ruthenium tetroxide is both a vigorous and versatile oxidant capable of oxidising alcohols (including

diols), alkanes, alkenes, amines, amides, ethers, sulphides and aromatics.243 However, its primary

utility is in the oxidation of alcohol groups on carbohydrates.249 It can act as a stoichiometric248 or

catalytic reagent with an excess of sodium periodate or bromate.250 Neither periodate nor bromate

effect the oxidation reaction directly in the absence of ruthenium and are therefore generally assumed

to be responsible for regenerating RuO4 from its reduced form.

The strengths of RuO4 are also its weaknesses. For example, its reactivity as an oxidising agent also

means it reacts violently with most organic solvents. This is noted somewhat alarmingly in the

original paper which assayed the solubility of RuO4:

“A small amount of ruthenium tetroxide (ca. 10 mg.) was tested with the following solvents:

anhydrous ether - small explosion, followed by yellow flame; benzene - vigorous explosion; pyridine

- no explosion, only flame”244

Therefore, despite its toxicity, carbon tetrachloride has been adopted as the organic solvent of

preference when working with ruthenium tetroxide because it is inert towards oxidation. The

reactivity of ruthenium tetroxide also leads to over-oxidation and poor selectivity when multiple

functional groups are present. In addition, RuO4 is volatile so special care must be taken when

handling these reactions; contact with the eyes and mucous membranes has serious effects, albeit

less-so than with its relative, osmium tetroxide.

Lower valent oxido complexes of ruthenium can be prepared as less reactive alternatives. The +7, d1

perruthenate anion ([RuO4]-) can be prepared in basic aqueous solution by oxidising RuCl3 with

93

NaBrO3251 or hypochlorite (NaClO)252 or by making an aqueous solution of RuO4 alkaline.253 In

alkaline aqueous solution, perruthenate equilibrates with the +6, d2 ruthenate anion

([RuO3(OH)2]2-).253 The position of the equilibrium is modulated by pH; strongly basic solutions (pH

12 – 14) favour ruthenate while perruthenate is stabilised by solutions with a pH of 10 – 12.

The structures of ruthenium tetroxide,254 perruthenate255 and ruthenate256 have been elucidated in both

the solid and solution phases by X-ray crystallography as well as spectroscopic techniques.253, 257-259

Figure 5-2 illustrates the near-tetrahedral and trigonal bipyramidal geometries which are observed for

each complex – in each case the geometry is conserved upon dissolution of the solid.

Figure 5-2 Tetrahedral ruthenium tetroxide/perruthenate and trigonal bipyramidal ruthenate.

Like ruthenium tetroxide, perruthenate and ruthenate have both been successfully employed as

stoichiometric oxidants260 or the reactions can be made catalytic by adding excess aqueous sodium

bromate or persulfate (Na2(S2O8)).261 Herein lies the first similitude to the Ley-Griffith protocol

whereby a catalytic amount of perruthenate facilitates the oxidation reaction (Figure 5-1) in the

presence of excess oxidising agent. While the experimental hazards associated with ruthenium

tetroxide are circumvented, these conditions still lead to problems of selectivity and over-oxidation.243

The very high pH requirements for stabilising the transition metal anions are also inimical to many

organic substrates.

In this regard, the synthesis of the tetrabutylammonium salt of perruthenate (TBAP) in 1985

represented a milestone achievement.262 The complex was originally prepared by adding potassium

perruthenate (K[RuO4]) to an aqueous solution of tetrabutylammonium hydroxide out of which the

insoluble TBAP precipitated as a green solid. Soluble in organic solvents, TBAB effected much

milder oxidations than its aqueous analogue; for example, primary alcohols could be converted to

aldehydes without over-oxidation to the acids. The reagent was also more tolerant to reactive

functional groups. However, no compatible organic-soluble co-oxidant was known of so TBAB was

used in stoichiometric quantities and this prevented widespread adoption of the method.

Two years later, the seminal paper of Griffith and Ley appeared where TBAP and its analogue –

tetrapropylammonium perruthenate (TPAP) were used as catalytic oxidants along with an excess of

94

the co-oxidant N-methylmorpholine N-oxide (NMO). The use of NMO in this capacity was first

reported by Sharpless et al. almost a decade earlier in conjunction with low-oxidation state ruthenium

catalysts RuCl2(PPh)3, Ru3(CO)12 or RuCl3·nH2O.263 The Sharpless paper reported that these

complexes could be used in catalytic quantities along with NMO or trimethylamine N-oxide (TMNO)

to oxidise a variety of alcohols in acetone or DMF. The amalgamation of the organic-soluble

perruthenate catalyst along with the co-oxidant from Sharpless led to the now widely utilised Ley-

Griffith method.

Before Griffith and Ley adopted the NMO co-oxidant, Sharpless stressed that its role was unknown

however he did make a series of key observations. For example while NMO and trimethylamine N-

oxide were useful co-oxidants, pyridine N-oxide was not. Furthermore, NMO was successful with

RuCl3 as a catalyst but not RuO2.nH2O.263 The Sharpless procedure, despite inspiring some ruthenium

catalyst development264-267 and protocol modification,268-272 received relatively little synthetic

deployment273-277 meaning the role of NMO was not clarified by 1987 and therefore this ambiguity

was inherited by the Ley-Griffith protocol.

In the 30 years since, only Lee278-279 and Swamy280 have made any serious attempt to elucidate the

mechanism of the reaction. The elegant work of Lee, however, only evaluated the catalyst itself (i.e.

in the absence of the co-oxidant) using one spectroscopic method (UV-vis). The turnover of alcohol

was not monitored, nor the role of the co-oxidant clarified. From the existing literature it is clear that

the solvent, catalyst and co-oxidant each impart unique mechanistic influences on the oxidation but

the origins of these effects remain unknown. Swamy undertook a kinetic analysis of the oxidation,

following the reaction by UV-vis however several of the assumptions of his report are not justified

as will be demonstrated throughout the following Chapters. Elucidating the mechanism of the Ley-

Griffith reaction, beginning with the role of the N-oxide, forms the basis for the work contained within

this and the subsequent Chapters.

95


The role of NMO

Like periodate or bromate, NMO does not directly oxidise alcohols in the absence of perruthenate.3

Its possible functions within the Ley-Griffith reaction include i) regenerating perruthenate or ii)

forming the true catalyst by some initial reaction with perruthenate.280 The direct oxidation of

alcohols by perruthenate in the absence of a co-oxidant refutes the second proposal that NMO is

necessary for formation of the true catalyst.262 Therefore, a tentative mechanism can be proposed as

shown in Scheme 5.2 where NMO regenerates perruthenate from its alcohol-reduced form. Little is

known about the fate of perruthenate during its reaction with an alcohol except for the ubiquitous

observation that in the absence of co-oxidant, a black ruthenium solid precipitates from solution.

Scheme 5-1 Proposed role of NMO in re-oxidising ruthenium.

In order to explore the redox chemistry of perruthenate in non-aqueous solution, cyclic voltammetry

was employed. By using the electrode in place of the alcohol to reduce perruthenate, the formation

of the black solid product (commonly referred to as ‘RuO2’) in the bulk solution is prevented; if it is

formed, it is only within the diffusion layer of the electrode. Cyclic voltammetry (CV) of TPAP

(herein referred to as n-Pr4N[RuO4]) was conducted in MeCN, which is a solvent often used in the

Ley-Griffith reactions.30, 240-241 The CV of freshly prepared solution of n-Pr4N[RuO4] in MeCN is

presented in Figure 5-3.

When the potential of the electrode was swept in the positive direction, a quasi-reversible redox

couple (ΔEp = 80 mV, ipc/ipa = 1) was observed at +370 mV vs. Fc+/0. The identity of the ruthenium

complex in the resting solution (i.e. at the start of the sweep), was determined by measuring the UV-

vis spectrum ‒ shown in Figure 5-4. The [RuO4]- anion has undergone widespread spectroscopic

analysis and in solution it retains the approximately tetrahedral geometry observed in the solid

state.253, 255, 258 Here, the spectrum of nPr4N[RuO4] in MeCN was consistent with that reported for the

96

[RuO4]- anion in alkaline aqueous solution253 and in dichloromethane (DCM),278 which confirms that

[RuO4]- is present at the start of the sweep.

Figure 5-3 Cyclic voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN. Sweep rate: 50 mV s-1. I = 0.1 M (Bu4N)(BF4).

Figure 5-4 Spectrum of n-Pr4N[RuO4] measured in MeCN.

Therefore, the electrochemical response at high potential is the [RuO4]0/- (RuVIII/VII) redox couple

which is consistent with published data.281 The voltammetry of [RuO4]- at lower potentials exhibited

a quasi-reversible redox couple at -1277 mV vs. Fc+/0 (ΔEp = 160 mV, ipa/ipc = 0.8) which was assigned

to a RuVII/VI response and a second irreversible response at -1677 mV; formally a RuVI/V couple. The

single electron stoichiometry of each couple was assigned based on their similar cathodic peak heights

-2500 -2000 -1500 -1000 -500 0 500

20 µA

RuVI / V

RuVII / VI

RuVIII / VII

E / mV vs Fc+/0

300 400 500 600 7000

500

1000

1500

2000

2500

3000

Mola

r A

bso

rptio

n / M

-1 c

m-1

Wavelength / nm

97

relative to the high potential [RuO4]0/- (RuVIII/VII) redox couple (Figure 5-3) and the redox potentials

were determined by simulating the voltammetry with Digisim82 (vide infra).

Quasi-reversibility of the RuVII/VI couple is linked with the appearance of an anodic peak at -850 mV

on the return sweep (Figure 5-5) from oxidation of an electrode-adsorbed RuVI species.73 The absence

of a corresponding cathodic peak at the same potential indicates that [RuO4]- does not adsorb to the

electrode (Scheme 5-2). The voltammetry of the RuVII/VI redox couple was successfully reproduced

with Digisim using this simple model. The magnitude of the peak around -850 mV was variable from

one experiment to the next and dependent on the degree of electrode polishing.

Figure 5-5 Experimental (A) and simulated (B) voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN in the vicinity of the

RuVII/VI couple. Sweep rate = 20 (red), 50 (green), 100 (yellow) and 200 mV s-1 (blue). I = 0.1 M (Bu4N)(BF4). The

potential was initially swept in the negative direction. The response from the adsorbed RuVI species is indicated with an

asterisk.

Scheme 5-2 The mechanism describing the RuVII/VI couple in Figure 5-5 including both diffusing and surface-adsorbed

redox components.

E (mV vs Fc+/0

)

-1800 -1600 -1400 -1200 -1000 -800 -600 -400 -200 0

cu

rre

nt

(A)

-4e-5

-3e-5

-2e-5

-1e-5

0e+0

1e-5

2e-5

3e-5

E (mV vs Fc+/0

)

-1800 -1600 -1400 -1200 -1000 -800 -600 -400 -200 0

cu

rre

nt

(A)

-4e-5

-3e-5

-2e-5

-1e-5

0e+0

1e-5

2e-5

* *

A B

98

The structure of the RuVI complex formed at -1277 mV is not known; RuVI has not been characterized

in organic solvent. Here it is tentatively assigned with the same trigonal bipyramidal geometry that is

observed in aqueous conditions (Figure 5-2) with acetonitrile coordinated in place of one of the axial

hydroxido ligands. More specific information regarding the nature of this complex is provided by

spectroelectrochemical measurements below.

With each of the observed redox couples from Figure 5-3 tentatively assigned, a UV-vis

spectroelectrochemical reduction of [RuO4]- was carried out by applying a constant potential of -2300

mV (vs Fc+/0) and monitoring changes in the spectrum over time (Figure 5-6). The applied potential

lies below both the RuVII/VI and RuVI/V redox potentials so the transition from RuVII to RuVI and then

RuV was followed spectroscopically. Two isosbestic points were observed (~440 nm and ~365 nm)

but during different time periods, suggesting two consecutive reactions. This is apparent from the

single wavelength kinetic profile (A350 nm, Figure 5-6 – inset).

Figure 5-6 UV-vis spectra of 1.0 mM [RuO4]- in MeCN during electrochemical reduction electrolysis at -2300 mV in an

anaerobic glovebox (O2 < 10 ppm). I = 0.1 M (Bu4N)(BF4). Spectra are plotted at 4.0 s intervals over the course of 200 s

and the arrows indicate the changes with time. Inset – absorbance changes at 350 nm during the reduction.

The initial spectrum is of perruthenate and the final spectrum after 200 seconds has a featureless

profile which has undergone a significant baseline shift. Global analysis of the data with the program

REACTLAB KINETICS188 using the mechanism A (RuVII) B (RuVI) C (RuV) allows

deconvolution of the spectral features of all three species (Figure 5-7 – Right). The rate constants

which are generated from the fit of the data are meaningless as their magnitude is dictated by the

applied potential. Applying a more negative overpotential would be coupled to larger rate constants

Wavelength (nm)

350 400 450 500

Ab

so

rba

nce

0.0

0.1

0.2

0.3

Time (Seconds)

0 50 100 150 200Ab

so

rba

nce

at

35

0 n

m

0.15

0.20

350 400 450 500

Wavelength / nm

Ab

so

rban

ce

0.0

0.1

0.2

0.3

99

and vice versa. The objective in fitting the data was to deconvolute the spectral changes and extract

the spectra of the three species generated during the reduction rather than to determine the rate

constants. The speciation profiles and spectra determined using this method are shown in Figure 5-7.

Figure 5-7 Time resolved speciation profiles (Left) and de-convoluted spectra (Right) from the spectroelectrochemical

reduction of [RuO4]-.

Spectrum A in Figure 5-7 (Right) is identical to the measured spectrum of [RuO4]- (Figure 5-4). The

profile of the intermediate ‘B’ exhibits a maximum at 414 nm and a shoulder at 370 nm. This spectrum

is different from genuinely tetrahedral [RuVIO4]2- ([RuO4]

2- doped into BaSO4) measured in the solid

state258 but is qualitatively similar to ruthenate(VI) measured in aqueous alkaline solution (max 460

nm, 1,820 M-1 cm-1 and 385 nm, 1,030).253 X-ray crystallographic studies of the barium and

potassium salts of ruthenate have shown it to be the trigonal bipyramidal complex [RuO3(OH)2]2-.256,

282 Therefore, by analogy, the optical spectrum of species ‘B’ was tentatively assigned as

[RuVIO4(NCMe)]2- (Scheme 5-2 and Scheme 5-3) and the differences in electronic maxima energies

may be due to replacement of OH- by MeCN. The presence of a hydroxido ligand instead of an oxido

ligand also cannot be ruled out (i.e. [RuO3(OH)(NCMe)]-).

The spectrum of the final product (species ‘C’) is featureless and the characteristic peaks from

perruthenate and ruthenate are absent. When the potential was poised at -2300 mV for longer than ca.

200 seconds the baseline absorbance continued to drift upwards due to the formation of a fine

precipitate. This was consistent with the voltammetry at low potential, which was irreversible and

sweep rate dependent (Figure 5-3 and Figure 5-9). Taken together, the voltammetric and

spectroelectrochemical data represent an initial quasi-reversible reduction of [RuO4]- to putative

[RuVIO4(NCMe)]2-, followed by a second reduction to an unstable RuV species that decomposes to

yield a solid product.

0 50 100 150 200 2500.0

0.2

0.4

0.6

0.8

1.0C.

B.

A.

Con

ce

ntr

atio

n /

mM

Time / s

350 400 450 5000

500

1000

1500

2000

2500

C.B.

A.

Mo

lar

Abso

rba

nce /

M-1 c

m-1

Wavelength / nm

100

It appears that disproportionation is the main cause of RuV instability. Bimolecular disproportionation

of RuV to RuVI and RuIV has been reported for oxido complexes of RuV supported by other organic

ligands.283 The instability of RuV appears to be a general phenomenon as only six structurally

characterized RuV complexes appear in the Cambridge Structural Database compared with more than

50 RuVI complexes and more than 250 RuIV complexes. Again, the exact identity of the RuV species

formed at low potential is unknown. Only one structurally relevant RuV complex has been reported284

which was prepared by adding two equivalents of 2-hydroxy-2-ethylbutyric acid to TPAP. The X-ray

structure reveals trigonal bipyramidal geometry with one axially-bound oxido ligand and two

bidentate carboxylate/alkoxido ligands occupying the remaining coordination sites (Figure 5-8).

Figure 5-8 Structure of Bis-2-hydroxy-2-ethylbutyrato(oxo)-ruthenate(V).

By analogy with this complex, the RuV species transiently formed below -1677 mV is assigned as

[RuVO3(NCMe)2]-. Although the solvent is dried before use and n-Pr4N[RuO4] is freshly prepared, a

stoichiometric equivalent of water must be present to facilitate the proton-coupled reduction of

ruthenate(VI) which liberates an oxido ligand as water and is followed by coordination of MeCN.

Upon formation, [RuVO3(NCMe)2]- undergoes rapid disproportionation to generate solid RuO2 and

[RuVIO4(NCMe)]2-; the latter of which is reduced again to [RuVO3(NCMe)2]- until RuO2 is the only

remaining species at low potential. The featureless spectrum of C is synonymous with suspended

RuO2 whose spectrum has been measured in DCM.278 The reduction of perruthenate involving each

of these key reactions is detailed below in Scheme 5-3.

Digisim82 was used to simulate the electrochemical behavior of n-Pr4N[RuO4] at low potentials

according to the mechanism of Scheme 5-3 (excluding the top reaction with NMO). The relevant

chemical and physical parameters obtained from the simulations are collected in Table 5.1 and the

results are shown in Figure 5-9. Including the RuV disproportionation (kdisp) reaction was essential for

reproducing the behavior of the low potential wave (E2).

101

The simulation parameters reveal that RuV rapidly disproportionates upon formation from RuVI (kdisp

3 × 106 M-1 s-1) and reinforce the point that the spectrum of species ‘C’ in Figure 5-7 is that of the

final decomposition product RuO2 (solid).

Scheme 5-3 Mechanism for simulation of the voltammetry of [RuO4]- in the presence of NMO. Reactions involving

surface adsorbed species were omitted for clarity (above).

Table 5-1 Key Physical and Chemical parameters obtained by simulating the sweep-rate dependent voltammetry of n-

Pr4N[RuO4] in MeCN.

Electrochemical Steps Chemical Reactions

E1 (mV vs Fc+/0) -1277 Kcat (kcat/k-cat) 200

α1 0.5 kcat-NMO (M-1s-1) 3 × 103

k0,1 (cm s-1) 5 × 10-2 kcat-TMNO (M-1s-1) 4 × 103

E2 (mV vs Fc+/0) -1677 Kdisp (kdisp/kcomp) 100

α2 0.5 kdisp (M-1 s-1) 3 × 106

k0,2 (cm s-1) 1 × 10-3 Do (cm2 s-1) 4 × 10-5

102

Figure 5-9 Experimental (A) and simulated (B) voltammetry of 0.8 mM n-Pr4N[RuO4] in MeCN. I = 0.1 M (Bu4N)(BF4).

Sweep rate: 50 mV s-1 (red), 100 mV s-1 (green) and 200 mV s-1 (yellow). The potential was initially swept in the negative

direction. Note the sloping baseline at very low potentials in A (< -2600 mV) was due to the onset of a reduction process

which was not modelled.

Figure 5-10 Experimental (A) and simulated (B) CVs of 0.8 mM [RuO4]- in MeCN with NMO concentrations of 0 mM

(red), 10 mM (green), 20 mM (yellow) and 30 mM (blue). Sweep rate = 100 mV s-1. I = 0.1 M (Bu4N)(BF4). Note the

sloping baseline at very low potentials in A (< -2600 mV) was due to the onset of a reduction process which was not

modelled.

Having examined the redox chemistry of perruthenate in the absence of a co-oxidant, NMO was then

titrated into the electrochemical cell (Figure 5-10). Scanning the same potential window revealed that

NMO has no effect on the reduction of [RuO4]- to RuVI demonstrating that neither RuVII or RuVI

undergo redox reactions with it. At low potentials however, the cathodic current is amplified with an

-2500 -2000 -1500 -1000 -500

10 µA

E / mV vs Fc+/0

-2500 -2000 -1500 -1000 -500

10 µA

E / mV vs Fc+/0

E1

E2

A B

E (mV vs Fc+/0

)

-3000 -2500 -2000 -1500 -1000 -500 0

E (mV vs Fc+/0

)

-3000 -2500 -2000 -1500 -1000 -500 0

10 A 10 A

A B

103

increasing concentration of NMO. This indicates that RuV, generated at low potential, is re-oxidized

by NMO to perruthenate so the reduction is catalytic (Scheme 5-3). Presumably this is an inner sphere

reaction which proceeds by coordination of NMO to RuV through its oxygen atom. Heterolytic

scission of the (RuV)O–N bond would yield perruthenate and N-methylmorpholine (NMM).

The capacity for N-oxides to act as ligands towards ruthenium has been reported for complexes of

RuIII.285-287 The intermediate, O-bound N-oxide complexes have been observed spectroscopically

however they are often short-lived, decomposing to the oxido RuV complex and the corresponding

deoxygenated N-containing compound. There is no evidence that the (RuIII)O–N bond cleaves

homolytically; like other oxido-donors, this is a heterolytic, two-electron reaction. Thus, where RuV

is complexed by an N-oxide, perruthenate rather than ruthenate will be regenerated.

The cyclic voltammetry of n-Pr4N[RuO4] with various equivalents of NMO and different sweep rates

was simulated, according to the mechanism in Scheme 5-3, to provide the single set of rate and

equilibrium constants for the reaction between RuV and NMO (Figure 5-10) which are collected in

Table 5.1. The most relevant value obtained from this analysis is the second order rate constant kcat =

3 × 103 M-1 s-1. The key finding is that at sufficiently high concentrations of NMO, RuV is oxidized

faster than it disproportionates and the active [RuO4]- catalyst is recovered. Conversely, in the absence

of NMO, RuV disproportionates to RuO2 and RuVI and the catalyst is lost.

The role of the NMO was clarified in a separate experiment where pyridine N-oxide was added to the

electrochemical cell (Figure 5-11). Pyridine N-oxide had no appreciable effect on the cathodic current

indicating that perruthenate was not regenerated on the voltammetric timescale. This is consistent

with the experimental observation that pyridine N-oxide is not a suitable co-oxidant.263

Figure 5-11 Cyclic voltammetry of 1.0 mM n-Pr4N[RuO4] in MeCN with pyridine N-oxide concentrations of 0.0 mM

(red), 10.0 mM (green), 20.0 mM (yellow) and 30 mM (blue). Sweep rate = 50 mV s-1. I = 0.1 M (Bu4N)(BF4).

-2500 -2000 -1500 -1000 -500

10 µA

E / mV vs Fc+/0

104

Apart from NMO, trimethylamine N-oxide (TMNO – Figure 5-12) is the only other organic-soluble

co-oxidant which has been successfully utilized in conjunction with perruthenate288 so the

electrochemistry was also measured with this N-oxide. Titrating TMNO into the cell elicited a similar

response to NMO with the cathodic current being catalytically amplified at E2 (Figure 5-13). The

concentration- and sweep rate-dependent voltammetry was modeled and the rate constant for the

homogeneous, catalytic chemical reaction was determined to be kcat = 4 × 103 M-1 s-1. Together these

results indicate that NMO and TMNO oxidise RuV at comparable rates leading to their suitability as

co-oxidants however pyridine N-oxide is unreactive towards RuV.

Figure 5-12 Various N-oxides relevant to this work and corresponding N+–O- bond lengths determined from the

corresponding crystal structures.

Figure 5-13 Experimental (A) and simulated (B) CVs of 0.8 mM [RuO4]- in MeCN with TMNO concentrations of 0 mM

(red), 10 mM (green), 20 mM (yellow) and 30 mM (blue). Sweep rate 100 mV s-1. I = 0.1 M (Bu4N)(BF4). Note the

sloping baseline in the experimental data, due to the onset of a low potential reduction below -2500 mV, was not modelled.

-2500 -2000 -1500 -1000 -500

20 µA

E / mV vs Fc+/0

-2500 -2000 -1500 -1000 -500

20 µA

E / mV vs Fc+/0

A B

105

Figure 5-14 Structural valance forms of pyridine N-oxide.

The crystal structures of each of these N-oxides have been reported and indicate that the N–O bond

is significantly shorter in pyridine N-oxide (1.35 Å)289 than either NMO (1.391 Å)290 or TMNO (1.388

Å).291 This is not a purely dative covalent bond in pyridine N-oxide but possess some double bond

character through resonance (Figure 5-14) which explains why it is shorter than in TMNO or NMO

which do not possess the same double bond character.292 It is tempting to correlate the O-donor

activity of each of the aforementioned N-oxides with the N–O bond length and say that pyridine N-

oxide is not an effective oxidant because this bond is much shorter, and therefore more difficult to

break. However, this is an oversimplified view of the bond dissociation enthalpy (∆H0298K) which is

actually the net enthalpy for the reaction shown in Scheme 5-4. The strength of the N–O bond is not

the only factor which contributes to the overall energetics of this reaction.

The stability of the corresponding tertiary amine product must also be considered. For example,

careful study of a series of substituted pyridine N-oxides revealed that, contrary to expectations,

∆H0298K is largely unaffected by electron withdrawing or donating functional groups on the pyridine

ring.293-294 The constancy of ∆H0298K was explained by the equal and opposite effects that the

functional groups have on the enthalplies of formation of the N-oxide and the free amine. Substituents

which increase the N–O bond length also increase the enthalpy of formation of the free amine and

the two effects cancel out one another. The same result was reported with a series of substituted

quinoxaline N-oxides.295-300 Therefore bond length is not a suitable, standalone estimate of the bond

dissociation enthalpy.

Scheme 5-4 Bond dissociation reaction and corresponding enthalpy.

In this regard, it is fortunate that ∆H0298K has been determined for TMNO (260 ± 5 kJ mol-1) and for

pyridine N-oxide (290 ± 10 kJ mol-1); 301-303 unfortunately NMO has not been evaluated.

Nevertheless, these results suggest that pyridine N-oxide is an inefficient oxidizing agent for RuV

because the energetic demands in forming perruthenate and pyridine, through heterolytic cleavage of

the N–O bond, are too high. Conversely TMNO is an effective O-donor because the bond dissociation

106

enthalpy is smaller. NMO, by virtue of its similar chemical structure around this bond, is expected to

have a comparable bond dissociation enthalpy to TMNO making it a suitable oxidant.

Overall this electrochemical method serves as a valuable tool for screening potential co-oxidants for

TPAP-catalyzed oxidations. For reactive N-oxides, the rate constant for the catalytic reaction (kcat)

can be isolated and used as a diagnostic to compare them against NMO. This should allow a more

rational approach to modifying the standard protocol. However, while some co-oxidants may be more

proficient at rescuing RuV, it will be shown later that this is not necessarily beneficial in terms of the

overall expediency of the alcohol oxidation reaction.

NMO and RuO4 electrochemistry

Although sweeping to low potentials unveiled the role of the NMO in the Ley-Griffith reaction,

sweeping to high potentials revealed that it also has the propensity to react with ruthenium tetroxide

(Figure 5-15). Increasing the concentration of NMO in the cell led to amplification of the anodic peak

current (ipa) at high potentials consistent with an alternative ECcat mechanism involving the oxidation

of the NMO ether.

Figure 5-15 Cyclic voltammetry of 1 mM n-Pr4N[RuO4] in MeCN with 0.0 mM (red), 10 mM (green), 20 mM (yellow)

and 30 mM (blue) NMO - A or pyridine N-oxide - B. I = 0.1 M (But4N)(BF4). Sweep rate = 100 mV s-1. Potentials were

swept initially in the forward (positive) direction.

RuO4 is the oxidant of choice for transforming relatively inert cyclic ethers into the corresponding

lactone products.304 The mechanism for this reaction has been investigated by Lee and Van Den Engh

as well as Bakke and Frøhaug260, 305 and several key observations were made. Firstly, the rate law for

the reaction was found to be first order in ruthenium tetroxide and ether. A large, negative enthalpy

of activation was also reported which suggests a highly ordered transition state. Changing the solvent

polarity and introducing electron donating or withdrawing substituents adjacent to the ether had a

0 200 400 600

20 µA

E / mV vs Fc+/0

0 200 400 600

20 µA

E / mV vs Fc+/0

A B

107

marginal effect on the rate of reaction (small Hammet constants). These latter observations indicate

that the reaction does not proceed via a carbocation intermediate. Finally, reacting RuO4 with

perdeuterio-ether (RCD2OCD2R) resulted in a large primary kinetic isotope effect (KIE). Together

these results suggest a mechanism in which RuO4 abstracts hydride from the ether and simultaneously

forms the alkoxide transition state. A second hydride transfer releases the lactone product and RuIV

and completes the reaction (Scheme 5-5). The exact identity of the RuIV product is unknown however

one oxido ligand from RuO4 is donated during the reaction and two hydrogens are abstracted from

the ether. Accordingly, this product is described as H2RuO3.

Scheme 5-5 Proposed mechanism for oxidation of NMO by RuO4 based off the work of Bakke and Frøhaug.

Scheme 5-6 Catalytic oxidation of NMO by RuO4. RuO4 is first produced at the electrode by oxidising perruthenate.

Applying the same principles to the system studied here, RuO4 (produced by electrochemically

oxidising the resting solution of [RuO4]-) reacts with NMO to give the corresponding lactone (N-

methyl-2-oxomorpholine N-oxide - NMOO) and RuIV (Scheme 5-6). Direct re-oxidation of H2RuO3

to RuO4 by the electrode is impossible; a preliminary O-atom transfer reaction between NMO (or

even the lactone product N-methyl-2-oxomorpholine N-oxide) and H2RuO3 must precede this step.

As already highlighted, N-oxides are two electron oxidants meaning RuVI will be produced by O-

108

atom transfer. A second oxidation of RuVI by NMO will not occur due to the insensitivity of this

oxidation state to the N-oxide. The remaining two electrons are provided by the electrode leading to

the catalytic voltammetry observed at high potential (Scheme 5-6). The two protons on H2RuO3 may

be removed either during oxidation to RuVI or during the subsequent oxidation to RuVII/VIII. Water or

N-methylmorpholine (pKa 7.4) likely act as the Brønsted base rather than the N-oxides (pKa ~ 4.5).306

When pyridine N-oxide was added instead of NMO the voltammetry was unaffected (Figure 5-15)

which confirms that oxidation of the ether is the key reaction. This electrochemical method for

producing RuO4 overcomes many of the hazards associated with classical ruthenium tetroxide

chemistry. A stable, inert resting solution of [RuO4]- is used, toxic CCl4 is avoided altogether and

RuO4 is produced at the electrode rather than in bulk. So whilst the electrochemistry of perruthenate

at high potentials is not directly relevant to the Ley-Griffith reaction, it reveals a unique possibility

for conducting oxidation reactions with the much more versatile ruthenium tetroxide. Exploiting the

electrochemically generated RuO4 has unique advantages and would be particularly interesting to

apply to the oxidation of carbohydrates, a reaction for which RuO4 is uniquely suited.

109

Conclusion

Cyclic voltammetry has revealed that NMO competitively regenerates the perruthenate catalyst in the

Ley-Griffith reaction from highly unstable RuV before it undergoes irreversible disproportionation to

RuIV and RuVI (kdisp ~ 106 M-1 s-1). Only a high concentration of the N-oxide ensures that this

regeneration reaction (kcat ~ 103 M-1 s-1) is competitive with disproportionation. The cyclic

voltammetry of perruthenate in the presence of NMO also revealed that RuVI is unreactive towards

the N-oxide.

In considering alternative N-oxide reagents, suitable candidates must have a sufficiently labile N–O

bond. In the case of pyridine N-oxide, the bond dissociation enthalpy is too high so it is not an

effective oxidant. However, trimethylamine N-oxide (TMNO), which has a similar chemical structure

to NMO within the vicinity of the N–O bond, is found to be a suitable substitute. The rates at which

NMO and TMNO oxidise RuV were determined as kcat = 3 × 103 M-1 s-1 and 4 × 103 M-1 s-1 respectively

by simulating the catalytic voltammetry of TPAP in the presence of each at low potentials.

Electrochemistry has therefore resolved which oxidation state of ruthenium is involved in the catalytic

regeneration reaction with NMO and also the rate constant for this reaction and for the

disproportionation reaction.

In addition, an unexpected second reaction was observed at high potentials between NMO and RuO4

whereby the ether moiety of NMO is catalytically oxidised to the corresponding lactone. Here NMO

acts as both the substrate and co-oxidant. While not directly relevant to the Ley-Griffith protocol, this

observation reveals a unique possibility for conducting specific RuO4 reactions, which are

conventionally hazardous, in a much safer manner.

110

Experimental

General

HPLC grade acetonitrile was distilled over calcium hydride onto 3Å molecular sieves (10% v/v) and

stored under Nitrogen before use. Tetrabutylammonium tetrafluoroborate was used as received from

Alfa Aesar. Ruthenium(III) trichloride hydrate was obtained from Precious Metals Online.

Synthesis

NMO

N-Methylmorpholine N-oxide monohydrate (NMO) was obtained by evaporating an aqueous solution

of commercial N-methylmorpholine N-oxide (50 wt%, obtained from Sigma Aldrich) under reduced

pressure.

n-Pr4N[RuO4]

A modified version of the original synthesis reported by Griffith et al. was used.3 Here two round

bottom flasks were connected through a short, tubular glass bridge (Figure 5-16). The first flask was

charged with tetra-n-propylammonium hydroxide solution (1.0 M, 1.25 mL), deionized water (2.5

mL) and sodium hydroxide solution (1.0 M, 10.0 mL). The second flask contained sodium periodate

(375 mg, 1.43 mmol) dissolved in deionized water (10.0 mL) to which was added a ruthenium

trichloride solution (374.0 mg in 2.5 mL of deionized water). The reaction flasks were stoppered, and

the reaction stirred at room temperature for 16 h. During the reaction, volatile RuO4, formed in the

second flask, diffused into the tetrapropylammonium solution and precipitated as a dark green solid

of n-Pr4N[RuO4], which was collected by filtration, washed with water (5 mL) and dried under

vacuum (147 mg, 33%).

M.p. (Dec): 165.2 ◦C ; HRMS: m/z for 102RuO4-, calcd: 165.8846, found: 165.8849; HRMS: m/z for

C12H28N+, calcd: 186.2216, found: 186.2213; UV-Vis: λmax (MeCN) = 316, 385 nm; IR: νmax 2969,

2940, 2879, 1476, 1457, 1390, 1039, 982, 968, 823, 750 cm−1; Anal. Calcd for C12H28NO4Ru: C,

41.01; H, 8.03; N, 3.99. Found: C, 41.12; H, 8.01; N, 3.98.

111

Figure 5-16 n-Pr4N[RuO4] synthesis, post-reaction. Flask A contains n-Pr4N[RuO4] in solution.

Physical methods

Cyclic Voltammetry


working electrode, platinum auxiliary electrode and a non-aqueous Ag/Ag+ reference electrode (ca.

0.01 M AgNO3) in acetonitrile. Ferrocene was used as an internal standard and all potentials are cited

versus Fc+/0. The supporting electrolyte was 0.1 M (But4N)(BF4) and all solutions were purged with

argon before measurement. Between each measurement the electrode was polished using aluminium

nanoparticles, washed and carefully dried – this was critical to avoid electrodeposition of ruthenium

onto the electrode surface.

For experiments in which N-oxide was added to the electrochemical solution, a concentrated solution

of the relevant N-oxide was prepared in acetonitrile and dried over 3Å molecular sieves for 12 hours

to remove waters of hydration before addition.

Spectro-Electrochemistry

The same setup described in Chapter 4 was used. The cuvette (1.7 mm pathlength) contained 2.0 mL

of 1.0 mM n-Pr4N[RuO4] in anhydrous MeCN with 0.1 M (But4N)(BF4) as the electrolyte.

Electrolysis was achieved by poising the potential at -2300 mV vs. Fc+/0 for 200 seconds; spectra were

recorded every two seconds throughout.

A B

112

The Mechanism of the Ley-Griffith Oxidation Part 2: The Role of Perruthenate

Introduction

With the role of NMO clarified by the preceding Chapter, the remaining unknown in the mechanism

of the Ley-Griffith oxidation is the function of the perruthenate catalyst. Two unique mechanisms

have been postulated by Lee, both of which rely on an inner sphere reaction between perruthenate

and alcohol to form a transient alkoxide complex (Scheme 6-1).278-279 In the first mechanism, this

intermediate decomposes by a rate limiting, hydride transfer reaction to release the organic product

along with RuV and water. The second, alternative mechanism involves a second perruthenate ion

which is proposed to adjoin to the alkoxide intermediate forming a ruthenium dimer. In this case the

products of hydride transfer are a RuVI dimer and the ketone/aldehyde.

Scheme 6-1 Mechanisms proposed by Lee for alcohol oxidation by perruthenate in organic solvent.

All of the work done by Lee in formulating these mechanisms followed the reaction between n-

Pr4N[RuO4] and alcohol using UV-vis spectroscopy; but because no co-oxidant was used, the spectral

profile over time only reflected the formation of the black solid ruthenium dioxide. The organic

product was not observable in the UV-vis region analysed and the baseline shifted spectrum of the

dioxide prevented the loss of perruthenate being followed. Nonetheless the rate of RuO2 formation

was monitored with various initial concentrations of catalyst and alcohol in order to formulae the rate

113

law which led to the mechanisms shown in Scheme 6-1. It is important to note that ruthenium dioxide

is not formed directly from the two electron reaction between perruthenate and alcohol; a minimum

of two sequential reactions are required for its formation (e.g. Scheme 6-2). Therefore, describing

the rate of formation of RuO2 solid is not synonymous with describing the rate of the alcohol oxidation

and this should not be used as a direct probe of the principle reaction mechanism. Furthermore, this

method ignores any influence of the co-oxidant and is therefore not truly representative of the Ley-

Griffith reaction.

In this regard the method of Swamy et al. is more appropriate; these authors monitored the oxidation

of aromatic secondary alcohols by TPAP in conjunction with NMO.280 The reaction velocity was

determined spectroscopically by following the appearance of the charge transfer transition from the

aromatic ketone product. The rate law which these authors report is first order in perruthenate and

alcohol and has a fractional order dependence on NMO. A mechanism involving a ternary

Ru/alcohol/NMO complex was invoked in order to account for these observations (Scheme 6-3). This

appears to be a satisfactory analysis of the mechanism until one considers that NMO, NMM and

TPAP also absorb in the region of the ketone and this was not taken into consideration when

formulating the rate law. There are other problems with this report which will be commented on later.

Scheme 6-2 Two-step mechanism produces RuIVO2 from perruthenate and alcohol.

Scheme 6-3 Mechanism proposed by Swamy for alcohol oxidation by perruthenate/NMO in organic solvent.

In each of the three mechanisms highlighted so far there are two further implicit assumptions. The

first is that the alcohol directly coordinates to ruthenium; the evidence for this will be considered

later. The second assumption is that there are no radical products; that is, the aldehyde/ketone is

produced directly from the alcohol by a concerted, two-electron oxidation. In aqueous solution this

114

assumption is not true; instead an organic radical is formed by the reaction between perruthenate and

alcohol.307

The presence of radical species is tested for by examining the product distribution after the oxidation

of cyclobutanol (Scheme 6.4). If the oxidation proceeds by single-electron steps, a radical

intermediate forms which is highly unstable towards cleavage in order to relieve the ring strain. In

this case a second equivalent of perruthenate is required to complete the oxidation and the major

products are the linear species shown on right hand side of Scheme 6.4 Thus, in aqueous conditions

cyclobutanone is only formed in 33% yield and the linear species are the major products; this confirms

the single-electron oxidation mechanism.307 However, in in dichloromethane cyclobutanone is

formed almost quantitatively indicating that the reaction involves a concerted, two-electron oxidation

.278

Scheme 6-4 Oxidation of cyclobutanol in H2O or CH2Cl2 by perruthenate.

Given the limitations of the aforementioned studies, a careful and comprehensive assessment of the

Ley-Griffith reaction mechanism was undertaken utilising not only visible absorption spectroscopy,

but also NMR and EPR to examine the organic and inorganic components of the system. Firstly, the

rate law for the reaction was determined in a manner similar to Swamy280 except that, instead of

measuring only the first fraction of the process (i.e. the initial rates method), the full reaction was

considered.

115


The mechanism of oxidation – bi-phasic kinetics

The oxidation of diphenylmethanol by n-Pr4N[RuO4] and NMO was followed by UV-vis

spectroscopy and revealed a clean conversion of the alcohol to the ketone (Scheme 6-5, Figure 6-1).

Diphenylmethanol is an ideal substrate because it does not absorb above 280 nm while the

corresponding ketone (benzophenone) absorbs at 336 nm (ε = 119.7 M-1 cm-1 – Figure 5.14 A) and

at 280 nm (ɛ = 2907 M-1 cm-1 - recorded by Swamy et. al).280

While the emerging charge transfer transition of benzophenone dominates the profile of Figure 6-1,

the distinctive spectral features of the perruthenate anion at 316 nm, which arise from vibronic

coupling, persist throughout the oxidation. This confirms that perruthenate is preserved throughout.

Subtracting the constant spectrum of perruthenate leads to Figure 6-2 B which, along with the

measured extinction coefficient of benzophenone at 336 nm, reveals that the alcohol is quantitatively

converted to the ketone (Figure 6-2 B – inset).

Scheme 6-5 Oxidation of diphenylmethanol to benzophenone using the Ley-Griffith reagents

Figure 6-1 Time-resolved spectra following the oxidation of 12.5 mM diphenylmethanol by 0.25 mM n-Pr4N[RuO4] and

67 mM NMO in MeCN (303 K). Spectra are displayed at ten minute intervals over the course of seven hours. Inset –

Single wavelength profile at 336 nm.

300 400 5000.0

0.5

1.0

1.5

2.0

0 10000 20000 30000

0.5

1.0

1.5

2.0

Ab

so

rba

nce

/ 3

36

nm

Time / s

Ab

so

rba

nce

/ a

.u.

Wavelength / nm

116

Figure 6-2 A) Spectrum of 1.00 × 10-2 M benzophenone in MeCN. ʎmax ~ 336 nm, ɛ = 119.7 M-1 cm-1. B) Time resolved

spectra from Figure 6-1 with the ‘constant’ spectrum of n-Pr4N[RuO4] subtracted. Inset – Benzophenone concentration

versus time profile.

The concentration versus time profile for the oxidation is unusual in showing two distinct phases. The

reaction initially proceeds slowly through an ‘induction phase’ in which a small amount of product is

formed at a slower rate; then at ~ 5000 seconds the reaction becomes much more rapid and the bulk

of the ketone is formed. This behaviour was not reported by Swamy because only the first few seconds

of the reaction were considered. Global analysis of the data using Specfit139 and REACTLAB

KINETICS188 was attempted but no combination of sequential or parallel reactions provided a

satisfactory fit. Therefore, a rate law analysis was conducted using an initial-rates approach in which

each phase of the reaction was considered independently.

The maximum rate of oxidation (vmax) was determined independently for both phases of the reaction

from the steepest tangent to the concentration curve of benzophenone (Figure 6-3). This was repeated

for a range of concentrations of perruthenate, NMO and alcohol and the reaction order with respect

to each reagent was determined by plotting the log of the slope versus the log of the concentration.

The resulting concentration dependence profiles and log-log plots are shown in Figure 6-4, Figure

6-5 and Figure 6-6. Note: The ‘noise’ observed in the profiles with high concentrations of [RuO4]- or

alcohol indicates that the absorbance at 336 nm has reached the detection limit.

300 350 400 450 500 550

0

50

100

150

200

Mola

r A

bso

rba

nce

/ M

-1 c

m-1

Wavelength / nm

300 350 400 450 500 5500.0

0.5

1.0

1.5

2.0

Abso

rba

nce

/ a

.u.

Wavelength / nm

0 10000 20000 300000

5

10

15

[Be

nzo

ph

en

on

e] / m

M

Time / s

A B

117

Figure 6-3 Maximum rate of oxidation during the induction and catalytic periods is determined by the slope of the steepest

tangent within each region. Experimental conditions are identical to Figure 6-1.

Figure 6-4 [RuO4]--dependent kinetics. Left) concentration profile for benzophenone during a reaction between n-

Pr4N[RuO4], 150 mM NMO and 12 mM diphenylmethanol in MeCN (T = 303 K). Black – 0.100 mM, Red – 0.125 mM,

Green – 0.150 mM, Blue – 0.250 mM, Aqua – 0.350 mM, Cyan – 0.500 mM n-Pr4N[RuO4]. Right) log-log plots for the

first (■ – slope = 1.1) and second (▲ – slope = 1.0) phases.

0 10000 20000 300000

5

10

15

[Be

nzo

ph

en

on

e] / m

M

Time / s

slope = vmax-cat

slope = vmax-ind

0 10000 20000 300000

5

10

15

[Ph

2C

O]

/ m

M

Time / s

-4.0 -3.8 -3.6 -3.4 -3.2-5.5

-5.0

-4.5

-4.0

-3.5

log(v

ma

x)

log[RuO4

-]

Increasing [RuO4

-]

118

Figure 6-5 [Diphenylmethanol]-dependent kinetics. Left) concentration profile for benzophenone during a reaction

between 0.25 mM n-Pr4N[RuO4], 150 mM NMO and diphenylmethanol in MeCN (T = 303 K). Black – 4.0 mM, Red –

8.0 mM, Green – 12.0 mM, Blue – 16.0 mM, Aqua – 24.0 mM, Cyan – 36 mM diphenylmethanol. Right) log-log plots

for the first (■ – slope = 1.1) and second (▲ – slope = 1.2) phases.

Figure 6-6 [NMO]-dependent kinetics. Left) concentration profile for benzophenone during a reaction between 0.25 mM

n-Pr4N[RuO4], 6.0 mM diphenylmethanol and NMO in MeCN (T = 303 K). Black – 5.0 mM, Red – 10.0 mM, Green –

20.0 mM, Blue – 30.0 mM, Aqua – 60.0 mM, Cyan – 120 mM NMO. Note: The different behavior with 5 mM NMO is

due to an excess of alcohol versus NMO. Right) log-log plots for the first (■) and second (▲) phases.

Increasing [Ph2CHOH]

0 10000 20000 300000

5

10

15

20

25

[Ph

2C

O]

/ m

M

Time / s

-2.5 -2.0 -1.5

-5.5

-5.0

-4.5

-4.0

-3.5

-3.0

log(v

ma

x)

log[Ph2COH]

-2.5 -2.0 -1.5 -1.0

-5.5

-5.0

-4.5

-4.0

log

(vm

ax )

log[NMO]

0 10000 200000

2

4

6

[Ph

2C

O]

/ m

M

Time / s

Increasing [NMO]

log[Ph2CHOH]

119

From these data it is apparent that the rate law for the oxidation is identical during both phases of the

reaction, being first order in perruthenate and alcohol and zero order with respect to the co-oxidant

(Equation 6.1a/b). As mentioned earlier, perruthenate is an effective oxidant in the absence of any

co-oxidant which indicates that NMO is not required to access the true catalyst. The zero-order

dependence of vmax on the concentration of NMO confirms this and contradicts the suggestions of

Swamy et al.280

vmax = k[RuO4-][Ph2CHOH] – slow (induction) phase (6.1a)

v’max = ‘k[RuO4-][Ph2CHOH] – fast phase (6.1b)

Equation 1 is consistent with the mechanism proposed by Lee in which the rate determining step

involves a second order reaction between the alcohol and perruthenate, which directly yields the

ketone product along with water and RuV.278 Chapter 5 demonstrated that NMO re-oxidises RuV so

the catalytic cycle would be completed.308 However reaction mechanism does not explain why there

are two distinct phases to the reaction i.e. there are two distinct values for the bimolecular rate

constants in Equations 6.1a and 6.1b. The observed behavior suggests that as the oxidation proceeds

one of the products, which slowly accumulates, accelerates the same reaction i.e. the reaction is

autocatalytic. Thus, where higher concentrations of alcohol and perruthenate are employed, the initial

rate of oxidation is faster (Equation 6.1a) and the induction period is shorter. Unfortunately, the rate

law does not describe which product is the ‘auto-catalyst’.

Adding benzophenone at t0 had no effect on the oxidation kinetics. Furthermore, RuV is short lived

and is rapidly re-oxidised by the large excess of NMO308 meaning neither of these are the product/s

responsible for autocatalysis. Any buildup of RuV would also be discernible spectroscopically as it

has a distinctive peak in the visible region.284 The effect of the remaining product, water, was studied

by adding different concentrations at t0 (Figure 6-7).

The near-zero slope of the log-log plots up to 20 mM H2O indicates that water is not directly involved

in the rate determining step. However, when much higher concentrations were added, the induction

period was truncated and the maximum rate was also slightly increased. One conceivable effect of

water is in disrupting hydrogen bonds between NMO and alcohol (Scheme 6-6). N-oxides are

ubiquitous hydrogen bond acceptors; in fact, NMO is used industrially in the Lyocell process to

disrupt hydrogen-bonded networks between cellulose sheets through its strongly electron-donating

oxygen atom.309-311 Crystal structures reveal that the O--atom of NMO bonds to waters of hydration

in the solid state (NMO.H2O290 and NMO.2.5H2O

312) as well as the hydroxyl protons of alcohols such

as 1,2-cyclohexandiol.313

120

Figure 6-7 [H2O]0-dependent kinetics. Left) concentration profile for benzohpenone during a reaction between 0.25 mM

n-Pr4N[RuO4], 6.0 mM diphenylmethanol and 20 mM NMO in MeCN (T = 303 K). Black – 2.5 mM, Red – 5.0 mM,

Green –10.0 mM, Blue – 20.0 mM, Aqua – 50.0 mM, Cyan – 100.0 mM H2O added at t0. Right) log-log plots for the first

(■) and second (▲) phases.

Scheme 6-6 Possible disruption of Hydrogen bonding by water.

The mechanism of oxidation – 1H NMR

The role of water as a hydrogen bond donor was explored by NMR spectroscopy in collaboration

with Mr. Peter Moore. Instead of diphenylmethanol, 1-octanol was selected as the ideal alcohol for

these measurements because it is much less hygroscopic. The spectra of NMO and 1-octanol are

presented in Figure 6-8 and the peaks match those reported in the literature in d-chloroform (See

Experimental/Table 6.1 and Table 6.2). Acetonitrile is an aprotic solvent so the hydroxyl peak of 1-

octanol appears as a triplet at 2.44 ppm. The adjacent methylene protons are clearly identified as a

triplet of doublets at 3.47 ppm. When a stoichiometric equivalent of NMO is added to 1-octanol, the

hydroxyl proton is shifted downfield to 4.78 ppm, where it appears as a broad singlet. The methylene

protons labeled 2, 4 and 5 along with the methyl protons at position 3’ also undergo a subtle shift (by

less than 0.05 ppm). Together these observations indicate that the hydroxyl proton of 1-octanol

hydrogen bonds to NMO in the manner shown (Figure 6-8).

0 10000 200000

2

4

6

8[P

h2C

O]

/ m

M

Time / s

Increasing [H2O]

-3.0 -2.5 -2.0 -1.5 -1.0-4.0

-3.5

-3.0

-2.5

log

(vm

ax)

log[H2O]

121

Figure 6-8 1H NMR spectra (500 MHz) of NMO (brown), 1-octanol (purple) and 1:1 NMO:1-octanol (green) in d3-

acetonitrile. Note: ax = axial, eq = equatorial.

Table 6-1 NMO 1H and 13C NMR data in d3-acetonitrile.

Position 1H NMR (ppm) 13C NMR (ppm)

1'eq 4.19 (td, J = 12.0, 2.2 Hz, 2H)

1'ax 3.66 (dd, J = 12.0, 3.7 Hz, 2H)

2'eq 3.34 (td, J = 11.5, 3.7 Hz, 2H)

2'ax 2.86 (d, J = 12.0 Hz, 2H)

3’ 3.08 (s, 3H) 61.8

OH 1

2

3-7 8

3’

1’eq

1’

ax

2’eq

2’ax

CH3CN

OH

1

122

Table 6-2 1-octanol 1H and 13C NMR data in d3-acetonitrile.

Position 1H NMR (ppm) 13C NMR (ppm)

OH 2.44 (t, J = 5.4 Hz, 1H)

1 3.47 (td, J = 6.6, 5.4 Hz, 2H) 62.5

2 1.47 (p, J = 6.6 Hz, 2H) 33.6

3 1.33-1.28 (m, 2H of 10H) 26.7

4/5 1.33-1.28 (m, 4H of 10H) 30.1/30.2

6 1.33-1.12 (m, 2H of 10H) 32.6

7 1.33-1.28 (m, 2H of 10H) 23.4

8 1.33-1.28 (t, J = 6.9 Hz, 3H) 14.4

Scheme 6-7 Hydrogen bonding between NMO and 1-octanol disrupted by water.

Hydrogen bonding was confirmed by diffusion-ordered NMR spectroscopy (DOSY) on the 1:1

NMO:1-octanol system. In this experiment diffusion coefficients are generated according to the

number of different molecules or systems present in the sample. Analysis of the 1:1 mixture revealed

a single system for which the diffusion coefficient was measured as 3.20-3.30 × 10-5 cm2 s-1. When

D2O was added to the mixture and the measurement repeated, two separate systems were formed: one

from the NMO signals with coefficients of 7.21-7.32 × 10-5 cm2 s-1 and another from the 1-octanol

signals with coefficients of 7.81-7.91 × 10-5 cm2 s-1. Together these results indicate the formation of

a strong hydrogen bond between NMO and the hydroxyl proton of the alcohol which is disrupted by

water – a product of the Ley-Griffith reaction (Scheme 6-7). Thus, in addition to its role as a co-

oxidant, NMO also modulates the concentration of the free alcohol.

Under standard Ley-Griffith conditions an excess of NMO versus alcohol is used and the reaction is

dry so the hydrogen-bonded species will be favoured. However, as the oxidation proceeds, the

concentration of NMO decreases and water is liberated. Together these lead to a higher concentration

of the ‘free’, non-bonded alcohol which reacts with perruthenate.

123

The behaviour observed in Figure 6-6 and Figure 6-7 can be reconciled with this proposal. Higher

concentrations of NMO shift the position of the equilibrium (Scheme 6-7) towards the adduct,

supressing the formation of water by oxidation. Thus the induction period is prolonged with a greater

excess of NMO. However, when water is added at the beginning of the reaction, this equilibrium is

shifted towards the free alcohol resulting in a shorter induction period. It is important to note that

even when five equivalents of water versus NMO were added, the induction period was still not

completely bypassed (Figure 6-7 – cyan) which suggests that the hydrogen bond equilibrium is not

the only contributor to the auto-catalytic behaviour.

The mechanism of oxidation – EPR

Here the key insight was provided by EPR measurements. EPR provides a lens through which the

inorganic component of the oxidation reaction is exclusively viewed as the organic components do

not contribute to the spectra (no organic radicals involved).

The continuous wave (CW) EPR spectrum of a frozen solution of n-Pr4N[RuO4] in MeCN was

measured at X-band frequency (Figure 6-9 A) and simulated with EPR50F to reveal a slightly

distorted d1 tetrahedral complex (gx,y = 1.937), gz = 1.910). These g-values are comparable with those

measured for other distorted d1 tetraoxoanions: [CrO4]3- gx = 1.84, gy = 1.85, gz = 1.94;314 [MnO4]

2-

gx = 1.98, gy = 1.97, gz = 1.94;315 [ReO4] g┴= 1.72, g═ = 1.85.316 Griffith and Gibson have previously

measured the EPR spectrum of n-Pr4N[RuO4] in frozen DCM reporting a rhombic spectrum (gx =

1.93, gy = 1.98, gz = 2.06).259 However, the g-values for paramagnetic transition metal complexes

whose d-orbitals are less than half-filled should not exceed 2.0. The authors invoke a strange

interaction between the single unpaired spin on perruthenate and a higher energy triplet state to

explain the anomalous value of gz which they report. A simpler explanation may suffice by

recognising that the raw spectrum from which these values were extracted is particularly noisy and

certainly inferior to the spectrum reported here. The effect of solvent on the spectrum was eliminated

by re-measuring a sample of perruthenate in a 50/50 v/v mixture of acetonitrile and toluene whereby

no differences were observed. Therefore, the data measured earlier should be discarded in favour of

the spectrum collected here.

124

Figure 6-9 X-band (νav = 9.7041 GHz) CW EPR spectra measured for n-Pr4N[RuO4] in acetonitrile, with differing

additives. A) n-Pr4N[RuO4], B) n-Pr4N[RuO4] + NMO, C) n-Pr4N[RuO4] + NMO + diphenylmethanol. T = 6 K for all.

Experimental spectra are solid lines and simulated spectra are the broken lines.

Table 6-3 Experimental Spin Hamiltonian Parameters for TPAP in acetonitrile, and after addition of NMO and ROH.

Sample gx gy gz Ru101 Hyperfine

couplings (MHz*)

Linewidths

(MHz)

n-Pr4N[RuO4] 1.937 1.937 1.910 ~90, ~90, — 200, 200, 200

n-Pr4N[RuO4] + NMO 1.938 1.938 1.910 ~110, ~110, — 75, 75, 180

n-Pr4N[RuO4] + NMO + ROH 1.939 1.936 1.918 ~110, ~110, — 75, 75, 180

* 101Ru, I = 5/2, abundance 17.06%; 99Ru, I = 5/2, abundance 12.76%. — indicates the value is undetermined.

When NMO was added to n-Pr4N[RuO4] the spin Hamiltonian parameters shifted slightly (gx,r =

1.938, gz = 1.910) and peaks from hyperfine coupling to 99Ru and 101Ru (I = 5/2) became resolved

(Figure 6-9 B). These observations indicate an interaction between the N-oxide and the ruthenium

ion. However, the spectrum lacks any discernible super hyperfine coupling to the I = 1, 14N nucleus

of NMO which suggests that the interaction is limited to the outer coordination sphere. Inner sphere

coordination of NMO would change the geometry of the complex and should therefore be

accompanied by a more significant effect on the spin Hamiltonian parameters than is observed in

Figure 6-9. The reluctance of perruthenate to expand its coordination number above four is also

known from the consistency of the optical spectrum in a range of solvents. When NMO was added to

[RuO4]- the optical spectrum was likewise unaffected (Appendix 5.1). The changes observed in the

3400 3500 3600 3700 3800

C

B

A

Field / Gauss

125

EPR spectrum must therefore arise from an outer-sphere rather than an inner-sphere interaction

between the two.

Further details of this interaction were provided by high-resolution pulsed EPR spectroscopy carried

out in d3-acetonitrile. This experiment measures 1H coupling to the unpaired d-electron. Figure 6-10

shows X-band HYSCORE spectra for [RuO4]- (A) and [RuO4]

- + NMO (B). Both spectra have the

expected peak at the deuterium Larmor frequency ((2H) = 2.2 MHz) from nearby solvent molecules.

The [RuO4]- + NMO sample (Figure 6-10 B) shows additional hyperfine couplings from 1H nuclei

with A(1H) ≤ 4 MHz. These results can only be explained by the presence of a nearby water molecule

as NMO does not have protons which are capable of H-bonding. As discussed earlier, NMO is a

ubiquitous hydrogen bond-acceptor and is isolated as a monohydrate. Despite drying a stock solution

of NMO to remove this water of hydration before use, spectrum 6.10 B indicates that at least a single

equivalent of water (versus ruthenium) is introduced by adding the N-oxide. Hydrogen bond

formation between an oxido ligand and a proton of the NMO-hydrate (Figure 6-11) explains the

appearance of the coupling observed in Figure 6-10 B as well as the change in the CW EPR upon

addition of NMO. It is also possible to envisage a second water molecule bridging the perruthenate

and NMO via a hydrogen bond to the ether oxygen. Earlier work by Swamy et al. suggested that

NMO directly coordinates [RuO4]- to form the ‘active catalyst’ for the Ley-Griffith reaction.280 The

results here confirm that the two do indeed associate but NMO is not bound as a ligand. Close

association of NMO to perruthenate may be important for the regeneration step after alcohol oxidation

but as mentioned earlier, [RuO4]- is the true oxidant.

Figure 6-10 X-band HYSCORE spectrum at 4K recorded in deuterated acetonitrile of A) n-Pr4N[RuO4] with no additives

(B0 = 352.9 mT, τ = 120 ns, ν = 9.561267 GHz), B) n-Pr4N[RuO4] + NMO (B0 = 358.65 mT, τ = 108 ns, ν = 9.733610

GHz). The anti-diagonal lines mark the deuterium ((2H) 2.3 MHz) and proton ((2H) 15.3 MHz) Larmor frequencies.

Inset: CW EPR spectrum with the positon of the HYSCORE experiment marked with the red line.

A B

126

Figure 6-11 Hydrogen-bonded NMO/[RuO4]- adduct.

When NMO and diphenylmethanol were added to [RuO4]- (i.e. under Ley-Griffith conditions) and

rapidly frozen the CW EPR spectrum was identical to [RuO4]- + NMO (Figure 6-9 C) indicating that

the alcohol does not form a stable complex with perruthenate.

An experiment following the oxidation of diphenylmethanol by concerted EPR/UV-vis was also

conducted. At various intervals throughout the oxidation, a small amount (100 µL) of the reaction

mixture in the cuvette was removed, frozen and the EPR spectrum measured. A lower concentration

of alcohol prolongs both the induction and main catalytic periods such that several measurements

could be taken during both phases; the results are shown in Figure 6-12. Throughout the induction

period, the only notable change to the EPR spectrum is a small decrease in the signal intensity. By

the time the oxidation reaches its maximum velocity (c.a. 8000 s), ~ 25% of the perruthenate signal

is lost. No new peaks appear during this time and the field positions of the existing peaks do not

change. Both RuVI (d2) and RuV (d3) are EPR-active and assume a different geometry from

perruthenate259, 284 so the results in Figure 6-12 are unlikely to be due to the formation of either of

these. Furthermore, the distinctive optical spectra of RuVI and RuV (λmax > 450 nm) are not observed

throughout the reaction. Together these observations indicate that perruthenate is slowly transformed

into an EPR-silent species during the induction period.

Figure 6-12 Left) UV-vis and concentration profile for benzophenone (inset) during a reaction between 0.25 mM n-

Pr4N[RuO4], 6.0 mM diphenylmethanol and 60 mM NMO in MeCN (T = 303 K). Right) frozen X-band spectra (νav =

9.7041 GHz) measured at the intervals indicated in the inset (T = 6K).

300 400 5000.0

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127

Figure 6-13 X-band (νav = 9.6766 GHz) CW EPR spectra measured at 6 K showing the decay of perruthenate EPR signal

over time after addition of substrate alcohol in the absence of co-oxidant NMO.

Formation of a small amount of insoluble ruthenium dioxide solid explains the changes observed in

the EPR. In the absence of a suitable co-oxidant, perruthenate is irreversibly reduced by alcohols,

solvent or adventitious water to solid ruthenium dioxide dihydrate (RuO2.2H2O).253, 262, 278-279, 317-318

Visibly this is discernible as a fine black precipitate which scatters incoming light to give a

featureless, baseline shifted optical spectrum; no EPR of this species has been reported. A separate

experiment was carried out in which a single equivalent of alcohol was added to a solution of n-

Pr4N[RuO4] in the absence of NMO. Every minute an aliquot of the solution was removed, frozen

and the CW EPR spectrum measured (Figure 6-13). The yellow-green solution turned black after two

minutes as RuO2.2H2O formed and this was coupled to a complete loss of the EPR signal.

Therefore, despite the high concentration of NMO which is present during the reaction shown in

Figure 6-12, a small but appreciable amount of RuO2.2H2O must be formed in order to explain the

decrease in the EPR signal intensity. The presence of RuO2.2H2O also accounts for the slight baseline

shift observed in the optical spectrum during the reaction (See Figure 6-12 A – shifting observable at

450 – 550 nm) and the small decrease in the perruthenate peak at 385 nm.

The mechanism of oxidation – RuO2.2H2O catalysis

The effect of RuO2.2H2O was explored by comparing parallel oxidations of diphenylmethanol with

and without added ruthenium dioxide (See Experimental). The results of the tandem reactions are

shown below in Figure 6-14. When RuO2.2H2O is added at the start of the reaction the induction

period is bypassed and the oxidation proceeds smoothly to completion (Figure 6-14 B).

3400 3500 3600 3700

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2 mins

1 min

0 mins

Field / Gauss

128

Figure 6-14 Time-resolved spectra following the oxidation of 6.0 mM diphenylmethanol by 0.25 mM n-Pr4N[RuO4] and

60 mM NMO in MeCN (303 K). A) – no added RuO2.2H2O, B) 16 µL of RuO2.2H2O stock solution added at t0. Spectra

are displayed at five minute intervals. Inset – Single wavelength profile at 336 nm.

The initial spectrum measured for Figure 6-14 B is baseline-shifted because of the insoluble

RuO2.2H2O. However, the change in absorbance at 336 nm (due to ketone formation) is similar in

both A and B. Any difference in the net absorbance change is easily reconciled by considering that a

small amount of the ketone is produced during the manual mixing phase at the start of reaction B

before the first spectrum is measured. Overall, these results indicate that RuO2.2H2O affects the rate

of oxidation but not the yield.

This key experiment reveals that RuO2.2H2O is the product which accelerates the Ley-Griffith

oxidation (Scheme 6-8). Ruthenium dioxide is formed by the concerted, two-electron reduction of

perruthenate (in this case by the alcohol) to the highly unstable RuV which rapidly disproportionates

(kdisp) to give RuO2.2H2O and ruthenate (RuVI).

Scheme 6-8 Catalytic cycle for the Ley-Griffith reaction in acetonitrile.

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129

Ruthenate is also an active oxidant and will undergo a second reaction with alcohol to produce more

of the dioxide.261, 307 Under the Ley-Griffith conditions, an excess of NMO prevents

disproportionation by competitively re-oxidising RuV to the active oxidant – perruthenate (kcat).

However, as the concentration of NMO diminishes throughout the oxidation, disproportionation

becomes competitive with the re-oxidation of RuV. It is ultimately the balance of these two reactions

which determines how quickly the autocatalytic region of the oxidation is accessed.

Higher concentrations of alcohol and perruthenate promote disproportionation (vdisp = 2kdisp[RuV]2)

by increasing the concentration of RuV (1st term of Equation 6.2). Accordingly, the main catalytic

region is accessed sooner with higher concentrations of these reagents (Figure 6-4 and Figure 6-5).

Conversely, higher concentrations of NMO prolong the induction period (Figure 6-6) by keeping

[RuV] low.

d[RuV

]

dt =

d[Ph2CO]

dt ‒ kcat[Ru

V][NMO] ‒ 2kdisp[RuV]

2

d[RuV

]

dt = 𝑘ox[RuO4

-][Ph2COH] ‒ kcat[Ru

V][NMO] ‒ 2kdisp[RuV]

2 (6.2)

The average turnover number of a single perruthenate anion before it irreversibly decomposes to the

dioxide can be estimated from Figure 6-12. At 8000 seconds, 3.0 mM benzophenone has been

produced with a loss of ~ 0.0625 mM catalyst (25%) so the turnover number is roughly 50.

It is important at this point to emphasise that while ruthenium dioxide accelerates the oxidation, it is

not the oxidant. A comprehensive study by Nobuko and Masakatsu demonstrated that RuO2 and

RuO22H2O are ineffective oxidants for un-activated alcohols.319 Unfortunately it is not possible to

conduct a kinetic analysis of the dependence of vmax on variable concentrations of added RuO2.2H2O

because even small amounts obscure the ketone peak. More importantly, it is not meaningful to talk

about the concentration of dioxide because it is a solid in its standard state. Nevertheless, its role in

the oxidation reaction can be postulated.

The rate law determined here during the initial phase of the oxidation of diphenylmethanol indicates

that the reaction proceeds through a transition state involving a single perruthenate anion and a single

molecule of alcohol. During this time, RuO2.2H2O is not present. As solid ruthenium dioxide forms,

the oxidation is accelerated into the main catalytic phase but still involves a single perruthenate anion

and alcohol molecule. The most likely role for ruthenium dioxide is that it acts as a surface catalyst

for the oxidation i.e. the system moves from homogeneous to heterogeneous catalysis.

Because perruthenate and diphenylmethanol both appear in the rate law during the RuO2.2H2O-

catalysed region, both species must be adsorbed onto the surface of the dioxide (i.e. a Langmuir-type

130

mechanism) which, possibly through favourable orientation of the reagents, increases the rate of

oxidation. This mechanism is analogous to the oxidations of alcohols and hydrocarbons by

permanganate which are catalysed by colloidal MnO2.320-321 Collectively, the data presented here,

along with previous reports, indicate that the Ley-Griffith oxidation is effected throughout by a

bimolecular reaction between [RuO4]- and alcohol, and that this reaction is accelerated by solid

ruthenium dioxide which forms when RuV is not rescued by NMO (Scheme 6-8).

The identity of the transition state involving perruthenate and alcohol (in non-aqueous solvent) is not

known. It is generally assumed that alcohols are oxidised by perruthenate through the formation of

an alkoxide or organometallic transition state complex. The classic example of this type of mechanism

is the chromic acid oxidation of isopropyl alcohol for which the proposed isopropyl chromate

intermediate was isolated.322-323 However, no such complex has been observed with perruthenate.

The only evidence for this type of transition state is the report from Lee which compares the rates of

alcohol and ether oxidation by n-Pr4N[RuO4].279 The authors found that THF was oxidised much

more slowly than isopropyl alcohol. Westheimer, who found a similar result using chromic acid,

suggested that the most straightforward explanation of this difference is that the alcohol is oxidised

by formation of the alkoxide intermediate whereas this is not possible for the ether.324 Accordingly,

Lee adopted this explanation. However, an alternative explanation was also offered by Westheimer

in the original paper. This mechanism involved outer-sphere transfer of the α-hydrogen of the alcohol,

as hydride, to one of the oxido ligands of the catalyst and subsequent loss of the hydroxyl proton (e.g.

Scheme 6-9). Further considerations showed this to be unlikely in the case of chromic acid however

no such comparison has been made for n-Pr4N[RuO4] which is deployed under very different

conditions (no acid, non-aqueous solvent).

Scheme 6-9 Outer sphere mechanism for oxidation

The mechanism for the oxidation of methanol by the related ferrate complex was examined by

Yoshizawa et al. using DFT calculations.325 They found that the most likely pathway involves

homolysis of the α-C–H bond to form a carbon radical which immediately coordinates to the metal

131

centre. According to their calculations, the resulting organometallic intermediate is energetically

stable and provides the driving force for this preferred pathway. A second, single electron transfer

completes the oxidation. This route was found to be energetically favourable compared to one

involving the formation of an alkoxide intermediate, a result which has been corroborated by Goddard

and Rappe.326 The comparable mechanism with perruthenate in place of ferrate is given in Scheme

6-10.

Scheme 6-10 Inner sphere mechanism for oxidation involving formation of an organometallic intermediate.

It seems more likely that n-Pr4N[RuO4] reacts with diphenylmethanol via an outer sphere mechanism

with the anion and alcohol forming a hydrogen bonded adduct prior to the rate determining step.

Hydride transfer coupled to hydrogen abstraction would then give the diprotonated complex

[H2RuVO4]- along with the ketone product (Scheme 6-9). Subsequent loss of water from [H2RuVO4]

-

and coordination by acetonitrile would form the unstable [RuV(O)3(NCMe)2] species which is re-

oxidised by NMO or disproportionates to give RuIV and RuVI. While no direct evidence for this

mechanism is provided, it is favoured because it is apparently very difficult to perturb the

coordination sphere of perruthenate; even highly concentrated, sterically unhindered ligands such as

acetonitrile or water do not coordinate. The bulky phenyl groups of diphenylmethanol should make

it even less likely to act as a ligand. The outer-sphere mechanism is also consistent with the EPR

results which show that perruthenate forms an adduct with nearby species which are capable of

hydrogen bonding (such as NMO – or an alcohol). A comparative oxidation of diphenylmethan-d-ol

revealed a small kinetic isotope effect (KIE = 1.7, Appendix 5.2) which is also expected for this

mechanism.

132

The mechanism of oxidation – synthetic relevance

There are hundreds (if not thousands) of reports detailing the synthetic application of the Ley-Griffith

protocol. Throughout the literature however, there is a curious lack of any reference to an induction

period like that observed here. Synthetic-scale oxidations typically utilise much higher concentrations

of the reagents ([ROH],[RuO4]- ~ 10 – 100 mM)327-329 than those examined here which may explain

the disparity; higher concentrations of alcohol and perruthenate being coupled to a short induction

period.

However, there is another explanation. The majority of these studies do not synthesise the catalyst,

instead using commercially available TPAP. When the oxidation of diphenylmethanol was repeated

using freshly purchased sample of the commercial catalyst instead of the synthesised product, no

induction period was observed and the oxidation proceeded smoothly to completion (Figure 6-15).

The baseline shift of the spectrum at t0 indicates that this behaviour is due to a small amount of

ruthenium dioxide present in the commercial catalyst. TPAP is known to be sensitive to moisture and

light and so must be kept sealed under argon in a brown bottle (preferably in the refrigerator). Even

with these precautions, over time it still degrades into a black mass with no oxidative activity. Here

these problems were largely bypassed by using a brand-new bottle of TPAP but nevertheless

ruthenium dioxide was still present. The solid was visibly discernible when identical masses of

synthetic and commercial TPAP were dissolved in MeCN and compared side by side (Appendix 6.1).

In practice, an induction period is undesirable for the sake of expediency and so by serendipitous

fortune, the commercial catalyst is actually preferable to the purified compound which is produced

by the technique described here. Therefore, contrary to conventional wisdom, the pure reagent is less

desirable.

Figure 6-15 Time-resolved spectra following the oxidation of 6.0 mM diphenylmethanol by 0.25 mM n-Pr4N[RuO4]

(commercial) and 60 mM NMO in MeCN (303 K). Spectra are displayed at five minute intervals. Inset – Single

wavelength profile at 336 nm.

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133

Conclusion

A kinetic study of the oxidation of diphenylmethanol using the Ley-Griffith reagents reveals that the

rate determining step involves a single alcohol molecule which is oxidised by a single perruthenate

anion. NMO is not required to form the active oxidant. The transition state was not observed by any

of the spectroscopic methods employed here and either decomposes back into the original reagents

or undergoes electron transfer releasing RuV, water and the organic product. A long-range, outer-

sphere association between perruthenate and NMO is however discernible by HYSCORE EPR.

1H NMR reveals that NMO also hydrogen bonds to the hydroxyl proton of the alcohol preventing its

reaction with the catalyst. This interaction is disrupted by water which is a product of the oxidation.

The major finding is that ruthenium dioxide, formed during the oxidation or by slow degradation of

the stored TPAP reagent, acts as surface catalyst for the reaction. When pure TPAP is utilised the

oxidation proceeds slowly at the beginning until a sufficient concentration of the dioxide is formed

by disproportionation of the RuV product. In this regard, commercial rather than freshly synthesised

TPAP is preferable because it already contains traces of ruthenium dioxide; thus the induction period

is bypassed.

134

Experimental

All common reagents and solvents were purified prior to use according to literature methods.330 All

experiments involving TPAP utilised the reagent produced by the method described in Chapter 5 –

except for the final experiment shown in Figure 6-15 which utilised a commercial sample, obtained

from Sigma Aldrich (97%). NMR spectra were recorded using either a Bruker AV300 (300 MHz, 75

MHz), AV400 (400 MHz,100 MHz) or AV500 (500 MHz, 125 MHz) instrument and all data was

processed using MestReNova software, version 10.0.1. Chemical shifts are given in parts per million

(ppm) and referenced to solvent signals: CD3CN (1.93 ppm for 1H and 1.30 ppm for 13C); CDCl3 (1H:

δ = 7.26 ppm; 13C: δ = 77.16 ppm). Coupling constants (J) are given in Hz. All manipulations

involving NMR (and EPR) solutions were conducted within a dry Belle Technology glovebox using

carefully dried reagents and solvent. In order to obtain a spectrum of NMO in the absence of water,

the hydrated solid was dissolved in d3-acetonitrile and dried over activated 4 Å sieves.

IR spectra were measured on a Perkin Elmer FT-IR spectrometer (Spectrum 2000) with Smiths

detection (DuraSamplerIR II). Melting points were recorded and uncorrected measurements repeated

three times using a Digimelt MPA161 SRS apparatus. GC-MS was recorded using a Shimadzu

GCMS-QP5000 machine using a Restek Rtx® -5MS column and analysed using GCMSsolutions

version 1.20. Microanalyses were performed by the University of Queensland Microanalytical

Service. EPR measurements were conducted at the Centre for Advanced Imaging at the University of

Queensland.

135

Synthesis

Benzophenone

Commercial Benzophenone was recrystallised from ethyl acetate, providing

benzophenone as a white crystalline solid.

Diphenylmethanol

Benzophenone (1.999 g, 11.0 mmol) was dissolved in methanol (25 mL) and

cooled to 0 °C (in an ice water bath). Sodium borohydride (951 mg, 22.0 mmol)

was added and the reaction stirred at 0 °C (in an ice water bath) for five minutes

before being allowed to warm to room temperature and stirred for 20 minutes. The reaction was

cooled to 0 °C (in an ice water bath) and diluted with water (25 mL), the reaction concentrated to half

volume, then extracted with diethyl ether (3 × 30 mL). The organic fractions were combined, washed

with brine, dried over sodium sulfate, passed through a silica plug (5 × 2 cm) and concentrated

providing diphenylmethanol in 97% yield (1.970 g).

1H-NMR (400 MHz): (CDCl3) δH 7.40 (m, 4H), 7.35 (m, 4H), 7.28 (m, 2H), 5.87 (s, 1H), 2.19 (brs,

1H); GC-MS: m/z (% relative intensity, ion): 184 (27, M+), 105 (100, [M−C6H7]+), 77 (63,

[M−C7H7O]+).

Diphenylmethan-d-ol

The same method as above was used except the reaction was performed on half

scale (benzophenone 1.007 g, 5.5 mmol) using NaBD4 (462 mg, 11.0 mmol)

providing diphenylmethan-d-ol in 88% yield (905 mg).

M.p. 68.4–69.2 ◦C; 1H-NMR (400 MHz): (CDCl3) δH 7.42-7.31 (m, 8H, Ar), 7.30-7.25 (m, 2H, Ar),

2.24 (s, OH); 2H-NMR (60 MHz): (CHCl3) δH 5.83 (s); 13C-NMR (100 MHz): (CDCl3) δC 143.9,

128.6, 127.7, 126.7, 76.2 (t, J = 22.0 Hz); IR: νmax 3262, 1491, 1446, 1190, 1045, 1024, 1002, 952,

758, 732, 697, 587 cm−1; GC-MS: m/z (% relative intensity, ion): 185 (47, M+), 184 (16, [M−H]+),

108 (13, [M−C7H5]+), 106 (37, [M−C7H5D]+), 105 (100, [M−C7H6D]+).

136

1-Octanol

1-Octanol was obtained commercially and distilled before use. NMR data

recorded in CD3CN.

1H-NMR (500 MHz): (CD3CN) δH 3.47 (td, J = 6.6, 5.4 Hz, 2H, CH2, H-1), 2.44 (t, J = 5.4 Hz, 1H,

OH), 1.47 (p, J = 6.5 Hz, 2H, CH2, H-2), 1.33-1.28 (m, 10H, CH2, H-3–H-7), 0.88 (t, J = 6.9 Hz, 3H,

CH3, H-8); 13C-NMR (125 MHz): (CD3CN) δC 62.5 (CH2, C-1), 33.6 (CH2, C-2), 32.6 (CH2, C-6),

30.2 (CH2, C-4 or C-5), 30.1 (CH2, C-4 or C-5), 26.7 (CH2, C-3), 23.4 (CH2, C-7), 14.4 (CH2, C-8).

N-Methylmorpholine N-oxide

NMO•H2O was obtained by concentrating under reduced pressure a 50

wt/wt% aqueous solution of NMO. Anhydrous NMO solutions in

CD3CN were obtained by drying the hydrate solution over 4 Å

molecular sieves for 24 hours. NMR data recorded in CD3CN.

1H-NMR (500 MHz): (CD3CN) δH 4.19 (td, J = 12.0, 2.2 Hz, 2H, CH2, H-1eq.), 3.66 (dd, J = 12.1, 3.6

Hz, 2H, CH2, H-1ax.), 3.34 (td, J = 11.5, 3.7 Hz, 2H, CH2, H-2eq.), 2.86 (d, J = 12.0 Hz, 2H, CH2, H-

2ax.), 3.08 (s, 3H, CH3, H-3); 13C-NMR (125 MHz): (CD3CN) δC 66.6 (CH2, C-2), 62.5 (CH2, C-1),

61.8 (CH2, C-3).

1:1 Mixture of 1-octanol : N-methylmorpholine N-oxide

Prepared by mixing 600 µL of 10 mg per mL NMO in CD3CN

and 12.5 µL of 1-octanol in a dry NMR tube under argon.

1H-NMR (500 MHz): (CD3CN) δH 4.74 (brs, 1H, OH), 4.19 (td,

J = 11.6, 2.2 Hz, 2H, CH2, H-1’eq.), 3.66 (dd, J = 12.3, 3.8 Hz, 2H, CH2, H-1’ax.), 3.43 (t, J = 6.6 Hz,

2H, CH2, H-1), 3.34 (td, J = 11.6, 3.7 Hz, 2H, CH2, H-2’eq.), 3.08 (s, 3H, CH3, H-3’), 2.87 (dd, J =

11.9, 2.3 Hz, 2H, CH2, H-2’ax.), 1.45 (p, J = 6.7 Hz, 2H, CH2, H-2), 1.33-1.28 (m, 10H, CH2, H-3–H-

7), 0.88 (t, J = 6.9 Hz, 3H, CH3, H-8); 13C-NMR (125 MHz): (CD3CN) δC 66.5 (CH2, C-2’), 62.5

(CH2, C-1’), 62.2 (CH2, C-1), 61.4 (CH2, C-3), 33.9 (CH2, C-2), 32.6 (CH2, C-6), 30.3 (CH2, C-4 or

C-5), 30.1 (CH2, C-4 or C-5), 26.8 (CH2, C-3), 23.4 (CH2, C-7), 14.4 (CH2, C-8).

137

Kinetics

Kinetic experiments were conducted using an Agilent 8453 diode array spectrophotometer equipped

with a multi-cell holder. The cell holder was coupled to a Bruker thermostat which maintained the

temperature at 303 K throughout. A catalytic concentration of n-Pr4N[RuO4] was always used

(typically 2.5 × 10-4 M) along with an excess of alcohol and NMO – consistent with the conditions

employed in a typical Ley-Griffith experiment. Diphenylmethanol was selected as the model alcohol

for kinetic analysis because 1) it does not absorb above 280 nm, 2) it is not susceptible to over-

oxidation through to an acid which would complicate kinetic analysis, and 3) the benzophenone

product has a distinctive peak at 336 nm (ɛ = 119.7 M-1 cm-1).

In a typical experiment, 2.0 mL of a fresh solution of n-Pr4N[RuO4] was added to the experimental

cuvette and termostatted at 303 K. After ca. 10 minutes, NMO was added from a concentrated stock

solution (0.85 M in MeCN). Finally, alcohol was added from a stock solution (0.1 M in MeCN) to

initiate the reaction.

The maximum rate (vmax) was obtained in each case from a tangent to the steepest portion of the

concentration profile of benzophenone (e.g. Figure 6-3). The slopes of the log-log plots of vmax against

the varied initial concentrations of n-Pr4N[RuO4], NMO and diphenylmethanol revealed the rate law

for the reaction.

NOTE: For alcohol-dependent kinetic measurements the blank cuvette contained 0.25 mM n-

Pr4N[RuO4] such that the constant spectrum of the catalyst was subtracted. This allowed a greater

range of alcohol concentrations to be exploited without flooding the detector. The detector was still

saturated by signal from the ketone for experiments which utilized a higher initial concentration of

alcohol; nevertheless, the conversion profiles could be analyzed to obtain vmax.

Effect of RuO22H2O

A stock solution of ruthenium dioxide was prepared by adding a stoichiometric amount of isopropanol

to a concentrated solution (25.0 mM) of perruthenate in acetonitrile. Within minutes the dark green

solution became brown but the reaction was allowed to continue at 303 K for one hour to ensure

quantitative formation of RuO2.2H2O. Previous work has demonstrated that in organic solvent,

secondary alcohols are oxidised by a stoichiometric quantity of perruthenate to give the ketone (in

this case acetone) and RuO2.2H2O.262 An aliquot of the stock solution was added to one of the two

pre-prepared cuvettes containing n-Pr4N[RuO]4 and NMO and the two reactions were initiated by a

final addition of diphenylmethanol. Acetone from the stock solution of the dioxide has no effect as it

is an inert solvent which can itself be used for these oxidation reactions.262

138

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Appendices

Appendix 2.1

General form of the Nernst equation:

E = E0 – RT

nFln

[O]

[R]

For EBr

E = EBr0 – 59.2ln

[CuII

LBr]

[CuILBr]

= EBr0 – 59.2ln

KII,Br[CuII][Br-]

KI,Br[CuII][Br-]

∴ E = E Br 0 – 59.2ln

KII,Br[CuII

L]

KI,Br[CuIL] 1

For ESol

E = ESolv0 – 59.2ln

[CuII

L]

[CuII

L] 2

Note: Charges are omitted to avoid confusion with the exponents of concentration. Combining 1 and

2 and expanding gives:

ESolv0 – 59.2ln (

[CuIIL]

[CuIL]

) = EBr0 – 59.2 (ln

KII,Br

KI,Br

+ ln[Cu

IIL]

[CuIL]

)

ESolv0 – EBr

0 ‒ 59.2ln ([Cu

IIL]

[CuIL]

) = ‒ 59.2 (lnKII,Br

KI,Br

+ ln[Cu

IIL]

[CuIL]

)

∴ ESolv0 – EBr

0 = -59.2lnKII,Br

KI,Br

(Eqn. 2.1)

154

Appendix 3.1

155

Appendix 3.2

156

Appendix 3.3

157

Appendix 3.4

158

Appendix 3.5

Values obtained for kobs as a function of the variables studied.

Reaction Solv T / ºC P /atm [CuII] / M [Br–];[Cl–];

[Solv'] / M kobs / s–1

[Cu(Me6tren)(Solv)]2+

+ Br-

MeCN 15 2.5×10-3 5.1

5.0×10-3 8.0

1.0×10-2 11.6

1.5×10-2 13.7

2.0×10-2 14.7

16 300 2.0×10–4 4.5×10-2 16

600 15

900 13

1200 13

1500 12

20 2.5×10-3 7.70

5.0×10-3 12.2

1.0×10-2 18.3

1.5×10-2 21.6

2.0×10-2 22.7

25 2.5×10-3 13

5.0×10-3 20

1.0×10-2 29

1.5×10-2 34

2.0×10-2 38

30 2.5×10-3 21

5.0×10-3 32

1.0×10-2 45

1.5×10-2 55

2.0×10-2 60

35 2.5×10-3 32

5.0×10-3 49

1.0×10-2 73

1.5×10-2 83

2.0×10-2 91


+ Cl-

MeCN 15 2.5×10-3 6.67

159

5.0×10-3 9.3

1.0×10-2 12.2

1.5×10-2 13.6

20 2.5×10-3 8.9

5.0×10-3 14.1

1.0×10-2 19.2

1.5×10-2 22.3

25 2.5×10-3 13.8

5.0×10-3 21.9

1.0×10-2 31

1.5×10-2 37

2.0×10-2 39

30 2.5×10-3 20

5.0×10-3 32

1.0×10-2 45

1.5×10-2 55

2.0×10-2 60

35 2.5×10-3 31

5.0×10-3 47

1.0×10-2 66

1.5×10-2 77

2.0×10-2 82

[Cu(Me6tren)Br]+ + Cl- MeCN 15 2.5×10-3 0.278

5.0×10-3 0.516

1.0×10-2 0.997

1.5×10-2 1.49

20 2.5×10-3 0.397

5.0×10-3 0.713

1.0×10-2 1.4

1.5×10-2 2.2

25 2.5×10-3 0.575

5.0×10-3 1.04

1.0×10-2 2.15

1.5×10-2 3.19

30 2.5×10-3 0.846

5.0×10-3 1.55

1.0×10-2 3.27

1.5×10-2 4.77

160

35 2.5×10-3 1.26

5.0×10-3 2.3

1.0×10-2 5.03

1.5×10-2 7.51

[Cu(Me6tren)Cl]+ + Br- MeCN 15 2.5×10-3 0.076

5.0×10-3 0.109

1.0×10-2 0.174

1.5×10-2 0.246

20 2.5×10-3 0.107

5.0×10-3 0.154

1.0×10-2 0.239

1.5×10-2 0.319

25 2.5×10-3 0.157

5.0×10-3 0.229

1.0×10-2 0.357

1.5×10-2 0.484

30 2.5×10-3 0.251

5.0×10-3 0.379

1.0×10-2 0.546

1.5×10-2 0.786

35 2.5×10-3 0.366

5.0×10-3 0.515

1.0×10-2 0.853

1.5×10-2 1.257


+ Br-

DMSO 25 2.0×10-3 18.6

2.5×10-3 23.1

5.0×10-3 37

1.0×10-2 47

1.5×10-2 58

2.0×10-2 73

1:1

DMSO:MMA 25 2.0×10-3

54

2.5×10-3 66

5.0×10-3 78

1.0×10-2 86

1.5×10-2 90

2.0×10-2 102

161

3:1

DMSO:MMA 2.0×10-3

44

2.5×10-3 49

5.0×10-3 65

1.0×10-2 75

1.5×10-2 81

2.0×10-2 95

1:1

DMSO:Sty 25 2.0×10-3

48

2.5×10-3 51

5.0×10-3 69

1.0×10-2 79

1.5×10-2 82

2.0×10-2 98

3:1

DMSO:Sty 2.0×10-3

29

2.5×10-3 38

5.0×10-3 49

1.0×10-2 60

1.5×10-2 65

2.0×10-2 84


+ Br-

DMF 15 2.0×10-3 68

2.5×10-3 76

5.0×10-3 106

1.0×10-2 138

1.5×10-2 151

[Cu(Me6tren)Br]+ + Cl- 15 2.5×10-3 4.94

5.0×10-3 12.1

1.0×10-2 21

1.5×10-2 28

2.0×10-2 36

20 2.5×10-3 7.9

5.0×10-3 18

1.0×10-2 32

1.5×10-2 42

2.0×10-2 54

25 2.5×10-3 12.2

5.0×10-3 28

162

1.0×10-2 50

1.5×10-2 65

2.0×10-2 85

30 2.5×10-3 19

5.0×10-3 43

1.0×10-2 79

1.5×10-2 102

2.0×10-2 133

35 2.5×10-3 29

5.0×10-3 64

1.0×10-2 120

1.5×10-2 162

2.0×10-2 -

[Cu(Me6tren)Cl]+ + Br- 15 2.5×10-3 0.32

5.0×10-3 0.37

1.0×10-2 0.4

1.5×10-2 0.49

2.0×10-2 0.46

20 2.5×10-3 0.44

5.0×10-3 0.46

1.0×10-2 0.54

1.5×10-2 0.62

2.0×10-2 0.7

25 2.5×10-3 0.6

5.0×10-3 0.76

1.0×10-2 0.76

1.5×10-2 0.98

2.0×10-2 1.05

30 2.5×10-3 0.7

5.0×10-3 1.02

1.0×10-2 1.27

1.5×10-2 1.58

2.0×10-2 1.74

35 2.5×10-3 0.98

5.0×10-3 1.34

1.0×10-2 1.95

1.5×10-2 2.62

2.0×10-2 2.83

163


+ Br-

EtOH 15 2.5×10-3 3

5.0×10-3 5

1.0×10-2 7.2

1.5×10-2 9

2.0×10-2 10

20 2.5×10-3 4.3

5.0×10-3 7.2

1.0×10-2 10.9

1.5×10-2 14

2.0×10-2 16.4

25 2.5×10-3 6.7

5.0×10-3 12

1.0×10-2 17.7

1.5×10-2 22.1

2.0×10-2 27

30 2.5×10-3 10.6

5.0×10-3 18.6

1.0×10-2 27.7

1.5×10-2 35.5

2.0×10-2 41

35 2.5×10-3 16.1

5.0×10-3 30.5

1.0×10-2 44.6

1.5×10-2 56

2.0×10-2 67


+ Cl-

EtOH 15 2.5×10-3 4.4

5.0×10-3 6.4

1.0×10-2 10.1

1.5×10-2 12

2.0×10-2 14.4

20 2.5×10-3 6.7

5.0×10-3 10.6

1.0×10-2 15.3

1.5×10-2 19.2

2.0×10-2 22.6

25 2.5×10-3 10.8

164

5.0×10-3 17.2

1.0×10-2 25.9

1.5×10-2 33

2.0×10-2 39

30 2.5×10-3 18.8

5.0×10-3 28

1.0×10-2 44.6

1.5×10-2 55

2.0×10-2 62

35 2.5×10-3 24.4

5.0×10-3 44.5

1.0×10-2 69

1.5×10-2 87

2.0×10-2 103


+ Br-

MeOH 15 2.5×10-3 2.78

5.0×10-3 4.3

1.0×10-2 6.4

1.5×10-2 7.9

2.0×10-2 8.6

20 2.5×10-3 3.9

5.0×10-3 6.4

1.0×10-2 9.5

1.5×10-2 12

2.0×10-2 13.5

25 2.5×10-3 5.8

5.0×10-3 10.2

1.0×10-2 15

1.5×10-2 19

2.0×10-2 22

30 2.5×10-3 8.4

5.0×10-3 16

1.0×10-2 24

1.5×10-2 31

2.0×10-2 34

35 2.5×10-3 14.8

5.0×10-3 24

1.0×10-2 37

165

1.5×10-2 47

2.0×10-2 53


+ Cl-

MeOH 15 2.5×10-3 3.1

5.0×10-3 5

1.0×10-2 7.3

1.5×10-2 8.9

2.0×10-2 10.6

20 2.5×10-3 4.4

5.0×10-3 7.2

1.0×10-2 11.2

1.5×10-2 13.4

2.0×10-2 16

25 2.5×10-3 7.2

5.0×10-3 11.3

1.0×10-2 18

1.5×10-2 23

2.0×10-2 27

30 2.5×10-3 11.6

5.0×10-3 18.8

1.0×10-2 29

1.5×10-2 37

2.0×10-2 43

35 2.5×10-3 18.7

5.0×10-3 31

1.0×10-2 45

1.5×10-2 56

2.0×10-2 70

[Cu(Et6tren)(Solv)]2+ +

Br-

MeCN 15 2.2×10-4 335

[Cu(Et6tren)Br]+ + Cl- MeCN 15 2.0×10-3 92

2.5×10-3 119

5.0×10-3 224

7.5×10-3 330

1.0×10-2 390

[Cu(Et6tren)Cl]+ + Br- MeCN 15 2.0×10-3 9.3

2.5×10-3 9.5

5.0×10-3 12.5

166

1.0×10-2 18.5

2.0×10-2 32.4

3.0×10-2 38.4

20 2.0×10-3 14.8

2.5×10-3 14

5.0×10-3 19.4

1.0×10-2 28.8

2.0×10-2 49.7

3.0×10-2 61.7

25 2.0×10-3 19.9

2.5×10-3 21.1

5.0×10-3 30.2

1.0×10-2 45.1

2.0×10-2 81

3.0×10-2 96

30 2.0×10-3 30.4

2.5×10-3 32.3

5.0×10-3 47.9

1.0×10-2 71.1

2.0×10-2 134

3.0×10-2 165

35 2.0×10-3 49.8

2.5×10-3 50.4

5.0×10-3 74.1

1.0×10-2 121

2.0×10-2 233

3.0×10-2 280

[Cu(tpa)Cl]+ + Br- MeCN 15 2.0×10-3 220

2.5×10-3 235

5.0×10-3 311

7.5×10-3 426

1.0×10-2 487

167

Appendix 4.1

Control experiment for a solution containing only DMSO and various concentrations of BrACN. I =

0.1 M (Et4N)(ClO4) Sweep rate = 500 mV s-1.

Appendix 4.2

2.0 mM CuBr2 in MeCN (0.1 M (Et4N)(ClO4)). Cell pathlength is 1.0 cm.

-1200 -1000 -800 -600 -400

0.0 mM

2.0 mM

4.0 mM

6.0 mM

20 A

E / mV vs. Fc+/0

400 600 8000.0

0.5

1.0

1.5

2.0

Ab

so

rban

ce

Wavelength / nm

168

Appendix 5.1

Time-resolved spectra of 0.25 mM n-Pr4N[RuO4] + 150 mM NMO in MeCN over the course of seven

hours (303 K). Spectra recorded every ten minutes. Spectrum throughout is identical to that of n-

Pr4N[RuO4].

Appendix 5.2

Single wavelength profile for benzophenone during a reaction between 0.25 mM n-Pr4N[RuO4], 150

mM NMO and 12 mM diphenylmethanol (Red) or 12 mM diphenylmethan-d-ol (Black) in MeCN

(T = 303 K). vmax-Ph2CHOH / vmax-Ph2CDOH = 1.7.

200 300 400 500 6000.0

0.5

1.0

1.5

2.0

2.5

3.0

Abso

rba

nce

/ a

.u.

Wavelength / nm

0 10000 20000 300000.0

0.5

1.0

1.5

2.0

2.5

Ab

s.

/ 3

36

nm

Time / s

169

Appendix 6.1

1.8 mg of n-Pr4N[RuO4] in MeCN (10 mL). Left – synthesised; Right ̶ commercial.

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