MECHANISM OF THE OUTER-SPHERE OXIDATION OF AQUEOUS L-CYSTEINE 
AND OF IODIDE IN ACETONITRILE BY A SERIES OF IRON(III) COMPLEXES 
 
Except where reference is made to the work of others, the work described in this 
dissertation is my own or was done in collaboration with my advisory committee. 
This dissertation does not include proprietary or classified information. 
 
 
 
                      ___________________________ 
Xiaoguang Wang 
 
 
 
 
 
 
Certificate of Approval: 
 
___________________________                ___________________________ 
Michael L. McKee                 David M. Stanbury, Chair  
Profesor           Profesor 
Chemistry and Biochemistry       Chemistry and Biochemistry 
 
___________________________                 ___________________________ 
Peter D. Livant           Holly R. Ellis  
Associate Professor          Assistant Professor 
Chemistry and Biochemistry        Chemistry and Biochemistry 
 
 
       ___________________________ 
                         George T. Flowers 
                    Interim Dean 
                    Graduate School 
 
 
 
 
 
 
 
MECHANISM OF THE OUTER-SPHERE OXIDATION OF AQUEOUS L-CYSTEINE 
AND OF IODIDE IN ACETONITRILE BY A SERIES OF IRON(III) COMPLEXES 
 
 
 
 
 
Xiaoguang Wang 
 
 
 
 
 
 
A Dissertation 
 
Submitted to 
 
the Graduate Faculty of 
 
Auburn University 
 
in Partial Fulfillment of the 
 
Requirements for the 
 
Degree of 
 
Doctor of Philosophy 
 
 
 
 
 
 
 
 
Auburn, Alabama 
May 10, 2007 
 
iii
 
 
 
 
 
 
 
MECHANISM OF THE OUTER-SPHERE OXIDATION OF AQUEOUS L-CYSTEINE 
AND OF IODIDE IN ACETONITRILE BY A SERIES OF IRON(III) COMPLEXES  
 
 Xiaoguang Wang 
 
 
Permission is granted to Auburn University to make copies of this dissertation at its 
discretion, upon request of individuals or institutions and at their expense. 
The author reserves all publication rights. 
 
 
 
 
 
 
 
 
 
 
                                                                        
                                          Signature of Author 
 
 
                                      
                                          Date of Graduation 
 
iv
 
 
 
 
 
 
 
VITA 
 
Xiaoguang Wang, son of Jiucheng Wang and Guirong Liu, was born on July 1, 1973, 
in Tonghua, Jilin Province, the People?s Republic of China. He graduated with a 
Bachelor?s degree in Chemical Engineering in 1994 from Jilin Institute of Technology. 
He received his M.E. degree in Chemical Engineering in 1997 from Jilin Institute of 
Technology. From 1997 to 2001, he worked as a research assistant in National 
Synchrotron Radiation Laboratory, P. R. China. In Fall 2001, he entered the Graduate 
School at Auburn University in the Department of Chemistry and Biochemistry and is 
currently pursuing the Ph.D degree. 
 
v
 
 
 
 
 
 
 
DISSERTATION ABSTRACT 
 
MECHANISM OF THE OUTER-SPHERE OXIDATION OF AQUEOUS L-CYSTEINE 
AND OF IODIDE IN ACETONITRILE BY A SERIES OF IRON(III) COMPLEXES  
 
Xiaoguang Wang 
Doctor of Philosophy, May 10, 2007 
(M.E., Jilin Institute of Technology, 1997) 
(B.E., Jilin Institute of Technology, 1994) 
 
238 Typed Pages 
 
Directed by David M. Stanbury 
 
 
Trace copper catalysis was observed in the oxidation of L-cysteine by 
[Fe(bpy)
2
(CN)
2
]
+
 and [Fe(bpy)(CN)
4
]
-
 in anaerobic aqueous solution. The copper 
catalysis was effectively inhibited with the addition of dipicolinate (abbreviated asdipic
2-
) 
for the reduction of [Fe(bpy)
2
(CN)
2
]
+
, and completely suppressed with the addition of 
EDTA (pH ? 10.0) or 1,4,8,11-tetraazacyclotetradecane (abbreviated as cyclam) (pH > 
10.0) for that of [Fe(bpy)(CN)
4
]
-
. 
1
H-NMR and UV-vis spectra show that the products of 
the reaction are L-cystine and the corresponding Fe(II) complexes, with the 
stoichiometric ratio 1:2 for ?[L-cystine]/?[Fe(II)]. The kinetics for the direct oxidation of 
L-cysteine by [Fe(bpy)(CN)
4
]
-
 were studied over the pH range 5.98 ~ 11.9, at ? = 0.10 M 
and 25.0 ?C. The kinetics for the direct oxidation of L-cysteine by [Fe(bpy)
2
(CN)
2
]
+
 were 
 
vi
 
studied over the pH range 3.48 ~ 7.89, at ? = 0.10 M and 25.0 ?C, with the addition of 
0.20 mM N-tert-butyl-?-phenylnitrone (PBN). The mechanism of the reactions was 
proposed. The rate-limiting step is electron transfer to form cysteine radicals and Fe(II). 
Only the thiolate forms of cysteine are reactive. Applying Marcus theory, the 
self-exchange rate constants (k
11
) of 
?
SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
 and 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
, were obtained. The relatively smaller k
11
 
value of 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 is ascribed to negative-negative 
charge repulsion.
 
The redox reactions may occur through one-electron outer-sphere 
transfer reactions.  
The oxidations of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
, [Fe
III
(dmbpy)
2
(CN)
2
]
+
, 
[Fe
III
(CH
3
Cp)
2
]
+
 and [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 at 25.0 
o
C and ? = 0.10 M in acetonitrile 
were studied. Trace copper catalysis was observed for the above four redox reactions. The 
kinetics of the direct oxidation of iodide were studied with the addition of 2,2?-bipyridine 
(bpy). According to 
1
H-NMR and UV-vis spectra, the products of the reaction are I
3
-
 and 
the corresponding Fe(II) complexes, with the stoichiometric ratio of 3:2 between I
-
 and 
Fe(III). Linear Free-Energy Relationships (LFER) were obtained for both log k
1
 and log 
k
2
 vs E
1/2
, implying that the reaction follows a one-electron outer-sphere mechanism. The 
standard-potential (E
1/2
) of I
?
/I
-
 was derived from the kinetic inhibition by 
Fe
II
(bpy)
2
(CN)
2
, with the value of (0.60 ? 0.01) V (vs [Fe(Cp)
2
]
+/0
). However, the 
exchange rate constant of I
?
/I
-
 could not be obtained due to the diffusion-control rate 
constant of k
-1
. 
 
vii 
 
 
 
 
 
 
ACKNOWLEDGMENTS 
 
 
Firstly, I would like to express my deepest appreciation to my advisor, Dr. David 
M. Stanbury, for his academically careful instruction and financial support. I am also 
grateful for his extensive patience and understanding during my Ph.D study. 
I want to thank my committee members for their kind help and good advice in my 
research. I thank their careful correction of my manuscript and instructive guidance in my 
preparation of final defense.  
I also would like to express my gratitude to Dr. German Mills for his generous 
help, especially in my first two years at Auburn University. I want to thank Thomas 
Carrington for his kind help in running Atomic Absorption Spectroscopy. I also 
appreciate that Dr. Leonard N. Bell allowing us to use his AF7 Coulometric Titrator for 
water content analysis. I thank my previous and current lab colleagues, Meiling Hung, 
Anastasia V, Makarycheva-Mikhailova, Thanasekaran Pounraj, and Na Song, for their 
generous help during my study at Auburn. I also appreciate that the Department of 
Chemistry and Biochemistry, Auburn University, and National Science Foundation 
provided funding support for my research. 
Last but not least, I am extremely grateful for my uncle, grandfather, and parents. 
Without their consistent and selfless support, I could not finish my Ph.D study at Auburn 
University.  
 
 viii
 
 
 
 
 
 
Style manual used: Journal of the American Chemical Society 
 
 
Computer software used: Microsoft Office 2001 Microsoft Word, ChemDraw Ultra 7.0.1, 
Datafit version 8.1, Prism 4 version 4.0a, SPECFIT/32 3.0.20. 
 ix
 
 
 
 
 
 
 
TABLE OF CONTENTS 
 
LIST OF TABLES AND FIGURES .................................................................................. xi 
CHAPTER ONE: LITERATURE REVIEW .......................................................................1 
General introduction .....................................................................................................1 
Basic concepts in kinetics ...........................................................................................20 
Electron-transfer mechanism and Marcus theory .......................................................24 
Techniques ..................................................................................................................27 
CHAPTER TWO: DIRECT OXIDATION OF L-CYSTEINE BY  
[Fe
III
(bpy)
2
(CN)
2
]
+
 AND [Fe
III
(bpy)(CN)
4
]
-
......................................................................29 
Introduction..................................................................................................................29 
Experimental section....................................................................................................31 
Reagents and solutions.........................................................................................31 
Preparation of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O ...............................................................33 
Preparation of [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O .......................................................34 
Preparation of K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O ............................................................34 
Preparation of Li[Fe
III
(bpy)(CN)
4
]?2.5H
2
O .........................................................35 
Methods................................................................................................................35 
Results..........................................................................................................................36 
Characterization of the iron complexes ...........................................................36 
Stability of [Fe
III
(bpy)
2
(CN)
2
]
+
 in aqueous solution.......................................44 
 x
 
 
Metal-ion catalysis and scavenger effect ...............................................................44 
The oxidation of EDTA and cyclam by [Fe
III
(bpy)(CN)
4
]
-
....................................53 
Effect of Fe(II) and L-cystine on the rate of the reaction.......................................54 
PBN effect on kinetics ...........................................................................................58 
Product identification and stoichiometry...............................................................64 
Kinetics ..................................................................................................................83 
Discussion....................................................................................................................97 
Conclusion .................................................................................................................114 
CHAPTER THREE: OXIDATION OF IODIDE BY A SERIES OF Fe(III)  
COMPLEXES IN ACETONITRILE...............................................................................116 
Introduction................................................................................................................116 
Experimental section..................................................................................................118 
Reagents and solutions.......................................................................................118 
Preparation of [Fe
II
(dmbpy)
2
(CN)
2
]?3H
2
O ........................................................119 
Preparation of [Fe
III
(dmbpy)
2
(CN)
2
]NO
3
?2H
2
O ................................................120 
Preparation of [Fe
III
(CH
3
Cp)
2
]PF
6
.....................................................................120 
Preparation of [Fe
II
(5-Cl-phen)
2
(CN)
2
]?2H
2
O ...................................................121 
Preparation of [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
?2H
2
O ...........................................121 
Methods..............................................................................................................121 
Results........................................................................................................................123 
Uv-visible spectra ................................................................................................123 
Electrochemistry and NMR spectra.....................................................................123 
Metal-ion catalysis and scavenger effect .............................................................137 
 xi
 
 
Effect of small amount of water on the kinetic study ..........................................140 
Product identification and stoichiometry.............................................................145 
Reaction of iodide with bpy.................................................................................151 
Kinetics ................................................................................................................152 
Discussion..................................................................................................................163 
Conclusion .................................................................................................................173 
REFERENCES ................................................................................................................174 
APPENDICES .................................................................................................................188 
A: Calculation of the equilibrium constants for the oxidation of I
-
 by 
[Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile ...........................................................................188 
B: Calculation of E
o
 (I
?
/I
-
) in acetonitrile...................................................................192 
C: Preliminary study of the oxidation of iodide by various outer-sphere oxidants in 
acetonitrile.................................................................................................................196 
D: Copper-catalyzed oxidation of iodide by [Fe(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O co-solvent  
(CH
3
CN/H
2
O: 99/1, v/v)............................................................................................202 
 
 xii
 
 
 
 
 
 
 
LIST OF TABLES AND FIGURES 
 
Table 2.1 UV-visible and electrochemical characteristics of the iron complexes in 
aqueous solution.................................................................................................................37 
Table 2.2 Effect of M
n+
 on the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
.................45 
Table 2.3 Effect of EDTA on the reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine ..............45 
Table 2.4 Effect of dipic
2-
 on the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 with L-cysteine ............46 
Table 2.5 Effect of dipic
2-
 on the reaction of 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 with 
L-cysteine at various pHs...................................................................................................49 
Table 2.6 Effects of Fe
3+
 and Cu
2+
 on reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine.......52 
Table 2.7 Effect of EDTA/cyclam on reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine in 
anaerobic aqueous solution................................................................................................52 
Table 2.8 The effect of Fe
II
(bpy)
2
(CN)
2
 and PBN on the oxidation of L-cysteine by  
[Fe
III
(bpy)
2
(CN)
2
]
+
.............................................................................................................57 
Table 2.9 The effect of [Fe
II
(bpy)(CN)
4
]
2-
 on the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98 .............................................................................................58 
Table 2.10 Effect of PBN on the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
..............61 
Table 2.11 [L-cysteine] dependence for the oxidation of L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 6.00 ...........................................................................................86 
Table 2.12 pH-dependent kinetic data for the oxidation of L-cysteine  
by [Fe
III
(bpy)
2
(CN)
2
]
+
........................................................................................................87 
 xiii
 
 
Table 2.13 [L-cysteine] dependence for the oxidation of L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 6.92 .............................................................................................92 
Table 2.14 pH-dependent kinetic data for the oxidation of L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 in aqueous solution...............................................................................94 
Table 2.15 Simulation results for the kinetic inhibition by Fe
II
(bpy)
2
(CN)
2
 in  
the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20 ........................................103 
Table 2.16 Mathematical models for the simulation of the oxidation of 1.25 ? 10
-3
 M 
L-cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20..............................................103 
Table 2.17 Simulation results for the kinetic inhibition by [Fe
II
(bpy)(CN)
4
]
2-
 in the  
oxidation of 2.6 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98 ......107 
Table 2.18 Mathematical models for the simulation of the oxidation of 2.6 ? 10
-3
 M 
L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98................................................107 
Table 3.1 UV-visible characteristics of the iron complexes in acetonitrile......................124 
Table 3.2 The half-wave potentials of the iron complexes in acetonitrile .......................135 
Table 3.3 Effect of metal ions on the oxidation of Et
4
NI by [Fe
III
(bpy)
2
(CN)
2
]
+
    
in acetonitrile ...................................................................................................................138 
Table 3.4 The effect of Cu
2+
 and bpy on the oxidation of Et
4
NI by 50 ?M Fe(III) 
complexes in acetonitrile .................................................................................................139 
Table 3.5 Kinetic comparison of NaI and Et
4
NI in the reduction by [Fe
III
(CH
3
Cp)
2
]
+
....144 
Table 3.6 [I
-
] dependence for the oxidation of Et
4
NI by [Fe
III
(bpy)
2
(CN)
2
]
+
..................154 
Table 3.7 [I
-
] dependence for the oxidation of Et
4
NI by [Fe
III
(dmbpy)
2
(CN)
2
]
+
.............157 
Table 3.8 [I
-
] dependence for the oxidation of Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
......................160 
Table 3.9 [I
-
] dependence for the oxidation of Et
4
NI by [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
........161 
 xiv
 
 
Table 3.10 The rate constants for the oxidation of I
-
 by Fe(III) complexes.....................162 
Table 3.11 The equilibrium constants for the oxidation of I
-
 by Fe(III) complexes ........164 
Table 3.12 Simulation results for the inhibition by [Fe
II
(bpy)
2
(CN)
2
] in the  
oxidation of I
-
 by [Fe
III
(bpy)
2
(CN)
2
]
+
..............................................................................169 
Table 3.13 Mathematical models for the simulation of the oxidation of I
-
  
by [Fe
III
(bpy)
2
(CN)
2
]
+
......................................................................................................169 
Table 3.14 The calculated rate constants (k
-1
) for the oxidation of I
-
 by Fe(III)  
complexes in acetonitrile .................................................................................................171 
Table C.1 UV-visible and electrochemical characteristics of the Ru and Fe  
complexes in acetonitrile .................................................................................................197   
Table C.2 Effect of Cu
2+
, bpy, and neocuproine on the oxidation of I
-
 by  
10
-3
 M [Fe
III
(Cp)
2
]
+
 in acetonitrile ...................................................................................201   
Table D.1 Water effect on the rate of the reaction between [Fe
III
(CH
3
Cp)
2
]
+
 and I
-
 in 
CH
3
CN/H
2
O co-solvent, in the presence of 5.0 ?M CuI.................................................205   
Table D.2 Effect of Fe
II
(CH
3
Cp)
2
 on the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O (99/1). .........................................................................................................206   
Table D.3 Effect of CuI on the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, 
v/v) ...................................................................................................................................208 
Table D.4 Comparison of CuI and [Cu(acn)
4
]PF
6
 for the oxidation of I
-
 by 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O(99/1)............................................................................208   
Table D.5 Water effect on the copper-catalyzed oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O co-solvent, in the presence of 5.0 ?M [Cu(acn)
4
]PF
6
................................209   
Table D.6 Effect of [Fe
II
(CH
3
Cp)
2
] on the oxidation of 5.0 mM Et
4
NI by 10 ?M 
 xv
 
 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1)...........................................................................211   
Table D.7 The water content in the products for the oxidation of 5.0 mM I
-
 by 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1)...........................................................................212   
Table D.8 The rate constants for the catalyzed oxidation of 5.0 mM Et
4
NI by 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), in the presence of various concentrations of 
[Cu(NCCH
3
)
4
]
+
 and Fe
II
(CH
3
Cp)
2
...................................................................................215   
Table D.9 Irreproducible kinetic data for the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O (99/1) ..........................................................................................................220   
 xvi
 
 
Figure 1.1 Principle of operation and energy level scheme of the DSSC..........................12 
Figure 2.1 CV of 1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 M NaCF
3
SO
3
................................38 
Figure 2.2 CV of 0.70 mM [Fe
III
(bpy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
................................39 
Figure 2.3 OSWV of 1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 M NaCF
3
SO
3
..........................40 
Figure 2.4 OSWV of 0.70 mM [Fe
III
(bpy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
..........................41 
Figure 2.5 
1
H-NMR spectrum of 0.20 mM Fe
II
(bpy)
2
(CN)
2
 in D
2
O .................................42 
Figure 2.6 
1
H-NMR spectrum of 1.0 mM [Fe
II
(bpy)(CN)
4
]
2-
 in D
2
O................................43 
Figure 2.7 
1
H-NMR spectrum of 12.0 mM sodium dipicolinate in D
2
O...........................47 
Figure 2.8 DSC of sodium dipicolinate dihydrate .............................................................48 
Figure 2.9 A reaction trace of oxidation of 4.7 ? 10
-4
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 7.80, with the addition of 5.0 mM dipic
2-
...................................50 
Figure 2.10 A reaction trace of oxidation of 1.25 ? 10
-3
 M L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20, with the addition of 1.0 mM dipic
2-
/dipic(1:1) ................56 
Figure 2.11 A reaction trace of oxidation of 1.25 ? 10
-3
 M L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20, with 1.0 mM dipic
2-
/dipic(1:1) and 0.10 mM PBN .........60 
Figure 2.12 A reaction trace of oxidation of 2.60 ? 10
-2
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 6.09, with 5.0 mM EDTA. ..........................................................63 
Figure 2.13 UV-vis spectra for the oxidation of 2.5 ? 10
-4
 M L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 6.02, with 1.0 mM dipic
2-
.........................................................65 
Figure 2.14 UV-vis spectra of 10
-4
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 and Fe
II
(bpy)
2
(CN)
2
 in  
aqueous solution.................................................................................................................66 
Figure 2.15 
1
H-NMR spectrum of the product for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
  
with L-cysteine in D
2
O, with 1.0 mM dipic
2-
 and 1.0 mM DSS, at pH 8.78.....................67 
 xvii
 
 
Figure 2.16 Spectrophotometric titration of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
, in the 
absence of PBN..................................................................................................................70 
Figure 2.17 
1
H-NMR spectrum of the product for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
  
with L-cysteine in D
2
O, with 1.0 mM dipic
2-
 and 0.10 mM PBN, at pH 3.10...................72 
Figure 2.18 
1
H-NMR spectrum of 10.0 mM PBN in D
2
O.................................................73 
Figure 2.19 
1
H-NMR spectrum of the product of [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in 
D
2
O, with 1.0 mM dipic
2-
, 0.50 mM PBN and 1.0 mM DSS, at pH 8.70 .........................74 
Figure 2.20 Spectrophotometric titration of L-cysteine with [Fe
III
(bpy)
2
(CN)
2
]
+
, in the 
presence of 1.0 mM PBN...................................................................................................78 
Figure 2.21 
1
H-NMR spectrum of the product of [Fe
III
(bpy)(CN)
4
]
-
 and L-cysteine in D
2
O, 
in the presence of 1.0 mM cyclam and 2.0 mM DSS, at pH 11.1......................................79 
Figure 2.22 Reaction trace of the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 
5.99, in the presence of 0.20 mM PBN and 1.0 mM dipic
2-
. .............................................84 
Figure 2.23 k
obs
 vs [RSH]
tot
 for the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 
6.00, with 1.0 mM dipic
2-
 and 0.20 mM PBN. ..................................................................85 
Figure 2.24 Plot of log(k
obs
/[L-cysteine]
tot
) vs pH for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
, in the presence of 0.20 mM PBN and 1.0 mM dipic
2-
.......................88 
Figure 2.25 A reaction trace of oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 at pH 7.04, 
with 5.0 mM EDTA.. .........................................................................................................91 
Figure 2.26 k
obs
 vs [RSH]
tot
 for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 at pH 
6.92, in the presence of 5.0 mM EDTA.. ...........................................................................93 
Figure 2.27 Plot of log(k
obs
/[L-cysteine]
tot
) vs pH for the oxidation of L-cysteine  
by [Fe
III
(bpy)(CN)
4
]
-
.. ........................................................................................................95 
 xviii
 
 
Figure 3.1 UV-vis spectra of 0.10 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 and Fe
II
(bpy)
2
(CN)
2
 in 
acetonitrile........................................................................................................................125 
Figure 3.2 The OSWV of 0.50 mM Fe
II
(Cp)
2
 and 0.40 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
...........................................................................................................126 
Figure 3.3 The CV of 1.0 mM Fe
II
(bpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
...........................................................................................................127 
Figure 3.4 The OSWV of 1.0 mM Fe
II
(bpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
...........................................................................................................128 
Figure 3.5 The CV of 1.0 mM Fe
II
(dmbpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
 
in 0.10 M Et
4
NBF
4
...........................................................................................................129 
Figure 3.6 The OSWV of 1.0 mM Fe
II
(dmbpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
...........................................................................................................130 
Figure 3.7 The CV of 1.0 mM Fe
II
(CH
3
Cp)
2
 and 1.0 mM Fe(Cp*)
2
 
in 0.10 M Et
4
NBF
4
...........................................................................................................131 
Figure 3.8 The OSWV of 1.0 mM Fe
II
(CH
3
Cp)
2
 and 1.0 mM Fe
II
(Cp*)
2
  
in 0.10 M Et
4
NBF
4
...........................................................................................................132 
Figure 3.9 The CV of 0.50 mM [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 in 0.10 M Et
4
NBF
 4
........133 
Figure 3.10 The OSWV of 0.20 mM Fe
II
(5-Cl-phen)
2
(CN)
2
 and 1.0 mM  
Fe(Cp*)
2
 in 0.10 M Et
4
NBF
4
...........................................................................................134 
Figure 3.11 
1
H-NMR spectrum of recrystallized Fe(CH
3
Cp)
2
 in CD
3
CN.......................136 
Figure 3.12 Kinetic traces for the oxidation of 2.0 ? 10
-3
 M NaI by 5.0 ? 10
-5
 M  
[Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile ....................................................................................141 
Figure 3.13 
1
H-NMR spectrum of product for the reaction of 1.0 mM NaI with 0.10  
 xix
 
 
mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile ............................................................................142 
Figure 3.14 Water effect on the rate constants for the oxidation of 80.0 mM NaI by 
0.25 mM [Fe
III
(CH
3
Cp)
2
]
+
...............................................................................................144 
Figure 3.15 UV-vis spectrum of product for the oxidation of 2.0 mM Et
4
NI by 50 ?M  
[Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, with 5.0 mM bpy ....................................................146 
Figure 3.16 UV-vis spectrum of product for the oxidation of 5.0 mM Et
4
NI by 47 ?M 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 in acetonitrile .............................................................................147 
Figure 3.17 
1
H-NMR spectrum of product for the oxidation of 80.0 mM NaI by 
2.5 mM [Fe
III
(CH
3
Cp)
2
]
+
 in CD
3
CN ..............................................................................148 
Figure 3.18 UV-vis spectra for the oxidation of 5.0 mM NaI by 0.10 mM  
[Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile ......................................................................................149 
Figure 3.19 
1
H-NMR spectrum of the product for the reaction between 3.0 mM Et
4
NI  
and 0.50 mM [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 in CD
3
CN, with 2.0 mM bpy .........................150 
Figure 3.20 UV-vis spectra of product of the reaction of 10.0 mM Et
4
NI with  
50 ?M [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 in acetonitrile ............................................................151 
Figure 3.21 Plot of 1/k
obs
 vs [Fe
II
(bpy)
2
(CN)
2
] for the reaction between 7.5 ? 10
-6
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 and 7.5 ? 10
-5
 M Et
4
NI in acetonitrile ..............................................153 
Figure 3.22 Reaction trace for the oxidation of 2.0 ? 10
-3
 M Et
4
NI by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
...........................................................................................................155 
Figure 3.23 k
obs
/[I
-
] vs [I
-
] for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 with Et
4
NI.................156 
Figure 3.24 Reaction trace for the oxidation of 4.01 ? 10
-2
 M Et
4
NI by 5.0 ? 10
-5
 M 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
......................................................................................................157 
Figure 3.25 Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by 5.0 ? 10
-5
 M  
 xx
 
 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
......................................................................................................158 
Figure 3.26 Reaction traces for the oxidation of 1.50 ? 10
-2
 M Et
4
NI by 2.50 ? 10
-4
 M 
[Fe
III
(CH
3
Cp)
2
]
+
...............................................................................................................159 
Figure 3.27 Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
..........160 
Figure 3.28 Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by  
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
.................................................................................................162 
Figure 3.29 Plot of log k vs E
f
 for the oxidation of I
-
 by Fe(III) complexes....................165 
Figure C.1 CV of 1.0 mM [Ru
II
(hfac)
3
]
-
 and 1.0 mM Fe(Cp*)
2
 in 0.10 M Et
4
NBF
4
. .....198 
Figure C.2 OSWV of 1.0 mM [Ru
II
(hfac)
3
]
-
 and 1.0 mM Fe(Cp*)
2
 in 0.10 M Et
4
NBF
4
199 
Figure C.3 Reaction trace for the oxidation of 8.0 ? 10
-5
 M Et
4
NI by 1.2 ? 10
-5
 M  
Ru
III
(hfac)
3
.......................................................................................................................200 
Figure D.1 Plot of k
obs
 vs [Fe(II)] for the oxidation of Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O (99/1, v/v)...................................................................................................210 
Figure D.2 
1
H-NMR spectrum of product for the oxidation of 8.0 mM Et
4
NI by 2.0 mM 
[Fe
III
(CH
3
Cp)
2
]
+
 in CD
3
CN..............................................................................................211 
Figure D.3 Kinetic trace of the oxidation of 5.0 mM Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
............214 
Figure D.4 Plot of k
obs
 vs [Cu(acn)
4
+
] for the oxidation of 5.0 mM Et
4
NI by 10 ?M 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v) ...................................................................216 
Figure D.5 Plot of 1/(k
obs
-0.0019) vs 1/[Cu(I)] for the oxidation of 5.0 mM Et
4
NI  
by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1)........................................................ 217 
Figure D.6 Plot of (k
obs
-0.0019)
-1
 vs [Fe(II)] for the oxidation of 5.0 mM Et
4
NI by 
10.0 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v)................................................... 217 
 1
 
 
 
 
 
CHAPTER ONE 
LITERATURE REVIEW 
 
GENERAL INTRODUCTION 
Sulfur, one of the chalcogen elements, is the source for the biosynthesis of sulfur-
containing amino acids that play a wide range of roles in essential biological processes.
1-6
 
Homocysteine, glutathione, and cysteine, the most important ?-aminothiols in biological 
chemistry, will not easily break their C-S bonds, and they can be regenerated for further 
use after oxidation to the disulfide.
7,8
 Under physiological conditions, ?-aminothiols can 
undergo a range of reactions, such as nucleophilic substitution, electron transfer, proton 
transfer, hydrogen atom transfer, hydride transfer, and oxygen atom transfer, etc.
2,9
 
Glutathione (GSH; a tripeptide formed by the amino acids glycine, cysteine, and 
glutamic acid), an abundant nonprotein extracellular and intracellular thiol present in 
animals, plants, and many bacteria, plays an important role in physiologically relevant 
antitoxic function; it participates in the maintenance of cellular redox potentials and 
protein thiol-disulfide ratios. The conversion of glutathione into its disulfide (GSSG) is 
readily reversible. There is suggestive evidence that the glutathione/glutathione disulfide 
redox couple modulates the regulation of DNA repair, gene expression and signal 
transduction in the immune system.
1,10
 Generally, the in vivo thiol/disulfide balance is 
tightly controlled, and the variations of its redox state represent proliferation, 
 2
differentiation, or apoptosis of the cells. The GSH/GSSG redox couple has the lowest 
intracellular reduction potential during the cells? proliferation, at -260 ? -230 mV; while 
it has the highest value during apoptosis, with values ranging between -170 and -150 
mV.
1
 The reductive capability of GSH/GSSR may be related to the concentration of GSH 
in the cell: there is the highest concentration of GSH during cell proliferation, while there 
is the lowest concentration of GSH during apoptosis.
1,11
 In differentiation and apoptosis 
of cells, the regeneration of GSH by cysteine may be suppressed, or the oxidation of GSH 
facilitated by free radicals. Studying the oxidation of GSH may help understand the role 
of enzymes in biological processes, or may help retard the apoptosis of cells. 
Cysteine, one of the simplest biological thiols, exists widely inside and outside of 
cells. The sulfur in cysteine is fully reduced (the oxidation state of -2 is assigned), and its 
thiol group can undergo a range of reactions to form a variety of redox-active sulfur-
containing compounds: thiyl radicals, sulfenic, sulfinic and sulfonic acids, sulfenyl-
amides, thiosulfinates and polysulfides.
5
 All reactions related to the formation of the 
above redox-active sulfur-containing compounds are shown in Scheme 1-1. Although its 
concentration is much less than that of glutathione,
4,12,13
 cysteine also fulfills various 
functions in proteins, including maintenance of the intracellular redox equilibrium, 
disulfide formation, metal-binding, electron donation, hydrolysis, biologic redox 
signaling, and redox-catalysis. All of the above functions are directly or indirectly related 
to the reductive capability of intracellular cysteine. The oxidation of cysteine can be 
reversible or irreversible, depending on the sulfur oxidation states involved. In many 
cases, oxidation is a reversible process resulting in disulfide formation, which is 
considered to be a common mechanism of protein stabilization and folding.  
 3
Scheme 1-1. Redox reactions of L-cysteine. (RSH: L-cysteine; GSSG: oxidized 
dimeric form of glutathione) 
 
RSH + GSSG RSSG + GSH                 thiol/disulfide exchange          (1-1)
2 RSH RSSR + 2 H
+ 
+ 2 e
-
 two-electron transfer            (1-2)
RSH + H
+ 
+ e
-
single-electron transfer        (1-3)
RSH RS
RS
+ H hydrogen-atom transfer       (1-4)
RSH + HNO
2
RSNO + H
2
O nitrosation (1-5)
RSH + H
2
O
2
RSOH + H
2
O two-electron transfer             (1-6)
RSOH + R'SH RSSR' + H
2
O nucleophilic substitution          (1-7)
RSOH + R'(H)NCR" R"CN(R')SR
O
O
nucleophilic substitution    (1-8)
RSOH + H
2
O
2
RSOH
O
+ H
2
O two-electron transfer             (1-9)
RSOH + R'SOH
RSSR'
O
+ H
2
O
two-electron transfer            (1-10)
RSSR'
O
+ H
2
ORSOH
O
+ R'SH
ATP
nucleophilic substitution         (1-11)
RSOH
O
RSOH
O
strong Ox
O
two-electron transfer           (1-13)
RSSR'
O
RSSR' + H
2
O
2
+ H
2
O
two-electron transfer           (1-12)
RSSR'
O
DecompositionRSS
x
SR'
(1-14)
+ H
2
O
 
 
 4
Cysteine also participates in the maintenance of the intracellular redox balance.
12,14
 
Generally, thiol/disulfide exchange reactions of cysteine/cystine involve balancing the 
intracellular redox potential. Thiol/disulfide exchange reactions, nucleophilic 
substitutions of a thiol or thiolate on a disulfide bond, lead to the formal oxidation of the 
nucleophile and reduction of the leaving group, as shown in Equation 1-1, in which there 
is no implication of involvement of electron transfer. However, thiol/disulfide exchange 
reactions can be accompanied by two-electron or two sequential one-electron transfers. 
For example, regeneration of GSH from GSSG is catalyzed by human glutathione 
reductase, which involves direct two-electron transfer from FADH
2
 to Cys 58-S-S-Cys 
63;
15
 the reduction of oxidized thioredoxins in plants and bacteria that is catalyzed by 
ferredoxin-thioredoxin reductase (FTR) occurs by two sequential one-electron transfers.
16
 
Many human diseases related to oxidative stress, such as rheumatoid arthritis, 
autoinflammatory and neurodegenerative disease and cancer, are associated with the 
imbalance of the intracellular redox equilibrium. It is assumed that disturbances in the 
intracellular redox equilibrium are due to a significant increase in the concentration of 
oxidizing species, which may oxidize cysteine to form a wide range of chemically 
distinct reactive sulfur species. Understanding the mechanism of the oxidation of cysteine 
may help provide a clue in the prevention and cure of potentially fatal diseases, such as 
leukemia and hepatocirrhosis. 
Cysteine-based radicals play important roles in biological systems as well.
17
 They 
are formed either by a long-range one-electron transfer from the thiol (Equation 1-3) or 
by short-range hydrogen atom abstraction from the thiol (Equation 1-4). Cysteine-based 
radicals are involved in protecting biological systems from ionizing radiation and 
 5
participating directly in enzyme-catalyzed reactions. For instance, it was found that 
sulfur-centered radicals take part in cycles catalyzed by pyuvate formate lyase (PFL).
18
 
Recent results have shown that redox-active sulfur, including sulfenic acids 
(RSOH), sulfinic acids [RS(O)OH], sulfonic acids [RS(O)
2
OH], sulfenyl-amides 
[RC(O)N(R
?
)SR?], thiosulfinates [RS(O)SR?] and polysulfides (RS
x
R?, x ? 3), play 
significant biological roles.
19-21
 Apart from glutathionylation, the formation of the cyclic 
sulfenyl-amide from sulfenic acid effectively prevents overoxidation of cysteine to 
sulfinic and sulfonic acid, as demonstrated by van Montfort
20
 and Salmeen.
21
 
Sulfiredoxin (Srx), a 13kDa yeast protein, was also shown to reduce the sulfinic acid in 
peroxiredoxin (Prx) back to sulfenic acid.
19
 So the reversible reaction between sulfinic 
and sulfenic acid in the presence of certain Prx enzymes contributes to redox-related cell 
signaling. As shown in Scheme 1-1, the formation of all of the sulfur-related compounds 
originates from the oxidation of cysteine. 
Increasingly, experimental studies have demonstrated that cysteine participates in 
electron-transfer process of proteins that modulate essential biological functions, such as 
photosynthesis and aerobic respiration.
22
 Cytochrome P 450, a class of mono-heme 
enzymes with cysteine as the proximal axial ligand, activates the conversion of O
2
 to 
water and catalyzes the oxidation of a wide range of organic substrates. Results from 
Sono et al.
22
 and Auclair et al.
23
 confirm that the thiolate ligand in P 450 enzymes 
facilitates the O?O bond cleavage, and electron transfer may participate in the process. 
More recently, Hirota, et al.
24
 proposed that the heme of myoglobin (Mb) is reduced by 
cysteine thiolate in cysteine-introduced Mbs, and the cysteine radical produced reacts 
with another cysteinyl radical of a different Mb molecule to form a Mb dimer, which 
 6
implies that intramolecular electron transfer from cysteine to Fe(III) is involved in the 
process. 
Erv2p, a yeast endoplasmic reticulum (ER) thiol oxidase, catalyzes the activation of 
O
2
 by the following two reactions (Equation 1-15 and 1-16). Firstly, the enzyme accepts 
electrons from reducing substrates, resulting in a reduction of the bound flavin cofactor. 
Secondly, the enzyme transfers the electrons to molecular oxygen to restore the bound 
cofactor to its initial state.
25
 Unfortunately, the mechanism of electron transfer from the 
reducing substrates to the enzyme remains unclear. 
 
E?FAD + 2 R-SH ? E?FADH
2
 + R-S-S-R                                          (1-15) 
 
E?FADH
2
 + O
2
 ? E?FAD + H
2
O
2 
                                                      (1-16) 
 
Azurin, a single blue copper protein (type 1) present in several bacteria, mediates 
electron transport between cytochrome C
551
 and nitrite reductase. Generally, copper ion is 
bound at one end of the ?-sandwich-shaped molecule, and the disulfide bridge is located 
at the opposite end from the copper ion with a distance of 2.6 nm in azurin molecule.
26
 
The cysteine disulfide radical could be the intermediate state of the disulfide, although no 
supportive evidence was provided. The disulfide radical that was generated from pulse 
radiolysis can transfer one electron to Cu(II) in azurin. Farver et al.
27
 reported that the 
rate constant increases by one order of magnitude upon decreasing pH from 8.0 to 4.0 in 
all azurins, and they proposed that this pH-dependent behavior is either due to changes in 
the electron-transfer distance or due to changes in the electronic coupling.  
 7
All of these biological roles of cysteine are directly or indirectly related to its 
reductive property, so it is of high importance to understand the mechanisms of the 
oxidation of cysteine. Many groups have focused on this interesting subject --- the 
oxidation of cysteine, and over 400 research papers have been published. A wide range of 
oxidants have been selected, from nonmetallic to metallic compounds.  
Of all nonmetallic oxidants involved in the reaction with cysteine, hydrogen 
peroxide, chlorine dioxide, and chlorite are particularly important. A significant increase 
in the concentration of intracellular oxidizing species, such as hydrogen peroxide and 
superoxide, is indicative of oxidative stress. It has triggered intensive research into 
understanding the reaction of cysteine with hydrogen peroxide and superoxide. For the 
reaction of hydrogen peroxide with excess cysteine, hydrogen peroxide reacts with 
cysteine to produce the corresponding sulfenic acid, and then the sulfenic acid reacts with 
excess cysteine to give cystine and water.
28-33
 At least two groups determined the rate 
constant of the reaction at pH 7.4, and both obtained similar results. The rate constant 
determined by Barton et al. 
32
 was around 1.4 M
-1
 s
-1
; while Winterbourn et al. 
33
 reported 
that the rate constant was 2.9 M
-1
 s
-1
. The difference in rate constants may be due to 
copper-catalysis effect.
28,32
 Unfortunately, no copper catalyst was tested in the above two 
experiments. Chlorine dioxide, an effective disinfectant of bacteria, may oxidize the ?SH 
group of glucose oxidase to an ?S-S- group. The oxidation of cysteine (CSH) by chlorine 
dioxide (ClO
2
) and chlorite (ClO
2
-
) was stimulated by such an assumption. Firstly, Lynch 
et al.
34
 reported that cystine is formed in the oxidation of cysteine by the mixture of ClO
2
 
and ClO
2
-
 at pH 7.0. Recently, Darkwa et al.
8
 reported that cysteic acid (CSO
3
H) is the 
product of the oxidation of cysteine by ClO
2
 and ClO
2
-
 in acidic solution (pH ? 1.0). For 
 8
the oxidation of cysteine by excess ClO
2
-
, a 28-step mechanism was proposed to fit the 
experimental data. More recently, the oxidation of cysteine by ClO
2
 was investigated in a 
rather wide pH range (pH: 2.7?9.5) by Margerum?s group.
35
 The reaction rate is first 
order in both CSH and ClO
2
, and increases rapidly with the increase of pH. The 
following mechanism was proposed: an electron transfers from CS
-
 to ClO
2
, with a 
subsequent rapid reaction between the CS
?
 radical and a second ClO
2
 to form a cysteinyl-
ClO
2
 adduct. In acidic solutions, the product of the reaction is cysteic acid; as the pH 
increases, the adduct reacts with CS
-
 by a second pathway to form cystine (CSSC) and 
chlorite ion (ClO
2
-
). Thus, the results from Margerum?s paper explain the above two 
seemingly contradictory reports. However, no copper catalytic effect was tested in that 
paper. 
Metallic cations also serve as good oxidants of cysteine, and this type of redox 
reaction has attracted much attention. The transition-metal cations of chromium(VI) and 
vanadium(V) are well known to be of high toxicity, while chromium(III) and 
vanadium(IV) are not detrimental to human health.
36
 It has been disclosed that exposure 
of cells to Cr(VI) and V(V) results in various forms of DNA damage.
36,37
 The reduction 
of Cr(VI) and V(V) by biothiols in the cells may contribute to the DNA damage. For the 
oxidation of cysteine by Cr(VI), a long lived cysteine-Cr ion (RS-Cr) was observed by 
McCann et al.
38
 and Kwong et al.,
39
 although different mechanisms were proposed by 
them. For the reaction of V(V) with cysteine, two different methods were employed by 
two groups. The UV-vis spectra of a bisulfate-vanadium complex was observed by 
Payasi et al.,
40
 while cysteinyl radical was detected by Shi et al.
37
 using EPR technique. 
The oxidations of cysteine by other metal ions, such as Tc(V),
41
 Fe
3+
,
42-44
 Mn
3+
,
45
 and 
 9
Ce(IV)
46
, also were reported. Cysteinyl radicals were found in the reduction of Mn
3+
 and 
Ce(IV), while only long lived cysteine-metal ions were observed with Tc(V) and Fe
3+
.  
Relative to metal ions, substitution-inert transition-metal complexes are considered 
as good model compounds in understanding the electron transfer (ET) process of 
metalloproteins. Several such metal complexes were selected for the reaction with 
cysteine, including [Co
III
(ox)
3
]
3-
 (ox: oxalate ligand),
47
 [Ir
IV
Cl
6
]
2-
,
48
 [Co
III
W
12
O
40
]
5-
,
49
 
[Fe
III
(CN)
6
]
3-
,
50,51
 [Mo
V
(CN)
8
]
3-
,
52
 [Co
2
(CN)
10
(O)
2
]
5-
,
53
 and 
[(bpy)
2
(H
2
O)Ru
III
ORu
III
(H
2
O)(bpy)
2
]
4+
.
54
 The products of the oxidation of cysteine by 
[Co
III
(ox)
3
]
3-
, [Fe(CN)
6
]
3-
, and [Co
III
W
12
O
40
]
5-
 are cystine and their corresponding 
reduced metal complexes; for the reactions of cysteine with [Co
2
(CN)
10
(O)
2
]
5-
 and 
[(bpy)
2
(H
2
O)Ru
III
ORu
III
(H
2
O)(bpy)
2
]
4+
, the products are cystine and their reduced 
mononuclear complexes; in the cases of [Mo
V
(CN)
8
]
3-
 and [Ir
IV
Cl
6
]
2-
, except for their 
reduced metal complexes, mixtures of cystine and cysteinesulfinate are formed for 
[Mo
V
(CN)
8
]
3-
, and only cysteic acid is produced for [Ir
IV
Cl
6
]
2-
. The above results seem to 
demonstrate that the products of sulfur-containing compounds are correlated with the 
oxidative capability of the metal complex: with weak oxidants, only cystine is formed; 
with strong oxidants, sulfur-containing compounds with higher oxidation states are 
obtained, such as cysteinesulfinate and cysteic acid. Moreover, the ubiquitous copper 
catalytic effect was tested only in the reactions of [Fe
III
(CN)
6
]
3-
, [Mo
V
(CN)
8
]
3-
, and 
[Co
2
(CN)
10
(O)
2
]
5-
. For the oxidations of L-cysteine by the other outer-sphere transition 
metal complexes, the copper-catalysis effect may be overlooked. So it seems reasonable 
to correlate the sulfur-containing products of the redox reaction to the oxidative 
capability of the transition metal complex. 
 10
Both [Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
 are substitution-inert metal 
complexes with mild oxidative capabilities. In Chapter Two, we report the results of 
studies of these two complexes as oxidants of cysteine. Copper catalysis is observed in 
the above reactions. The direct oxidations of cysteine by the Fe(III) complexes are 
achieved by the addition of chelating reagents: dipic
2-
 (dipicolinate) for reaction of 
[Fe(bpy)
2
(CN)
2
]
+
; EDTA (pH ? 10.0) or 1,4,8,11-tetraazacyclotetradecane (abbreviated 
as cyclam) (pH > 10.0) for the reduction of [Fe(bpy)(CN)
4
]
-
. The mechanism for the 
direct oxidation of cysteine is proposed. 
Moreover, the energy crisis is another concern of mankind. In Chapter Three, the 
oxidation of iodide by a series of Fe(III) complexes in acetonitrile is studied. 
Understanding the mechanism of the redox reaction may help improve the efficiency of 
dye-sensitized solar cells (DSSC) in the near future. 
Oil, gas, and coal are three major energy resources on which humans rely. In 1998, 
the mean global energy consumption rate was about 12.8 Trillion watts (TW) per year, 
from which the three major sources account for 80%.
55
 The global energy consumption 
rate is expected to reach 28-35 TW per year by 2050. So there would be an estimated 15-
22 TW shortage of energy supply in the planet if the same amount of fuel were consumed. 
Renewable energy is a good candidate to fill the 15-22 TW gap, although it currently 
constitutes less than 1%. Furthermore, the earth?s oil reserves will last for 50 years; 
natural gas is adequate for 70 years; coal reserves are available for 200 years.
55
 In the 
long run, the substitution of fossil energy resources by renewable energy is inevitable. 
The five most common renewable sources of primary power in the world are 
hydroelectric, geothermal, wind, biomass, and solar. Of all of them, solar energy is one of 
 11
the best choices for noncarbon primary power.
56
 The supply of energy from the sun to the 
Earth is about 10
4
 times more than what mankind consumes today. It indicates that solar 
cells with an efficiency of 10 % that cover only 0.1 % of the Earth would suffice to meet 
our current needs. In 1991, Gr?tzel and coworkers initially made a nanoncrystalline dye-
sensitized solar cell (DSSC), a low-cost alternative to traditional inorganic photovoltaic 
devices.
57
 
The DSSCs consist of five major components: (1) conductive mechanical support, 
(2) nanocrystalline semiconductor film, (3) dye sensitizer, (4) redox couple, and (5) Pt 
counter-electrode.
58,59
 Among them, the nanocrystalline semiconductor film is at the heart 
of the DSSCs system. Figure 1-1 shows the principle of operation and energy level 
scheme of the DSSC. The operating principles of the DSSC are as follows: generally, a 
monolayer of the dye sensitizer is attached to the surface of the nanocrystalline film. The 
injection of an electron into the conduction band of the oxide (generally, TiO
2
 is selected) 
follows photoexcitation of the sensitizer. The dye is regenerated by electron transfer from 
the electrolyte, such as the I
-
/ I
3
-
 couple. Then iodide is regenerated by the reduction of 
triiodide at the counter electrode. The circuit is completed via electron migration through 
the external load.
58,59
 All in all, with the employment of DSSCs, electric power is 
generated from light without suffering any permanent chemical transformation. Further 
extensive investigations of this type of cell are stimulated by its increasing academic and 
industrial importance. Most studies are focusing on modifications of the nanocrystalline 
semiconductor, the redox electrolyte and the dye.
58-63
 
 
 12
 
Figure 1-1. Principle of operation and energy level scheme of the dye-sensitized 
nanocrystalline solar cell. Photoexcitation of the sensitizer (S) is followed by electron 
injection into the conduction band of an oxide semiconductor film. The dye molecule is 
regenerated by the redox system, which itself is regenerated at the counter-electrode by 
electrons passed through the load. Potentials are referred to the normal hydrogen 
electrode (NHE). The open circuit voltage of the solar cell corresponds to the difference 
between the redox potential of the mediator and the Fermi level of the nanocrystalline 
film indicated with a dashed line.  
 13
Wide-band gap semiconductor oxides, such as ZnO, SnO
2
, Nb
2
O
5
 or TiO
2
, have 
been employed in DSSCs. TiO
2
 is the most widely used oxide material among them.
58,60
 
The morphology of the nanocrystalline film plays a key role in photovoltaic conversion. 
A bipyramidal TiO
2
 nanoparticle film, with average particle size of 20 nm and (101) 
orientation on the exposed facets, is successfully prepared by a hydrothermal method.
58
 
The high surface roughness of the mesoporous film slows down charge carrier loss by 
recombination, the efficiency of DSSCs therefore being improved. In case of the redox 
electrolyte, it has been demonstrated that DSSCs with I
-
/I
3
-
 redox electrolyte solution 
display the highest efficiency, although several drawbacks impede its wide 
application.
64,65
 A number of alternative redox mediators and electrolyte systems have 
been explored to redress these limitations, including I
-
/I
3
-
 in either solid polymer, gel, 
ionic liquid, plastic crystal system, solid inorganic materials, Co
II
/Co
III
, or  
SeCN
-
/(SeCN)
3
-
 redox couples.
66-70
 Unfortunately, all of them have a decreased 
efficiency relative to that of I
-
/I
3
-
 in volatile solvent. The inefficient performance of the 
cells prepared with all other redox electrolytes is ascribed to their poorer electron-transfer 
kinetics, and poor contact at the dye-HTM (hole-transport material) interface. The Br
-
/Br
3
-
 redox couple is an exception. Very recently, Sugihara and his coworkers 
demonstrated that DSSCs with Br
-
/Br
3
-
 redox couple have higher overall energy 
conversion efficiency and higher open-circuit photovotage (V
oc
) production than I
-
/I
3
-
 if 
the dye sensitizer has a more positive potential than that of Br
-
/Br
3
-
.
71
 However, their 
allegation awaits further verification. Currently, I
-
/I
3
-
/volatile solvent cells are still the 
first option. 
 14
Selection of dyes is also of high importance in improving the efficiency of DSSCs. 
Basically, the dye must bind strongly to TiO
2
 by means of an anchoring group to 
guarantee efficient electron injection into the TiO
2
 conduction band. The dye must absorb 
solar energy in the visible or near-IR range. Electron injection from the dye to the TiO
2
 
must be much faster than decay to ground state of the dye. Following these essential 
requirements, a variety of mononuclear and polynuclear metal complexes have been 
tested, such as Ru
II
, Os
II
, Pt
II
, Re
I
, Cu
I
, and Fe
II
.
64,72-76
 Progress has been made in the 
optimization of the dye component of the cell by systematically varying the ligands, 
metal and other substituent groups in candidate transition-metal complexes. The family of 
complexes [{(4,4?-CO
2
H)
2
 (bpy)
2
}RuX
2
] (bpy = 2,2?-bipyridyl; X = Cl, Br, I, CN, NCS) 
have shown good performance. DSSCs with the highest solar-to-electrical energy 
conversion efficiency of 11 % are achieved by applying the dye [{(4,4?-
CO
2
H)
2
(bpy)}
2
Ru(NCS)
2
] (cis-RuL
2
(NCS)
2
 is known as N3 dye) now.
64
 
Furthermore, the kinetics of electron transfer at the semiconductor/dye/electrolyte 
interface is critical to the efficiency of the device. Electron-transfer processes and the 
associated rate constants at the sensitized nanocrystalline titanium dioxide ?electrolyte 
interface are shown in Scheme 1-2.
72
 To generate electric power efficiently and 
permanently, injection of an electron from the dye excited state into the semiconductor 
conduction band must be much faster than decay of the excited state to the ground state, 
and faster than the reverse charge recombination between the injected electrons in the  
 15
Scheme 1-2: Electron-Transfer Processes and the Associated Rate Constants k
i
 at DSSCs 
interface
a
  
 
S?TiO
2
 + h? ? S
*
?TiO
2
 
k
1
(1-17)
S
*
?TiO
2
 + Q ? S?TiO
2
 + Q
* 
k
-1
(1-18)
S
*
?TiO
2
 ? S
+
?TiO
2
 + e
-
cb
(TiO
2
) 
k
2
(1-19)
S
+
?TiO
2
 + e
-
cb
(TiO
2
) ? S?TiO
2
 
k
3
(1-20)
S
+
?TiO
2
 + 3/2 I
-
 ? S?TiO
2
 + 1/2 I
3
-
 
k
4
(1-21)
 
conduction band and the oxidized form of the dye, i.e. k
2
 >> k
-1
, k
3
; the rate of reduction 
of the dye cation by a redox couple in liquid electrolyte must be greater than the rate of 
recombination of this dye cation with electrons injected into the semiconductor, i.e. k
4
 > 
k
3
 (see Scheme 1-2). This general rule arouses great interest in the kinetics of both the 
photoinduced electron injection and recombination reactions. Increasing evidence has 
shown that, with favorable energetics and strong dye binding to the film surface, the 
electron injection in the N3-sensitized TiO
2
 and similar system proceeds largely on 
ultrafast time scales (10
12
 ? 10
13
 s
-1
).
77-79
 For instance, Tachibana et al.
77
 used 
subpicosecond transient absorption spectroscopy to study the rate of electron injection 
from an excited state of N3 to conduction band of TiO
2
, with a rate constant close to 10
12
 
s
-1
. Ellingson et al.
78
 measured the near- and mid-IR transient absorption for N3 adsorbed 
to films of nanocrystalline TiO
2
, and an electron injection time with an upper limit of 
about 50 fs was obtained; Hannappel et al.
79
 measured a transient absorption signal of the 
 16
injected hot electrons with a rise time of less than 25 fs. More importantly, Durrant and 
his coworkers
77
 found that the electron injection kinetics are insensitive to the solvent 
environment, applied electrical potential, and the redox potential of the excited-state dye, 
which is consistent with the activationless electron injection into the TiO
2
 conduction 
band acceptor state.
77
 However, the recombination kinetics between the N3 dye cation 
and electron injected into the TiO
2
 film are strongly dependent on the excitation intensity, 
solvent/electrolyte composition, and the applied electrical bias, with recombination time 
constants from tens of picoseconds to milliseconds.
80
 A rapid acceleration of the charge 
recombination kinetics is observed provided that the excitation intensity is above a 
certain level or the applied potential is more negative than a threshold potential V
kin
; the 
recombination rate constant in CH
3
CH
2
OH/0.1 M TBAT (tetrabutylammonium triflate) is 
about 10
6
 faster than that in dry acetonitrile/0.1 M TBAP (tetrabutylammonium 
perchlorate)/0.1 M LiClO
4
/0.5 M 4-tert-butypyridine if the applied potential is -400 mV 
vs Ag/AgCl, while it is weakly dependent upon electrolyte composition if no or positive 
applied potentials are applied, with the recombination time constants of milliseconds.
80
 A 
very interesting result was also obtained by Tachibana et al. in their comparison of 
ruthenium bipyridyl with porphyrin sensitizer dyes.
77
 They found that the three different 
dyes have similar electron injection and recombination rate constants, while only Ru(2,2?-
bipyridyl-4,4?-dicarboxylate)
2
-cis-(NCS)
2
 (Ru(dcbpy)
2
(SCN)
2
) dye shows high 
photoelectrochemical conversion. They ascribed the lower efficiency of porphyrin 
sensitizers to an increasing probability of their excitation annihilation, which may slow 
down the electron transfer from the iodide to the dye cation. 
 17
The important role of the interfacial oxidation of iodide by the dye cation in DSSC 
cells prompted several groups to study its kinetics. The complexity of the interfacial 
interaction impedes our understanding of its kinetic behavior. So far only a few 
inconsistent results have been given. Meyer and his coworkers found that cis-
Ru(dcb)
2
(CN)
2
/TiO
2
 (dcb: 4,4?-(COOH)
2
-2,2?-bipyridine) converts absorbed photons to 
electrons nearly quantitatively while cis-Os(dcb)
2
(CN)
2
/TiO
2
 does not.
81
 Then they 
measured the rate constants of the oxidation of iodide by these two dye cations using in-
situ time-resolved diffuse reflectance, with 2 ? 10
7
 M
-1
 s
-1
 (with an assumed 1
st
-order 
iodide dependence) for cis-Ru(dcb)
2
(CN)
2
/TiO
2
 and less than 9.0 ? 10
6
 M
-1
 s
-1
 for cis-
Os(dcb)
2
(CN)
2
/TiO
2
; they ascribed this poor photon conversion of cis-
Os(dcb)
2
(CN)
2
/TiO
2
 to the sluggish oxidation of iodide by Os(III), the rate-limiting step 
in the sensitization of the cis-Os(dcb)
2
(CN)
2
/TiO
2
 photoelectrochemical cell. However, 
for cis-Ru(dcb)
2
(CN)
2
/TiO
2
, the rate-limiting step is the recombination of injected 
electron with dye cation. Iodide can efficiently regenerate the sensitizer ground state 
within about 200 ns. So the efficient photocurrents are realized in cis-
Ru(dcb)
2
(CN)
2
/TiO
2
 solar cells. Kuciauskas et al.
72
 also determined the rate constants (k
4
 
in Scheme 1-2) for the oxidation of 0.50 M iodide by Ru/Os dye cations by transient 
absorption spectroscopy, with 8.0 ? 10
5
 s
-1
 for Ru(dcb)
2
(CN)
2
/TiO
2
 and 6.5 ? 10
5
 s
-1
 for 
Os(dcb)
2
(CN)
2
/TiO
2
. Unfortunately, no iodide-dependent kinetic results were provided 
by these two reports, so the rate constants obtained are incomparable. The regeneration of 
the Ru(II) dye by iodide in [Ru(bpy)
2
(dcb)]
2+
/SnO
2
 photoelectrochemical solar cells was 
studied by Nasr, et al.,
60
 from which an intermediate state of I
2
?-
 was observed. The 
reaction is first order in iodide and Ru(III) with the rate constant of 1.2 ? 10
10
 M
-1
 s
-1
, and 
 18
the mechanism of the reaction was proposed. The higher oxidation potential of 
[Ru(bpy)
2
(dcb)]
2+
 and the interfacial interaction between [Ru(bpy)
2
(dcb)]
2+
 and SnO
2
 
may contribute to the higher rate constant of the reduction of [Ru
III
(bpy)
2
(dcb)]
3+
. 
The kinetics and the mechanism of the oxidation of iodide by substitution-inert 
transition metal coordination complexes in aqueous solution are well understood. 
82-85
 
The stoichiometry of the reaction is: 
 
2 M
ox
 + 2I
-
 ? 2 M
red
 + I
2
                                                                            (1-22) 
 
with the rate law being  
 
]M)[]I[]I[(2
dt
]d[M
ox
2
21
ox ??
+=? kk                                                         (1-23) 
 
The reversibility of Equation 1-22 is dependent on the oxidative capability of the 
transition metal complexes. With a strong oxidants, the reverse reaction is negligible; 
otherwise, the forward and reverse reactions are included. The oxidation of iodide by 
outer-sphere transition metal complexes in aqueous solution proceeds by two parallel 
paths, with the formation of I
?
 and I
2
?-
 intermediate states, as shown in Equations 1-24 
and 1-25. Both reactions are first order in transition metal complex, but are first order and 
second order in iodide, respectively. 
 
M
ox
 + I
-
 ? M
red
 + I
?
                                                           k
1
                      (1-24) 
 19
 
M
ox
 + 2 I
-
 ? M
red
 + I
2
?-
                                                       k
2
                     (1-25) 
 
The standard potentials of I
?
/I
-
 and I
2
?-
/2I
-
 in aqueous solution have been determined, 
with E
?
 (I
?
/I
-
) = 1.33 V, and E
?
 (I
2
?-
/2I
-
) = 0.94 V vs NHE.
86
 Moreover, no trace metal-ion 
catalysis has been reported in the oxidation of iodide in aqueous solution, with one 
exception for the reduction of [Fe(CN)
6
]
3-
.
87
 Unfortunately, the oxidation of iodide by 
outer-sphere transition metal complexes in a genuine nonaqueous solvent has not been 
reported. Our research is stimulated by this paradox ?  the important role of the 
interfacial oxidation of iodide by dye cation in acetonitrile, and the lack of understanding 
of the kinetics and mechanism of such redox reactions.  
In Chapter Three, the oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]NO
3
, 
[Fe
III
(CH
3
Cp)
2
]PF
6
, [Fe
III
(dmbpy)
2
(CN)
2
]NO
3
, and [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 in 
acetonitrile is described. The general copper catalysis is inhibited with the addition of 
2,2?-bipyridine. The kinetics of the direct oxidation of iodide was studied with the 
addition of 2,2?-bipyridine. Linear Free-Energy Relationships (LFER) were obtained for 
both log k
1
 and log k
2
 vs E
1/2
. The reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 is mildly inhibited by 
Fe(II), and a quantitative study of this effect leads to a value of 0.60 ? 0.01 V (vs 
[Fe(Cp)
2
]
+/0
) for the reduction potential (E?) of I
?
/I
?
. Deriving from this work, two papers 
have been published: J. Phys. Chem. A, 2004, 108, 7637-7638; Inorg. Chem. 2006, 45, 
3415-3423. 
Next, in this chapter, we briefly introduce some of the basic concepts, theories and 
techniques that are related to this dissertation. 
 20
 
BASIC CONCEPTS IN KINETICS 
Chemical kinetics is the study of the rates of reactions. The overall goal is to 
understand reactions in terms of simple steps involving atomic, molecular, or ionic 
interactions and collisions. The reaction rate is the increase in molar concentration of the 
product of a reaction per unit time, or the decrease in molar concentration of the reactant 
per unit time. For general reaction 
 
        nNmMbBaALL+?+                                                              (1-26) 
 
the reaction rate, the rate of disappearance of reactants and rate of formation of products 
are related by 
dt
d[N]
n
1
dt
d[M]
m
1
dt
d[B]
b
1
dt
d[A]
a
1
Rate ==?=?=                                    (1-27) 
 
Another important concept is the rate law. The rate law is the experimentally 
determined dependence of the reaction rate on reagent concentrations. In simple cases it 
has the following general form: 
 
??
[B][A]Rate k=                                                                                    (1-28) 
 
where k is a proportionality constant called the rate constant. The rate of a reaction is 
dependent on the concentrations of reactants and products, while the rate constant is not. 
 21
The rate constant is affected by temperature, pressure, catalyst, and properties of solvent. 
The exponents m and n are determined experimentally from kinetic studies. The sum of 
the exponents is the overall order of the rate law. However, the order of the reaction is 
usually represented with respect to individual reagents. The exponents in the rate law 
have no necessary relationship to the stoichiometric coefficients in the balanced chemical 
reaction.  
Combining Equation 1-27 with 1-28, the following equation is obtained: 
 
??
[B][A]
dt
d[N]
n
1
dt
d[M]
m
1
dt
d[B]
b
1
dt
d[A]
a
1
k===?=?                            (1-29) 
 
If the reaction were monitored by the disappearance of A, then the experimental 
rate constant would be ak, and it must be divided by a to get the numerical value of k as 
defined by Equation 1-28. 
The half-life is one more important concept in kinetics. The half-life, t
1/2
, is the time 
required for a reactant concentration to reach half of its initial value. This term has a 
quantitative relationship to the rate constant in simple cases, for example, t
1/2
 = 1n2/k for 
first-order reactions. When the kinetic traces are not well fit by mathematical models, the 
half-life is a good way to represent the rate of the reactions. 
Reactions can be of zero-order, first-order or second-order. Here we are interested 
in first-order and pseudo-first-order reactions. For the irreversible elementary reaction: 
 
BA
1
???
k
                                                                   (1-30) 
 22
 
the rate of disappearance of A and appearance of B are given by: 
 
]A[
dt
d[B]
dt
d[A]
Rate
1
k==?=                                                                   (1-31) 
 
Suppose the concentrations of A at time of zero and t are [A]
0
 and [A]
t
, integration 
of Equation 1-31 gives the following equations. 
 
t
0t
1
e[A][A]
k?
=                                                                                         (1-32) 
 
tln[A]ln[A]
10t
k?=?                                                                              (1-33) 
 
Generally, some physical properties (P) proportional to the concentration of 
reactants/products are followed in kinetics. In the simplest situation of first-order 
reactions, if only the reactant A contributes to this property P, it is easy to see that the 
fractional change in P is proportional to [A], and may be substituted for it in Equation  
1-32 and 1-33. It has been demonstrated by Espenson
88
 that the same result also holds 
when the product contributes to the value of P for first-order reactions. Therefore, in our 
kinetic studies, the reaction could be monitored by the absorbance at any wavelength 
provided it follows first order kinetic behavior.  
Pseudo-first-order reaction conditions are widely used in kinetic studies. The 
pseudo-first-order condition is that the concentration of the limiting reactant is much less 
 23
than that of any of other reactants, from which the reaction is first-order in the limiting 
reactant. For example: 
 
5 Br
-
 + BrO
3
-
 + 6 H
+
 = 3 Br
2
 + 3 H
2
O                                                   (1-34) 
 
The rate law is:
88
 
 
2
3
3
]][H][Br[BrO
dt
]d[BrO
+??
?
=? k                                                       (1-35) 
 
When the concentrations of H
+
 and Br
-
 are at least ten times that of BrO
3
-
, the rate 
of the reaction would become: 
 
]BrO[]][H][Br[BrO
dt
]d[BrO
3
'2
3
3
?+??
?
==? kk                                 (1-36) 
 
where k' is: 
 
2'
]][H[Br
+?
= kk                                                                                   (1-37) 
 
In our research, the concentration of cysteine (in Chapter Two) or iodide (in 
Chapter Three) is at least ten times greater than that of transition metal complexes. The 
 24
reactions are monitored by the loss of Fe(III) or the increase of Fe(II). Applying pseudo-
first-order conditions simplifies our kinetic studies, and good results are obtained. 
 
ELECTRON-TRANSFER REACTION MECHANISM AND MARCUS THEORY 
Generally, inorganic reactions can be divided into either substitution reactions or 
oxidation-reduction (redox) reactions. Redox reactions, in which one or more electrons 
transfer from the reducing reagent to the oxidant, are involved in many biological and 
environmental processes. According to the mechanism, Taube classified the electron-
transfer processes into inner-sphere or outer-sphere reactions.
89,90
 
In inner-sphere redox reactions, there exists a bridged transition state such that a 
chemical bond is formed between one of the reactants and a ligand that is still bonded to 
the other reactant. Any inner-sphere reaction must include three steps: 1. formation of the 
bridged complex; 2. the redox process; and 3. decomposition of the successor complex to 
final products. The first thoroughly investigated inner-sphere reaction, the oxidation of 
Cr
2+
 by [CoCl(NH
3
)
5
]
2+
, is a good example, as shown in Scheme 1-3.
91
 
From their central electronic structures, the complexes of cobalt(III) (low spin d
6
) 
are inert, of chromium (II)(high-spin d
4
) are labile, of cobalt (II)(high-spin d
7
) are labile, 
and of chromium(III) (d
3
) are inert.
91
 This combination of conditions allowed the reaction 
products to establish definitively an inner-sphere mechanism for the reaction in Scheme 
1-3. The metal ions keep their original oxidation states in the precursor complex. The 
relative labilities of Co(II) and Cr(III) result in the successor complex decomposing to the 
Cr(III) product in which the Cl
-
 ligand originally on cobalt(III) has been transferred to 
chromium(III).
91
 The inner-sphere mechanism of the above reaction was also 
 25
experimentally proven in the radioactive 
36
Cl
-
 tracer studies,
90
 from which it was revealed 
that over 99% of the chloride incorporated into the product arises from the 
[CoCl(NH
3
)
5
]
2+
 oxidant. 
 
Scheme 1-3  
(NH
3
)
5
CoCl
2+
 + Cr(OH
2
)
6
2+
-H
2
O
Co
III
Cl
H
3
N
H
3
N
H
3
N
NH
3
H
3
N
Cr
II
OH
2
OH
2
OH
2
H
2
O
OH
2
electron
transfer
5H
+
5H
2
O
Co
II
Cl
H
3
N
H
3
N
H
3
N
NH
3
H
3
N
Cr
III
OH
2
OH
2
OH
2
H
2
O
OH
2
6H
2
O 5H
+
Co(OH
2
)
6
2+
 + 5NH
4
+
 + (H
2
O)
5
CrCl
2+
Precursor complex Successor complex
 
 
For outer-sphere electron transfer reactions, the coordination spheres of the oxidant 
and reductant remain intact during the electron transfer. Generally, both reductant and 
oxidant have nonlabile coordination spheres, and no ligand substitution occurs during the 
reaction. There are three steps in outer-sphere reactions. The first step is the formation of 
precursor. The second step is the electron transfer itself. The third step is dissociation of 
the successor complex, which is generally fast. In most cases, the electron-transfer step is 
rate limiting. 
In 1954, Marcus theory was developed to better describe the behavior of outer-
sphere reactions. The Franck-Condon principle is fundamental to the theory. The Franck-
 26
Condon principle states that electron movement is much faster than nuclear motion. 
According to the Franck-Condon principle and Marcus theory, the time required for 
electron transfer between donor and acceptor nuclei is markedly less than the time needed 
for those nuclei to change their positions. Before the electron can be transferred, however, 
the free energies of the donor and acceptor species must become identical. This can only 
be achieved if the inner-sphere coordination shells and their solvent environments both 
simultaneously come into new and appropriate configurations. The reaction rate constant 
will depend largely on the reorganization energy (?G
IS
?
 and ?G
OS
?
). Moreover, the 
Coulombic work (?G
CO
?
) between the donor and acceptor contributes to the free energy 
of activation as well. The self-exchange rate constants of outer-sphere reaction are 
obtained by applying the following equation:
91
 
 
)
RT
?G
exp(??
?
osnelcalc
?
=?k                                                         (1-38) 
 
?
CO
?
OS
?
IS
?
?G?G?G?G ++=                                                     (1-39) 
 
where ?
el
 is the transmission coefficient, ?
n
 represents the tunneling constant (unit: s
-1
), 
K
os
 is the pre-exponential factor (unit: M
-1
), ?G
?
 represents the overall activation free 
energy.  
However, the unavailability of ?G
IS
?
, ?G
OS
?
, and ?G
co
?
 impedes the wide 
application of Equation 1-38. To overcome this limitation, the following simplified 
equation was developed by Marcus to calculate the self-exchange rate constant.
92
 
 27
 
12
2/1
1212221112
)( WfKkkk =                                                                (1-40) 
 
where k
12
 is the observed cross electron-transfer rate constant, k
11
 and k
22
 are the self-
exchange rate constants of redox couple 1 and 2, K
12
 is the cross electron-transfer 
equilibrium constant, W
12
 is work term, f
12
 is a factor, usually close to one. 
In this research, the Marcus cross relationship was used to calculate the self-exchange 
rate constants for 
?
SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
,and 
?
SCH
2
CH(NH
2
)CO
2
?
 
/
?
SCH
2
CH(NH
2
)CO
2
?
. 
 
TECHNIQUES 
Several techniques were used in this research, including NMR, UV-vis spectroscopy, 
cyclic voltammetry, atomic absorption spectroscopy, and stopped-flow 
spectrophotometry. NMR spectroscopy was used to determine the purity of synthesized 
ligands, identify the products of the reaction, and determine stoichiometric reaction ratios; 
cyclic voltammetry was employed to determine the standard potential of the outer-sphere 
Fe(III) complexes; atomic absorption spectroscopy was used to determine the copper 
concentration in the reactants and products; UV-vis spectroscopy was used to identify the 
reaction products, determine the stoichiometric ratio, and perform kinetic studies with 
slow rate constants. Most of the kinetic studies were carried out on a stopped-flow 
instrument. Here, we will present the basic principle of the stopped-flow technique. 
The basic principle of the stopped-flow technique is to mix equal volumes of two 
reactant solutions in the flow cell simultaneously. Solutions are transferred into the 
 28
driving syringes and stored in a thermostated tank to keep the temperature of the 
solutions constant. Solutions are then injected into the optical flow cell by applying 5 bar 
of compressed air on the driving plate to push the plungers of the driving syringes. The 
solution mixture in the cell is then expelled to the stopping syringe and the plunger is 
pushed against the stopping block. The trigger on the stopping block then initiates the 
computing system to collect the signal from the detector at the same time.
93,94
  
 29
 
 
 
 
 
CHAPTER TWO 
Direct Oxidation of L-Cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
  
and [Fe
III
(bpy)(CN)
4
]
-
 
 
Introduction 
Thiol-containing compounds play important roles in biochemical process, including 
maintaining cellular redox potentials, protecting cells against oxidative stress, and 
regulating metabolism and gene expression.
1-3,10,12
 All of these processes involve 
thiol/disulfide exchange, hydrogen-atom transfer, or electron transfer.
2
 The most 
important thiols in biological chemistry are cysteine, homocysteine, and glutathione. 
Cysteine, one of the essential ?-amino acids and one of the simplest biological thiols with 
the oxidation state of -2 on element of sulfur, will not easily cleave at the C-S bond, and 
can undergo oxidation either by metal ions or nonmetallic oxidants to form sulfinic acid 
or disulfide, depending on the oxidative capability of the oxidant. The electron-transfer 
reaction of cysteine may be the mechanistic basis for the anti-oxidation of DNA, lipids 
and proteins.
95,8
 Furthermore, the concentration of cysteine in living cells is very much 
indicative of a human being?s physical condition: depletion of cysteine can accompany 
leukemia and several types of cancer; abnormally high levels of cysteine can be 
associated with brain disorders, chromosome aberrations, etc. Understanding the 
mechanism of the oxidation of cysteine by metal complexes will be beneficial to 
 30
developing rapid and reliable sensing techniques for detection of cysteine and other 
biological thiols, and therefore it will contribute to the prevention of some fatal 
disease.
3,96-98
 
The oxidation of cysteine by metal cations or metal complexes has been widely 
investigated, and a variety of mechanisms have been proposed.
38,40,44,48-53,99-101
 For metal-
ion oxidants, such as V(V),
40
 Cr(VI),
38
 Fe(III),
44
 a long lived cysteine-metal ion (RS-M) 
complex was observed in the reactions. To simplify their kinetic behavior, metal ions 
were replaced by some typical outer-sphere transition metal complexes, including 
[Ir
IV
Cl
6
]
2-
,
48
 [Co
III
W
12
O
40
]
5-
,
49
 [Fe
III
(CN)
6
]
3-
,
50
 [Co
2
(CN)
10
(O)
2
]
5-
,
53
 [Mo
V
(CN)
8
]
3-
,
52
 and 
[{(bpy)
2
H
2
ORu
III
}
2
O]
4+
.
54
 The reductions of [Fe(
III
CN)
6
]
3-
, [Co
2
(CN)
10
(O)
2
]
5-
 and 
[Mo
V
(CN)
8
]
3-
 are trace copper catalyzed, while no catalysis was reported for the reactions 
of [Co
III
W
12
O
40
]
5-
, [Ir
IV
Cl
6
]
2-
, and [{(bpy)
2
H
2
ORu
III
}
2
O]
4+
. It is possible that the copper 
catalysis was overlooked by the investigators of these reactions. 
The ubiquitous copper catalysis was also observed in our group for the oxidation of 
thioglycolate (TGA) by [Ir
IV
Cl
6
]
2-
,
102
 [Mo
V
(CN)
8
]
3-
,
103
 and [Os
III
(phen)
3
]
3+
,
104
 and for the 
oxidation of L-cysteine by [Mo
V
(CN)
8
]
3-
.
52
 The kinetic complication due to copper 
catalysis was overcome with the addition of 2,6-pyridinedicarboxylic acid (dipicolinic 
acid, abbreviated as dipic hereafter). For the oxidation of TGA by [Ir
IV
Cl
6
]
2-
, 
[Mo
V
(CN)
8
]
3-
 and [Os
III
(phen)
3
]
3+
, the weakest oxidant, [Mo
V
(CN)
8
]
3-
, gives only 
disulfide as product; the strongest oxidant, [Ir
IV
Cl
6
]
2-
, yields a mixture of the disulfide 
and the sulfonic acid; and the reduction of [Os
III
(phen)
3
]
3+
, with intermediate standard 
potential between [Mo
V
(CN)
8
]
3-
 and [Ir
IV
Cl
6
]
2-
, shows dramatically different products: 
disulfide, [Os
III
(phen)
3
]
2+
, and [Os(phen)
2
(phen-TGA)]
2+
. The different products of the 
 31
above direct outer-sphere redox reactions may be closely related to the oxidative 
capability of substitution inert transition-metal complexes. In the case of oxidation of L-
cysteine by [Mo
V
(CN)
8
]
3-
, the reaction also follows a similar mechanism?electron 
transfer from thiolate to form cysteinyl radical. The products of the reaction of 
[Mo
V
(CN)
8
]
3-
 are a mixture of cystine and cysteinesulfinate. 
To extend our understanding of the mechanism of oxidation of cysteine, two 
relatively weaker oxidants, [Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
, were selected in 
this study. The direct reduction of [Fe
III
(bpy)
2
(CN)
2
]
+
 is achieved with the addition of 1.0 
mM dipic
2-
, while that of [Fe
III
(bpy)(CN)
4
]
-
 is achieved by the addition of EDTA (pH ? 
10.0) or cyclam (pH > 10.0). No cysteinesulfinate was detected in the above two 
reactions. The rate of the reaction is pH-dependent. The mechanism for the direct 
oxidation of cysteine is proposed, revealing the rate limiting step of the reaction?
electron transfer from the thiolate of cysteine to Fe(III). The self-exchange rate constants 
of ?SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
 and 
?
SCH
2
CH(NH
2
)CO
2
?
 
/
?
SCH
2
CH(NH
2
)CO
2
?
 are obtained by applying Marcus? equations.  
 
Experimental section 
Reagents and solutions. Ferrous ammonium sulfate hexahydrate (Fisher), 
potassium cyanide (Fisher), 2,2?-bipyridyl (Aldrich), L-cysteine (Fluka), L-cystine 
(Aldrich), N-tert-butyl-?-phenylnitrone (PBN, Aldrich), cacodylic acid (Sigma), sodium 
phosphate monobasic monohydrate (J. T. Baker), sodium phosphate dibasic (MCB), 
sodium hydroxide (Fisher), D
2
O (Aldrich), 3-(trimethylsilyl)-1-propanesulfonic acid, 
sodium salt (DSS, Aldrich), 5,5?-dithiobis (2-nitrobenzoic acid) (DTNB, Aldrich), 
 32
ammonium chloride (Fisher), sodium bicarbonate (J. T. Baker), sodium carbonate 
(Fisher), cupric sulfate (Fisher), nickel sulfate hexahydrate (Fisher), ferric nitrate 
nonahydrate (Fisher), ferrous sulfate heptahydrate (J. T. Baker), cobalt sulfate 
heptahydrate (Fisher), manganese sulfate monohydrate (Fisher), zinc sulfate heptahydrate 
(J. T. Baker), ammonium molybdate (Aldrich), ethylenediaminetetraacetic acid disodium 
salt dihydrate (abbreciated as EDTA hereafter, MCB), 2,6-pyridinedicarboxylic acid 
(abbreviated as dipic hereafter, Aldrich), 1,4,8,11-tetraazacyclotetradecane (cyclam, 
Aldrich), tetraphenylphosphonium chloride (Aldrich), chlorine gas (Matheson Gas 
Products, Inc), acetic acid (Fisher), sulfuric acid (Fisher), nitric acid (Fisher), 
hydrochloric acid (Fisher), chloroform (Fisher), dimethyl sulfoxide (DMSO, Fisher), and 
Dowex 50W-X8 resin (J. T. Baker) were used without further purification. The structure 
of several compounds is shown in Scheme 2-1. 
 
Scheme 2-1 
 
Thioglycolic acid L-cysteine
 
N CO
2
HHO
2
C
Dipicolinic acid
CH N
+
C(CH
3
)
3
O
?
PBN
NH HN
NH HN
Cyclam
HS
CO
2
H HS
CO
2
H
NH
3
+
 
 33
Sodium perchlorate (Fisher), lithium perchlorate (GFS chemical Co.), and sodium 
triflate (Aldrich) were recrystallized from hot water; their stock solutions were 
standardized by running an aliquot through a cation-exchange column packed with 
Dowex 50W-X8 resin, then titrating by standardized NaOH aqueous solution. Sodium 
acetate trihydrate (Fisher) was recrystallized from hot water and dried in a vacuum 
desiccator. Chloroacetic acid (Fisher) was recrystallized from hot water, and dried in a 
vacuum desiccator; stock solutions were titrated with standardized NaOH aqueous 
solution. 
Distilled deionized water was obtained from a Barnstead NANO pure infinity 
ultrapure water system. All Fe(III) and L-cysteine solutions were prepared in de-ionized 
water just prior to use, and purged with Ar or N
2
 for at least 15 minutes prior to the 
reaction to prevent possible complications caused by O
2
. Fe(III) solutions were kept in 
the dark to prevent any photochemical changes. The concentration of L-cysteine was 
determined spectrophotometrically with Ellman?s reagent at pH 7.46.
105,106
 
Preparation of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O. [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O was prepared 
following a standard procedure.
107
 0.47 g of 2,2?-bipyridyl (3.0 mmol) and 0.39 g of 
ferrous ammonium sulfate hexahydrate (1.0 mmol) were added to 40.0 mL of water. The 
solution became dark red. Then it was heated to just below the boiling point. A freshly 
prepared solution of 1.00 g of KCN (15.4 mmol) in 2.0 mL of water was immediately 
added to the heated solution. After brief stirring, the solution was left to cool at room 
temperature for 4 hours. The dark violet product was obtained by vacuum filtration, 
rinsed with cool water, and dried in vacuum at room temperature. Yield of crude 
[Fe
II
(bpy)
2
(CN)
2
]?3H
2
O: 0.36 g (76%). For purpose of electronic spectra and 
 34
electrochemical measurements, a portion of the initial products was recrystallized from 
concentrated sulfuric acid (96.4%). The yield of recrystallized [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O: 
70%. Anal. Calcd for C
22
FeH
22
N
6
O
3
: C, 55.71; H, 4.68; N, 17.72. Found: C, 55.66; H, 
4.41; N, 17.67.
 1
H-NMR (400 MHz/D
2
O): 9.40 (d, J = 6.5, 2H); 8.36 (d, J = 7.7, 2H); 
8.32 (d, J = 8.0, 2H); 8.09 (t, J = 7.8, 2H); 7.92 (t, J = 7.8, 2H); 7.60 (t, J = 5.8, 2H); 7.32 
(d, J = 6.0, 2H); 7.21 (t, J = 7.0, 2H). 
Preparation of [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O. [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O was 
synthesized by oxidation of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O with nitric acid.
107
 1.60 mL of 
concentrated nitric acid (70%) was added to 0.40 g of the crude [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O. 
A deep purple solution was formed, with the evolution of nitrogen dioxide. The mixture 
was slightly warmed to make the reaction complete. It was then diluted with 40.0 mL of 
distilled deionized water to give a purple precipitate. A purple solution was obtained by 
heating the mixture. Then the filtrate was obtained by filtering the hot clear solution. A 
purple precipitate was recovered after putting the filtrate in a refrigerator overnight. The 
product was obtained by vacuum filtration, rinsed with cool water, and dried in vacuum 
at room temperature. Yield of recrystallized [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O: 0.28 g (65%). 
Anal. Calcd for C
22
FeH
20
N
7
O
5
: C, 50.98; H, 3.89; N, 18.92. Found: C, 51.02; H, 3.58; N, 
19.11. 
Preparation of K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O. Schilt?s method
107
 was used to prepare 
K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O. 0.474 g of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O (1.00 mmol) and 19.5 g of 
KCN (0.299 mol) were mixed in 150.0 mL of water, and the solution was refluxed for 24 
hrs. It was extracted with 50.0 mL of chloroform to remove the remaining 
[Fe
II
(bpy)
2
(CN)
2
]?3H
2
O. 90.0 mL of water was removed by heating the solution on a hot 
 35
plate at 70 
o
C. The remaining heated solution was placed in a refrigerator overnight. The 
dark brown precipitate was obtained by vacuum filtration. Yield of 
K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O: 0.425 g (95%). For purpose of electronic spectra and 
electrochemical measurements, a portion of the initial products was recrystallized from 
hot water. The yield of recrystallized K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O: 23%. 
1
H-NMR (400 
MHz/D
2
O): 9.32 (d, J = 4.9, 2H), 8.20 (d, J = 7.6, 2H), 7.94 (t, J = 7.8, 2H), 7.47 (t, J = 
6.7, 2H) 
Preparation of Li[Fe
III
(bpy)(CN)
4
]?2.5H
2
O. Li[Fe
III
(bpy)(CN)
4
]?2.5H
2
O was 
prepared according to published procedures.
108
 Chlorine gas was bubbled through a warm 
(T ? 50 
o
C) dark red aqueous solution (40.0 mL) of K
2
[Fe
II
(bpy)(CN)
4
]?3H
2
O (1.10 
mmol) for 30.0 minutes under continuous stirring, leading to a color change from red to 
orange. A concentrated warm aqueous solution of tetraphenylphosphonium chloride 
(0.449 g/6.0 mL H
2
O, 1.20 mmol) was added to the above solution, and the solution was 
kept heating and stirring for 10 minutes. An orange solid (PPh
4
[Fe
III
(bpy)(CN)
4
]?H
2
O) 
was obtained by vacuum filtration. To prepare Li[Fe
III
(bpy)(CN)
4
]?2.5H
2
O, 0.251 g of 
lithium perchlorate (1.50 mmol) was slowly added to an acetonitrile solution (0.805 
g/13.0 mL) of PPh
4
[Fe
III
(bpy)(CN)
4
]?H
2
O (1.20 mmol). An orange precipitate was 
collected by vacuum filtration and dried in a vacuum desiccator for 12 hrs. Yield of 
Li[Fe
III
(bpy)(CN)
4
]?2.5H
2
O: 0.292 g (66%). Anal. Calcd for C
14
FeH
12.4
LiN
6
O
2.2
: C, 
46.32; H, 3.39; N, 23.17. Found: C, 46.37; H, 3.54; N, 22.90. 
Methods. All single UV-vis spectra were recorded on a HP-8453 or HP-8452 
diode-array spectrophotometer equipped with 1 cm cells and a Brinkman Lauda RM6 
thermostated water bath to maintain the temperature at 25.0 ? 0.1
 o
C. Cyclic 
 36
voltammograms and Osteryoung Square Wave Voltammograms (OSWV) were recorded 
at room temperature on a BAS 100 B electrochemical analyzer equipped with BAS cell 
stand C3 with purging and stirring system, using a glassy carbon working electrode (d = 
3.1 mm), Ag/AgCl (3.0 M NaCl) reference electrode, and a Pt wire auxiliary electrode. 
The scan rate was 100 mV/s. A 450 Corning pH/ion meter with a Mettler Toledo InLab 
421 pH electrode was used for pH measurement. 
1
H NMR spectra were obtained on a 
multi-pulse Bruker AV 400 spectrometer. Kinetic studies were performed by mixing 
equal volumes of the two reactants on a Hi-Tech Scientific model SF-51 stopped-flow 
apparatus that was equipped with a SU-40 spectrophotometer unit and a 2095 circulatory 
water bath. Temperature was maintained at 25.0 ? 0.1
o
C. An Olis 4300S system was used 
for data acquisition and analysis. All apparent rate constants were the average of at least 
five runs with ? 5% error or less. Reactions were monitored at 522 nm (Fe
II
(bpy)
2
(CN)
2
, 
?
522
 = 5115 M
-1
 cm
-1
) and 482 nm ([Fe
II
(bpy)(CN)
4
]
2-
, ?
482
 = 2825 M
-1
 cm
-1
), and the rate 
constants were obtained by fitting the data with OLIS-supplied first-order functions. A 
nonlinear-least squares computer program with relative weighting (weighting by 1/Y
2
) 
was used to fit the overall rate law to the values of k
obs
. SPECFIT/32 version 3.0.15 
global analysis system was applied to simulate the reaction traces. For all kinetics studies, 
the concentration of L-cysteine was at least ten times greater than that of Fe(III) 
complexes.  
 
Results 
Characterization of the iron complexes. The UV-vis absorbance characteristics of 
the iron complexes are shown in Table 2-1, in which the extinction coefficients are very 
 37
close to previous reports.
107,109
 The cyclic voltammograms (CV) of [Fe
III
(bpy)
2
(CN)
2
]
+
 
and [Fe
III
(bpy)(CN)
4
]
-
 are quasi-reversible. The CV of 1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 
M NaCF
3
SO
3
 is shown in Figure 2-1, with ?E
p/p 
= 67 mV, and E
1/2
 = 554 mV (E
1/2
 = (E
pc
 
+ E
pa
)/2) versus Ag/AgCl
(s)
 (E
o
Ag/AgCl(s)
 = 0.205 V vs NHE
110
) that is, 0.759 V versus a 
normal hydrogen electrode, the same as Terrettaz?s report.
111
 The CV of 0.70 mM 
[Fe
III
(bpy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
 is shown in Figure 2-2, with ?E
p/p 
= 64 mV, and 
E
1/2
 = 332 mV versus Ag/AgCl
(s)
 (E
o
Ag/AgCl(s)
 = 0.205 V vs NHE
110
), that is, 0.537 V 
versus a normal hydrogen electrode, the same as previous reports.
111,112
 The OSWVs of 
1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 and 0.70 mM [Fe
III
(bipy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
 are 
shown in Figure 2-3 and Figure 2-4, both of them consistent with their CV results. Figure 
2-5 shows the 
1
H-NMR spectrum of 0.20 mM Fe
II
(bpy)
2
(CN)
2
 with 1.0 mM DSS as 
reference in D
2
O, the same as Agrawala?s NMR report.
113
 The 
1
H-NMR spectrum of 1.0 
mM [Fe
II
(bipy)(CN)
4
]
2-
 in D
2
O is shown in Figure 2-6, which is consistent with its 
structure, except that one very small singlet at 8.43 ppm is unexplainable.  
 
Table 2-1. UV-visible absorbance and electrochemical characteristics of the iron 
complexes in aqueous solution 
?, M
-1
 cm
-1
 
Compounds Band ?
max
, nm
our work references 
E
1/2
, mV 
a
I 352 5258  
Fe
II
(bpy)
2
(CN)
2
 
II 522 5522 5800 
 
I 394 1224 1200 
[Fe
III
(bpy)
2
(CN)
2
]NO
3
 
II 544 241 200 
759 ? 2
I 346 3437 3470 
K
2
[Fe
II
(bpy)(CN)
4
] 
II 482 2825 2880 
 
I 375 1474 1450 
Li[Fe
III
(bpy)(CN)
4
] 
II 416 938 920 
537 ? 2 
a
 E
1/2
 vs NHE, at 22.0 
o
C and ? = 0.10 M. E
1/2
 = E
f
 
 38
 
 
 
Figure 2-1. The CV of 1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 M NaCF
3
SO
3
. With a glassy 
carbon disc as working electrode, an Ag/AgCl(s) reference electrode, and a Pt wire as a 
counter electrode. 
 39
 
Figure 2-2. The CV of 0.70 mM [Fe
III
(bpy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
. With a glassy 
carbon disc as working electrode, an Ag/AgCl(s) reference electrode, and a Pt wire as a 
counter electrode. 
 40
 
 
Figure 2-3. The OSWV of 1.0 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 M NaCF
3
SO
3
. With a 
glassy carbon disc as working electrode, an Ag/AgCl(s) reference electrode, and a Pt  
wire as a counter electrode, with E
p
 = 556 mV. 
 
 41
 
 
Figure 2-4. The OSWV of 0.70 mM [Fe
III
(bpy)(CN)
4
]
-
 in 0.10 M NaCF
3
SO
3
. With a 
glassy carbon disc as working electrode, an Ag/AgCl(s) reference electrode, and a Pt wire 
as a counter electrode, with E
p
 = 335 mV. 
 42
 
 
Figure 2-5. 
1
H-NMR spectrum of 0.20 mM Fe
II
(bpy)
2
(CN)
2
 in D
2
O, with 1.0 mM DSS as 
reference. 
 
 
 43
 
 
Figure 2-6. 
1
H-NMR spectrum of 1.0 mM [Fe
II
(bpy)(CN)
4
]
2-
 in D
2
O. 
 44
Stability of [Fe
III
(bpy)
2
(CN)
2
]
+
 in aqueous solution. Although cyanide ligands 
stabilize the +3 oxidation state of iron in [Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
, 
Papula et al. have found that [Fe
III
(bpy)
2
(CN)
2
]
+
 is photo-reduced to Fe
II
(bpy)
2
(CN)
2
 with 
a slow rate under a mercury arc lamp at pH 7.0.
109
 In our kinetic studies, the pH ranges 
from 2.95 to 7.89 for the reduction of [Fe
III
(bpy)
2
(CN)
2
]
+
. So it is of importance to know 
the stability of the Fe(III) complex in the dark under the pH studied. The UV-vis spectra 
of 0.50 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in 10.0 mM chloroacetate buffer (pH = 3.05) and 10.0 
mM phosphate buffer (pH = 7.98) were recorded on a HP 8453 diode-array 
spectrophotometer with a cycle time of 30 s in the range of 3600 s, respectively. The 
absorbance at 394 nm for the above solutions was constant over 3600 s. Therefore, it is 
safe to run kinetic experiments in the dark. 
Metal-ion catalysis and scavenger effect. The copper catalysis is ubiquitous in the 
oxidation of thiols by inert transition metal complexes.
50,52,53,102-104,114
 For the oxidation 
of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
, some preliminary 
experiments were performed on the SF-51 stopped-flow apparatus. For the oxidation of 
1.20 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 4.76 (10.0 mM acetate 
buffer) and 25.0 
o
C, in the presence of 0.092 M NaCF
3
SO
3
 (? = 0.10 M), the half-life is 
0.36 s. With the addition of 5.0 ?M CuSO
4
, the rate of the reaction is so fast that it occurs 
within the deadtime of the instrument (? 2 ms). Then 5.0 ?M Zn
2+
, Ni
2+
, Fe
2+
, Co
2+
, Mn
2+
 
and Mo(VI) were deliberately added to the above reactants (no Cu
2+
), respectively, the 
half-life is about 0.22 s for each reaction (Table 2-2). For the oxidation of 2.60 ? 10
-3
 M 
L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.96 (10.0 mM cacodylate buffer) and 
25.0 
o
C, in the presence of 0.093 M NaClO
4
, the half-life is 0.373 s; With the addition of 
 45
1.0 ?M CuSO
4
, the half-life of the reaction is only 0.0250 s, as shown in Table 2-3. The 
above results clearly demonstrate that traces of copper ions are good catalysts for the 
oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
 in aqueous solution. 
 
Table 2-2. Effect of metal cations on the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. 
With the addition of 0.092 M NaCF
3
SO
3
 and 10.0 mM acetate buffer, at 25.0 
o
C.
a 
[M
n+
], ?M 
t
1/2
, s pH of product  
0 0.36 4.76 
[Cu
2+
] = 5.0 
< 0.002
b 
4.76 
[Ni
2+
] = 5.0 0.25 4.75 
[Fe
2+
] = 5.0 0.20 4.78 
[Co
2+
] = 5.0 0.21 4.74 
[Mn
2+
] = 5.0 0.22 4.78 
[Mo
VI
] = 5.0 0.24 4.76 
[Zn
2+
] = 5.0 0.24 4.77 
a
 [L-cysteine]
0
 = 1.20 ? 10
-3
 M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 5.0 ? 10
-5
 M. 
b
 the dead time for 
SF-51 apparatus is 2 ms. 
 
 
 
Table 2-3. Effect of EDTA on the reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine. In the 
presence of 0.093 M NaClO
4
 and 10.0 mM cacodylate buffer (pH = 5.96), at 25.0 
o
C. 
 [Cu
2+
], ?M [EDTA], mM t
1/2
, s 
#1 0.0 0.0 0.373 
#2 1.0 0.0 0.0250 
#3 0.0 2.0 20.3 
#4 0.0 5.0 20.3 
#5 1.0 2.0 19.7 
#6 1.0 5.0 19.5 
a
 [L-cysteine]
0
 = 2.60 ? 10
-3
 M, [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M. 
 
 46
The copper-catalysis effect was successfully inhibited with the addition of 2,6-
pyridinedicarboxylic acid (dipic) in the oxidation of TGA and L-cysteine by  
[Mo(CN)
8
]
3-
.
52,103
 For the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
, the copper-
catalysis effect was completely suppressed by adding sodium dipicolinate dihydrate 
(Na
2
(C
7
H
3
NO
4
)?2H
2
O), which was synthesized by mixing dipic with 2 equivalent NaOH, 
then drying by rotary evaporation. The 
1
H-NMR spectrum of sodium dipicolinate is 
shown in Figure 2-7, in which the doublet and triplet are overlapped. The Differential 
Scanning Calorimetry (DSC) of sodium dipicolinate is shown in Figure 2-8. 
Thermodynamic calculation based on the enthalpy value at 100 
o
C confirms that it has 
two molecules of water. For the oxidation of 6.0 ? 10
-4
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.60 (10.0 mM chloroacetate buffer) and 25.0 
o
C, in the 
presence of 0.092 M NaCF
3
SO
3 
(? = 0.10 M), 1.0 mM dipic
2-
/dipic(1:1) is sufficient to 
control the pH and inhibit the copper ion catalysis, as shown in Table 2-4. For the later 
kinetic studies, 1.0 mM dipic
2-
/dipic (1:1, for pH < 4.0), and 1.0 mM dipic
2-
 (for pH > 4.0) 
were added to inhibit the copper-catalysis effect. 
 
Table 2-4. Effect of dipic
2-
 on the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 with L-cysteine. In the 
presence of 10.0 mM chloroacetate buffer (pH = 3.60), at ? = 0.10 M (0.092 M 
NaCF
3
SO
3
) and 25 
o
C.
a 
 [Cu
2+
], ?M [dipic
2-
/dipic], mM t
1/2
, s 
#1 0.0 0.0 0.770 
#2 0.0 1.0 20.0 
#3 5.0 0.0 0.006 
#4 5.0 1.0 20.0 
#5 5.0 2.0 20.0 
a
 [L-cysteine]
0
 = 6.0 ? 10
-4
 M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 5.0 ? 10
-5
 M. 
 47
 
Figure 2-7. 
1
H-NMR spectrum of 12.0 mM sodium dipicolinate in D
2
O. 
 
 48
 
 
Figure 2-8. DSC (Differential Scanning Calorimetry) of sodium dipicolinate dihydrate  
 
 49
However, for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
, the copper-catalysis 
is still dominant with the addition of 1.0 mM dipic
2-
. At pH 6.0, when the concentration 
of dipic is 2.0 mM or higher is the copper-catalysis well masked. When the pH of the 
solution is 7.80 or greater, copper-catalysis can?t be inhibited even with the addition of 
5.0 mM dipic
2-
, as shown in Table 2-5 and the highly non-first-order kinetics shown in 
Figure 2-9. Compared with [Mo
V
(CN)
8
]
3-
(E
1/2
 = 0.771 V
103
 vs NHE) and 
[Fe
III
(bpy)
2
(CN)
2
]
+
 (E
1/2
 = 0.759 V
111
 vs NHE), [Fe
III
(bpy)(CN)
4
]
-
 is a much weaker 
oxidant. To completely inhibit strong copper catalysis, EDTA (pH ? 10.0), and cyclam 
(pH > 10.0) were used in our kinetic studies. 
 
Table 2-5. Effect of dipic
2-
 on the reaction of 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 with L-
cysteine at various pHs. In the presence of dipic
2-
 and 10.0 mM buffer solution, at 25.0 
o
C. 
 [L-cysteine]
0
, 
mM 
[dipic
2-
], 
mM 
[Cu
2+
]
added
, 
?M 
[NaCF
3
SO
3
], 
M 
pH t
1/2
, s k
obs
, s
-1
 
2.54 1.0 0.0 0.093 5.99 14.9 ?
a 
2.54 2.0 0.0 0.093 6.00 17.0 0.0384 
2.51 2.0 1.0 0.093 6.01 16.3 ?
a 
2.54 5.0 0.0 0.093 6.02 17.0 0.0419 
2.50 5.0 1.0 0.093 6.03 18.0 0.0393 
2.50 7.5 0.0 0.093 6.04 16.0 0.0420 
2.50 10.0 0.0 0.093 6.05 18.0 0.0421 
0.470 5.0 0.0 0.056 7.80 1.63
b 
?
a 
1.02 2.0 0.0 0.075 9.89 0.017 40.2 
1.02 5.0 0.0 0.075 9.87 0.021 32.2 
1.02 5.0 0.25 0.075 9.88 0.0085 78.2 
1.02 5.0 1.0 0.075 9.88 0.0031 207 
a 
Non-first order; 
b
 irreproducible. 
 50
 
 
Figure 2-9. A reaction trace of oxidation of 4.7 ? 10
-4
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 7.80 (10.0 mM phosphate buffer), with the addition of 5.0 mM 
dipic
2-
. At ? = 0.10 M (0.056 M NaClO
4
) and 25.0 
o
C. 
 51
For the oxidation of 2.60 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at 
pH 5.96 (10.0 mM cacodylate buffer) and 25.0 
o
C, in the presence of 0.093 M NaClO
4
, 
2.0 mM EDTA is able to inhibit the copper ion catalysis, as shown in Table 2-3. 
Furthermore, iron cation is an omnipresent impurity in the chemical reagents, and it was 
reported that Fe-EDTA complex is a good catalyst for the oxidation of hydroxylamine by 
hexacyanoferrate(III).
115
 The above catalytic feature prompted us to test whether the Fe-
EDTA complex acts as a catalyst in the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. For 
the oxidation of 2.50 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.90 
(10.0 mM cacodylate buffer) and 25.0 
o
C, in the presence of 0.093 M NaCF
3
SO
3
, a series 
of experiments was performed on SF-51 apparatus. First, 1.0 mM EDTA was added; 
second, 1.0 mM EDTA and 1.0 ?M CuSO
4
 were added; third, 1.0 mM EDTA and 1.0 
?M Fe(NO
3
)
3
 were added. The rate constants for these three reactions shown in Table 2-6 
are about 0.030 s
-1
, ruling out catalysis by Fe-EDTA. To completely inhibit copper 
catalysis, 5.0 mM EDTA was employed in later kinetic studies at pH values no higher 
than 9.0. For the oxidation of 5.0 ? 10
-4
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
, 
the data in Table 2-7 indicate that copper ion catalysis is fully masked with the addition 
of 10.0 mM EDTA (pH = 10.0). However, our preliminary kinetic results demonstrate 
that EDTA fails to suppress copper catalysis at pH values higher than 10.0. The higher 
stability constant of 1,4,8,11-tetra-azacyclotetradecane (cyclam) with copper (II) than that 
of EDTA makes it a possible alternative chelating reagent in the above reaction.
116,117
 It is 
assumed that, at very high pH, the fully deprotonated cyclam are coordinated to copper(II) 
cation. The relativly high pK
a
 values of cyclam and its poor solubility at low pH limits its 
chelating efficiency at lower pH.
118,119
 Our preliminary kinetic studies (see Table 2-7) 
 52
confirm that copper-catalysis is inhibited with the addition of 1.0 mM cyclam at pH 
higher than 10.0. In our later kinetic studies, 5.0 mM EDTA (pH ? 9.00), 10.0 mM 
EDTA (9.00 < pH ? 10.0) and 1.0 mM cyclam (pH > 10.0) were added to inhibit the 
copper-catalysis. 
 
Table 2-6. Effects of Fe
3+
 and Cu
2+
 on reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine. In 
the presence of 0.093 M NaCF
3
SO
3
 and 10.0 mM cacodylate buffer, at 25.0 
o
C 
 [Fe
3+
], ?M 
[EDTA], 
mM 
pH t
1/2
, s k
obs
, s
-1
 
#1 0.0 1.0 5.89 26.0 0.0278 
#2 1.0 1.0 5.91 24.5 0.0299 
#3 0.0
b 
1.0 5.91 24.0 0.0301 
a
 [L-cysteine]
0
 = 2.50 ? 10
-3
 M, [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M, 
b
 [Cu
2+
] = 1.0 ?M 
 
 
Table 2-7. Effect of EDTA/cyclam on reaction of [Fe
III
(bpy)(CN)
4
]
-
 with L-cysteine. 
In the presence 10.0 mM EDTA or 1.0 mM cyclam, at 25.0 
o
C 
[EDTA], 
mM 
[Cu
2+
]
added
, 
?M 
[NaClO
4
], 
M 
[buffer], 
mM 
pH t
1/2
, s k
obs
, s
-1
 
10.0 0.0 0.0 0.0 10.21 0.28 2.21
 
10.0 0.25 0.0 0.0 10.22 0.25 ?
b
 
0.0
c 
0.0 0.099 0.0
d 
10.98 0.16 4.30
 
0.0
c
 0.50 0.099 0.0
d 
10.91 0.15 4.35 
0.0
c
 0.0 0.090 0.0
e 
11.90 0.28 5.19
 
0.0
c
 0.50 0.090 0.0
e 
11.85 0.24 5.19 
a
 [L-cysteine]
0
 = 5.0 ? 10
-4
 M, [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M, 
b
 Non pseudo-first order; 
c
 [cyclam] = 1.0 mM; 
d
 [NaOH] = 1.0 mM; 
e
 [NaOH] = 10.0 mM. 
 53
Note the concentration of copper cation in products? solution was determined by 
AA 240 Atomic Absorption Spectrometer. For the oxidation of 2.0 ? 10
-3
 M L-cysteine 
by 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at ? = 0.10 M and pH = 6.94, the concentration of 
copper cation in the product solution is 0.25 ?M; for the oxidation of 2.0 ? 10
-3
 M L-
cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at ? = 0.10 M and pH = 9.94, the 
concentration of copper cation in the product solution is 0.22 ?M, of which most of the 
copper is from sodium perchlorate and sodium carbonate buffer. According to the above 
preliminary kinetic studies, the copper cation can be completely scavenged even with the 
deliberate addition of 1.0 ?M CuSO
4
 (for reduction of [Fe
III
(bpy)
2
(CN)
2
]
+
) and 0.50 ?M 
CuSO
4
 (for reduction of [Fe
III
(bpy)(CN)
4
]
-
). However, the concentration of copper cation 
from AA measurements is much less than 1.0 ?M. So the copper catalysis is definitely 
inhibited with the addition of 1.0 mM dipic
2-
 for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
, 5.0 mM EDTA (pH ? 9.00), 10.0 mM EDTA (9.00 < pH ? 10.0) and 
1.0 mM cyclam (pH > 10.0) for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. 
The oxidation of EDTA and cyclam by [Fe
III
(bpy)(CN)
4
]
-
. According to the 
above preliminary studies, the addition of EDTA and cyclam slows down the copper-
catalyzed reaction dramatically. Does [Fe
III
(bpy)(CN)
4
]
-
 react with EDTA or cyclam? 
How fast is it? To investigate these questions, 1.0 mL of 1.0 ? 10
-4
 M [Fe
III
(bpy)(CN)
4
]
-
 
was mixed with 1.0 mL of 2.0 ? 10
-3
 M EDTA in one cuvette, in the presence of 10.0 
mM cacodylate (pH = 6.04), and the absorbance at 482 nm was recorded over time on 
HP-8453 diode-array spectrophotometer (optical cutoff filter: 455 nm). There is no 
reaction between [Fe
III
(bpy)(CN)
4
]
-
 and EDTA at pH 6.04 within 6000 s. Then the above 
 54
reaction was retested at higher pH. For the reaction between 1.0 ? 10
-3
 M EDTA and 5.0 
? 10
-5
 [Fe
III
(bpy)(CN)
4
]
-
, in the presence of 10.0 mM sodium carbonate buffer  
(pH = 10.20), the rate of the reaction is fairly slow, with a rate constant of 3.0 ? 10
-3 
 s
-1
. 
Finally, the reaction between 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 and 1.0 ? 10
-3
 M cyclam 
was tested at pH 10.1. The reaction is very slow, with the half-life longer than 2400 s. So 
it is safe to add 5.0 mM EDTA (pH < 9.00), 10.0 mM EDTA (9.0 < pH ? 10.0) and 1.0 
mM cyclam (pH > 10.0) to the flask containing cysteine for the oxidation of L-cysteine 
by [Fe
III
(bpy)(CN)
4
]
-
. 
Effect of Fe(II) and L-cystine on the rate of the reaction. 
(1). The oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. According to our study (in the 
product identification section), the products of the copper-catalysis inhibited reaction are 
Fe
II
(bpy)
2
(CN)
2
 and L-cystine. Figure 2-10 shows a typical kinetic trace for the oxidation 
of 1.25 ? 10
-3
 M L-cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, with 10.0 mM 
chloroacetate buffer (pH 3.20), ? = 0.10 M (0.093 M NaCF
3
SO
3
) and 1.0 mM dipic
2-
/dipic(1:1) at 25.0 
o
C. In Figure 2-10, it shows that the rate constant of the reaction is 
(1.391 ? 0.012) ? 10
-2
 s
-1
, with final absorbance of 0.1277, standard deviation (SD) of 
2.17 ? 10
-3
, and Durbin-Watson (DW) factor of 0.03/1.6. Numerically, a good fit is 
characterized by a low standard deviation and a relatively high value for the Durbin-
Watson (DW) factor. Therefore, the above fitting result demonstrates a significant 
deviation from pseudo-first-order kinetics, suggesting that product inhibition exists in the 
above reaction. Possible inhibition by Fe(II) was tested by adding different 
concentrations of Fe
II
(bpy)
2
(CN)
2
 to the reactants. For the oxidation of 1.25 ? 10
-3 
M L-
 55
cysteine by 2.5 ? 10
-5 
M [Fe
III
(bpy)
2
(CN)
2
]
+
, with 10.0 mM chloroacetate buffer (pH = 
3.20), 1.0 mM dipic
2-
/dipic (1:1) and at ? = 0.10 M (NaCF
3
SO
3
), 1.0 ? 10
-4 
M and 2.5 ? 
10
-5 
M Fe
II
(bpy)
2
(CN)
2
 were deliberately added to the above reactants, separately. The 
half-lives were 90.0 and 52.2 s, respectively, much longer than that without the addition 
of Fe
II
(bpy)
2
(CN)
2
, as shown in Table 2-8. The above results indicate that a step in the 
reaction is reversible, and not favorable in the forward direction at such a pH. In order to 
obtain perfect pseudo-first-order fitting, PBN was added to the reactants for all kinetic 
studies, which will be discussed in detail. Most importantly, the addition of 0.20 mM 
PBN counteracts the inhibition by Fe
II
(bpy)
2
(CN)
2
, as shown in Table 2-8. The reaction 
between L-cystine and [Fe
III
(bpy)
2
(CN)
2
]
+
 was tested. 2.5 ? 10
-4
 M L-cystine and 2.5 ? 
10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, in the presence of 10.0 mM phosphate buffer (pH = 7.94), 
were mixed with equal volume in the cuvette. The rate of the reaction is very slow, with 
the half-life at least 900 s. So the reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cystine is 
neglectable. 
 56
 
 
Figure 2-10. A reaction trace of oxidation of 1.25 ? 10
-3
 M L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20 (10.0 mM chloroacetate buffer), with the addition of 1.0 
mM dipic
2-
/dipic(1:1). At ? = 0.10 M (0.093 M NaCF
3
SO
3
) and 25.0 
o
C. 
 57
Table 2-8. The effect of Fe
II
(bpy)
2
(CN)
2
 and PBN on the oxidation of L-cysteine by  
[Fe
III
(bpy)
2
(CN)
2
]
+
 in anaerobic aqueous solution.
a 
[Fe
II
(bpy)
2
(CN)
2
]
0
, M [PBN], mM t
1/2
, s k
obs
, s
-1 
0.0 0.00 36.8 
?
c 
2.5 ? 10
-5
 
0.00 52.2 
?
c 
1.0 ? 10
-4
 
0.00 90.0 
?
c 
0.0 0.10 36.1
b 
1.92 ? 10
-2 
2.5 ? 10
-5
 
0.10 39.8
b 
1.74 ? 10
-2
 
1.0 ? 10
-4
 
0.10 52.5
b 
1.32 ? 10
-2
 
0.0 0.20 35.2
b 
1.97 ? 10
-2
 
2.5 ? 10
-5
 
0.20 38.1
b 
1.82 ? 10
-2
 
1.0 ? 10
-4
 
0.20 46.8
b 
1.48 ? 10
-2
 
a
 [L-cysteine]
0
 = 1.25 ? 10
-3
 M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 2.5 ? 10
-5
 M, pH = 3.20, ? = 0.10 
M (NaCF
3
SO
3
), [dipic
2-
/dipic] = 1.0 mM; ? = 0.10 M (NaCF
3
SO
3
); 
b
 t
1/2
 is calculated 
from k
obs
, t
1/2 
= ln2/k
obs
.
c
 non-first-order. 
 
(2). The oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. According to our study (in the 
product identification section), the products of the direct oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 are [Fe
II
(bpy)(CN)
4
]
2-
 and L-cystine. It is known that [Fe
III
(bpy)(CN)
4
]
-
 
has weaker oxidative capabilities than [Fe
III
(bpy)
2
(CN)
2
]
+
. Therefore, it is of importance 
to test the effect of [Fe
II
(bpy)(CN)
4
]
2-
 on the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. 
For the oxidation of 2.60 ? 10
-3 
M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
, with 
10.0 mM cacodylate buffer (pH = 5.98), 5.0 mM EDTA and 0.093 M NaClO
4
, 1.0 ? 10
-4 
M and 5.0 ? 10
-5 
M [Fe
II
(bpy)(CN)
4
]
2-
 were deliberately added to the above reactants. 
The half-lives were 28.0 and 25.4 s, very close to that without the addition of 
[Fe
II
(bpy)(CN)
4
]
2-
, as shown in Table 2-9. The inhibition by [Fe
II
(bpy)(CN)
4
]
2-
 at pH 5.98 
 58
will be discussed in detail (see discussion section). Then can L-cystine, one of the 
products of the reaction, react with [Fe
III
(bpy)(CN)
4
]
-
? Due to the poor solubility of L-
cystine 
120
 from pH values between 3.0 and 7.0, the following experiment was carried out 
at pH 8.00. 1.0 ? 10
-3
 M L-cystine and 1.0 ? 10
-4
 M [Fe
III
(bpy)(CN)
4
]
-
, in the presence of 
10.0 mM phosphate buffer (pH = 8.00), were mixed in equal volume, and the kinetic 
experiment was monitored on HP-8452 diode-array spectrophotometer. There is no 
reaction between L-cystine and [Fe
III
(bpy)(CN)
4
]
-
over 4800 seconds. 
 
Table 2-9. The effect of [Fe
II
(bpy)(CN)
4
]
2-
 on the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98. With 5.0 mM EDTA, and ? = 0.10 M (0.093 M NaClO
4
) 
[Fe(II)]
0
, ?M t
1/2
, s 
0.0 20.3 
50 25.4 
100 28.0 
a
 [L-cysteine]
0
 = 2.60 ? 10
-3
 M, [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M  
 
N-tert-butyl-?-phenylnitrone (PBN) effect on kinetics. 
(1). The oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. Our preliminary results 
have demonstrated that, at pH values less than 4.76, the rate of the reaction between L-
cysteine and [Fe
III
(bpy)
2
(CN)
2
]
+
 is non-first order even when 1.0 mM dipic was added to 
the reactants. The non-first-order kinetics was also observed by Hung and Stanbury in 
their studying the kinetics of the oxidation of thioglycolate by [Os(phen)
3
]
3+
. Perfect 
pseudo-first-order kinetics was obtained by adding 1.0 mM PBN to the reactants to 
scavenge the TGA radicals.
104
 The efficient scavenging of cysteine thiyl radical by PBN 
 59
was also observed by Graceffa.
121
 Here, a series of kinetic experiments was carried out 
with the addition of various concentrations of PBN at different pHs. For the oxidation of 
1.25 ? 10
-3 
M L-cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, with 10.0 mM chloroacetate 
buffer (pH 3.20), ? = 0.10 M (0.093 M NaCF
3
SO
3
) and 1.0 mM dipic
2-
/dipic(1:1) at 25.0 
o
C, the rate of the reaction is non-pseudo-first-order, as shown in Figure 2-10; with the 
addition of 0.10 mM PBN, the rate of reaction follows pseudo-first-order kinetic behavior, 
as shown in Figure 2-11, and the rate constant increases by a factor of 1.4. Then 0.20, 
0.50 and 1.0 mM PBN was added to the above reactants separately. The rate constants 
were the same as that for the addition of 0.10 mM PBN. The reaction has the same 
behavior at pH 4.62. For the oxidation of 1.00 ? 10
-3 
M L-cysteine by 2.5 ? 10
-5 
M 
[Fe
III
(bpy)
2
(CN)
2
]
+
, with 10.0 mM acetate buffer (pH 4.62), ? = 0.10 M (0.092 M 
NaCF
3
SO
3
) and 1.0 mM dipic
2-
/dipic(1:1) at 25 
o
C, the rate of the reaction is non-pseudo-
first-order, and the rate constant increases from 0.310 to 0.360 s
-1
 after the addition of 1.0 
mM PBN. However, the rate of the reaction decreases with the addition of PBN at pH 
7.62. For the oxidation of 2.50 ? 10
-4
 M L-cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, 
with 10.0 mM phosphate buffer (pH = 7.62), ? = 0.10 M (0.075 M NaCF
3
SO
3
) and 1.0 
mM dipic
2-
 at 25.0 
o
C, the rate of the reaction is pseudo-first-order; the rate constant 
decreases from 75.0 to 71.0 s
-1
after the addition of 0.10 mM PBN. Unexpectedly, the rate 
constant decreases with the increase of PBN. When 7.5 mM PBN was added to the above 
reaction system, the rate constant decreased by 32.5 %. All of the results are shown in 
Table 2-10.  
 
 60
 
Figure 2-11. A reaction trace of oxidation of 1.25 ? 10
-3
 M L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20 (10.0 mM chloroacetate buffer), with 1.0 mM dipic
2-
/dipic(1:1) and 0.10 mM PBN. At ? = 0.10 M (0.093 M NaCF
3
SO
3
) and 25.0 
o
C. 
 61
Table 2-10. The effect of PBN on the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
.
a 
[L-cysteine]
0
, 
mM 
[PBN], 
mM 
Buffer 
solution 
pH of 
products
k
obs
, s
-1
 t
1/2
, s 
1.25 0.0
 
chloroacetate 3.20 0.0139
b 
36.8 
1.25 0.10
 
chloroacetate 3.22 0.0194 35.7
c 
1.25 0.20
 
chloroacetate 3.21 0.0196 35.7
c 
1.25 0.50
 
chloroacetate 3.21 0.0199 35.4
c 
1.25 1.0
 
chloroacetate
 
3.24 0.0205 33.8
c
 
1.00 0.0 acetate 4.61 0.310
b 
2.00 
1.00 1.0 acetate 4.61 0.360 1.92
c 
1.00 2.0 acetate 4.62 0.367 1.89
c 
0.25 0.0 phosphate 7.62 75.0 0.0092
 c 
0.25 0.10 phosphate 7.62 71.0 0.0098
 c 
0.25 0.50
 
phosphate 7.62 63.0 0.011
 c
 
0.25 1.0
 
phosphate 7.62 56.0 0.012
 c
 
0.25 2.0 phosphate 7.61 51.3 0.014
 c
 
0.25 5.0 phosphate 7.62 51.5 0.013
 c
 
0.25 7.5 phosphate 7.61 50.6 0.014
 c
 
a 
[Fe(III)]
0
 = 25.0 ?M; [dipic
2-
] = 1.0 mM; ? = 0.10 M (NaCF
3
SO
3
); [buffer] = 10.0 mM. 
b 
Fitting is not perfect. 
c
 t
1/2
 is calculated from k
obs
, t
1/2
 = ln2/k
obs
. 
 
In the absence of PBN, it has been found that the rate constant decreases with the 
addition of Fe
II
(bpy)
2
(CN)
2
 at lower pH vlaues. With the addition of PBN, the inhibition 
by Fe
II
(bpy)
2
(CN)
2
 becomes insignificant. Here, for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20, in the presence of various concentrations of Fe(II), 0.10 
and 0.20 mM PBN was added to the reaction systems ([L-cysteine]
0
 = 1.25 ? 10
-3 
M; 
[Fe(III)]
0
 = 2.5 ? 10
-5
 M; [chloroacetate] = 10.0 mM; [dipic
2-
/dipic] = 1.0 mM; ? = 0.10 
 62
M (NaCF
3
SO
3
); [Fe(II)]
0
 = 2.5 ? 10
-5
 or 1.0 ? 10
-4 
M), respectively. The results shown in 
Table 2-8 indicate that 0.20 mM PBN is high enough to counteract the inhibition by 
Fe
II
(bpy)
2
(CN)
2
. (In the presence of 0.20 mM PBN, for [Fe(II)]
0
/[Fe(III)]
0
 = 0, the rate 
constant is 1.97 ? 10
-2
 s
-1
; for [Fe(II)]
0
/[Fe(III)]
0
 = 1, the rate constant is 1.82 ? 10
-2
 s
-1
.) 
Can PBN react with [Fe
III
(bpy)
2
(CN)
2
]
+
 or L-cysteine? The kinetic experiments 
were run by mixing equal volume of 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 with 1.0 ? 10
-2
 M 
PBN, in the presence of 1.0 mM dipic
2-
 and 10.0 mM phosphate buffer (pH = 7.56). The 
rate of the reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and PBN is rather slow, with the half-life 
over 900 s. Recently, Potapenko et al. reported that 5-diethoxyphosphoryl-5-methyl-1-
pyrroline N-oxide (DEPMPO) is nucleophilically attacked by L-cysteine to form 
DEPMPO hydroxylamine derivatives with equilibrium constant of 0.03 M
-1
.
122
 Then 1.50 
mL 0.050 mM PBN was spectrophotometrically titrated by 10.0 mM L-cysteine at pH 
7.60. The absorbance of PBN at 287 nm decreases with the dilution of PBN by L-cysteine. 
Therefore, it is safe to add small amounts of PBN to the reaction to scavenge the cysteine 
radicals that are produced in the reaction.  
Based on the above experimental results, 0.20 mM PBN was added to the reactants 
for all kinetic studies. 
(2). The oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. For the oxidation L-
cysteine by [Fe
III
(bpy)(CN)
4
]
-
, the pH for the kinetic experiments was selected from 6.00 
to 12.0. Figure 2-12 shows that it follows perfect pseudo-first-order behavior for the 
reaction between 2.60 ? 10
-2 
M L-cysteine and 5.0 ? 10
-5 
M [Fe
III
(bpy)(CN)
4
]
-
 (5.0 mM 
EDTA and 0.066 M NaClO
4
), at ? = 0.10 M and pH 6.09, which demonstrates that no 
PBN need be added to the reactants. 
 63
 
 
Figure 2-12. A reaction trace of oxidation of 2.60 ? 10
-2
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 6.09 (10.0 mM cacodylate buffer), with 5.0 mM EDTA. At ? = 
0.10 M (0.066 M NaClO
4
) and 25.0 
o
C. 
 64
Product identification and stoichiometry.  
(1). The oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. Without the addition of 
PBN. The products of the non-copper-catalyzed reaction were identified by UV-vis and 
1
H-NMR spectra. Figure 2-13 shows UV-vis spectra for the oxidation of 2.5 ?10
-4
 M L-
cysteine by 2.5 ?10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 with 10.0 mM cacodylate buffer (pH = 6.02), 
? = 0.10 M (0.093 M NaCF
3
SO
3
) and 1.0 mM dipic
2-
 at room temperature. At the 
beginning of the reaction, there are two peaks at 394 and 544 nm, which correspond to 
the UV-vis absorbance of [Fe
III
(bpy)
2
(CN)
2
]
+
. Within the time of the reaction, these two 
peaks disappeared, and two new peaks appeared. After 2.5 s of reaction, the shape of the 
curve did not change. Comparing it with the UV-vis spectrum of pure Fe
II
(bpy)
2
(CN)
2
 
(Figure 2-14), it very strongly suggests that Fe
II
(bpy)
2
(CN)
2
 is one of the products of the 
reaction. Furthermore, the well-defined isosbestic point at 339 nm in Figure 2-13 
suggests that there is no accumulation of any long-lived intermediate during the reaction. 
The identity of the sulfur-containing product was determined by 
1
H-NMR spectroscopy, 
from which the formation of Fe
II
(bpy)
2
(CN)
2
 was also observed
.
 A sample for 
1
H-NMR 
spectra was prepared as follows: 0.36 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 and 2.0 mM L-cysteine 
were equally mixed in anaerobic D
2
O, in the presence of 1.0 mM dipic
2-
 and 1.0 mM 
DSS at pH 8.78 (adjusted by 1.0 M NaOD). The 
1
H-NMR spectra shown in Figure 2-15 
clearly reveals that L-cystine (3.91 (dd, J = 7.9, 4.1; 2H), 3.28, (dd, J = 14.6, 4.1; 2H), 
3.07, (dd, J = 14.5, 8.0; 2H)) and Fe
II
(bpy)
2
(CN)
2
 (four triplets and four doublets, 
chemical shift from 7.20 to 9.40 ppm) are the products of the reaction. Unlike the 
reaction of L-cysteine with octacyanomolybdate(V),
52
 the 
1
H-NMR spectra for the 
oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 provided no evidence for the formation of 
 65
L-cysteine-sulfinate. The stoichiometry of the reaction was determined by the 
quantitative analysis of the
 1
H-NMR spectra of the product, with ?[Fe(II)]/?[L-cystine]
,
 
is 1.5 ? 0.1. 
300 400 500 600 700
0.00
0.05
0.10
0.15
0.20
0 s
4 s
Wavelength, nm
A
b
s
o
r
b
an
ce
 
Figure 2-13. The UV-vis spectra for the oxidation of 2.5 ? 10
-4
 M L-cysteine by  
2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 6.02 (10.0 mM cacodylate buffer), with 1.0 mM 
dipic
2-
. At ? = 0.10 M (0.093 M NaCF
3
SO
3
) and room temperature. (time interval: 0.10 s) 
 66
 
300 400 500 600 700
0.0
0.4
0.8
1.2
1.6
0.10 mM Fe(III)
0.10 mM Fe(II)
Wavelength (nm)
A
b
s
o
r
b
an
ce
 
Figure 2-14. UV-vis spectra of 1.0 ? 10
-4
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 (---) and Fe
II
(bpy)
2
(CN)
2
 
(? ) in aqueous solution. 
 67
 
 
Figure 2-15a. 
1
H-NMR spectrum of the product for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
  
with L-cysteine in D
2
O, with 1.0 mM dipic
2-
 and 1.0 mM DSS, at pH 8.78 (adjusted by 
1.0 M NaOD). [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.18 mM, [L-cysteine]
0
 = 1.0 mM. 
 
dipic
2-
 68
 
 
Figure 2-15b. 
1
H-NMR spectrum of the product for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
  
with L-cysteine in D
2
O, with 1.0 mM dipic
2-
 and 1.0 mM DSS, at pH 8.78 (adjusted by 
1.0 M NaOD). [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.18 mM, [L-cysteine]
0
 = 1.0 mM. 
 
 69
The stoichiometry of the reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in the 
presence of dipic
2- 
was also determined by spectrophotometric titration. Titration of 2.0 
mL of 1.04 ? 10
-4
 M L-cysteine by 5.09 ? 10
-3
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 7.61 was 
monitored at 522 nm, with 1.0 mM dipic
2-
 and 10.0 mM phosphate buffer. The 
consumption ratio of ?[Fe(II)]/?[L-cysteine] was calculated from the titration curve that 
is shown in Figure 2-16, with 1.12 ? 0.02 for ?[Fe(II)]/?[L-cysteine]. 
The stoichiometry of the reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in the 
presence of dipic
2- 
was determined by spectrophotometric analysis as well. Equal 
volumes of 0.40 mM deaerated L-cysteine and 0.20 mM [Fe
III
(bpy)
2
(CN)
2
]
+
, with 1.0 
mM dipic
2-
 and 10.0 mM phosphate buffer (pH = 6.94), were mixed together in a 
bubbling flask for 10.0 minutes, then the product solution was transferred to a stopcocked 
cuvette. The concentration of Fe
II
(bpy)
2
(CN)
2
 was determined from its characteristic 
absorbance at 522 nm, and the concentration of L-cysteine before and after the reaction 
was determined by Ellman?s reagent.
105,106
 The stoichiometric ratio,  
?[Fe(II)]/?[L-cysteine], was calculated as 1.002 ? 0.002, similar to results obtained from 
spectrophotometric titrations. The above results imply that the overall reaction is: 
 
2 [Fe
III
(bpy)
2
(CN)
2
]
+
 + 2 HSCH
2
CH(NH
3
+
)CO
2
-
 ? 
         2 Fe
II
(bpy)
2
(CN)
2
 + 
-
O
2
CCH(NH
3
+
)CH
2
SSCH
2
CH(NH
3
+
)CO
2
-
 + 2H
+
    (2-1-a) 
 70
 
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
End point
Volume of Fe(III), mL
A
522
 
Figure 2-16. Spectrophotometric titration of L-cysteine with [Fe
III
(bpy)
2
(CN)
2
]
+
, in the 
absence of PBN. Absorbance at 522 nm is shown as a function of the volume of Fe(III) 
added. pH =7.61, [L-cysteine]
0
 = 1.04 ? 10
-4
 M, n
0
(L-cysteine) = 0.00208 mmol, [Fe(III)] 
= 5.09 ? 10
-4
 M, and the initial volume of L-cysteine is 2.0 mL.  
 71
In the presence of PBN. The products for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
 in anaerobic aqueous solution, in the presence of PBN, are identified 
by their 
1
H-NMR spectra in D
2
O. Figure 2-17 shows the 
1
H-NMR spectrum of the 
product of 0.20 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 and 1.0 mM L-cysteine in D
2
O, in the presence 
of 1.0 mM dipic and 0.10 mM PBN at pH 3.10. From 7.19 to 9.40 ppm, there are eight 
peaks with the same intensity, confirming the formation of Fe
II
(bpy)
2
(CN)
2
. The chemical 
shifts at 4.00 and 3.09 ppm are ascribed to L-cysteine. Comparing Figure 2-17 with 
Figure 2-18, the chemical shifts at 8.61 and 8.41 ppm in Figure 2-17 are assigned as PBN 
and overlap of PBN with dipic
2-
. The ratio of the integrated intensity of the chemical shift 
at 7.20 (2H in Fe
II
(bpy)
2
(CN)
2
) and 1.56 ppm (9H in tert-butyl in PBN) in the above 
1
H-NMR spectrum demonstrates that around 40 % of PBN was consumed in the reaction. 
The consumption of PBN was thought to scavenge the cysteine thiyl radical, and it is 
found that the PBN adduct of cysteine thiyl radical is not very stable.
121
 So the two 
singlets (1.33 and 2.05 ppm) are ascribed to the decomposition of the PBN-cysteine 
radical adduct. Moreover, no L-cystine was detected in the 
1
H-NMR spectrum. The 
stoichiometric ratio for the reaction was obtained from the integrated intensities of the 
1
H-NMR spectrum, from which ?[L-cysteine]/?[Fe(III)] = 1.0 ? 0.2. 
According to PBN effect on kinetics, the rate for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 7.61 (see Table 2-10) decreases with the addition of PBN. 
Furthermore, L-cystine is more soluble in alkaline solution,
120
 so it is significant to 
identify the products and determine the stoichiometric ratio in alkaline media. The 
1
H-
NMR spectrum for the oxidation of 1.0 mM L-cysteine by 0.20 mM [Fe
III
(bpy)
2
(CN)
2
]
+
 
in D
2
O, in the presence of 1.0 mM dipic
2-
, 0.50 mM PBN and 1.0 mM DSS at pH 8.70, is 
 72
shown in Figure 2-19, which indicates that the products of the reaction are 
Fe
II
(bpy)
2
(CN)
2
 and L-cystine. The concentration of PBN decreases by 16 % during the 
reaction, much less than that of PBN in acidic solution. Combined with the above kinetic 
results, it implies that PBN scavenges more cysteine thiyl radical at lower pH values. The 
stoichiometric ratio was obtained from the integrated intensities of the 
1
H-NMR spectrum, 
from which ?[Fe(II)]/?[L-cystine] = 1.82 ? 0.07. 
 
 
Figure 2-17. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine  
in D
2
O, with 1.0 mM dipic
2-
 and 0.10 mM PBN, at pH 3.10 (adjusted by 1.0 M NaOD). 
[Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.20 mM, [L-cysteine]
0
 = 1.0 mM. 
 
 73
 
Figure 2-18. 
1
H-NMR spectrum of 10.0 mM PBN in D
2
O 
 
CH N
+
C(CH
3
)
3
O
?
ac
c
c a
b
CH N
+
C(CH
3
)
3
O
?
ac
c
c a
b
a
b
c
 74
 
 
Figure 2-19a. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in 
D
2
O, with 1.0 mM dipic
2-
, 0.50 mM PBN and 1.0 mM DSS, at pH 8.70 (adjusted by 1.0 
M NaOD). [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.20 mM, [L-cysteine]
0
 = 1.0 mM. 
 
 75
 
 
Figure 2-19b. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in 
D
2
O, with 1.0 mM dipic
2-
, 0.50 mM PBN and 1.0 mM DSS, at pH 8.70 (adjusted by 1.0 
M NaOD). [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.20 mM, [L-cysteine]
0
 = 1.0 mM. 
? : Fe(bpy)
2
(CN)
2
; ? : PBN; ?: overlap of dipic
2-
 and PBN;  
? : overlap of PBN and Fe(bpy)
2
(CN)
2
. 
 
?
?
? ? ? ??
?
?
 76
 
 
Figure 2-19c. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in 
D
2
O, with 1.0 mM dipic
2-
, 0.50 mM PBN and 1.0 mM DSS, at pH 8.70 (adjusted by 1.0 
M NaOD). [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 0.20 mM, [L-cysteine]
0
 = 1.0 mM. 
? : RSSR; ? : RSH; ?: DSS; ?: PBN; ? : RSH/PBN adduct; ? : overlap of RSH and 
DSS. 
 77
The stoichiometry of the reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and L-cysteine in the 
presence of 1.0 mM dipic
2-
 and 1.0 mM PBN was also determined by spectrophotometric 
titration. Titration of 2.0 mL of 1.04 ? 10
-4
 M L-cysteine by 4.92 ? 10
-3
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 7.61 was monitored at 522 nm, with 1.0 mM dipic
2-
, 1.0 mM 
PBN and 10.0 mM phosphate buffer. The stoichiometric ratio was determined from the 
titration curve (Figure 2-20), with 1.13 ? 0.09 for ?[Fe(II)]/?[L-cysteine], consistent with 
the results from 
1
H-NMR. Therefore, the products and stoichiometric ratio of the reaction 
are the same even when PBN is added.  
(2). The oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. The products of the non-
copper-catalyzed reaction were identified by 
1
H-NMR spectra. A sample for running 
1
H-
NMR spectra was prepared as follows: equal volume of 0.80 mM [Fe
III
(bpy)(CN)
4
]
-
 and 
4.0 mM L-cysteine were mixed in anaerobic D
2
O, in the presence of 1.0 mM cyclam and 
2.0 mM DSS at pH 11.1 (adjusted by 0.10 M NaOD). The 
1
H-NMR spectrum shown in 
Figure 2-21 clearly shows that L-cystine (two doublets of doublets, (3.62, dd, J = 7.6, 4.4; 
2H), (3.12, dd, J = 16, 4.4; 2H), and [Fe
II
(bpy)(CN)
4
]
-
 (two triplets and two doublets,
113
 
chemical shift from 7.47 to 9.34 ppm) are the products of the reaction, consistent with 
that of the oxidation L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. Furthermore, the third doublet of 
doublets of 
1
H-NMR spectra of L-cystine overlap with that of cyclam. Unlike the 
oxidation of L-cysteine by octacyanomolybdate(V)
52
, the 
1
H-NMR spectra for the 
oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 provided no evidence for the formation of 
L-cysteine-sulfinate. The stoichiometry of the reaction was determined by the 
quantitative analysis of the
 1
H-NMR spectra of the product, with ?[Fe(II)]/?[L-cystine] = 
1.97 ? 0.03. 
 78
 
0.00 0.25 0.50 0.75 1.00
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
End point
Volume of Fe(III), mL
A
522
 
Figure 2-20. Spectrophotometric titration of L-cysteine with [Fe
III
(bpy)
2
(CN)
2
]
+
, in the 
presence of 1.0 mM PBN. Absorbance at 522 nm is shown as a function of the volume of 
Fe(III) solution added. pH = 7.57, [L-cysteine]
0
 = 1.04 ? 10
-4
 M, n
0
(L-cysteine) = 2.09 ? 
10
-4 
mmol, [Fe(III)] = 0.492 mM, and the initial volume of L-cysteine was 2.0 mL. 
 79
 
 
Figure 2-21a. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)(CN)
4
]
-
 and L-cysteine in 
D
2
O, in the presence of 1.0 mM cyclam and 2.0 mM DSS, at pH 11.1 (adjusted by 0.10 
M NaOD). [Fe
III
(bpy)(CN)
4
-
]
0
 = 0.40 mM, [L-cysteine]
0
 = 2.0 mM. 
 
 80
 
 
Figure 2-21b. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)(CN)
4
]
-
 and L-cysteine in 
D
2
O, in the presence of 1.0 mM cyclam and 2.0 mM DSS, at pH 11.1 (adjusted by 0.10 
M NaOD). [Fe
III
(bpy)(CN)
4
-
]
0
 = 0.40 mM, [L-cysteine]
0
 = 2.0 mM. 
 
 81
 
 
Figure 2-21c. 
1
H-NMR spectrum of the product of [Fe
III
(bpy)(CN)
4
]
-
 and L-cysteine in 
D
2
O, in the presence of 1.0 mM cyclam and 2.0 mM DSS, at pH 11.1 (adjusted by 0.10 
M NaOD). [Fe
III
(bpy)(CN)
4
-
]
0
 = 0.40 mM, [L-cysteine]
0
 = 2.0 mM. 
 82
The stoichiometry of the reaction between [Fe
III
(bpy)(CN)
4
]
-
 and L-cysteine in the 
presence of EDTA was determined by spectrophotometric analysis as well. Two 
experiments were performed. Equal volume of 0.50 mM deaerated L-cysteine and 0.30 
mM [Fe
III
(bpy)(CN)
4
]
-
, with 5.0 mM EDTA and 10.0 mM ammonia buffer (pH = 8.78), 
was mixed together in a bubbling flask for 10.0 minutes, then the product solution was 
transferred to a stopcocked cuvette. The concentration of [Fe
II
(bpy)(CN)
4
]
2-
 was 
determined from its characteristic absorbance at 482 nm, and the concentration of L-
cysteine before and after reaction was determined by Ellman?s reagent.
105,106
 The 
stoichiometric ratio, ?[Fe(II)]/?[L-cysteine], was calculated as 1.28 ? 0.02. The above 
result implies that L-cysteine-sulfinate could be formed when the concentration of L-
cysteine is slightly higher than that of [Fe
III
(bpy)(CN)
4
]
-
. But in later kinetic experiments, 
the concentration of L-cysteine is at least ten times that of [Fe
III
(bpy)(CN)
4
]
-
. Equal 
volumes of 1.0 mM deaerated L-cysteine and 0.30 mM [Fe
III
(bpy)(CN)
4
]
-
, with 5.0 mM 
EDTA and 10.0 mM ammonia buffer (pH = 8.78), were mixed together in a bubbling 
flask for 10.0 minutes, and the product solution was transferred to a stopcocked cuvette. 
The concentration of [Fe
II
 (bpy)(CN)
4
]
2-
 was determined from its characteristic 
absorbance at 482 nm, and the concentration of L-cysteine before and after the reaction 
was determined by Ellman?s reagent.
105,106
 The stoichiometric ratio was calculated as 
1.05 ? 0.04 for ?[Fe(II)]/?[L-cysteine], the same as the results from 
1
H-NMR spectra.  
Based on the above results, the overall reaction is as follows: 
 
2 [Fe
III
(bpy)(CN)
4
]
-
 + 2 HSCH
2
CH(NH
3
+
)CO
2
-
 ? 
            2 [Fe
II
(bpy)(CN)
4
]
2-
 + 
-
O
2
CCH(NH
3
+
)CH
2
SSCH
2
CH(NH
3
+
)CO
2
-
 + 2H
+
     (2-1-b) 
 83
Kinetics.  
(1). The oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. The reduction of 2.5 ?  
10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 by various concentrations of L-cysteine, in the presence of 1.0 
mM dipic
2-
, 0.20 mM PBN, and pH from 3.0 to 7.87, was investigated at 25.0 
o
C and at 
constant ionic strength (? = 0.10 M, NaCF
3
SO
3
). The pseudo-first-order rate constants 
(k
obs
, s
-1
) were evaluated by exponential fits of the kinetic curves, and replicate runs 
agreed to better than ? 5%. Figure 2-22 shows a kinetic trace of oxidation of 4.7 ? 10
-4
 M 
L-cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 5.99 (10.0 mM cacodylate buffer), 
with 1.0 mM dipic
2-
, 0.20 mM PBN, ? = 0.10 M (0.092 M NaCF
3
SO
3
), and at 25.0 
o
C. 
The rate of the reaction follows pseudo-first-order behavior, with a rate constant (k
obs
)
 
of 
3.84 s
-1
. A series of kinetic experiments was run by changing [L-cysteine]
tot
 (the total 
concentration of L-cysteine), with 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, 1.0 mM dipic
2-
, 10.0 
mM cacodylate buffer (pH = 6.01), 0.20 mM PBN, and ? = 0.10 M at 25.0 
o
C. The rate 
of the reaction follows pseudo-first-order behavior over 3 half-lives. The rate constant 
dependence on [L-cysteine]
tot
 is summarized in Table 2-11, and a plot of k
obs
 versus [L-
cysteine]
tot
 is shown in Figure 2-23. It is linear with slope of (7.25 ? 0.09) ? 10
3 
M
-1
s
-1
 
and an intercept of (0.30 ? 0.05) s
-1
. The intercept value is statistically significant. To 
determine if it is chemically important, lower concentrations of L-cysteine should be 
added in the kinetic experiments. The small non-zero intercept value will be omitted in 
our kinetic studies. So the rate of the reaction has a first-order dependence on [L-
cysteine]
tot.
 
k
obs
 = k[L-cysteine]
tot
                                                                                   (2-2) 
 84
 
 
Figure 2-22. Reaction trace of the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 
5.99 (10.0 mM cacodylate buffer), in the presence of 0.20 mM PBN and 1.0 mM dipic
2-
. 
At ? = 0.10 M (0.093 M NaCF
3
SO
3
) and 25.0 
o
C, using first-order fitting. 
[Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 2.5 ? 10
-5
 M, [L-cysteine]
0
 = 4.7 ? 10
-4
 M. 
 85
0.0000 0.0025 0.0050 0.0075
0
10
20
30
40
50
[L-cysteine]
tot
, M
k
obs
, s
-1
 
Figure 2-23. k
obs
 vs [RSH]
tot
 for the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 
6.00 (10.0 mM cacodylate buffer), with 1.0 mM dipic
2-
 and 0.20 mM PBN. At ? = 0.10 
M (0.093 M NaCF
3
SO
3
), and 25.0 
o
C, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 2.5 ? 10
-5
 M. (Weight by 
1/Y
2
, 1/Y
2
 = ?(Y
data
- Y
curve
)
2
/Y
data
2
). 
 86
Table 2-11. [L-cysteine] dependence for the oxidation of L-cysteine by 2.5 ? 10
-5
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 6.00 (10.0 mM cacodylate buffer). In the presence of 1.0 mM 
dipic
2-
, at ?= 0.10 M (NaCF
3
SO
3
) and 25.0 
o
C. 
[L-cysteine]
0
, M k
obs
, s
-1
 
0.000236 1.99 
0.000471 3.82 
0.00188 13.4 
0.00288 21.0 
0.00481 35.2 
0.00576 42.8 
0.00721 53.3 
 
The kinetic pH dependence was revealed in a series of kinetic runs with the pH 
varied from 2.95 to 7.89 (the rate of the reaction is too fast when pH is higher than 7.89, 
and out of limit of the SF-51 instrument), the values of k
obs
 are summarized in Table 2-12. 
A plot of (log k
obs
/ [L-cysteine]
tot
) versus pH shown in Figure 2-24 indicates that the rates 
increase with the increase of pH. The rate pH dependence is clearly related to the three 
proton dissociation equilibrium constants of L-cysteine as follows,
117
 and it implies that 
sulfur is the reactive center in the oxidation of L-cysteine. 
 
HSCH
2
CH(NH
3
+
)CO
2
H          HSCH
2
CH(NH
3
+
)CO
2
-
 +
 
H
+
pK
a1
 = 1.90      (2-3)
HSCH
2
CH(NH
3
+
)CO
2
-
-
SCH
2
CH(NH
3
+
)CO
2
-
 + H
+ pK
a2
 = 8.18      (2-4)
-
SCH
2
CH(NH
3
+
)CO
2
- -
SCH
2
CH(NH
2
)CO
2
-
 + H
+ pK
a3
 = 10.30    (2-5)
 
 
 
 
 
 87
Table 2-12. pH-dependent kinetic data for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
. At ?= 0.10 M (NaCF
3
SO
3
) and 25.0 
o
C.
a
 
[L-cysteine]
0
, 
mM 
[Buffer], mM 
[CF
3
SO
3
Na], 
M 
pH k
obs
, s
-1
 
k
obs
/[L-cysteine]
0
, 
M
-1
 s
-1
 
10.3 [chloroacetate]=10 0.095 2.95 0.0953 9.23 
1.25 [chloroacetate]=10 0.092 3.21 0.0197 15.8 
10.3 [chloroacetate]=10 0.090 3.62 0.366 35.5 
2.06 [acetate]=10.0 0.096 4.06 0.203 98.5 
2.06 [acetate]=10.0 0.092 4.44 0.467 2.26  ? 10
2
 
2.48 [acetate]=10.0 0.092 4.85 1.42
 
5.73 ? 10
2
 
2.48 [acetate]=10.0 0.089 5.23 3.26 1.31 ? 10
3
 
2.48 [cacodylate]=10.0 0.094 5.69 8.74 3.52 ? 10
3
 
0.470 [cacodylate]=10.0 0.093 5.99 3.82 8.11 ? 10
3
 
0.990 [cacodylate]=10.0 0.093 6.46 19.1 1.93 ? 10
4
 
0.488 [phosphate]=10.0 0.075 6.98 34.7 7.11 ? 10
4
 
0.244 [phosphate]=10.0 0.075 7.40 41.6 1.70 ? 10
5
 
0.244 [phosphate]=10.0 0.075 7.57 57.3 2.35 ? 10
5
 
0.244 [phosphate]=10.0 0.075 7.89 106
 
4.33 ? 10
5
 
a 
[Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 25.0 ?M; [PBN] = 0.20 mM; [dipic
2-
] = 1.0 mM. 
 
In principle, the four different forms of L-cysteine can react with [Fe
III
(bpy)
2
(CN)
2
]
+
 
through kinetically distinguishable terms, as shown in Equation 2-6-a: 
 
tot
32121
2
1
3
3214213
2
a12
3
1
 obs
RSH][
]H[]H[]H[
]H[]H[]H[
?
?
?
?
?
?
+++
+++
=
+++
+++
aaaaaa
aaaaa
KKKKKK
KKKkKKkKkk
k              (2-6-a) 
 
where k
1
-k
4
 represent the reactivity of protonated, neutral, monoanionic, and dianionic L-
cysteine species, respectively. A nonlinear least-squares fit of the data in Table 2-12 to  
 88
 
Figure 2-24. Plot of log(k
obs
/[L-cysteine]
tot
) vs pH for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
, in the presence of 0.20 mM PBN and 1.0 mM dipic
2-
. At ? = 0.10 M  
and 25.0 
o
C. Solid line shows the fit to eq 2-8. (Weight by 1/Y
2
, 1/Y
2
 = ?(Y
data
- 
Y
curve
)
2
/Y
data
2
) 
 
 89
 
Equation 2-6-a yields (-21.3 ? 9.2) M
-1
 s
-1
 for k
1
. In the oxidation of L-cysteine by 
octacyanomolybdate, Hung et al.
52
 reported the rate constants, k
1
 to k
4
, with such order: 0 
? k
1
< k
2
 < k
3 
> k
4
, and the smaller k
4
 is thought due to the greater repulsion between 
dianionic L-cysteine and [Mo(CN)
8
]
3-
. Since [Fe
III
(bpy)
2
(CN)
2
]
+
, with + 1 charge, has 
very close standard potential as 
[Mo(CN)
8
]
3-
, the rate constants may be in such order: 0 ? 
k
1
< k
2
 < k
3 
< k
4
. Then Equation 2-6-a can be converted to: 
 
tot
32121
2
1
3
3214213
2
a12
 obs
]RSH[
]H[]H[]H[
]H[]H[
?
?
?
?
?
?
+++
++
=
+++
++
aaaaaa
aaaaa
KKKKKK
KKKkKKkKk
k                      (2-7-a) 
 
The data in Table 2-12 were fit to Equation 2-7-a by using nonlinear least-squares model, 
with pK
a1
, pK
a2
 and pK
a3
 held at the literature values of 1.90, 8.18 and 10.30.
117
 The rate 
constants k
2
, k
3
 and k
4
 that were obtained from the above fitting are (2.94 ? 0.24) M
-1
 s
-1
, 
(1.22 ? 0.02) ? 10
6
 M
-1
 s
-1
, and (6.30 ? 32.3) ? 10
6
 M
-1
 s
-1
, respectively. The large 
statistical error in the value of k
4
 implies that the contribution of dianionic L-cysteine is 
kinetically insignificant in such pH range. Then Equation 2-7-a can be simplified as: 
 
tot
32121
2
1
3
213
2
a12
 obs
]RSH[
]H[]H[]H[
]H[]H[
?
?
?
?
?
?
+++
+
=
+++
++
aaaaaa
aa
KKKKKK
KKkKk
k                        (2-8) 
 
Applying a nonlinear least-squares model, the data in Table 2-12 were refit by Equation 
2-8, as shown in Figure 2-24. The obtained rate constants k
2
 and k
3
 are (2.93 ? 0.23) and 
 90
(1.22 ? 0.02) ? 10
6
 M
-1
 s
-1
. These results demonstrate that Equation 2-8 accurately 
represents the data. 
(2). The oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. The reduction of 5.0 ?  
10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 by various concentrations of L-cysteine, in the presence of 5.0 
mM EDTA (pH ? 9.00), 10.0 mM EDTA (9.00 < pH ? 10.0) and 1.0 mM cyclam (pH > 
10.0), was investigated at 25.0 
o
C and at constant ionic strength (? = 0.10 M, NaClO
4
). 
The rate of the reaction follows pseudo-first-order behavior over 3 half-lives. The 
pseudo-first-order rate constants (k
obs
, s
-1
) were evaluated by exponential fits of the 
kinetic curves, and replicate runs agreed to better than ? 5%. Figure 2-25 shows a 
reaction trace of oxidation of 5.05 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 
at pH 6.94 (10.0 mM cacodylate buffer), with 5.0 mM EDTA, at ionic strength of 0.10 M 
(0.061 M NaClO
4
) and 25.0 
o
C. The rate of the reaction follows pseudo-first-order 
behavior, with a rate constant (k
obs
) of 0.556 s
-1
. A series of kinetic experiments was run 
by changing [L-cysteine]
tot 
(the total concentration of L-cysteine), with 5.0 ? 10
-5
 M 
Fe(III), 5.0 mM EDTA, and 10.0 mM cacodylate buffer (pH = 7.04), at ? = 0.10 M and 
25.0 
o
C. The rate constant dependence on [L-cysteine]
tot
 is summarized in Table 2-13, 
and plot of k
obs
 versus [L-cysteine]
tot
 is shown in Figure 2-26. It is linear with slope of 
(1.11 ? 0.02) ? 10
2 
M
-1
s
-1
 and intercept of (8.44 ? 26.1) ? 10
-4
 s
-1
. The intercept value is 
statistically zero. So the rate of the reaction has a first-order dependence on [L-cysteine]
tot. 
 
 91
 
 
Figure 2-25. A reaction trace of oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 at pH 7.04 
(10.0 mM cacodylate buffer), with 5.0 mM EDTA. At ? = 0.10M (0.061 M NaClO
4
) and 
25.0 
o
C. [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M, [L-cysteine]
0
 = 5.05 ? 10
-3
 M 
 
 92
Table 2-13. [L-cysteine] dependence for the oxidation of L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 6.92 (10.0 mM cacodylate). In the presence of 5.0 mM EDTA, at 
? = 0.10 M (0.061 M NaClO
4
) and 25.0 
o
C. 
[L-cysteine]
0
, mM pH k
obs
, s
-1
 
0.50 6.91 0.058 
0.98 6.87 0.102 
2.02 6.91 0.230 
3.64 6.90 0.428 
5.05 6.90 0.555 
7.35 6.92 0.819 
9.80 6.94 1.07 
12.3 6.95 1.36
 
 93
 
0.000 0.003 0.006 0.009 0.012 0.015
0.0
0.5
1.0
1.5
[RSH]
tot
, M
k
obs
, s
-1
 
Figure 2-26. k
obs
 vs [RSH]
tot
 for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 at pH 
6.92 (10.0 mM cacodylate buffer), in the presence of 5.0 mM EDTA. At ? = 0.10 M 
(0.061 M NaClO
4
), and 25.0 
o
C, [Fe
III
(bpy)(CN)
4
-
]
0
 = 5.0 ? 10
-5
 M.  
(Weight by 1/Y
2
, 1/Y
2
 = ?(Y
data
- Y
curve
)
2
/Y
data
2
). 
 94
The kinetic pH dependence was revealed in a series of kinetic runs with pH values 
from 6.09 to 11.9, the values of k
obs
 being summarized in Table 2-14. A plot of (log k
obs
/ 
[L-cysteine]
tot
) versus pH is shown in Figure 2-27. The rate of the reaction increases with 
the increase of pH. The rate pH-dependence is clearly related to the three proton 
dissociation equilibrium constants of L-cysteine.
117
 
 
Table 2-14. pH-dependent kinetic data for the oxidation of L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 in aqueous solution.
a
 
[L-cysteine]
0
, 
mM 
[EDTA], 
mM 
[NaClO
4
], 
M 
[buffer], 
mM 
pH k
obs
, s
-1
 
k
obs
/[L-cysteine]
0
, 
s
-1
 M
-1 
24.6 5.0 0.075 10.0 6.09 0.371 15.1 
9.70 5.0 0.063 10.0 6.44 0.406 41.8 
10.4 5.0 0.063 10.0 6.61 0.469 45.1 
5.05 5.0 0.061 10.0 6.90 0.555 110 
1.04 5.0 0.046 10.0 7.37 0.305 293 
0.47 5.0 0.041 10.0 7.72 0.300 638 
0.50 5.0 0.041 10.0 8.06 0.521 1.04 ? 10
3
 
0.52 5.0 0.060 10.0
 
8.59 0.939 1.80 ? 10
3
 
0.51 5.0 0.050 10.0
 
9.04 1.21 2.37 ? 10
3
 
0.49 10.0 0.034 0.0 9.56 1.75 3.57 ? 10
3
 
0.49 10.0
 
0.026 0.0 9.86 2.01
 
4.10 ? 10
3
 
0.52 0.0
b
 0.076 10.0 10.6 3.83 7.36 ? 10
3
 
0.52 0.0
b
 0.099 0.0
c 
11.0 4.30 8.26 ? 10
3
 
0.52 0.0
b
 0.090 0.0
d 
11.9 5.19 9.98 ? 10
3
 
a
 at 25.0 
o
C, ? = 0.10 M; 
b
 [cyclam] = 1.0 mM; 
c
 [NaOH] = 1.0 mM; 
d
 [NaOH] = 10.0 mM. 
 95
 
 
Figure 2-27. Plot of log(k
obs
/[L-cysteine]
tot
) vs pH for the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
. With 5.0 mM EDTA (pH from 6.0 to 9.0), 10.0 mM EDTA (pH from 
9.0 to 10.0) or 1.0 mM cyclam (pH > 10.0), at ionic strength of 0.10 M and 25.0 
o
C. Solid 
line shows the fit to eq 2-9. (Weight by 1/Y
2
, 1/Y
2
 = ?(Y
data
- Y
curve
)
2
/Y
data
2
) 
 96
tot
32121
2
1
3
321'421'3
2
a12'
3
1'
 obs
RSH][
]H[]H[]H[
]H[]H[]H[
?
?
?
?
?
?
+++
+++
=
+++
+++
aaaaaa
aaaaa
KKKKKK
KKKkKKkKkk
k        (2-6-b) 
 
tot
32121
2
1
3
321'421'3
2
a12'
 obs
]RSH[
]H[]H[]H[
]H[]H[
?
?
?
?
?
?
+++
++
=
+++
++
aaaaaa
aaaaa
KKKKKK
KKKkKKkKk
k                        (2-7-b) 
 
In principle, the four different forms of L-cysteine can react with [Fe
III
(bpy)(CN)
4
]
-
 
through kinetically distinguishable terms, as shown in Equation 2-6-a, 
where k
1
-k
4
 
represent the reactivity of protonated, neutral, monoanionic, and dianionic L-cysteine 
species, respectively. In the oxidation of L-cysteine by octacyanomolybdate, Hung et al. 
reported the rate constants, k
1
 to k
4
, with such order: 0 ? k
1
< k
2
 < k
3 
> k
4
, the smaller k
4
 
value is thought due to the greater repulsion between dianionic L-cysteine and 
[Mo(CN)
8
]
3-
.
52
 In our case, [Fe
III
(bpy)(CN)
4
]
-
 is a much weaker oxidant than 
[Mo(CN)
8
]
3-
,
 
so k
1?
 probably be close to zero and this term is omitted in the fitting. Then Equation 2-6-
b is simplified into Equation 2-7-b. The data in Table 2-14 were fit to Equation 2-7-b by 
using nonlinear least-squares model, with pK
a1
, pK
a2
 and pK
a3
 held at the literature values 
of 1.90, 8.18 and 10.30.
117
 The rate constants k
2?
, k
3?
 and k
4?
 were obtained from the above 
fitting are (-2.83 ? 2.74) M
-1
 s
-1
, (2.21 ? 0.22) ? 10
3
 M
-1
 s
-1
, and (1.04 ? 0.22) ? 10
4
  
M
-1
 s
-1
 respectively. Excluding both k
1?
 and k
2?
 terms still yields good fitting results, as 
shown in Figure 2-27, with k
3?
, k
4?
 being (2.04 ? 0.16) ? 10
3
 M
-1
 s
-1
, and (1.07 ? 0.25) ? 
10
4
 M
-1
 s
-1
. Therefore, Equation 2-9 represents the experimental data best. 
 
 97
tot
32121
2
1
3
321'421'3
 obs
]RSH[
]H[]H[]H[
]H[
?
?
?
?
?
?
+++
+
=
+++
+
aaaaaa
aaaaa
KKKKKK
KKKkKKk
k                 (2-9) 
 
Alkaline cation effect. The alkaline cation catalysis was recently observed for the 
oxidation of L-cysteine by [Mo
V
(CN)
8
]
3-
 in alkaline condition.
52
 In that reaction, there is 
strong electrostatic repulsion between the ?3 charge in [Mo
V
(CN)
8
]
3-
 and the negative 
charge in the dianion form of L-cysteine. Since [Fe
III
(bpy)(CN)
4
]
-
 has ?1 charge, does it 
have alkaline cation effect in the oxidation of L-cysteine? For the oxidation of 5.1 ? 10
-4
 
M L-cysteine by 5.01 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 11.0, in the presence of 1.0 mM 
cyclam and 1.0 mM NaOH, 0.099 M LiClO
4
 and NaClO
4
 were added separately. The 
pseudo-first-order constants are 4.45 and 4.30 s
-1
. It implies that no alkaline cation 
catalysis exists for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 at pH 11.0. 
 
Discussion 
In anaerobic aqueous solution, the oxidation of L-cysteine by the typical outer-
sphere oxidants, [Fe
III
(bpy)
2
(CN)
2
]
+
 and 
[Fe
III
(bpy)(CN)
4
]
-
, are strongly catalyzed by 
trace amount of Cu
2+
. The copper-catalysis is well inhibited with the addition of 1.0 mM 
dipic
2-
 for oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
. The direct oxidation of L-
 
cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 was studied with the addition of 1.0 mM dipic
2-
. For 
[Fe
III
(bpy)(CN)
4
]
-
, 
dipic
2-
 does not completely inhibit copper catalysis, but EDTA 
(pH ? 
10.0) and cyclam (pH > 10.0) successfully suppress copper catalysis. The direct oxidation 
of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 was studied with the addition of 5.0 mM EDTA (pH ? 
 98
9.00), 10.0 mM EDTA (9.00 < pH ? 10.0) and 1.0 mM cyclam (pH > 10.0)
. The rates of 
the above two reactions are pH dependent, and they increase with an increase in pH. 
The copper catalyzed auto oxidation of L-cysteine at pH 13 was carefully studied by 
Cavallini et al.
123
 Their UV-vis, EPR spectra and kinetic results indicate that the cysteine-
Cu(II) complex represents the real intermediate catalyst. The cysteine-Cu(II) is slowly 
reduced to cysteine-Cu(I) when oxygen is removed. Recently, the oxidation of L-cysteine 
by ferricytochrome, nitro blue tetrazolium and oxygen was investigated by Pecci et al. at 
pH 7.4.
124
 Their UV-vis and EPR results show that the cuprous bis-cysteine complex 
(RS-Cu
I
-SR) is the catalytic species involved in the oxidation of cysteine. The two totally 
different results may be ascribed to their inconsistent pH environment. In our case, the 
kinetics was studied from pH 6.95 to 12.0 under anaerobic condition. So bis-cysteine 
complex (RS-Cu
I
-SR) may be the catalytic species for the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
. Due to its complexity, understanding the mechanism of the copper 
catalytic oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
 could be an area for future studies. 
Generally, cuprous cation, soft acid, is coordinated to fairly soft base S
-
 and NH
2
 to 
form five-member-ring cuprous cysteinate (RSCu(I)) and cuprous dicysteinate 
(RSCu(I)SR).
96,124-126
 There are only a few reports about the stability constant of cuprous 
cysteinate (RSCu(I)) in the literature. Firstly, Kolthoff et al. 
126
 determined the stability 
constant of cuprous cysteinate (K
1
) by polarographic titration in ammonia buffer (pH = 
9.25), with a value of 1.5 ? 10
19
 M
-1
. Then Knoblock et al. 
127
 mistakenly reported the 
overall stability constant (?) of cupric dicysteinate (RSCu(II)SR) at pH 7.40 and ionic 
strength of 1.0 M, which should correspond to the overall stability constant of cuprous 
dicysteinate (RSCu(I)SR), with a value of 1.0 ? 10
16
 M
-1
. The above reported stability 
 99
constants show that it decreases with a decrease in pH, consistent with the theoretical 
estimation from the coordination chemistry of cysteine with cuprous cation. To break the 
strong binding of Cu(I) to the cysteine, strong chelating reagent EDTA or cyclam was 
added in the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. Its copper-catalysis was 
inhibited with the addition of 5.0 mM EDTA (pH ? 9.00), 10.0 mM EDTA (9.00 < pH ? 
10.0) and 1.0 mM cyclam (pH > 10.0). 
Our
 
1
H-NMR and UV-vis spectra results indicate that the products of the reaction 
are L-cystine and the corresponding Fe(II) complexes; while in the oxidation of L-
cysteine with octacyanomolybdate(V), L-cysteine-sulfinate was formed as well as 
[Mo(CN)
8
]
4-
 
and L-cysteine.
52
 This is attributed to the lower oxidative capability of 
[Fe
III
(bpy)
2
(CN)
2
]
+
 and [Fe
III
(bpy)(CN)
4
]
-
 (E
1/2
, [Fe(bpy)
2
(CN)
2
]
+/0
 = 0.759 V; E
1/2
, 
[Fe(bpy)(CN)
4
]
-/2-
 = 0.537 V; E
1/2
, [Mo(CN)
8
]
3-/4-
 = 0.771 V 
52
). The same phenomena 
were observed in the oxidation of TGA: dithiodiglycolic acid (RSSR) and sulfoacetate 
(RSO
3
-
) were formed when hexachloroiridate(IV) was used as the oxidant; while only 
dithiodiglycolic acid (RSSR) was produced if octacyanomolybdate(V) was used as an 
oxidant (E
1/2
, [IrCl
6
-
]
2-/3-
 = 0.892 V, E
1/2
, [Mo(CN)
8
]
3-/4-
 = 0.771 V).
103,102
 
The results 
from 
1
H-NMR spectra and spectrophotometric analysis demonstrate that the 
stoichiometric ratio of ?[L-cysteine]/?[Fe(III)] is around 1:1. Based on the products? and 
stoichiometric ratio results, it is presumed that Fe(III) first reacts with L-cysteine to 
form 
L-cysteine thiyl radical, then cysteine combines with cysteine thiyl radical to form cystine 
radical, and finally cystine radical is oxidized by Fe(III) to produce Fe(II) and L-cystine 
(see mechanism, page 101 and 106). 
 100
The observance of non-pseudo-first-order kinetics in the presence of 
Fe
II
(bpy)
2
(CN)
2
 at lower pH for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
 is 
ascribed to the forward rate constant pH-dependence in Equation 2-10. Recently,  
Mezyk 
128
 measured the rate constants of radical disulfide anion formation for cysteine at 
different pH values: at pH 7.00, the rate constant (k
10
) is (3.77 ? 0.33) ? 10
8
 M
-1
 s
-1
, and 
the rate constant increases by a factor of nearly 2 at pH 8.00. Although the measurements 
were not performed in acidic media, the rate constants are assumed to decrease at pH 
values less than 7.00. So the slower forward rate constant of Equation 2-10 in acidic 
media may shift Equation 2-12 to backward direction, therefore the kinetic inhibition by 
Fe
II
(bpy)
2
(CN)
2
 at lower pH was observed. 
Perfect pseudo-first-order kinetics was obtained with the addition of 0.20 mM PBN 
to the reactants: at lower pH values, the rate constant increases by 15%; at higher pH 
values, the rate constant decreases by 6.7%. Earlier EPR results showed that PBN 
scavenges L-cysteine radical efficiently. 
121
 Although no rate constant for the reaction 
between PBN and L-cysteine was reported, the rate constant for formation of the adduct 
of cysteine thiyl radical to DMPO was determined by Davies et al., with a value of 2.10 ? 
10
8
 M
-1
 s
-1
.
129
 It is also known that PBN is less reactive than DMPO, so the rate constant 
k
11 
in Equation 2-11 may be smaller than 2.10 ? 10
8
 M
-1
 s
-1
. The addition of PBN induces 
the competition between Reaction 2-10 and 2-11. At lower pH values, the forward rate 
constant (k
11
) in Equation 2-11 is higher than that (k
10
) in Equation 2-10, which enables 
PBN to scavenge cysteine thiyl radical efficiently. So the reaction in Equation 2-12 shifts 
to the right, increasing the forward rate constant, k
12
. At higher pH values, k
11
 is very 
close to or smaller than k
10
, so the addition of PBN will not change k
12
. Unfortunately, the 
 101
decrease of k
obs
 with the addition of PBN at pH 7.61 remains a puzzle. With the addition 
of PBN, the 
1
H-NMR spectra show that the products of the reaction are Fe
II
(bpy)
2
(CN)
2
, 
L-cystine and some unidentified compound with singlets. Since the adduct of 
PBN/cysteine thiyl radical is reportedly very unstable,
121
 the singlets are possibly due to 
the decomposition of the adduct. The stoichiometric ratio of ?[L-cysteine]/?[Fe(III)] is 
still 1:1 in the presence of PBN. Furthermore, according to our results, PBN reacts with 
[Fe
III
(bpy)
2
(CN)
2
]
+
 very slowly at pH 7.56, and no reaction takes place between PBN and 
L-cysteine. Thus, it is safe to add PBN to the reactants to scavenge the cyteine thiyl 
radicals that were formed in the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
. 
 
RS  + RSH RSSR 
-
 + H
+
     k
10
, k
-10
, K
rad
                     (2-10)
 
 
RS  + PBN PBN/RS                                                  k
11
, k
-11  
                          (2-11)
 
 
 + [Fe(bpy)
2
(CN)
2
]
+
RS  + Fe(bpy)
2
(CN)
2
              k
12
, k
-12
                          (2-12)
RS
-
 
 
The mechanism of oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 is proposed as: 
 
HSCH
2
CH(NH
3
+
)CO
2
H          HSCH
2
CH(NH
3
+
)CO
2
-
 +
 
H
+
pK
a1
 = 1.90             (2-3)
HSCH
2
CH(NH
3
+
)CO
2
-
-
SCH
2
CH(NH
3
+
)CO
2
-
 + H
+ pK
a2
 = 8.18            (2-4)
 
 
HSCH
2
CH(NH
3
+
)CO
2
- -
SCH
2
CH(NH
3
+
)CO
2
H                pK
i
 = 5.5            (2-13)
 
-
SCH
2
CH(NH
3
+
)CO
2
H + [Fe(bpy)
2
(CN)
2
]
+
Fe(bpy)
2
(CN)
2
  +   SCH
2
CH(NH
3
+
)CO
2
H           k
a
, k
-a
            (2-14)
 
 102
-
SCH
2
CH(NH
3
+
)CO
2
-
 + [Fe(bpy)
2
(CN)
2
]
+
 Fe(bpy)
2
(CN)
2
 +   SCH
2
CH(NH
3
+
)CO
2
-
                 k
b
, k
-b
      (2-15)
RS  + RSH RSSR 
-
 + H
+
k
10
 k
-10
 K
rad
,    (2-10)
RS  + PBN
PBN/RS                                                                         k
11
, k
-11       
(2-11)
RSSR 
- 
+ [Fe(bpy)
2
(CN)
2
]
+
RSSR + Fe(bpy)
2
(CN)
2
                  k
d
        (2-16)
 
 
 
The distinct kinetic inhibition by Fe
II
(bpy)
2
(CN)
2
 at pH 3.20, in the absence of PBN, 
was well simulated by using Specfit/32 software package, as shown in Table 2-15. The 
mathematical model and one of the simulated results are listed in Table 2-16. The 
inhibition by Fe
II
(bpy)
2
(CN)
2
 is insignificant if the rate constant of k
-3a
 (see Table 2-16) is 
set at lower than 5.0 ? 10
8
 M
-1
 s
-1
. The above simulations assist us to understand the 
contribution of PBN to the lower pH?s kinetic behavior. At lower pH, efficient 
scavenging of the cysteine radial by PBN slows down the backward reaction of Equation 
2-14 dramatically. At higher pH, the cysteine radical is scavenged by cysteine efficiently. 
So it is reasonable to consider Equations 2-14 and 2-15 as irreversible in derivation of the 
rate law.  
Applying steady-state for RS
?
 and RSSR
?-
, the rate law of the reaction is derived: 
 
[Fe(III)]]RSH[
]H[]H[]H[
]H[]H[
2
dt
d[Fe(III)]
tot
32121
2
1
3
21
2
a1
?
?
?
?
?
?
+++
+
=
?
+++
++
aaaaaa
aabia
KKKKKK
KKkKKk
     (2-17-a) 
 
where k
obs
 is: 
 
 103
Table 2-15. Simulation results for the kinetic inhibition by Fe
II
(bpy)
2
(CN)
2
 in  
the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20. In the presence of 1.0 mM 
dipic/dipic
2-
(1:1), at ? = 0.10 M and 25.0 
o
C, using Specfit/32 software package. 
t
1/2
, s 
[Fe
II
(bpy)
2
(CN)
2
]
0
, ?M
Experimental
b
Theory
c
 
0 36.8 37.7 
25 52.2 52.6 
100 90.0 90.0 
a
 [L-cysteine]
0
 = 1.25 ? 10
-3
 M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 2.5 ? 10
-5
 M, 
b
 Calculated from 
k
obs
.
c
 Obtained from simulation 
 
Table 2-16. Mathematical models for the simulation of the oxidation of 1.25 ? 10
-3
 M L-
cysteine by 2.5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 at pH 3.20. In the presence of 1.0 mM 
dipic/dipic
2-
(1:1) and 2.5 ? 10
-5
 
M Fe
II
(bpy)
2
(CN)
2
, at ? = 0.10 M and 25.0 
o
C, using 
Specfit/32 software package. 
Model Reaction 
Equation 
Name Rate constants
Initial 
Concentration, M 
Half-life, 
s 
A ? B + C 
k
1a
 10.0
 
A 
1.25 ? 10
-3
 
B + C ? A 
k
-1a
 79.4 B 0.0 
B ? D 
k
2a
 1.0 C 
6.31 ? 10
-4
 
D ? B 
k
-2a
 
3.1? 10
5
 
D 0.0 
D + E ? F + G 
k
3a
 
2.80 ? 10
6
 
E 
2.5 ? 10
-5
 
F + G ? D + E 
k
-3a
 
4.00 ? 10
9
 
F 0.0 
B + F ? C + H 
k
4a
 
2.20 ? 10
8
 
G 
2.5 ? 10
-5
 
C + H ? B + F 
k
-4a 1.38 ? 10
7
 
H 0.0 
E + H ? G + I 
k
5a 2.4 ? 10
9
 
I 0.0 
52.6 
A: HSCH
2
CH(NH
3
+
)CO
2
H; B: HSCH
2
CH(NH
3
+
)CO
2
?
; C: H
+
; D: 
-
SCH
2
CH(NH
3
+
)CO
2
H; 
E: [Fe
III
(bpy)
2
(CN)
2
]
+
; F: 
?
SCH
2
CH(NH
3
+
)CO
2
H; G: Fe
II
(bpy)
2
(CN)
2
; H: RSSR
?-
; I: 
RSSR 
 104
tot
32121
2
1
3
21
2
a1
obs
]RSH[
]H[]H[]H[
]H[]H[
2
?
?
?
?
?
?
+++
+
=
+++
++
aaaaaa
aabia
KKKKKK
KKkKKk
k                          (2-18-a) 
 
Comparing Equation 2-18-a with Equation 2-8, 2k
a
K
i
 = k
2
 = (2.93 ? 0.23) and 2k
b
 = 
k
3
 = (1.22 ? 0.02) ? 10
6
 for the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
4
]
+
. 
Very recently, a substantial kinetic isotope effect was observed in the oxidation of 
hydroxylamine by hexachloroiridate(IV) by Stanbury and his coworkers.
130
 They 
ascribed the kinetic isotope effect to concerted proton-coupled electron transfer (PCET) 
from the neutral form of hydroxylamine to hexachloroiridate(IV). Therefore, PCET from 
the neutral form of L-cysteine to [Fe
III
(bpy)
2
(CN)
2
]
+
 (Equation 2-19) needs to be 
considered. If Equation 2-19 is the correct mechanism, then its rate constant, k
c
, would 
correspond to half of k
2
 in Equation 2-8, with (1.46 ? 0.12) M
-1
 s
-1
 for k
c
. The standard 
potential of RS
?
/RSH was calculated by using the standard potential of RS
?
/RS
-
 (E
1/2
 = 
0.88 V vs NHE, refer to page 111) and the dissocation constant of neutral cysteine (pK
a2
 
= 8.18), with E
1/2
, RS
?
/RSH = 1.36 V vs NHE. Based on the half-wave potential of 
[Fe(bpy)
2
(CN)
2
]
+/0
, and that of RS
?
/RSH, ?G
o
 for Equation 2-19 is 58 kJ mol
-1
. Although 
this PCET mechanism can?t be ruled out, the high ?G
o
 value of Equation 2-19 implies 
that this process is unfavorable. The resonable self-exchange rate constants of 
?
SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
 and 
?
SCH
2
CH(NH
2
)CO
2
?
/ 
?
SCH
2
CH(NH
2
)CO
2
?
 that were calculated from Marcus equations also imply that this 
process occurs through electron transfer, not PCET. In contrast to electron transfer being 
favored for cysteine oxidation, PCET may be the favored pathway in oxidation of 
H
2
NOH because ?G
o
 is not unfavorable. 
 105
 
HSCH
2
CH(NH
3
+
)CO
2
-
 + [Fe(bpy)
2
(CN)
2
]
+
SCH
2
CH(NH
3
+
)CO
2
-
 + Fe(bpy)
2
(CN)
2
+ H
+
(2-19)k
c
 
 
The oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 shows rather different behavior.
 
For the oxidation of 2.60 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M 
[Fe
III
(bpy)(CN)
4
]
-
 at pH 
5.98 (0.093 M NaClO
4
, 10.0 mM cacodylate, and 5.0 mM EDTA), non-first-order 
kinetics were obtained; when the concentration of L-cysteine increases by 10 times, first-
order kinetics were successfully achieved. At pH values higher than 8.0, even when only 
5.0 ? 10
-4
 M L-cysteine were added, perfect first-order kinetics was obtained. Most 
importantly, the rate of the reaction does not change with the addition of PBN. 
The above kinetic behavior is ascribed to the pH-dependence on the equilibrium of 
Equation 2-10. The equilibrium constant of Equation 2-10 was measured by Mezyk
131
 
using two techniques: integrated yields and measurements of kinetics. Both methods gave 
consistent results: at pH 6.0, the equilibrium constant (K
10
) is 24.6 ? 5.4; at pH 7.0, K
10
 is 
228 ? 4.3; at pH 8.5, it reaches a maximum, with a value of 1451 ? 16; then it decreases 
slowly with the increase of pH, it is 394.6 ? 8.6 at pH 13.0.
131
 At lower pH and lower 
concentration of L-cysteine, less cysteine thiyl radical that was formed in the oxidation of 
L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 combines with L-cysteine to form radical disulfide anion; 
with an increase of concentration of L-cysteine, more cysteine thiyl radical is consumed 
to form radical disulfide anion, favoring the achievement of first-order kinetics. With the 
increase of pH, the equilibrium constant of Equation 2-10 increases dramatically, so 
 106
Equation 2-10 shifts to the forward direction even at very low concentrations of L-
cysteine. Therefore, first-order kinetics were observed in the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 complex at higher pH. 
Based on above information, the mechanism for the oxidation of L-cysteine by 
[Fe
III
(bpy)(CN)
4
]
-
 
is proposed:  
-
SCH
2
CH(NH
3
+
)CO
2
-
 + [Fe(bpy)(CN)
4
]
-
 [Fe(bpy)(CN)
4
]
2-
 +   SCH
2
CH(NH
3
+
)CO
2
-
              k
b'
, k
-b'
   (2-20)
 
-
SCH
2
CH(NH
2
)CO
2
- 
+ [Fe(bpy)(CN)
4
]
-
[Fe(bpy)(CN)
4
]
2-
 +  SCH
2
CH(NH
2
)CO
2
- 
                 k
c'
,  k
-c'
  (2-21)
 
RS  + RSH RSSR 
-
  + H
+
k
10
, k
-10
, K
rad
,        (2-10)
RSSR 
- 
+ [Fe(bpy)(CN)
4
]
-
RSSR + [Fe(bpy)(CN)
4
]
2-
                       k
d'
       (2-22)
 
 
Furthermore, the mild kinetic inhibition by 
[Fe
II
(bpy)(CN)
4
]
2-
 at pH 5.98 was well 
simulated by using Specfit/32 software package,
132
 as shown in Table 2-17. The 
mathematical model and one of the simulated results are listed in Table 2-18. The 
simulations indicate that the forward rate constant of Equation 2-10 is one of the rate 
limiting steps, not the typical forward rate constant of Equation 2-22. Since the inhibition 
is insignificant for the reaction without the addition of [Fe
II
(bpy)(CN)
4
]
2-
, the reactions 
from 2-20 to 2-21 are safely considered irreversible in the following derivation of rate 
law.  
 107
Table 2-17. Simulation results for the kinetic inhibition by [Fe
II
(bpy)(CN)
4
]
2-
 in the  
oxidation of 2.6 ? 10
-3
 M L-cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98. In the 
presence of 5.0 mM EDTA, at ? = 0.10 M and 25.0 
o
C, using Specfit/32 software. 
t
1/2
, s 
[[Fe
II
(bpy)(CN)
4
]
-
]
0
, ?M 
Experimental
a
 Theory
b
 
0 20.2 21.0 
50 25.4 24.6 
100 28.0 28.2 
a
 Calculated from k
obs
.
b
 Obtained from simulation 
 
 
 
Table 2-18. Mathematical models for the simulation of the oxidation of 2.6 ? 10
-3
 M L-
cysteine by 5.0 ? 10
-5
 M [Fe
III
(bpy)(CN)
4
]
-
 at pH 5.98. In the presence of 5.0 mM EDTA 
and 5.0 ? 10
-5
 
M [Fe
II
(bpy)(CN)
4
]
2-
, at ? = 0.10 M and 25.0 
o
C, using Specfit/32 software. 
Model Reaction 
Equation 
Name 
Rate 
constants 
Initial 
Concentration, M 
Half-life,
s 
A ? B + C 
k
1b
 10.0
 
A 
2.6 ? 10
-3
 
B + C ? A 
k
-1b
 
1.50 ? 10
9
 
B 0.0 
B ? C + D 
k
2b
 1.0 C 
1.06 ? 10
-6
 
C + D ? B 
k
-2b
 
2.0 ? 10
10
 
D 0.0 
B + E ? F + G 
k
3b
 
1.05 ? 10
3
 
E 
5.0 ? 10
-5
 
F+G ? B + E 
k
-3b
 
1.75 ? 10
9
 
F 0.0 
A + F ? C + H 
k
4b
 
2.20 ? 10
8
 
G 
5.0 ? 10
-5
 
C + H ? A + F 
k
-4b 1.38 ? 10
7
 
H 0.0 
E + H ? G + I 
k
5b 2.4 ? 10
8
 
I 0.0 
24.6 
 
A: HSCH
2
CH(NH
3
+
)CO
2
-
; B: 
-
SCH
2
CH(NH
3
+
)CO
2
-
; C: H
+
; D: 
-
SCH
2
CH(NH
2
)CO
2
-
; E: 
[Fe
III
(bpy)(CN)
4
]
-
; F: 
?
SCH
2
CH(NH
3
+
)CO
2
-
; G: [Fe
II
(bpy)(CN)
4
]
2-
; H: RSSR
?-
; I: RSSR 
 108
Applying steady-state for RS
?
 and RSSR
?-
, the rate law of the reaction is derived:  
 
[Fe(III)]]RSH[
]H[]H[]H[
]H[
2
dt
d[Fe(III)]
tot
32121
2
1
3
321'21'
?
?
?
?
?
?
+++
+
=
?
+++
+
aaaaaa
aaacaab
KKKKKK
KKKkKKk
    (2-17-b) 
 
where k
obs
 is: 
tot
32121
2
1
3
321'21'
obs
]RSH[
]H[]H[]H[
]H[
2
?
?
?
?
?
?
+++
+
=
+++
+
aaaaaa
aaacaab
KKKKKK
KKKkKKk
k                        (2-18-b) 
 
The rate law has the same form as the empirical rate law (Equation 2-9). Comparing 
Equation 2-18-b with Equation 2-9, 2k
b?
 = k
3?
 = (2.04 ? 0.16) ? 10
3
 M
-1
 s
-1
, 2k
c?
 = k
4?
? = 
(1.07 ? 0.25) ? 10
4
 M
-1
 s
-1
 for the oxidation of L-cysteine by [Fe
III
(bpy)(CN)
4
]
-
. 
In the above mechanism, the rate-limiting steps correspond to Equations 2-14 and 
2-15 for the reduction of [Fe
III
(bpy)
2
(CN)
2
]
+
, and Equations 2-20 and 2-21 for the 
reduction of [Fe
III
(bpy)(CN)
4
]
-
. They involve one-electron transfer from the thiolate 
forms of cysteine to the corresponding Fe(III) complex. The cysteine thiyl radicals 
generated in the above reactions react with cysteine rapidly to form disulfide radical 
anion, as shown in Equation 2-10.
133,134
 Then the disulfide radical anion is oxidized to 
form L-cystine, as shown in Equation 2-16 and 2-22. 
 
HSCH
2
CH(NH
3
+
)CO
2
-
         (2-23)HSCH
2
CH(NH
2
)CO
2
H                       K
23
 
Cysteine can exist as different tautomers in its neutral and monoanionic forms.
135,136
 
HSCH
2
CH(NH
3
+
)CO
2
-
, the zwitterion, predominates in the neutral form, as well as a 
 109
small amount of 
-
SCH
2
CH(NH
3
+
)CO
2
H and HSCH
2
CHNH
2
CO
2
H (pK
i
 = 5.5, pK
23
 = 
4.88).
136
 
-
SCH
2
CH(NH
3
+
)CO
2
-
 is the main form in monoanionic cysteine, and the minor 
tautomers are HSCH
2
CHNH
2
CO
2
-
 and 
-
SCH
2
CHNH
2
CO
2
H. The equilibrium constant for 
the tautomerization of the monoanion to the thiol form is 0.40, while the equilibrium 
constant to the carboxylic form is quite small (K = 6.02 ?10
-6
).
136
 Since the thiolate is the 
reactive form in the oxidation of L-cysteine by Fe(III), and the concentration of  
-
SCH
2
CHNH
2
CO
2
H is rather low, only
 
-
SCH
2
CH(NH
3
+
)CO
2
-
 needs to be included. 
Combining the estimated K
i
 with the measured value for 2k
a
K
i 
= k
2
 = (2.93 ? 0.23) M
-1
 s
-1
 
gives k
a
 value of (4.63 ? 0.36) ? 10
5
 M
-1
 s
-1
 and 2k
b
 = k
3
 = (1.22 ? 0.02) ? 10
6
 gives k
b
 
(6.10 ? 0.01) ? 10
5
 M
-1
 s
-1
 for the oxidation of L-cysteine by [Fe
III
(bpy)
2
(CN)
2
]
+
; 2k
b?
 = 
k
3?
 = (2.04 ? 0.16) ? 10
3
 M
-1
 s
-1
 gives k
b?
 (1.02 ? 0.08) ? 10
3
 M
-1
 s
-1
 and 2k
c?
 = k
4?
 = (1.07 
? 0.25) ? 10
4
 M
-1
 s
-1
 gives k
c?
 (5.35 ? 1.20) ? 10
3
 M
-1
 s
-1
 for the oxidation of L-cysteine 
by [Fe
III
(bpy)(CN)
4
]
-
.  
 
RS 
-
 + RS
2-
RSSR 
3-
         k
24
, k
-24
                         (2-24)
 
         E
25
o
                               (2-25)
 
+ e
-
2 RS
2-
RSSR 
3-
RS  
-
  +
         E
26
o
 
      
                          (2-26)e
-
RS
2-
 
 + Fe(III) RS 
-
 + Fe(II)                                        k
27
, k
-27
                           (2-27)RS
2-
 
 
RS  + Fe(II)                                        k
28
, k
-28
                           (2-28)RS
-
 + Fe(III)
 
Here RS
?
, RS
-
, RS
?-
, RS
2-
 and RSSR
?3-
 refer to 
?
SCH
2
CH(NH
3
+
)CO
2
?
,  
 110
-
SCH
2
CH(NH
3
+
)CO
2
?
, 
?
SCH
2
CH(NH
2
)CO
2
?
, 
?
SCH
2
CH(NH
2
)CO
2
?
 and  
-
O
2
C(NH
2
)CCHCH
2
S     SCH
2
CH(NH
2
)CO
2
-
. 
 
The driving force for the one-electron transfer reaction can be evaluated if the 
standard potentials of [Fe(bpy)
2
(CN)
2
]
+/0
, [Fe(bpy)(CN)
4
]
-/2-
 and Cys
?-
/Cys
2-
 are known. 
The first reported standard potential of Cys
?-
/Cys
2-
 was (0.73 ? 0.05) V, which was 
obtained by measurement of the electron-transfer equilibrium constants between the 
cystine radical anion and tyrosine at pH 13.0, with a value of 0.73 ? 0.05 V.
137
 Recently, 
Mezyk
131
 and Zhao et al.
133
 determined the equilibrium constant for the formation of the 
cysteine radical disulfide anion (K
24
), and both of them gave very close results. The 
equilibrium constant determined by Mezyk is 514 ? 13 M
-1
, 466 ? 20 M
-1
and 414.5 ? 6.2 
M
-1
 at pH 10.0, 11.0 and 12.0, respectively; while that from Zhao is 706 M
-1
 at pH 10.5. 
Then Mezyk and Armstrong theoretically calculated the standard potential of RSSR
?3-
/2RS
2-
, with E
1/2,
 (RSSR
?3-
/2RS
2-
) = 0.60 V vs NHE.
134
 At pH 12.0, nearly all of L-
cysteine exists in dianion form. Based upon the standard potential (E
25
o
 = 0.60 V vs NHE) 
of Equation 2-25 and the equilibrium constant of Equation 2-24 at pH 12.0 (K
24
 = 414.5 ? 
6.2 M
-1
), the standard potential of RS
?-
/RS
2-
 was calculated, with E
1/2
, RS
?-
/RS
2-
 = 0.755 
V vs NHE. The standard potential of Cys
?-
/Cys
2-
 was also calculated by Hung et al.
52
, 
with E
1/2
, (Cys
?-
/Cys
2-
) = (0.76 ? 0.02) V at pH 13, consistent with our results. The 
standard potential of Cys
?
/Cys
-
 was calculated by using the standard potential of Cys
?-
/Cys
2-
( E
1/2
 = 0.76 V), microscopic dissociation constant of cysteine radical 
(
?
SCH
2
CH(NH
3
+
)CO
2
?
, pK
S?
 = 8.26 ? 0.04 
134
) and the macroscopic dissociation constant 
 111
of cysteine monanion (
-
SCH
2
CH(NH
3
+
)CO
2
?
, pK
SNN
 = 10.36 
134
), with E
1/2
, Cys
?
/Cys
-
 = 
0.88 V vs NHE. 
Then ?G
o
, (21.5 ? 1.9 kJ mol
-1
), is derived from the corrected standard potential of 
Cys
?-
/Cys
2-
 and [Fe(bpy)(CN)
4
]
-/2-
 (E
1/2
, [Fe(bpy)(CN)
4
]
-/2-
 = 0.537 V, E
1/2
, Cys
?-
/Cys
2-
 = 
0.76 V), corresponding to the equilibrium constant of 1.70 ? 10
-4
 for Equation 2-27. The 
?G
o
 of Equation 2-28 is derived from the corrected standard potentials of Cys
?
/Cys
-
, 
[Fe(bpy)
2
(CN)
2
]
+/0
 and [Fe(bpy)(CN)
4
]
-/2-
 (E
1/2
, Cys
?
/Cys
-
 = 0.88 V, E
1/2
,  
[Fe(bpy)(CN)
4
]
-/2-
 = 0.537 V, E
1/2
, [Fe(bpy)
2
(CN)
2
]
+/0
 = 0.759 V), which are 33.1 kJ mol
-1
, 
11.7 kJ mol
-1
 for the reduction of [Fe(bpy)
2
(CN)
2
]
+
 and [Fe(bpy)(CN)
4
]
-
, corresponding 
to the equilibrium constants of 9.0 ? 10
-3
 and 1.59 ? 10
-6
. 
The cross-relationship of Marcus theory is applicable to the self-exchange rate 
constant of the 
?
SCH
2
CH(NH
2
)CO
2
-
/
-
SCH
2
CH(NH
2
)CO
2
-
 couple provided reaction 2-27 
follows an outer-sphere mechanism. The following equations from Marcus theory were 
used to calculate the self-exchange rate constant of the 
?
SCH
2
CH(NH
2
)CO
2
-
/
-
SCH
2
CH(NH
2
)CO
2
-
 couple:
92
 
 
12
2/1
1212221112
)( WfKkkk =                                                                                    (2-29) 
 
)/RT)]()/Z4[ln(
]RT/)([ln
ln
2211
2
2211
2
211212
12
wwkk
wwK
f
++
?+
=                                                         (2-30) 
 
RT2/)exp(
2211211212
wwwwW ++??=                                                             (2-31) 
 
 112
)]328.01(/[23.4 ?rrZZw
jiij
+=                                                                      (2-32) 
 
In these equations, k
12
 is the observed cross electron-transfer rate constant, (5.20 ? 
0.11) ? 10
3
 M
-1
 s
-1
, and k
11
 and k
22
 are the self-exchange rate constants of the 
?SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 and [Fe(bpy)(CN)
4
]
-/2-
 redox couples. The 
self-exchange rate constant (k
22
) of [Fe(bpy)(CN)
4
]
-/2-
 redox couple was reported by 
Stasiw and Wilkins in the study of the oxidation of [Fe(bpy)(CN)
4
]
2-
 by [Fe(CN)
6
]
3-
, with 
the value of 4 ? 10
7
 M
-1
 s
-1
.
138
 However, the charge effect was not included in the 
calculation, and the self-exchange rate constant of [Fe(CN)
6
]
3-/4-
(k
11
 = 5 ? 10
3
 M
-1
 s
-1
) is 
not accurate. It is known that the electron-transfer reaction between [M(CN)
6
]
3-
 and 
[M(CN)
6
]
4-
(M = Fe, Os) in homogeneous solution is catalyzed by alkali metal ion.
139,140
 
The uncatalyzed self exchange rate constant of [Fe(CN)
6
]
3-/4-
 was determined by NMR 
line broadening method using either crypt-2,2,2 or 18-crown-6 to inactivate the 
potassium counterion catalysis, with k
0
 of 2.4 ? 10
2
 M
-1
 s
-1
, much smaller than Stasiw and 
Wilkins? report.
140
 To avoid the high charge effect of [Fe(CN)
6
]
3-/4-
, the self-exchange 
rate constant (k
22
) of [Fe(phen)(CN)
4
]
-/2-
 redox couple was recalculated by us using 
Marcus? equations from the oxidation L-ascorbic acid by [Fe(phen)(CN)
4
]
-
.
92,111,141,142
 K
33
 
is the electron-transfer equilibrium constant of Equation 2-33, with a value of 0.02. Z is 
the collision frequency that the value of 1 ? 10
11
 M
?1
 s
?1
 is used in our calculation. Z
i
 and 
Z
j
 are ionic charges on the reactants, R is the ideal gas constant (1.987 ? 10
?3
 kcal mol
?1
), 
and r is the center-to-center distance between reactants while in contact. The radii of 
[Fe(phen)(CN)
4
]
-
 and HA
-
 are 5.33 
111
 and 3.00 ?
141
, respectively. Using the above 
 113
known parameters, the self-exchange rate constant of [Fe(phen)(CN)
4
]
-/2-
 was 
recalculated with 8.39 ? 10
5
 M
-1
 s
-1
. 
 
[Fe(phen)(CN)
4
]
-
 + HA
-
 ? [Fe(phen)(CN)
4
]
2-
 + HA
?
                     K
33
                (2-33) 
 
Since [Fe(bpy)(CN)
4
]
-
 has similar size as that of [Fe(phen)(CN)
4
]
-
, the self-
exchange rate constant (k
22
) of [Fe(bpy)(CN)
4
]
-/2-
 redox couple is assumed as 8.39 ? 10
5
 
M
-1 
s
-1
, and it is applied for the following calculation. K
12
 is the electron-transfer 
equilibrium constant of Equation 2-27, with a value of 1.70 ? 10
-4
. Z is the collision 
frequency that the value of 1 ? 10
11
 M
?1
 s
?1
 is used in our calculation. Z
i
 and Z
j
 are ionic 
charges on the reactants, R is the ideal gas constant (1.987 ? 10
?3
 kcal mol
?1
), and r is the 
center-to-center distance between reactants while in contact. The radii of [Fe(bpy)(CN)
4
]
-
 
and 
-
SCH
2
CH(NH
2
)CO
2
?
 are 5.33 ?
111
 and 3.00 ?
52
, respectively. Using the above 
known parameters, the self-exchange rate constant of 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
, k
11
,
 
is calculated as 3.66 ? 10
5
 M
?1
 s
?1
, which is higher than that 
derived from the reaction of [Mo(CN)
8
]
3-
 with L-cysteine (k
11
 = 5.4? 10
3
 M
?1
 s
?1
).
52
 The 
smaller self-exchange rate constant of ?SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 that 
was derived from the oxidation of L-cysteine by [Mo(CN)
8
]
3-
 is ascribed to the high 
negative charge of Mo(V). Although alkaline metal cation catalysis exists for the 
oxidation of L-cysteine by [Mo(CN)
8
]
3-
,
52
 the lower self-exchange rate constant implies 
that repulsion between Mo(V) and RS
2-
 is dominant at higher pH values. However, less 
negative charge in [Fe(bpy)(CN)
4
]
-
 endows its weaker repulsion with L-cysteine dianion 
 114
(RS
2-
), and no alkaline metal cation effect was observed in the oxidation L-cysteine by 
[Fe(bpy)(CN)
4
]
-
 at pH 11.0. Therefore, the higher k
11
 value of 
?
SCH
2
CH(NH
2
)CO
2
?
/ 
?
SCH
2
CH(NH
2
)CO
2
?
 that was derived from the oxidation of L-cysteine by 
[Fe(bpy)(CN)
4
]
-
 is reasonable. Using Marcus? Equations (2-29 to 2-32) and the known 
parameters, the self-exchange rate constant of 
?
SCH
2
CH(NH
3
+
)CO
2
?
/ 
?
SCH
2
CH(NH
3
+
)CO
2
?
, k
11
, is calculated as 2.62 ? 10
6
 M
?1
 s
?1
. 
Marcus? theory is also applicable for the oxidation of L-cysteine by 
[Fe
III
(bpy)
2
(CN)
2
]
+
.
 There is no report about the self-exchange rate constant (k
22
) of the 
[Fe(bpy)
2
(CN)
2
]
+/0
 redox couple. It was calculated using Marcus? equations from the 
oxidation L-ascorbic acid by [Fe(bpy)
2
(CN)
2
]
+
, 
92,111,141,142
 with the value of 
1.30 ? 10
7
 
M
?1
 s
?1
. The radii of [Fe(bpy)
2
(CN)
2
]
+
 and 
-
SCH
2
CH(NH
3
+
)CO
2
?
 estimated from CPK 
models are 6.77 ? 
111
 and 3.00 ? 
52
, respectively. Using the above known parameters, the 
self-exchange rate constant of ?SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
, k
11
,
 
is 
calculated as 2.50 ? 10
6
 M
?1
 s
?1
, very close to that obtained from the oxidation of L-
cysteine by [Fe(bpy)(CN)
4
]
-
 (k
11
= 2.62 ? 10
6
 M
?1
 s
?1
). Comparing the self-exchange rate 
constant of 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 with ?SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
, the smaller self-exchange rate constant of 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 is ascribed to their negative-negative charge repulsion. The rapid 
self-exchange rate constants of 
?
SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
 and 
?
SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
) are consistent with the small internal 
reorganizational energy for the formation of cysteine radical.
143
 This implies that the 
 115
oxidation of L-cysteine by [Fe(bpy)(CN)
4
]
-
 may occur through one-electron outer-sphere 
transfer reactions.  
 
Conclusion 
 
Copper catalysis was observed in the oxidation of L-cysteine by [Fe(bpy)
2
(CN)
2
]
+
 
and [Fe(bpy)(CN)
4
]
-
 in anaerobic aqueous solution. With the addition of 1.0 mM dipic
2-
, 
5.0 mM EDTA/1.0 mM cyclam, the copper-catalysis was effectively inhibited, and the 
direct oxidation of L-cysteine by outer-sphere oxidants ([Fe(bpy)
2
(CN)
2
]
+
 and 
[Fe(bpy)(CN)
4
]
-
) was first reported by us. The rate-limiting step is electron transfer to 
form a cysteine radical and the corresponding Fe(II) complex
. Only the thiolate forms of 
cysteine are reactive. The products of the reaction are L-cystine and the corresponding 
Fe(II) complex. The rate constants increase with increasing pH. For the oxidation of L-
 
cysteine by [Fe(bpy)(CN)
4
]
-
, 
the self-exchange rate constants of 
?
SCH
2
CH(NH
3
+
)CO
2
?
/ 
?
SCH
2
CH(NH
3
+
)CO
2
?
 
and 
?SCH
2
CH(NH
2
)CO
2
?
/
?
SCH
2
CH(NH
2
)CO
2
?
, k
11
, were obtained 
by applying Marcus theory with 2.62 ? 10
6
 M
?1
 s
?1
and 3.66 ? 10
5
 M
?1
 s
?1
, respectively
. 
The self-exchange rate constants of 
?
SCH
2
CH(NH
3
+
)CO
2
?
/
?
SCH
2
CH(NH
3
+
)CO
2
?
 was 
also obtained from the reaction between L-cysteine and [Fe(bpy)
2
(CN)
2
]
+
 by applying 
Marcus theory, with 2.50 ? 10
6
 M
?1
 s
?1
. No observance of a long lived intermediate state 
in the reaction and the rapid self-exchange rate constants of cysteine/cysteine radical 
imply that the oxidation of L-cysteine by the Fe(III) complexes may occur through a one-
electron outer-sphere transfer reaction.  
 
 116
 
 
 
 
 
CHAPTER THREE  
OXIDATION OF IODIDE BY A SERIES OF Fe(III)  
COMPLEXES IN ACETONITRILE 
 
Introduction 
The general features of the kinetics and mechanisms of oxidation of excess iodide 
by typical outer-sphere oxidants in aqueous solution are well understood. Currently, there 
are insufficient experimental data about the kinetics of oxidation of iodide in nonaqueous 
solvent. All of the relevant studies are carried out in aqueous-nonaqueous cosolvent 
systems, and are limited to three oxidants: [Fe
III
(phen)
3
]
3+
, [Ir
IV
Cl
6
]
2-
 and [Co
III
W
12
O
40
]
5-
.
144-146
 The reaction of [Fe
III
(phen)
3
]
3+
 was studied in water/methanol with up to 20% 
alcohol.
144
 In the case of [Ir
IV
Cl
6
]
2-
, a variety of aqueous solvent mixtures were used, but 
the aqueous fraction was never < 20%.
145
 With [Co
III
W
12
O
40
]
5-
, aqueous solvent mixtures 
again were used in this case, with up to 40% methanol, 40% acetonitrile, and 60% 
dimethyl sulfoxide (DMSO).
146
 Hence, little is revealed about the oxidation of iodide in 
genuine nonaqueous media. Moreover, the association constant of I
-
 with I
2
 in acetonitrile 
is much higher than that in aqueous solution,
147
 which introduces a strong driving force 
for the oxidation of iodide in acetonitrile. Therefore, this redox reaction in acetonitrile 
may have some new kinetic features. 
 117
The development of dye-sensitized solar cell (DSSC) prompts us to understand the 
oxidation of iodide in acetonitrile. In 1990s, Gr?tzel and his coworkers initially invented 
nanoncrystalline dye-sensitized solar cell (DSSC) with a solar-to-electric energy 
conversion efficiency of 11%.
57,64
 They are typically fabricated with the two electrodes 
immersed in a nonaqueous solvent with an I
-
/I
3
-
 electrolyte that mediates the redox 
reaction. One of the electrodes is composed of TiO
2
 with a surface-adsorbed dye, 
(Ru(4,4?-dicarboxylic acid-2,2?-bipyridine)
2
(NCS)
2
.
64,148
 It is generally believed that one 
function of the iodide is to reduce the photochemically oxidized dye back to its resting 
state.
60,72
 A variety of [Ru
II
(bpy)
2
(SCN)
2
] derivatives have been used in the Gr?tzel 
photoelectrochemical cell over the last decade. Unfortunately, the poor stability of 
[Ru
III
(bpy)
2
(SCN)
2
]
+
 derivatives in acetonitrile impedes studies of its reduction by iodide. 
Recently, it was observed that electron transfer from the excited states of the 
adsorbed sensitizers to the conduction band of TiO
2
 occurs within hundreds of 
femtoseconds.
78,79,149
 With its near unity quantum efficiency, it implies that this electron 
transfer occurs from initially populated, nonrelaxed excited states, not from the lowest 
excited states.
150
 This new theory makes it possible to introduce iron(II) polypyridyls to 
replace current [Ru
II
(bpy)
2
(SCN)
2
] derivatives. More recently, Ferrere and Gregg 
demonstrated that [Fe(4,4?-dicarboxylic acid-2,2?-bipyridine)
2
(CN)
2
], an iron(II) 
polypyridyl, can sensitize nanocrystalline TiO
2
 in acetonitrile/3-methyl-2-oxazolidinone 
(9/1).
151
 So iron(II) polypyridyls can be potentially effective and inexpensive dyes in 
DSSCs. Therefore, four different substitution-inert Fe(III) complexes were selected, 
including [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
, [Fe
III
(bpy)
2
(CN)
2
]
+
, [Fe
III
(dmbpy)
2
(CN)
2
]
+
 and 
[Fe
III
(CH
3
Cp)
2
]
+
. 
 118
Herein, the kinetics and mechanism of the oxidation of excess iodide by a series of 
substitution-inert Fe(III) complexes in acetonitrile are studied. The copper catalysis was 
effectively inhibited with the addition of 5.0 mM 2,2?-bipyridine (bpy). The rate of the 
copper-catalysis inhibited reaction is closely related to the half-wave potentials of the 
Fe(III) complexes: the higher of the half-wave potential, the faster of the reaction it is. 
The rate law of the direct oxidation of iodide has the same form as that in aqueous media. 
The standard potential of I
?
/I
-
 is derived from the kinetic inhibition by Fe
II
(bpy)
2
(CN)
2
. 
 
Experimental Section 
Reagents and Solutions. 2, 2?-bipyridyl (Aldrich), 4, 4?-dimethyl-2, 2?-bipyridyl 
(Aldrich), 5-chloro-phenanthroline (GFS Chemicals), ferrous ammonium sulfate 
hexahydrate (Fisher), potassium cyanide (Fisher), sodium hexafluorophosphate (Aldrich), 
ferric nitrate (Fisher), decamethylferrocene (Aldrich), cupric nitrate trihydrate (Fisher), 
iodine (J. T. Baker), silver nitrate (Fisher), nitric acid (Fisher), sulfuric acid (Fisher), 
acetonitrile (Fisher), acetonitrile-d
3
 (Aldrich), acetone (Fisher), n-hexane (Fisher), 
methanol (Fisher), ethanol (Fisher), acetic acid (Fisher), diethyl ether (Fisher), and 
Sephadex LH-20-100 resin (Sigma) were used without further purification. 1,1?-
dimethylferrocene (Aldrich) was recrystallized from ethanol. Et
4
NBF
4
 (Aldrich) was 
recrystallized three times from a mixture of methanol and n-hexane (4:1) and dried under 
vacuum at 96 
o
C for 12 hours. Et
4
NI (Aldrich) was recrystallized from water and dried 
under vacuum at 100 
o
C for 12 hours. The concentration of Et
4
NI in acetonitrile was 
standardized by titration with standard aqueous AgNO
3
, using Eosin as indicator.
152
 
 119
Distilled deionized water was obtained from a Barnstead NANO pure infinity 
ultrapure water system. Solutions of Fe(III) complexes and iodide in acetonitrile were 
prepared just prior to use, and kept in the dark to prevent any photochemical change. For 
reaction between [Fe
III
(CH
3
Cp)
2
]
+
 and iodide conducted in the stopped-flow 
spectrophotometer, both Et
4
NI and [Fe
III
(CH
3
Cp)
2
]
+
 were prepared maintaining the 
appropriate concentrations of the other reagents, purged with argon gas, and transferred 
via gastight glass syringe. For the oxidation of iodide by the other three Fe(III) 
complexes, the reactants are not sensitive to O
2
. However, all solutions were purged with 
Ar or N
2
 before reaction to prevent potential complications caused by O
2
. 
Preparation of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O and [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O. 
[Fe
II
(bpy)
2
(CN)
2
]?3H
2
O and [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O were prepared by following 
Schlit?s method.
107
 The detailed procedures for the synthesis of these two iron complexes 
are described in Chapter Two. 
1
H-NMR (400 MHz/CD
3
CN) of [Fe
II
(bpy)
2
(CN)
2
]?3H
2
O: 
8.94 (s, 2H); 8.34 (d, J = 8.1, 2H), 8.19 (d, J = 8.0, 2H), 7.96 (t, J = 11.6, 2H), 7.70 (d, J = 
5.8, 2H), 7.31 (s, 2H), 7.15 (s, 2H), 6.50 (s, 2H). 
Preparation of [Fe
II
(4,4?-dimethyl-bipyridine)
2
(CN)
2
]?3H
2
O. [Fe
II
(4,4?-
dimethyl-bipyridine)
2
(CN)
2
]?3H
2
O ([Fe
II
(dmbpy)
2
(CN)
2
]?3H
2
O) was prepared as 
described in the literature.
153
 0.378 g (0.96 mmol) (NH
4
)
2
Fe(SO
4
)
2
?6H
2
O, 0.535 g (2.90 
mmol) 4,4?-dimethyl-2,2?-bipyridyl (dmbpy), and 30.0 mL deionized water were mixed 
together. 5.0 mL acetone was added to promote dissolution and reaction. Then the 
solution was heated to remove acetone. A solution of 0.995 g KCN in 4 mL water was 
added all at once. The solution was heated to near boiling for 30 minutes, and allowed to 
cool overnight. The precipitate was collected by vacuum filtration, rinsed with water and 
 120
diethyl ether. Pure [Fe
II
(dmbpy)
2
(CN)
2
]?3H
2
O was obtained by passing a saturated 
ethanol solution through a column of Sephadex LH-20-100 resin, eluting with ethanol, 
and removing ethanol by rotary evaporation. Anal. Calcd for C
26
FeH
30
N
6
O
3
: C, 58.88; H, 
5.70; N, 15.84. Found: C, 58.86; H, 5.58; N, 15.47. 
1
H-NMR (400 MHz/CD
3
CN): 9.52 (d, 
J = 5.8, 2H); 8.14 (s, 2H), 8.11 (s, 2H), 7.36 (d, J = 5.8, 2H), 7.12 (d, J = 5.8, 2H), 7.00 (d, 
J = 5.8, 2H), 2.55 (s, 6H), 2.39 (s, 6H). 
Preparation of [Fe
III
(dmbpy)
2
(CN)
2
]NO
3
?2H
2
O. The synthesis of this Fe(III) 
compound was similar to that of [Fe
III
(bpy)
2
(CN)
2
]NO
3
?2H
2
O. Fe
II
(dmbpy)
2
(CN)
2 
was 
oxidized by concentrated HNO
3
 to form [Fe
III
(dmbpy)
2
(CN)
2
]NO
3
. Then it was purified 
by recrystallization from hot distilled deionized water. Yield of 
[Fe
III
(dmbpy)
2
(CN)
2
]NO
3
?2H
2
O: 73%. Anal. Calcd for C
26
FeH
28
N
7
O
5
: C, 54.37; H, 4.91; 
N, 17.07. Found: C, 54.39; H, 4.79; N, 17.24. 
1
H-NMR (400 MHz/CD
3
CN): 8.94 (s, 2H); 
8.34 (d, J = 8.1, 2H), 8.18 (d, J = 7.9, 2H), 7.95 (t, 2H), 7.69 (d, 2H), 7.31 (s, 2H), 7.15 (s, 
6H), 6.50 (s, 2H). 
Preparation of [Fe
III
(CH
3
Cp)
2
]PF
6
. [Fe
III
(CH
3
Cp)
2
]PF
6 
was prepared by 
following Pladziewicz?s procedure.
154
 0.604 g (2.73 mmol) 1,1?-dimethylferrocene was 
dissolved in 5.0 mL ether. 2.27 g (5.62 mmol) ferric nitrate was added to 5.0 mL of 0.04 
N HCl. The two solutions were mixed together with the appearance of blue color 
immediately. Then 0.724 g (4.31 mmol) sodium hexafluorophosphate was added to the 
above blue solution. The precipitate was collected by vacuum filtration, rinsed with cold 
dilute acid, water, diethyl ether, and dried in a vacuum desiccator. Yield of 
[Fe
III
(CH
3
Cp)
2
]PF
6
: 61%. For UV-vis spectra and kinetic study, [Fe
III
(CH
3
Cp)
2
]PF
6
 was 
 121
recrystallized from hot water, and dried under vacuum for 12 hours. Anal. Calcd for 
C
12
FeH
14
PF
6
: C, 40.14; H, 3.93; F, 31.75. Found: C, 39.89; H, 3.80; F, 31.47. 
Preparation of [Fe
II
(5-Cl-phen)
2
(CN)
2
]?2H
2
O. [Fe
II
(5-Cl-phen)
2
(CN)
2
]?2H
2
O 
was prepared according to Schilt?s method.
155
 0.744 g (3.0 mmol) 5-chloro-1,10-
phenanthroline and 0.399 g (1.0 mmol) ammonium ferrous sulfate hexahydrate were 
added to 250 mL water. The red mixture was stirred and heated at nearing boiling point 
for 30 minutes. A solution of 1.0 g KCN in 4.0 mL H
2
O was added, and it was continued 
heating for 5 more minutes. Then the solution was cooled to room temperature. 3 hours 
later, the solvent was removed by vacuum suction. The residue was rinsed with 15.0 mL 
0.5 N ammonia and 20.0 mL cool water, and dried in vacuum for 4 hours. 
1
H-NMR (400 
MHz/MeOD): 10.04 (dd, 1H); 9.55 (dd, 1H), 8.94 (dd, 1H), 8.70 (dd, 1H), 8.66 (d, 1H), 
8.43 (dd, 2H), 8.36 (s, 1H), 8.14 (dd, 1H), 8.04 (dd, 1H), 7.66 (dd, 1H), 7.62 (dd, 1H), 
7.57 (dd, 1H), 7.51 (dd, 1H). 
Preparation of [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
?2H
2
O. [Fe
III
(5-Cl-
phen)
2
(CN)
2
]NO
3
?2H
2
O was prepared by oxidizing the crude [Fe
II
(5-Cl-
phen)
2
(CN)
2
]?2H
2
O with 70% nitric acid. The Fe(III) complex was precipitated out with 
the addition of some amount of water. It was collected by vacuum suction, rinsed with 
water to remove acid, and dried in vacuum. Yield of [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
?2H
2
O: 
51%. Anal. Calcd for C
26
FeH
18
N
7
Cl
2
O
5
: C, 49.16; H, 2.86; N, 15.43. Found: C, 49.37; H, 
2.66; N, 15.47. 
 
Methods. All single UV-vis spectra were recorded on a HP-8453 diode-array 
spectrophotometer with 1.0 cm rectangular quartz cells, equipped with Brinkman Lauda 
 122
RM6 thermostated water bath to maintain the temperature at 25.0 ? 0.1
 o
C. Cyclic 
voltammograms (CV) and Osteryoung Square Wave Voltammograms (OSWV) were 
recorded at room temperature on a BAS 100 B electrochemical analyzer with scan rate of 
100 mV/s, equipped with BAS cell stand C3 with purging and stirring system, using a 
glassy carbon working electrode, Ag/AgCl (saturated aqueous KCl) reference electrode, 
and a Pt wire auxiliary electrode. In acetonitrile, the standard potential of [Fe(Cp)
2
]
+/0
 is 
close to that of [Fe(dmbpy)
2
(CN)
2
]
+/0
, so decamethylferrocene was used as internal 
reference in all CV and OSWV experiments. Since the standard potential of [Fe(Cp)
2
]
+/0
 
versus [Fe(Cp
*
)
2
]
+/0
 is known,
156,157
 it is of no difficulty to present the standard potentials 
of the four Fe(III) complexes with reference to [Fe(Cp)
2
]
+/0
. 
1
H NMR spectra were 
recorded on a Bruker AV 400 spectrometer. 
For fast reactions, kinetic studies were performed by mixing equal volumes of the 
two reactants on a Hi-Tech Scientific model SF-51 stopped-flow apparatus that is 
equipped with a SU-40 spectrophotometer unit and a C-400 circulatory water bath. For 
slow reaction, equal volumes of the two reactants were mixed in cuvette, then kinetic 
studies were immediately performed on HP-8453 diode-array spectrophotometer that is 
equipped with a RMS Lauda circulatory water bath, from which the rate constants (k
obs
) 
were obtained by fitting the experimental data with the software that was provided by HP. 
Temperature was maintained at 25.0 ? 0.1
 o
C. On Hi-Tech Scientific model SF-51 
stopped-flow apparatus, an Olis 4300S system was used for data acquisition and analysis. 
Reactions were monitored at fixed wavelength, and the rate constants were obtained by 
fitting the data with OLIS-supplied first-order functions. All apparent rate constants were 
the average of at least five runs with ? 5% error or less. A nonlinear-least squares 
 123
computer program with relative weighting (weighting by 1/Y
2
) was used to fit the overall 
rate law to the values of k
obs
. The products were identified by 
1
H NMR spectra and UV-
visible spectra. 
 
Results
 
1. UV-visible spectra. Figure 3-1 shows the UV-visible spectra of 0.10 mM 
Fe
II
(bpy)
2
(CN)
2
 and 0.10 mM [Fe
III
(bpy)
2
(CN)
2
]NO
3
 in acetonitrile. For Fe
II
(bpy)
2
(CN)
2
, 
there are two characteristic absorbance peaks at 388 and 605 nm, for which the extinction 
coefficients (?) are 6988 and 7360 M
-1 
cm
-1
; for [Fe
III
(bpy)
2
(CN)
2
]NO
3
, the absorbance at 
605 nm is nearly zero. Iodide does not absorb in the visible range, and triiodide (I
3
-
), one 
of products of the reaction that was confirmed by our experimental results, has 
characteristic absorbance at 365 nm (?
365
 = 26,250 M
-1
 cm
-1
; ?
605 
?10 M
-1 
cm
-1
). Thus, for 
the oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
, the stopped-flow experiments were 
monitored at the wavelength of 605 nm. The UV-visible absorbance characteristics of all 
iron complexes in acetonitrile are shown in Table 3-1, in which the extinction coefficients 
are very close to previous reports.
153,158-160
 The wavelength for the oxidization of iodide 
by [Fe
III
(dmbpy)
2
(CN)
2
]
+
, [Fe
III
(CH
3
Cp)
2
]
+
 and [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 are monitored 
at 608, 650, and 609 nm, respectively. 
2. Electrochemistry and NMR spectra. The CV and OSWV of 0.50 mM 
ferrocene in 0.10 M Et
4
NBF
4(AN)
, with 0.40 mM decamethylferrocene as internal 
reference, were recorded. The OSWV graph (Figure 3-2) shows that the half-wave 
potential, E
1/2
, of [Fe(Cp)
2
]
+/0
 is 504 mV higher than that of [Fe(Cp*)
2
]
+/0
, the same as 
previous reports.
156,157
 
 124
 
Table 3-1. UV-visible absorbance characteristics of the iron complexes in acetonitrile 
Compounds Band 
?
max
, nm ?, M
-1
 cm
-1
I 388 6988 [Fe
II
(bpy)
2
(CN)
2
] 
II 605 7360 
I 301 25037 
[Fe
III
(bpy)
2
(CN)
2
]NO
3
 
II 503 226 
I 383 7370 
[Fe
II
(dmbpy)
2
(CN)
2
] 
II 608 6930 
[Fe
III
(dmbpy)
2
(CN)
2
]NO
3
 I 356 3884 
I 324 76 
[Fe
II
(CH
3
Cp)
2
] 
II 438 110 
[Fe
III
(CH
3
Cp)
2
]PF
6
 I 650 354 
[Fe
II
(5-Cl-phen)
2
(CN)
2
] I 609 10700 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 I 362 3740 
 
 125
 
300 400 500 600 700 800 900
0.0
0.2
0.4
0.6
0.8
Fe(II)
Fe(III)
Wavelength, nm
Abs
o
r
b
a
n
c
e
 
Figure 3-1. UV-visible spectra of 0.10 mM Fe
II
(bpy)
2
(CN)
2
 (? ) and 0.10 mM  
[Fe
III
(bpy)
2
(CN)
2
]NO
3
 (---) in acetonitrile 
 
 126
 
 
Figure 3-2. The OSWV of 0.50 mM Fe
II
(Cp)
2
 and 0.40 mM Fe(Cp*)
2
 
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an 
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 127
Fe
II
(bpy)
2
(CN)
2
. The CV (Figure 3-3) of 1.0 mM Fe
II
(bpy)
2
(CN)
2
 in 0.10 M 
Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal reference, is quasi-reversible, 
with ?E
p/p
 = 72 mV, E
1/2
 = 71 mV vs. [Fe(Cp)
2
]
+/0
. The OSWV (Figure 3-4) of 1.0 mM 
Fe
II
(bpy)
2
(CN)
2
 in 0.10 M Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal 
reference, has the half-potential, E
1/2
, of 70 mV, nearly the same as the result from CV. 
 
 
Figure 3-3. The CV of 1.0 mM Fe
II
(bpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an  
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 128
 
 
Figure 3-4. The OSWV of 1.0 mM Fe
II
(bpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an  
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 129
Fe
II
(dmbpy)
2
(CN)
2
. The CV (Figure 3-5) of 1.0 mM Fe
II
(dmbpy)
2
(CN)
2
 in 0.10 M 
Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal reference, is quasi-reversible, 
with ?E
p/p
 = 62 mV, E
1/2
 = -35 mV. The OSWV (Figure 3-6) of 1.0 mM 
Fe
II
(dmbpy)
2
(CN)
2
 in 0.10 M Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal 
reference, has the half-potential of -36 mV, nearly the same as that obtained from CV. 
 
 
Figure 3-5. The CV of 1.0 mM Fe
II
(dmbpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
 
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an 
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 130
 
 
Figure 3-6. The OSWV of 1.0 mM Fe
II
(dmbpy)
2
(CN)
2
 and 1.0 mM Fe(Cp*)
2
  
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an Ag/AgCl
(s)
  
electrode as reference, and a Pt wire as a counter electrode.
 
 
 131
Fe
II
(CH
3
Cp)
2
. The CV (Figure 3-7) of 1.0 mM Fe
II
(CH
3
Cp)
2
 in 0.10 M 
Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal reference, is quasi-reversible, 
with ?E
p/p
 = 82 mV, E
1/2 
 = -110 mV. The OSWV (Figure 3-8) of 1.0 mM Fe
II
(CH
3
Cp)
2
 
in 0.10 M Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal reference, has the 
half-wave potential of -108 mV. The half-wave potential from OSWV is 2 mV higher 
than that from CV. 
 
 
Figure 3-7. The CV of 1.0 mM Fe
II
(CH
3
Cp)
2
 and 1.0 mM Fe(Cp*)
2
 
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an 
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 132
 
 
Figure 3-8. The OSWV of 1.0 mM Fe
II
(CH
3
Cp)
2
 and 1.0 mM Fe
II
(Cp*)
2
  
in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, an  
Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 133
[Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3 
and Fe
II
(5-Cl-phen)
2
(CN)
2
. No decamethylferrocen-
e was added in the CV of [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 due to the redox reaction between 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 and decamethylferrocene. The CV (Figure 3-9) of 0.50 mM 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 in 0.10 M Et
4
NBF
4(AN)
 is quasi-reversible, with ?E
p/p
 = 91 
mV, E
1/2 
= 610 mV (vs. Ag/AgCl
(s)
), i.e. 176 mV vs [Fe(Cp)]
+/0
.
161-163
 The OSWV 
(Figure 3-10) of 0.20 mM Fe
II
(5-Cl-phen)
2
(CN)
2
 in 0.10 M Et
4
NBF
4(AN)
, with 1.0 mM 
decamethylferrocene as internal reference, has the half-wave potential of 168 mV. 
 
 
Figure 3-9. The CV of 0.50 mM [Fe
III
(5-Cl-phen)
2
(CN)
2
]NO
3
 in 0.10 M  
Et
4
NBF
4
. With a glassy carbon disc as working electrode, an Ag/AgCl
(s)
  
electrode as reference, and a Pt wire as a counter electrode. 
 134
 
 
Figure 3-10. The OSWV of 0.20 mM Fe
II
(5-Cl-phen)
2
(CN)
2
 and 1.0 mM  
Fe(Cp*)
2
 in 0.10 M Et
4
NBF
4
. With a glassy carbon disc as working electrode, 
an Ag/AgCl
(s)
 electrode as reference, and a Pt wire as a counter electrode. 
 
 135
The half-wave potentials of the four iron complexes versus [Fe(Cp)
2
]
+/0
 are shown 
in Table 3-2. In the case of Fe
II
(CH
3
Cp)
2
, it has the same half-wave potential as previous 
report;
157
 while half-wave potential of [Fe(dmbpy)
2
(CN)
2
]
+/0
 is 58 mV higher than the 
literature.
153
 However, nobody reported the half-wave potentials of [Fe(bpy)
2
(CN)
2
]
+/0
 
and [Fe(5-Cl-phen)
2
(CN)
2
]
+/0
 in acetonitrile. The half-wave potentials in Table 3-2 
demonstrate that [Fe
III
(CH
3
Cp)
2
]
+
 is the weakest oxidant, and [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 is 
the strongest oxidant, with a difference of 276 mV between them; [Fe
III
(bpy)
2
(CN)
2
]
+
, 
one mild oxidant, has 108 mV higher than [Fe
III
(dmbpy)
2
(CN)
2
]
+
. The rate constants of 
the oxidation of the iodide are closely related to the half-wave potential of the iron 
complexes. This will be discussed in kinetics section.  
 
Table 3-2. The half- wave potentials of the iron complexes in acetonitrile
 
 
Compounds E
1/2
, mV 
a
 
[Fe(CH
3
Cp)
2
]
+/0 
-108 ? 2 
[Fe(dmbpy)
2
(CN)
2
]
+/0 
- 36 ? 2 
[Fe(bpy)
2
(CN)
2
]
+/0 
+ 70 ? 2 
[Fe(5-Cl-phen)
2
(CN)
2
]
+/0
 168 ? 2 
a
 E
1/2 
vs [Fe(Cp)
2
]
+/0
, at 22.0 
o
C and ? = 0.10 M. E
1/2
 = E
f
 
 
1
H-NMR spectra. The 
1
H-NMR spectra of Fe
II
(bpy)
2
(CN)
2
 and Fe
II
(dmbpy)
2
(CN)
2
 
show that both of them have cis geometries without any impure component. The 
1
H-
NMR spectrum (Figure 3-11) of recrystallized Fe(CH
3
Cp)
2
 has two characteristic 
chemical shift at 3.94 and 1.95 ppm (overlap with CD
3
CN), with the intensity of 4:3. For 
cis-Fe
II
(5-Cl-phen)
2
(CN)
2
, it has three isomers, as shown in Scheme 3-1. The 
1
H-NMR 
 136
spectrum of 1.0 mM Fe
II
(5-Cl-phen)
2
(CN)
2
 in MeOD has 14 peaks with the same 
intensity, confirming that only isomer (C) exists in the compound. So the 
1
H-NMR 
spectra of the four Fe(II) complexes are consistent with their structure. 
 
 
Figure 3-11.
1
H-NMR spectrum of recrystallized Fe(CH
3
Cp)
2
 in CD
3
CN 
 137
Scheme 3-1. The structure of cis-Fe
II
(5-Cl-phen)
2
(CN)
2
 with three different isomers.  
ClN represents the Cl-substituted part of the 5-Cl-phen ligand. 
N
NCN
Fe
NCl
NCl
CN ClN
ClN CN
Fe
N
N
CN N
ClN CN
Fe
NCl
N
CN
(A)
(B) (C)
 
3. Metal-ion catalysis and scavenger effect. We initially studied the kinetics of 
the oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile due to its mild oxidative 
capability. Our preliminary kinetic studies were plagued by its irreproducibility. For a 
typical reaction of 2.0 ? 10
-3
 M NaI with 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 in 0.10 M 
Et
4
NBF
4
, the half-life ranges from 0.31 to 0.81 s. Then 5.0 ?M Ni
2+
, Fe
3+
 and Cu
2+
 were 
deliberately added to the reactant of iodide. After the addition of 5.0 ?M Cu
2+
, the rates 
of reaction are much faster, with the half-life of 3 ms; while there is no distinct change 
for Ni
2+
 and Fe
3+
, as shown in Table 3-3. Therefore, trace of Cu ions are good catalyst for 
the redox reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and iodide in acetonitrile.
164
 Due to the 
copper-catalysis effect, the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 with excess iodide in 
acetonitrile does not follow first-order behavior, so the rate of the reaction is represented 
by half-life, rather than by k
obs
. The copper-catalysis effect was also observed for the 
other three Fe(III) complexes, although it is not distinct for [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 and 
[Fe
III
(CH
3
Cp)
2
]
+
. For the oxidation of 2.0 ? 10
-3
 M NaI by 5.0 ? 10
-5
 M 
 138
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 in 0.10 M Et
4
NBF
4
 at 25.0 
o
C, the half-life for the formation of 
Fe
II
(dmbpy)
2
(CN)
2
 decreases by 33 times after the addition of 1.0 ?M Cu(NO
3
)
2
. For the 
oxidation of 8.0 ? 10
-2
 M NaI by 2.5 ? 10
-4
 M [Fe
III
(CH
3
Cp)
2
]
+
 in 0.02 M Et
4
NBF
4
 at 
25.0 
o
C, the half-life for the formation of Fe
II
(CH
3
Cp)
2
 decreases by 7 times when 1.0 
?M Cu(NO
3
)
2
 was deliberately added. For the oxidation of 0.20 mM I
-
 by 10.0 ?M 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 in 0.10 M Et
4
NBF
4
 at 25.0 
o
C, the half-life for the formation of 
Fe
II
(5-Cl-phen)
2
(CN)
2
 decreases by a factor of four due to the addition of 1.0 ?M 
Cu(NO
3
)
2
. All of the results are shown in Table 3-4. 
 
Table 3-3. Effect of metal ions on the oxidation of Et
4
NI by [Fe
III
(bpy)
2
(CN)
2
]
+
  
in acetonitrile. At ? = 0.10 M (Et
4
NBF
4
), and 25.0 
o
C. 
Molarity of cation, ?M 
t
1/2
, s 
0 0.60 
[Cu
2+
] = 5.0 ?M 0.003 
[Ni
2+
] = 5.0 ?M 0.11 
[Fe
2+
] = 5.0 ?M 0.20 
a
 [Et
4
NI]
0
 = 2.0 mM, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 50 ?M. 
 
To inhibit its copper-catalysis effect, some chelating reagents should be added to 
the reactants. In our group, Saha and Hung
103,52
 have found that 2,2?-bipyridine (bpy), 
2,6-pyridinedicarboxylic acid (dipic), and 1,10-phenanthroline (phen) can effectively 
mask the copper-catalysis effect in the reduction of octacyanomolybdate(V) by 
thioglycolic acid and cysteine in aqueous media. 
 139
Table 3-4. The effect of Cu
2+
 and bpy on the oxidation of Et
4
NI by 50 ?M Fe(III) 
complexes in acetonitrile. At ? = 0.10 M (Et
4
NBF
4
), and 25.0 
o
C. 
Compounds [I
-
]
0
, mM [Et
4
NBF
4
], M [Cu
2+
], ?M [bpy], mM t
1/2
, s
80.0 0.02 0.0 0.0 0.375
80.0 0.02 5.0 0.0 0.044[Fe
III
(CH
3
Cp)
2
]
+
 
a 
80.0 0.02 5.0 5.0 1.16 
2.0 0.10 0.0 0.0 15.50 
2.0 0.10 1.0 0.0 0.46 [Fe
III
(dmbpy)
2
(CN)
2
]
+
 
2.0 0.10 1.0 5.0 86.48 
2.0 0.10 0.0 0.0 0.60 
2.0 0.10 1.0 0.0 0.015[Fe
III
(bpy)
2
(CN)
2
]
+
 
2.0 0.10 1.0 5.0 4.20 
0.2 0.10 0.0 0.0 0.037
0.2 0.10 1.0 0.0 0.011
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+ 
b
 
0.2 0.10 1.0 5.0 0.969
 
a
 [[Fe
III
(CH
3
Cp)
2
]
+
]
0
 = 0.25 mM; 
b
 [[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
]
0
 = 10 ?M 
 
Firstly, the effect of bpy on the oxidation of NaI by [Fe
III
(bpy)
2
(CN)
2
]
+
 was tested. 
1.0 mM and 5.0 mM bpy was added to the reactants (2.0 mM NaI, 50.0 ?M 
[Fe
III
(bpy)
2
(CN)
2
]
+
, 0.098 M Et
4
NBF
4
), respectively. It was found that both reactions 
have the same half-life, 4.20 s. Ethylenediamine (en) also was tested. 1.0 mM 
ethylenediamine (en) was added to the same concentration of the above reactants, the 
half-life is 4.52 s. It indicates that en has the same function as bpy in the inhibition of the 
Cu
2+
 ion catalysis. It was also found that the half-life for the reaction between 5.0 ? 10
-5
 
M [Fe
III
(bpy)
2
(CN)
2
]
+
 and 1.0 ? 10
-3
 M en in acetonitrile is about 700 s, while there is no 
 140
reaction between [Fe
III
(bpy)
2
(CN)
2
]
+
 and bpy. To avoid the complication caused by the 
reaction of en with [Fe
III
(bpy)
2
(CN)
2
]
+
, bpy, rather than en, was selected to inhibit the 
copper catalysis. Direct oxidation of iodide by the four outer-sphere Fe(III) complexes 
was achieved by the addition of 5.0 mM bpy to the reactants. Figure 3-12 shows a kinetic 
trace of reaction between 2.0 ? 10
-3
 M I
-
 and 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
. The inset 
in Figure 3-12 shows that, with the addition of 5.0 mM bpy, the rate of the reaction is 
much slower and it follows pseudo-first-order kinetic behavior; however, it does not 
follow pseudo-first-order kinetics in the absence of bpy. The same kinetic features were 
observed for the reduction of [Fe
III
(CH
3
Cp)
2
]
+
, [Fe
III
(dmbpy)
2
(CN)
2
]
+
 and [Fe
III
(5-Cl-
phen)
2
(CN)
2
]
+
. 
4. Effect of small amount of water on the kinetics study. Our preliminary kinetic 
results demonstrate that the non-catalyzed oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
 is 
much slower in water than in acetonitrile. For the oxidation of 2.0 ? 10
-3
 M NaI by 5.0 ? 
10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
, at ? = 0.10 M and 25.0 
o
C, the half-lives of the reaction are 
0.60 s in acetonitrile and 58.7 s in water, respectively. Therefore, the trace amount of 
water in acetonitrile and reactants may affect the rate of the reaction. The water-content 
effect on the rate of the oxidation of iodide was studied by selecting [Fe
III
(bpy)
2
(CN)
2
]
+
 
as oxidant.  
 
 141
0.0 0.1 0.2 0.3
0.0
0.1
0.2
0.3
Time, s
A
b
sor
b
an
ce
0 10 20 30 40 50 60
0.0
0.1
0.2
0.3
Time, s
A
b
s
o
r
b
an
ce
 
Figure 3-12. Kinetic traces for the oxidation of 2.0 ? 10
-3
 M NaI by 5.0 ? 10
-5
 M  
[Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile. At ? = 0.10 M (Et
4
NBF
4
), and 25.0 
o
C. Data was  
recorded at 605 nm. Main trace with 1.0 ?M Cu(NO
3
)
2
?3H
2
O. Inset shows data  
with 1.0 mM bpy but no added Cu
2+
. Note the vastly differing time scales for the  
two traces. 
 
Firstly, the concentration of water for the nominally anhydrous stopped-flow 
experiment was determined by 
1
H-NMR spectra. 2.0 mM NaI and 0.20 mM 
[Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, in the presence of 0.10 M Et
4
NBF
4
, was mixed with 
equal volume in the syringes of SF-51 instrument. Analysis of the integrate intensities of 
Et
4
NBF
4
 and H
2
O in the 
1
H-NMR spectrum (Figure 3-13) demonstrates that it contains 
0.20% (v/v) water in the product solution. Then 1% water was deliberately added to the 
 142
reactants that contained 1.0 mM bpy, the rate constants decrease from 0.57 to 0.53 s
-1
. In 
the absence of bpy but with 1.0 ?M Cu(NO
3
)
2
, the addition of 1% water to the reactants 
does not change the rate of the reaction, either. Therefore, the effect of small amount of 
water (0.2 %) in nominally anhydrous conditions on the rate of the reaction is negligible. 
 
Figure 3-13. 
1
H-NMR spectrum of product for the reaction of 1.0 mM NaI with 0.10  
mM [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile. In the presence of 0.10 M Et
4
NBF
4
. 
?: ?CH
2
 in Et
4
NBF
4
; ?: H
2
O; ?: ?CH
3
 in Et
4
NBF
4 
 
Since the structure of [Fe
III
(CH
3
Cp)
2
]
+
 is quite different from that of the other three 
Fe(III) complexes, it is of high importance to test the water effect on the oxidation of 
 143
iodide by [Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile. Figure 3-14 shows that k
obs
 decreases linearly 
as the water content increases from 0.2 to 1.0%. When 0.20 % of water was deliberately 
added to the reaction system (80.0 mM NaI, 0.25 mM [Fe
III
(CH
3
Cp)
2
]
+
, 10.0 mM bpy 
and 20.0 mM Et
4
NBF
4
, at 25.0 
o
C), the rate constant decreases from 0.28 to 0.26 s
?1
. For 
all kinetic experiments, acetonitrile was used directly, without any further drying process. 
Furthermore, a series of kinetic experiments was performed with the same concentrations 
of NaI and Et
4
NI. At lower concentration of iodide, the rate constant is nearly the same; 
at higher concentration of iodide, the rate of the reaction is much slower for NaI, as 
shown in Table 3-5. This can be ascribed to the more hygroscopic character of NaI. With 
the increase of NaI, significant great amount of water can be introduced to the reaction 
system. Accordingly, non-hygroscopic Et
4
NI was used, rather than NaI, in all later kinetic 
studies.  
 
 144
0.2 0.4 0.6 0.8 1.0
0.0
0.1
0.2
0.3
H
2
O %
k
ob
s
, s
-1
 
Figure 3-14. Water effect on the rate constants for the oxidation of  
80.0 mM NaI by 0.25 mM [Fe
III
(CH
3
Cp)
2
]
+
. In the presence of 10.0  
mM bpy and 20.0 mM Et
4
NBF
4
, at 25.0 
o
C.  
 
 
Table 3-5. Kinetic comparison of NaI and Et
4
NI in the reduction by  
[Fe
III
(CH
3
Cp)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C. 
Iodide [I
-
], mM k
obs
, s
-1
 
5.0 0.0018 
10.0 0.0058 
20.0 0.0211 
NaI 
40.0 0.0822 
5.0 0.0022 
10.0 0.0066 
20.0 0.0247 
Et
4
NI 
40.0 0.0895 
 145
5. Product identification and stoichiometry. The products of the four copper-
catalysis inhibited reactions were identified by 
1
H-NMR and UV-visible spectra. The 
stoichiometry of the reactions was determined by 
1
H-NMR and spectrophotometric 
analysis. 
Oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
. The products of the non-copper 
catalyzed reaction for the oxidation of Et
4
NI by [Fe
III
(bpy)
2
(CN)
2
]
+
 were identified by 
spectrophotometric analysis. The UV-vis spectrum (Figure 3-15) of the product solution 
of 30.0 mM Et
4
NI and 50 ?M [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, with 0.10 M Et
4
NBF
4
 
and 5.0 mM bpy, and at 25.0 
o
C, shows that there are two maxima peaks at 369 and 605 
nm. Compared with the UV-vis spectra of pure Fe
II
(bpy)
2
(CN)
2
 in acetonitrile, it is 
confirmed that Fe
II
(bpy)
2
(CN)
2
 is one of the products. To assign other product of the 
reaction, UV-vis spectra of triiodide (I
3
-
) were recorded, and it was observed that the I
3
-
 
has characteristic absorbance at 363 nm with the extinction coefficient of 26,250 M
-1 
cm
-1
, 
which is very close to Isci and Mason?s result.
165
 So the absorbance peak at 369 nm for 
the product solution is ascribed to the overlap of absorbance of I
3
-
 and Fe
II
(bpy)
2
(CN)
2
. 
Thus, the products of the copper-catalysis inhibited reaction are I
3
-
 and Fe
II
(bpy)
2
(CN)
2
. 
Quantitative dual-wavelength analysis at 605 and 369 nm (?
605
 = 7360; ?
369
 = 5713 M
-1
 
cm
-1
 for Fe
II
(bpy)
2
(CN)
2
) demonstrates that the yield of Fe
II
(bpy)
2
(CN)
2
 is 99.0%, and 
that the product ratio of ?[I
3
-
]/?[Fe(II)] equals to 0.48 ? 0.03. The above results imply 
that the overall reaction is:  
 
2 Fe(III)
 
 
+ 3 I
-
2 Fe(II) + I
3
-
(3-1)K
tot
 
 146
 
400 500 600 700 800
0.0
0.2
0.4
0.6
0.8
Wavelength, nm
Ab
s
o
r
b
a
n
c
e
 
Figure 3-15. UV-vis spectrum of product for the oxidation of 30.0 mM Et
4
NI by  
50 ?M [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, with 5.0 mM bpy. 
 
Oxidation of iodide by [Fe
III
(dmbpy)
2
(CN)
2
]
+
. The products of the direct 
oxidation of iodide by [Fe
III
(dmbpy)
2
(CN)
2
]
+
 were also identified by spectrophotometric 
analysis. The UV-vis spectrum (Figure 3-16) of the product solution of 5.0 mM Et
4
NI 
and 47 ?M [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, in the presence of 5.0 mM bpy, at ? = 0.10 
M and at 25.0 
o
C, shows that there are two maxima peaks at 368 and 605 nm. The 
spectrophotometric analysis of the spectrum indicates that the products of the reaction are 
I
3
-
 and Fe
II
(dmbpy)
2
(CN)
2
, and that the product ratio of ?[I
3
-
]/?[Fe(II)] is 0.43 ? 0.03, 
consistent with Equation 3-1. 
 147
400 500 600 700 800
0.00
0.25
0.50
0.75
1.00
Wavelength, nm
A
b
s
o
r
b
an
ce
 
Figure 3-16. UV-vis spectrum of product of the reaction between 47 ?M 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 and 5.0 mM Et
4
NI in acetonitrile. In the presence  
of 5.0 mM bpy and 0.095 M Et
4
NBF
4
, at 25.0 
o
C. 
 
Oxidation of iodide by [Fe
III
(CH
3
Cp)
2
]
+
. The products of the direct oxidation of 
iodide by [Fe
III
(CH
3
Cp)
2
]
+
 were identified by 
1
H-NMR spectra and spectrophotometric 
analysis. The 
1
H-NMR spectrum of product solution of the oxidation of 80.0 mM NaI by 
2.5 mM [Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile, in the presence of 10.0 mM bpy and 20.0 mM 
Et
4
NBF
4
, is shown in Figure 3-17, demonstrates that Fe
II
(CH
3
Cp)
2
 is one of the 
products,
166,167
 and the yield of Fe(CH
3
Cp)
2
 is (99 ? 1)%. To identify other product of the 
reaction, UV-vis spectra of the product solution were recorded. Figure 3-18 is the UV-vis 
spectra for the oxidation of 5.0 mM NaI by 0.10 mM [Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile, in 
 148
the presence of 10.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. The product solution 
has strong absorbance at 363 nm that is assigned to that of I
3
-
 (?
363
 ? 8 M
-1
 cm
-1 
for 
Fe
II
(CH
3
Cp)
2
; ?
363
 = 26,250 M
-1
 cm
-1
 for I
3
-
). Thus, the products of the reaction are 
Fe
II
(CH
3
Cp)
2
 and I
3
-
. 
 
 
Figure 3-17. 
1
H-NMR spectrum of product for the oxidation of 80.0 mM NaI  
by 2.5 mM [Fe
III
(CH
3
Cp)
2
]
+
 in CD
3
CN. In the presence of 10.0 mM bpy and  
20.0 mM Et
4
NBF
4
. ? : bpy; ? : Et
4
NBF
4
; ?: Fe(CH
3
Cp)
2
; ?: H
2
O;  
?: overlap of CD
3
CN and -CH
3
 in Fe(CH
3
Cp)
2
. 
 149
 
 
Figure 3-18. The UV-vis spectra for the oxidation of 5.0 mM NaI by 0.10 mM  
[Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile. In the presence of 10.0 mM bpy, at ? = 0.10 M  
(Et
4
NBF
4
) and 25.0 
o
C, 60 seconds interval between two spectra.  
(Inset shows the enlarged spectra from 480 to 950 nm). 
 
The stoichiometric ratio of the reaction was determined from UV-vis spectra. The 
product ratio of ?[I
3
-
]/?[Fe
II
(CH
3
Cp)
2
] is 0.45 (yield of Fe(CH
3
Cp)
2
 is 99%; ?
363
 ? 8  
M
-1
 cm
-1 
for Fe
II
(CH
3
Cp)
2
; ?
363 
= 26,250 M
-1
 cm
-1
 for I
3
-
). Therefore, the stoichiometric 
ratio of above reaction is consistent with Equation 3-1. 
Oxidation of iodide by [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
. The products of the direct 
oxidation of iodide by [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 were identified by 
1
H-NMR spectra. The 
1
H-NMR spectrum (Figure 3-19) of the product of the reaction between 0.50 mM [Fe
III
(5-
 150
Cl-phen)
2
(CN)
2
]
+
 and 3.0 mM Et
4
NI in CD
3
CN has 14 peaks with the same intensity, 
which implies that only one isomer of Fe
II
(5-Cl-phen)
2
(CN)
2
 was formed. The UV-vis 
spectrum of the product for the oxidation of 10.0 mM Et
4
NI by 50 ?M [Fe
III
(5-Cl-
phen)
2
(CN)
2
]
+
 in acetonitrile, with 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C, is shown in 
Figure 3-20. The spectrophotometric analysis indicates that the product ratio of  
?[I
3
-
]/?[Fe
II
(5-Cl-phen)
2
(CN)
2
] is 0.51 ? 0.02. Thus, the stoichiometric ratio of the 
reaction complies with Equation 3-1 as well. 
 
 
Figure 3-19. 
1
H-NMR spectrum of the product of the reaction between  
0.50 mM [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 and 3.0 mM Et
4
NI in CD
3
CN, with 2.0  
mM bpy. 
 151
300 400 500 600 700 800
0.00
0.25
0.50
0.75
Wavelength, nm
A
b
s
o
r
bnc
e
 
Figure 3-20.UV-vis spectra of product of the reaction of 10.0 mM Et
4
NI with  
50 ?M [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 in acetonitrile. In the presence of 5.0 mM bpy,  
at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 
6. Reaction of iodine with bpy. For the oxidation of iodide by Fe(III) complexes in 
acetonitrile, I
2
 was firstly formed in the reaction. Spectroscopic studies have shown that 
iodine reacts with ?-donors to form DI
+
 and I
3
-
.
168,169
 Bpy, one good ?-donor, is known to 
react with iodine very fast to form bpy
2
I
+
 and I
3
-
, as shown in reaction 3-2. Thus, I
2
 may 
combine with bpy to form bpy
2
I
+
 and I
3
-
 in present reaction system. The UV-vis spectra 
of 0.60 mM I
2
 in air-saturated acetonitrile show that it is stable more than 60 minutes. 
Then a solution containing 0.20 mM I
2
 and 5.0 mM bpy was prepared in anaerobic 
acetonitrile, and UV-vis spectra were recorded. These spectra show the slow formation of 
0.027 mM I
3
-
 with a half-life of about 580 s. It implies that the equilibrium constant for 
 152
reaction 3-2 is unfavorable under such dilute conditions. The large associate constant 
(pK
eq
 = -7.18
147
) of triiodide from iodine and iodide in acetonitrile ensures that reaction 
3-2 will shift to backward direction, excluding the possibility of the formation of bpy
2
I
+
 
in the direct oxidation of iodide by Fe(III) complexes. Furthermore, the evidence that the 
Fe(III)/I
-
 rate constants independent of the concentration of bpy supports the insignificant 
role of I
2
/bpy in the current study. 
 
2bpy + 2I
2
 ? bpy
2
I
+
 + I
3
-
                                                                                (3-2) 
 
7. Kinetics. 
 
Oxidation of iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
. According to above results, the 
products of the reaction are Fe
II
(bpy)
2
(CN)
2
 and I
3
-
. It is of importance to know the effect 
of products on the rate of the reaction. For the reaction between 7.5 ? 10
-6
 M 
[Fe
III
(bpy)
2
(CN)
2
]
+
 
and 7.5 ? 10
-5
 M Et
4
NI, in the presence of 0.10 M Et
4
NBF
4
 and 5.0 
mM bpy, various concentrations of Fe
II
(bpy)
2
(CN)
2
 were added to the above reaction 
system. The rate constant decreases with the addition of Fe
II
(bpy)
2
(CN)
2
. Plot of 1/k
obs
 vs 
[Fe
II
(bpy)
2
(CN)
2
] is linear with the slope of (3.07 ? 0.30) ? 10
6
 s/M
 
and intercept of 200 
? 8.7 s, as shown in Figure 3-21. 
 153
0 10 20 30 40 50 60 70
150
200
250
300
350
400
[Fe
II
(bpy)
2
(CN)
2
], ?M
1/
k
obs
, s
 
Figure 3-21. Plot of 1/k
obs
 vs [Fe
II
(bpy)
2
(CN)
2
] for the reaction between  
7.5 ? 10
-6
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 and 7.5 ? 10
-5
 M Et
4
NI in acetonitrile. In  
the presence of 0.10 M Et
4
NBF
4
 and 5.0 mM bpy, at 25.0 
o
C. 
 
Then the reduction of [Fe
III
(bpy)
2
(CN)
2
]
+
 by various concentrations of Et
4
NI in 
acetonitrile, in the presence of 5.0 mM bpy, was investigated under 25.0 
o
C and constant 
ionic strength (? = 0.10 M, Et
4
NBF
4
). A pseudo-first-order dependence of the rate 
constant was established by fitting the experimental curves with one exponential equation. 
The pseudo-first-order rate constants, k
obs
, were obtained from the fits, and replicate runs 
agree to better than 5%. Figure 3-22 is the reaction trace of oxidation of 2.0 ? 10
-4
 M 
Et
4
NI by 5 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
 in acetonitrile, with 5.0 mM bpy and ? = 0.10 M, 
at 25.0 
o
C. It indicates that the rate of the reaction follows pseudo-first-order kinetic 
 154
behavior. The rate constants for various concentrations of I
-
 are shown in Table 3-6. Plots 
of k
obs
/[I
-
] vs [I
-
] (Figure 3-23) are linear, in which the intercept is 58.4 M
-1
 s
-1
, and slope 
is 7760 M
-2
 s
-1
. Based on the above results, the rate law of the reaction can be expressed 
by Equation 3-3, with k
1 
and k
2
 being (2.9 ? 0.1) ? 10
1
 M
-1
 s
-1 
and (3.88 ? 0.14) ? 10
3
 M
-2
 
s
-1
, respectively. 
 
Rate = - d[Fe(III)]/dt = 2 (k
1
[I
-
] + k
2
[I
-
]
2
) [Fe(III)]                                                 (3-3) 
 
Table 3-6. [I
-
] dependence for the oxidation of Et
4
NI by 5.0 ? 10
-5
 M  
[Fe
III
(bpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C. 
 
 [Et
4
NI], mM k
obs
, s
-1
 
0.5 0.028 
1.0 0.0672 
2.0 0.151 
5.0 0.50 
10.0 1.51 
20.0 4.18 
25.0 6.46 
30.0 8.73 
40.0 13.6 
 
 155
 
 
Figure 3-22. Reaction trace for the oxidation of 2.0 ? 10
-3
 M Et
4
NI  
by 5.0 ? 10
-5
 M [Fe
III
(bpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at  
? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 156
0.00 0.01 0.02 0.03 0.04 0.05
0
100
200
300
[I
-
], M
k
obs
/[
I
-
],
s
-1
.M
-1
 
Figure 3-23. k
obs
/[I
-
] vs [I
-
] for the reaction of [Fe
III
(bpy)
2
(CN)
2
]
+
 with  
Et
4
NI. In the presence of 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C,  
[Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 5.0 ? 10
-5
 M. 
 
Oxidation of iodide by [Fe
III
(dmbpy)
2
(CN)
2
]
+
. The rate of reduction of 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 by various concentrations of Et
4
NI in acetonitrile, in the presence 
of 5.0 mM bpy, was investigated under 25.0 
o
C and constant ionic strength (? = 0.10 M, 
Et
4
NBF
4
). The rate constants follow pseudo-first-order kinetics, as shown in Figure 3-24. 
The iodide-dependent rate constants are shown in Table 3-7. Plots of k
obs
/[I
-
] vs [I
-
] 
(Figure 3-25) are linear with intercept of 1.30 M
-1
 s
-1
 and slope of 692 M
-2
 s
-1
. The rate 
law conforms to Equation 3-3, with k
1
 of (0.65 ? 0.05) M
-1
 s
-1
 and k
2 
of (3.46 ? 0.16) ? 
10
2
 M
-2
 s
-1
. 
 157
Table 3-7. [I
-
] dependence for the oxidation of Et
4
NI by 5.0 ? 10
-5
 M  
[Fe
III
(dmbpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C. 
[Et
4
NI], mM k
obs
, s
-1
 
0.5 
7.81 ? 10
-4 
1.0 
1.94 ? 10
-3 
2.0 
5.80 ? 10
-3 
5.0 
2.39 ? 10
-2
 
10.0 
1.07 ? 10
-1
 
20.0 
3.16 ? 10
-1
 
30.1 
6.08 ? 10
-1
 
40.1 1.10 
50.2 1.64 
 
 
Figure 3-24. Reaction trace for the oxidation of 4.01 ? 10
-2
 M Et
4
NI  
by 5.0 ? 10
-5
 M [Fe
III
(dmbpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy,  
at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 158
 
0.00 0.01 0.02 0.03 0.04 0.05
0
10
20
30
[Et
4
NI], M
k
obs
/[E
t
4
NI
]
,
 s
-1
M
-1
 
Figure 3-25. Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by 5.0 ? 10
-5
 M  
[Fe
III
(dmbpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
)  
and 25.0 
o
C. 
 
Oxidation of iodide by [Fe
III
(CH
3
Cp)
2
]
+
. [Fe
III
(CH
3
Cp)
2
]PF
6
 is very air sensitive 
in acetonitrile. The half-life for the decomposition of 1.0 mM [Fe
III
(CH
3
Cp)
2
]PF
6
 in air-
saturated acetonitrile is comparable to that of the reduction by lower concentration of 
iodide. To prevent its decomposition, all of solutions were purged with argon or N
2
 for all 
of kinetic studies.  
The kinetic experiments for the reduction of [Fe
III
(CH
3
Cp)
2
]
+
 by various 
concentrations of Et
4
NI in acetonitrile, in the presence of 5.0 mM bpy, was investigated 
using both SF-51 and HP 8453 diode-array spectrophotometer, under 25.0 
o
C and 
 159
constant ionic strength (? = 0.10 M, Et
4
NBF
4
). Figure 3-26 shows that the rate constants 
follow pseudo-first-order kinetic behavior. The rate constants for various concentrations 
of I
-
 are shown in Table 3-8. Plot of k
obs
/[I
-
] vs [I
-
] (Figure 3-27) are linear with intercept 
of 0.10 M
-1
 s
-1
 and slope of 54.2 M
-2
 s
-1
. The rate law is consistent with Equation 3-3, in 
which k
1
 and k
2
 are (0.05 ? 0.01) M
-1
 s
-1
 and (27.1 ? 1.0) M
-2
 s
-1
, respectively. 
 
0 250 500 750 1000
0.000
0.025
0.050
0.075
0.100
0.125
time, s
A
650
 
Figure 3-26. Reaction traces for the oxidation of 1.50 ? 10
-2
 M Et
4
NI  
by 2.50 ? 10
-4
 M [Fe
III
(CH
3
Cp)
2
]
+
. In the presence of 5.0 mM bpy,  
at ? = 0.10M and 25.0 
o
C, using first-order fitting (k
obs
 = 0.0125 s
-1
). 
(Red dash line: pseudo-first-order fitting) 
 160
Table 3-8. [I
-
] dependence for the oxidation of Et
4
NI by 2.50 ? 10
-4
 M [Fe
III
(CH
3
Cp)
2
]
+
.  
In the presence of 5.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 [Et
4
NI], mM k
obs
, s
-1
 
0.0050 0.00208 
0.0073 0.00370 
0.010 0.00611 
0.0142 0.0125 
0.020 0.0247 
0.030 0.055 
0.040 0.0895 
0.050 0.143 
 
0.00 0.01 0.02 0.03 0.04 0.05 0.06
0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
[I
-
], M
k
ob
s
/[I
-
],
 M
-1
s
-1
 
Figure 3-27. Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by  
[Fe
III
(CH
3
Cp)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M  
and 25.0 
o
C. [Fe
III
(CH
3
Cp)
2
+
]
0
 = 2.50 ? 10
-4
 M 
 161
Oxidation of iodide by [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
.
 The reaction of [Fe
III
(5-Cl-
phen)
2
(CN)
2
]
+
 with excess iodide in acetonitrile, in the presence of 5.0 mM bpy, at ? = 
0.10 M and 25.0 
o
C, follows pseudo-first-order kinetics. The rate constants for various 
concentrations of I
-
 are shown in Table 3-9. Plots of k
obs
/[I
-
] vs [I
-
] are shown in Figure 3-
28, with the intercept of 2900 M
-1
 s
-1
 and slope of 3.36 ? 10
5
 M
-2
 s
-1
. The experimental 
data are not exactly linear in Figure 3-28, which implies that k
1
 is dominant in the above 
reaction, especially at the concentration of iodide higher than 5.0 mM. The rate law is 
consistent with Equation 3-3, in which k
1
 and k
2 
are (1.45 ? 0.03) ? 10
3
 M
-1
 s
-1
 and (1.68 
? 0.11) ? 10
5
 M
-2
 s
-1
, respectively.  
 
Table 3-9. [I
-
] dependence for the oxidation of Et
4
NI by 10
-5
 M [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
. 
In the presence of 5.0 mM bpy, at ? = 0.10 M and 25.0 
o
C. 
 [Et
4
NI], mM k
obs
, s
-1
 
0.20 0.57 
0.50 1.61 
1.00 3.28 
2.60 9.36 
4.00 17.8 
5.00 23.1 
6.00 29.2 
7.00 36.7 
8.00 40.0 
9.00 45.4 
10.0 52.2 
 
 
 162
0.0000 0.0025 0.0050 0.0075 0.0100 0.0125
0
1000
2000
3000
4000
5000
6000
                                    [I
-
], M
k
ob
s
/[
I
-
],
 s
-1
M
-1
 
Figure 3-28. Plot of k
obs
/[I
-
] vs [I
-
] for the oxidation of Et
4
NI by  
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M  
and 25.0 
o
C. [Fe
III
(5-Cl-phen)
2
(CN)
2
+
]
0
 = 10
-5
 M 
 
Table 3-10. The rate constants for the oxidation of iodide by Fe(III) complexes in 
acetonitrile. In the presence of 5.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
Oxidants E
1/2
, mV 
a
k
1
, M
-1 
s
-1
 k
2
, M
-2  
s
-1
 
[Fe
III
(CH
3
Cp)
2
]
+
 -108 ? 2 0.05 ? 0.01 27.1 ? 1.0 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 - 36 ? 2 0.65 ? 0.05 (3.46 ? 0.16) ? 10
2
 
[Fe
III
(bpy)
2
(CN)
2
]
+
 + 70 ? 2 29.2 ? 0.82 (3.88 ? 0.14) ? 10
3
 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 168 ? 2 (1.45 ? 0.03) ? 10
3
 (1.68 ? 0.11) ? 10
5
 
a
 The E
1/2 
vs [Fe(Cp)
2
]
+/0  
 
The rate constants for the above reactions are summarized in Table 3-10. For the 
reduction of [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
, it has the highest rate constants; while in the case 
 163
of [Fe
III
(CH
3
Cp)
2
]
+
, it has the smallest values. The kinetic experimental results indicate 
that the rate constants of the above reactions is proportional to the oxidative capability of 
the four Fe(III) complexes. 
 
Discussion 
The half-wave potentials of Fe(III) complexes are obtained from CV and OSWV, 
as shown in Table 3-10. The equilibrium constants (K
tot
, Equation 3-1) for the oxidation 
of iodide by Fe(III) complexes in acetonitrile can be calculated provided that the standard 
potential of I
3
-
/I
-
 is known.  
The standard potential of I
3
-
/I
-
 was reported by several groups in their voltammetric 
studies.
163,170-172
 However, the employment of various reference electrodes in their 
determination yields seemingly incomparable values. The above impedance was 
overcome by converting them relative to the same reference, ie. [Fe(Cp)
2
]
+/0
. The results 
from our calculation (see appendix A) indicate that their determination gave similar 
standard-potential value of I
3
-
/I
-
: from Popov and Geske, E
?
(I
3
-
/I
-
) = -0.36 V
163
 vs 
[Fe(Cp)
2
]
+/0
; from Desbarres et al., E
?
(I
3
-
/I
-
) = -0.337 V
170
 vs [Fe(Cp)
2
]
+/0
; from Benoit et 
al., E
?
(I
3
-
/I
-
) = -0.315 V 
171
 vs [Fe(Cp)
2
]
+/0
; from Nelson et al., E
?
(I
3
-
/I
-
) = -0.329 V 
172
 vs 
[Fe(Cp)
2
]
+/0
; from Datta et al., E
?
(I
3
-
/I
-
) = -0.372 V 
172
 vs [Fe(Cp)
2
]
+/0
. The average value, 
E
?
(I
3
-
/I
-
) = - (0.35 ? 0.02) V vs [Fe(Cp)
2
]
+/0
, was used in the calculation of equilibrium 
constants in Equation 3-1. Furthermore, the known equilibrium constant for the 
association of I
-
 with I
2
 to form I
3
-
 in acetonitrile ensures us to calculate the equilibrium 
constants in Equation 3-5. All of the equilibrium constants for the reaction of I
-
 with the 
 164
four Fe(III) complexes are shown in Table 3-11. Even for the oxidation of I
-
 by the 
weakest oxidant, [Fe
III
(CH
3
Cp)
2
]
+
, it is still favorable for the forward direction due to the 
higher equilibrium constants of K
I2
 and K
tot
. Theoretically, the yield of Fe(II) for these 
four redox reactions can reach 100%. Our results in product identification and 
stoichiometry section support this thermodynamic calculation. 
 
I
2
 
      
+ (3-4)K
eq
 =10
7.18
 I
3
-
 I
-
2 Fe(III)
 
 
+ 2 I
-
2 Fe(II) + I
2
(3-5)K
I2
 
 
Table 3-11. The equilibrium constants (K
tot
 and K
I2
) for the oxidation of iodide by Fe(III)  
complexes in acetonitrile. 
Compounds K
tot
, M
-2
 K
I2
, M
-1
 
[Fe
III
(CH
3
Cp)
2
]
+
 
7.0 ? 10
7
 
4.6 
[Fe
III
(dmbpy)
2
(CN)
2
]
+
 
1.9 ? 10
10
 
1.2 ? 10
3
 
[Fe
III
(bpy)
2
(CN)
2
]
+
 
7.3 ? 10
13
 
4.8 ? 10
6
 
[Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
 
1.5 ? 10
17
 1.0 ? 10
10
 
 
The results from CV and OSWV show that [Fe(5-Cl-phen)
2
(CN)
2
]
+/0
 has the 
highest standard potential (E
?
), and that [Fe(CH
3
Cp)
2
]
+/0
 has the lowest value. The kinetic 
results for the oxidation of I
-
 by four Fe(III) complexes are closely related to their 
standard potential values: the rate constant is the fastest for the reaction of [Fe
III
(5-Cl-
phen)
2
(CN)
2
]
+
, while it has the slowest rate constant for the reaction of [Fe
III
(CH
3
Cp)
2
]
+
. 
 165
Plots of log k versus E
f
 (LFER curves) in Figure 3-29 are linear with the slope of 16.1 
(log k
1
) and 13.3 (log k
2
), respectively. The higher slope of log k
1
 than log k
2
 implies that 
the rate constant k
1
 will be dominant when the half-wave potential of substitution-inert 
transition metal complex reaches to some value. Moreover, Linear Free-Energy 
Relationship (LFER) curves imply that the reaction may follow one-electron outer-sphere 
mechanism. 
-0.2 -0.1 0.0 0.1 0.2
-2
-1
0
1
2
3
4
5
6
log k
1
log k
2
E
f
(vs [Fe(Cp)
2
]
+/0
), V
lo
g
k
 
Figure 3-29. Plot of log k vs E
f
 for the oxidation of I
-
 by [Fe
III
(5-Cl-phen)
2
(CN)
2
]
+
, 
[Fe
III
(bpy)
2
(CN)
2
]
+
, [Fe
III
(dmbpy)
2
(CN)
2
]
+
 and [Fe
III
(CH
3
Cp)
2
]
+
 in acetonitrile. 
 
The oxidation of iodide by substation-inert transition metal complexes in aqueous 
solution has been well studied.
83,85
 It shows that no copper-catalysis effect is present in 
the reaction. For the copper-catalysis inhibited oxidation of I
-
 by the Fe(III) complexes in 
acetonitrile, it is possible that the reaction follows the similar mechanism. Here, the 
 166
mechanism of the copper-catalysis inhibited oxidation of I
-
 by Fe(III) complexes in 
acetonitrile is proposed as:  
 
Fe(III)   +   I
-
Fe(II)  +  I (3-6)
Fe(III)  +  2I
-
Fe(II)
 
+ I
2
(3-7)
I   +   I
-
              I
2
                                                K
rad
= k
3
 / k
-3
(3-8)
Fe(III)  +  I
2
Fe(II) +  I
2 
                                   k
4
                               (3-9)
I
-
 + I
2 
            I
3
-
                                                           K
eq 
                           ( 3-10)
k
1
, k
-1
k
2
, k
-2
-
 
 
 
-
-
 
Applying steady-state approximation for I
?
 and I
2
?-
, the rate law of the reaction is 
represented by Equation 3-11: 
 
]Fe[]I][Fe[]Fe][I[]Fe][Fe[]Fe[
]I[]Fe]}[I[]Fe[{2
rate
II
31
III
43
II
32
IIIII
14
2II
21
22III
432
II
142431
??
??
????
??
?
++++
++
=
kkkkkkkkkk
kkkkkkkkk
   (3-11) 
 
In the above equation, the rate constants k
1
, k
-1
, k
2
, k
-2
, k
3
 and k
-3
 are not 
independent of each other. They are correlated with the following equation: 
 
2
2
3
3
1
1
???
=
?
?
?
?
?
?
?
?
?
?
?
?
?
?
?
?
k
k
k
k
k
k
                                                                                  (3-12) 
 
 167
In the copper-catalysis inhibited oxidation of excess iodide by [Fe
III
(bpy)
2
(CN)
2
]
+
 
in acetonitrile, it was observed that the rate of the reaction decreases with the addition of 
Fe
II
(bpy)
2
(CN)
2
. Supposes that k
4
[Fe(III)] is much higher than k
-2
[Fe(II)], and k
4
[Fe(III)] 
is much higher than k
-3
, Equation 3-11 is simplified into such form: 
2III
2
3
II
1
31
]I][Fe[2
]I[]Fe[
2
rate
?
?
?
?
?
?
?
?
?
?
?
+
+
= k
kk
kk
                                                   (3-13) 
 
Where k
obs
 is:  
 
2
2
3
II
1
31
obs
]I[2
]I[]Fe[
2
?
?
?
?
?
?
?
?
?
?
?
+
+
= k
kk
kk
k                                                            (3-14) 
 
Furthermore, Equation 3-14 can be converted into another form: 
][I2
1
][Fe
]I[2]I[2
1
1
II
2
31
1
2
2obs
??
?
?
+=
? kkk
k
kk
                                               (3-15) 
 
k
obs
 will be much higher than 2 k
2
 [I
-
]
2
 provided that the concentration of I
-
 is about 
10 times as that of Fe(III) complexes. Under such experimental condition, Equation 3-15 
can be simplified into the following form: 
 
][I2
1
][Fe
]I[2
1
1
II
2
31
1
obs
??
?
+=
kkk
k
k
                                                               (3-16) 
 
 168
The theoretical linearity of 1/k
obs
 relative to [Fe(II)] is confirmed by our Fe
II
(bpy)
2
(CN)
2
 
kinetic inhibition results in the direct oxidation of 75 ?M Et
4
NI by 7.5?M 
[Fe
III
(bpy)
2
(CN)
2
]
+
. Combing the slope value in Figure 3-21 with the slope term in 
Equation 3-16, k
-1
/k
3
 was obtained with the value of 1.06. The rate constant (k
3
, see the 
mechanism) of the reaction between iodide and iodine radical in acetonitrile was reported 
as 2.3 ? 10
10
 M
-1
 s
-1
.
173
 Then k
-1
 was derived, with the value of 2.4 ? 10
10
 M
-1
 s
-1
. The 
derived k
-1
 value implies that the reverse reaction of Equation 3-6 is diffusion-
control.
174,175
 
The validity of the approximations in the derivation of Equation 3-16 is confirmed 
by simulating the complete proposed mechanism (Equation 3-6 to 3-10) using the 
Specfit/32 software package.
132
 Note that in such simulation the rate constants are 
constrained by Equation 3-12. Using the value of k
1
 and k
2
 in Table 3-10, k
-1
 and k
3
 given 
above, and working with the above constraint, an excellent simulation of the observed 
kinetic inhibition by Fe
II
(bpy)
2
(CN)
2
 was achieved, as shown in Table 3-12. A full 
description of the simulation model is shown in Table 3-13. 
Applying the diffusion-control reverse rate constants (k
-1
), the experimental 
forward rate constant (k
1
 = 29.2 ? 0.82 M
-1
 s
-1
), and the standard potential of 
[Fe(bpy)
2
(CN)
2
]
+/0
 (E
1/2
 = 0.070 V vs [Fe(Cp)
2
]
+/0
), the half-wave potential of I
?
/I
-
 was 
calculated by us with the employment of the following equations: 
 
?E
rxn
 = (RT Ln K
1
)/nF                                                                                (3-17) 
 169
Table 3-12. Simulation results for the inhibition by [Fe
II
(bpy)
2
(CN)
2
] in the oxidation of 
I
-
 by [Fe
III
(bpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 
25.0 
o
C, using Specfit/32 software package. 
t
1/2
, s 
 [Fe
II
(bpy)
2
(CN)
2
]
0
, ?M k
obs 
? 10
3
, s
-1 
Experimental Theory
b
 
0 4.94 140 142 
15 4.23 164 171 
30 3.19 217 204 
45 3.06 226 237 
60 2.58 269 270 
a
 [Et
4
NI]
0
 = 75.0 ?M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 7.5 ?M; 
b
 Obtained from simulation 
 
 
 
Table 3-13. Mathematical models for the simulation of the oxidation of I
-
 by 
[Fe
III
(bpy)
2
(CN)
2
]
+
. In the presence of 5.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C, 
using Specfit/32 software package. [Et
4
NI]
0
 = 75.0 ?M, [Fe
III
(bpy)
2
(CN)
2
+
]
0
 = 7.5 ?M, 
[Fe
II
(bpy)
2
(CN)
2
]
0
 = 60.0 ?M. 
Model Reaction 
Equation 
Name Rate constants
Initial 
Concentration, M 
Half-life,
 s 
A + B ? C + D 
k
1c
 35.2 A 
7.5 ? 10
-6
 
C + D ? A + B 
k
-1c
 
2.4 ? 10
10
 
B 
7.5 ? 10
-5
 
A + 2*B ? C + E 
k
2c
 
3.88 ? 10
3
 
C 
6.0 ? 10
-5
 
C + E ? A + 2*B 
k
-2c
 
8.0 ? 10
7 
 
D 0.0 
B + D ? E 
k
3c
 
2.4 ? 10
10
 
E 0.0 
E ? B + D 
k
-3c
 
1.0 ? 10
4
 
F 0.0 
A + E ? C + F 
k
4c
 
2.4 ? 10
10
 
  
270 
A: [Fe
III
(bpy)
2
(CN)
2
]
+
; B: I
-
; C: Fe
II
(bpy)
2
(CN)
2
; D: I
?
; E: I
2
?-
; F: I
2
. 
 
 170
?E
rxn
 = E
f
 (Fe
III
/Fe
II
) - E
f 
(I
?
/I
-
)                                                                   (3-18) 
 
Where K
1
, equilibrium constant for reaction 3-6, equals to k
1
/k
-1
; T: the temperature of 
reaction; R: gas law constant, 8.314 m
2
 kg s
-2
 K
-1
 mol
-1
; n, number of electron transfer in 
reaction 3-6; F, Faraday constant, 96,485 C/mol. The half-wave potential of I
?
/I
-
 is 
obtained, with 0.597 ? 0.002 V vs [Fe(Cp)
2
]
+/0
. Then the reverse rate constants (k
-1
) for 
the oxidation of I
-
 by the other three Fe(III) complexes were calculated using the standard 
potential of I
?
/I
-
 (0.597 ? 0.002 V vs [Fe(Cp)
2
]
+/0
), of Fe
III/II
, and the experimental 
forward rate constants k
1
. All of the calculated reverse rate constants (k
-1
) shown in Table 
3-14 are close to diffusion-control, which further supports that the validity of our 
derivation of the half-wave potential of I
?
/I
-
. 
The half-wave potential of I
?
/I
-
 was also estimated by using thermodynamic cycle. 
Non-close values were obtained with the employment of different assumptions. Firstly, it 
is assumed that iodine radical has the same free transfer energy from water to acetonitrile 
as Xe, the half-wave potential of I
?
/I
-
 was obtained, with E
1/2
 = (0.54 ? 0.04) V vs 
([Fe(Cp)
2
]
+/0
; secondly, suppose that iodine radical has the same free transfer energy from 
gas phase to acetonitrile as Xe, E
1/2
 was obtained, with a value of (0.61? 0.04) V vs 
([Fe(Cp)
2
]
+/0
. The detail calculation was described in Appendix B. Actually, the 
calculated half-wave potential of I
?
/I
-
 from the second assumption is more close to that 
obtained from our experimental results. It implies that Xe is a better model for iodine 
radical in acetonitrile than in water. 
 171
Table 3-14. The calculated rate constants (k
-1
) for the oxidation of I
-
 by Fe(III)  
complexes in acetonitrile.
a
 
Oxidants E
1/2
, 
mV 
b
 
k
1
, 
M
-1
 s
-1
 
K
1
c
 
k
-1 
? 10
-10
,
c
M
-1
 s
-1
 
[Fe(CH
3
Cp)
2
]
+
 
-108 ? 2 0.05 ? 0.01 1.2 ? 10
-12
 4.2 ? 0.8 
[Fe(dmbpy)
2
(CN)
2
]
+
 
- 36 ? 2 0.65 ? 0.05 2.0 ? 10
-11
 3.3 ? 0.3 
[Fe(bpy)
2
(CN)
2
]
+
 
+ 70 ? 2 29.2 ? 0.82 1.2 ? 10
-9
 2.4 ? 0.1 
[Fe(5-Cl-phen)
2
(CN)
2
]
+
 
168 ? 2 1454 ? 30 5.6 ? 10
-8
 2.6 ? 0.1 
a
 [bpy] = 5.0 mM, ? = 0.10 M (Et
4
NBF
4
), T = 25.0 
o
C; 
b 
E
1/2
 vs [Fe(Cp)
2
]
+/0
, 
c
 Calculated with E
1/2
(I
?
/I
-
) = 0.597 V 
 
For a typical outer-sphere reaction, it is generally believed that the oxidative and 
reducing reagents combine to form ion pair (precursor); then electron transfer occurs 
within the ion pairs to form successor; finally, the successor dissociates to obtain the 
products, as shown in Scheme 3-2.
85,176
 The theoretical rate constant, k
1
, can be 
represented by:
85
 
 
k
1
 = K
1ip
k
1et
k
1d
/(k
-1et
 + k
1d
)                                                                   (3-19) 
 
If the rate-limiting step is diffusion-control, not electron transfer, Equation 3-19 can 
be simplified into such form: 
 
k
1
 = K
1ip
k
1et
k
1d
/k
-1et
 = K
1eq
k
-1d
                                                              (3-20) 
 
 172
Then the theoretical slope for a plot of logk
1
 versus E
f
, where k
-1
 is diffusion-
control, was obtained from Equation 3-20, with the value of 16.9, close to the 
experimental data (with slope of 16.1 in Figure 3-29). Therefore, Marcus theory is not 
applicable to calculate the exchange rate constant of I
?
/I
-
. 
 
Scheme 3-2 
 
Ox + X
-
Ox,X
-
K
1ip
Ox,X
-
Red, X k
1et
, k
-1et
Red, X Red + X
k
1d
, k
-1d
 
 
 
 
The exchange rate constant of I
2
?-
/I
-
 could be estimated if the reverse reaction of 3-7 
is not diffusion-control and the half-wave potential of I
2
?-
/I
-
 is known. The equilibrium 
constant, K
rad
, for Equation 3-8 is reported with a value higher than 10
5
 M
-1
.
173
 
Combining the equilibrium constant (K
rad
 ? 10
5
 M
-1
) with the half-wave potential of I
?
/I
-
 
(0.597 V vs [Fe(Cp)
2
]
+/0
), the half-wave potential of I
2
?-
/I
-
 was calculated, with upper 
limit of 0.301 V (E
f
, I
2
?-
/I
-
) vs [Fe(Cp)
2
]
+/0
. Then the rate constant, k
-2
, was obtained with 
upper limit of 2.2 ? 10
8
 M
-1
 s
-1
, much smaller than the diffusion control rate constant in 
acetonitrile.
174,175
 Furthermore, the non-diffusion-control of k
-2
 is further confirmed from 
the smaller slope value of log k
2
 in Figure 3-29. Unfortunately, the unavailability of 
standard potential of I
2
?-
/I
-
 impedes us to obtain its exchange rate constants. 
 
 173
Conclusion 
Trace copper-ion catalysis appears to be a general phenomenon in the oxidation of 
iodide by substitution-inert transition metal complex in acetonitrile. The direct oxidation 
of iodide has the same stoichiometry and two-term rate law as in aqueous solution. 
Analysis of the slopes in LFERs implies that the reverse rate constants for the two paths 
are diffusion-controlled and activation controlled, respectively. The intermediate state of 
I
?
 and I
2
?-
 is inferred, and the standard potential of I
?
/I
-
 is derived from the kinetic 
inhibition by Fe
II
(bpy)
2
(CN)
2
. 
 174
 
 
 
 
 
 
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 188
 
 
 
 
 
 
APPENDIX A  
CALCULATION OF THE EQUILIBIUM CONSTANTS FOR THE OXIDATION 
OF I
-
 BY [Fe
III
(bpy)
2
(CN)
2
]
+ 
 IN ACETONITRILE 
 
                     2 Fe(III) +3 I
-
  2 Fe(II) + I
3
-
                  K
eq1
      (A-1) 
                     2 Fe(III) + 2 I
-
  2 Fe(II) + I
2
                  K
eq2
       (A-2) 
                     I
-
 + I
2
 
 
I
3
-
                                                pK
0
       (A-3) 
E
1/2
, [Fe
III
(bpy)
2
(CN)
2
]
+/0
 = 0.072 V vs [Fe(Cp)
2
]
+/0
 
 
1. E
1/2
, I
2
/I
3
-
 = 0.39 V
163
 vs Ag/AgNO
3
 (0.01M, AN) 
E
1/2
, I
3
-
/I
-
 = -0.27 V
163
 vs Ag/AgNO
3
 (0.01M, AN) 
E
1/2
, [(Fe(Cp)
2
]
+/0
 = 0.089 V
177
 vs Ag/AgNO
3
 (0.01M) 
Thus, E
1/2
, I
2
/I
3
-
 = 0.30 V vs [Fe(Cp)
2
]
+/0
, E
1/2
, I
3
-
/I
-
 = -0.36 V vs [Fe(Cp)
2
]
+/0
 
E
1/2
, I
2
/I
-
 = -0.14 V vs [Fe(Cp)
2
]
+/0
 
For reaction (A-1), 
lnK
eq1
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.36))/(8.314 ? 298) =33.65 
K
eq1
 = 4.1 ? 10
14
 M
-2
 
For reaction (A-2), 
lnK
eq2
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.14)/(8.314 ? 298) = 16.51 
K
eq2
 = 1.5 ? 10
7 
M
-1 
 189
(A-1) - (A-2): 
I
-
 + I
2
 
 
 
I
3
-
                                          (A-3) 
K
0
 = K
eq1
/K
eq2
 = 4.1 ? 10
14
 M
-2
/1.5 ? 10
7
 M
-1
 = 2.73 ? 10
7
 M
-1
 
pK
0
 = -7.44 
 
2. E
1/2
, I
2
/I
3
-
 = 0.396 V
170
 vs Ag/AgNO
3
 (AN), E
1/2
, I
3
-
/I
-
 = -0.248 V
170
 vs Ag/AgNO
3
 
(AN) 
E
1/2
, [Fe(Cp)
2
]
+/0
 = 0.089 V
177
 vs Ag/AgNO
3
 (AN) 
Thus, E
1/2
, I
2
/I
3
-
 = 0.309 V vs [Fe(Cp)
2
]
+/0
, E
1/2
, I
3
-
/I
-
 = -0.337 V [Fe(Cp)
2
]
+/0
 
E
1/2
, I
2
/I
-
 = -0.122 V vs [Fe(Cp)
2
]
+/0
 
For reaction (A-1), 
lnK
eq1
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.337))/(8.314 ? 298) = 31.85 
K
eq1
 = 6.84 ? 10
13
 M
-2
 
For reaction (A-2), 
ln K
eq2
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.122))/(8.314 ? 298) = 15.11 
K
eq2
 = 3.6 ? 10
6
 M
-1
 
(A-1) - (A-2): 
I
-
 + I
2
 
 
I
3
-
                                          (A-3) 
K
0
 = K
eq1
/K
eq2
 = 6.84 ? 10
13
 M
-2
/3.6 ? 10
6 
M
-1
 = 1.9 ? 10
7
 M
-1
 
pK
0
 = -7.27 
 
3 E
1/2
, I
2
/I
3
-
 = 0.65 V
172
 vs S.C.E, E
1/2
, I
3
-
/I
-
 = 0.06 V
172,178
 vs S.C.E 
E
1/2
, Ag/AgNO
3
 (AN) = 0.30 V
163
 vs S.C.E (aq) 
 190
E
1/2
, [Fe(Cp)
2
]
+/0
 = 0.089 V
177
 vs Ag/AgNO
3
 (AN) 
Thus, E
1/2
, I
2
/I
3
-
 = 0.261 V vs [Fe(Cp)
2
]
+/0
, E
1/2
, I
3
-
/I
-
 = -0.329 V vs [Fe(Cp)
2
]
+/0
 
E
1/2
, I
2
/I
-
 = -0.132 V vs [Fe(Cp)
2
]
+/0
 
For reaction (A-1),  
lnK
eq1
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.329))/(8.314 ? 298) =31.23 
K
eq1
 = 3.7 ? 10
13
 M
-2
 
For reaction (A-2),  
lnK
eq2
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.132))/(8.314 ? 298) = 15.9 
K
eq2
 = 8.0 ? 10
6
 M
-1 
(A-1) - (A-2): 
I
-
 + I
2
 
 
I
3
-
                                          (A-3) 
K
0
 = K
eq1
/K
eq2
 = 3.7 ? 10
13
 M
-2
/8.0 ? 10
6
 M
-1
 = 4.6 ? 10
6
 M
-1
 
pK
0
 = -6.66
 
 
4. E
1/2
, I
2
/I
3
-
 = 0.33 V
171
 vs [Fe(Cp)
2
]
+/0
, E
1/2
, I
3
-
/I
-
 = -0.315 V
171
 vs [Fe(Cp)
2
]
+/0
 
?E
1/2
, I
2
/I
-
 = -0.10 V vs [Fe(Cp)
2
]
+/0
 
For reaction (A-1), 
lnK
eq1
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.315))/(8.314 ? 298) = 30.14 
K
eq1
 = 1.2 ? 10
13
 M
-2
 
For reaction (A-2), 
lnK
eq2
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.10))/(8.314 ? 298) = 13.4 
K
eq2
 = 6.6 ? 10
5
 M
-1 
(A-1) - (A-2): 
 191
I
-
 + I
2
 
 
 
I
3
-
                                          (A-3) 
K
0
 = K
eq1
/K
eq2
 =1.2 ? 10
13
 M
-2
/6.6? 10
5
 M
-1
 = 1.8 ? 10
7
 M
-1 
pK
0
 = -7.26 
 
5. E
1/2
, I
2
/I
3
-
 = 0.617 V
179
 vs S.C.E, E
1/2
, I
3
-
/I
-
 = 0.017 V
179
 vs S.C.E 
E
1/2
, Ag/AgNO
3
 (AN) = 0.30 V
163
 vs S.C.E (aq) 
E
1/2
, [Fe(Cp)
2
]
+/0
 = 0.089 V
177
 vs Ag/AgNO
3
 (AN) 
Thus, E
1/2
, I
2
/I
3
-
 = 0.228 V vs [Fe(Cp)
2
]
+/0
, E
1/2
, I
3
-
/I
-
 = -0.372 V vs [Fe(Cp)
2
]
+/0
 
E
1/2
, I
2
/I
-
 = -0.172 V vs [Fe(Cp)
2
]
+/0
 
For reaction (A-1),  
lnK
eq1
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.372))/(8.314 ? 298) =34.58 
K
eq1
 = 1.0 ? 10
15
 M
-2
 
For reaction (A-2),  
lnK
eq2
 = (nF?E
rxn
)/(RT) = 2 ? 96485 ? (0.072-(-0.172))/(8.314 ? 298) = 19.0 
K
eq2
 = 1.8 ? 10
8
 M
-1 
(A-1) - (A-2): 
I
-
 + I
2
 
 
I
3
-
                                          (A-3) 
K
0
 = K
eq1
/K
eq2
 = 1.0 ? 10
15
 M
-2
/1.8 ? 10
8
 M
-1
 = 5.6 ? 10
6
 M
-1
 ? pK
0
 = -6.74
 
 
Reported pK
0
 from literature: 
pK
0
 = -7.18
147
 
Thus, the standard potentials from Popov et al.
163
, Desbarres et al.
170
 and Benoit et al. 
171
 
reports look more accurate. 
 192
 
 
 
 
 
 
APPENDIX B 
CALCULATION OF E
o 
(I
?
/I
-
) IN ACETONITRILE  
 
1. METHOD ONE 
 
I
?
 (aq) + ? H
2
 (g) ? I
-
 (aq) + H
+
 (aq), E
o
 = (1.33 ? 0.02) V
83,85,180
 vs NHE               (B-1) 
Ag
+
 (aq) + ? H
2
 (g) ? Ag (s) + H
+
 (aq), E
o
 = (0.7996 ? 0.001) V
181
 vs NHE           (B-2) 
Ag
+
 (aq) + I
-
 (aq) ? Ag
+
 (AN) + I
-
 (AN),                                                                   (B-3) 
I
?
 (aq) ? I
?
 (AN)                                                                                                           (B-4) 
[Fe(Cp)
2
]
+
 (AN) + Ag (s) ? Fe(Cp)
2
 (AN) + Ag
+
 (AN)                                              (B-5) 
(B-1) - (B-2) ?  
I
?
 (aq) + Ag (s) ? I
-
 (aq) + Ag
+
 (aq)                                                                             (B-6) 
?G
?
6
 = ?G
?
1
 - ?G
?
2
 ? ?E
?
6
 = (1.33 ? 0.02) V - (0.7996 ? 0.001) V = (0.53 ? 0.02) V 
?G
?
6
 = -n*F*?E
?
6
 = -1* (96485 C)* ((0.53 ? 0.02) V) = - (51.14 ? 2.00) kJ mol
-1 
(B-6) + (B-3):  
I
?
 (aq) + Ag (s) ? I
-
 (AN) + Ag
+
 (AN)                                                                       (B-7) 
?G
?
7
 = ?G
?
6
 + ?G
?
3
, ?G
?
3
 = ?G
?
tr
 (I
-
) + ?G
?
tr
 (Ag
+
), 
?G
?
tr
 (I
-
) = - (23.2 ? 2.00) kJ mol
-1
,
182
 ?G
?
tr
 (Ag
+
) = (16.8 ? 2.00) kJ mol
-1
,
182
 
?G
?
7
 = (-51.14 ? 2.00) + (-23.2 ? 2.00) + (16.8 ? 2.00) = - (57.5 ? 3.5) kJ mol
-1
 
 193
I
?
 (aq) ? I
?
 (AN), ?G
?
tr
 (I
?
)  
Suppose ?G
?
tr
 (I
?
) = ?G
?
tr
 (Xe),  
ln K
H
 (AN) = (- 4.98 ? 0.01) 
183
, ln K
H
 (H
2
O) = (- 9.46 ? 0.03) 
183,184
 
 
52.8ml
)0.777(g/ml
41.05)(g)1131.29(e
d
MnMn
V
4.98)(
AN
ANANXeXe
AN
=
?+?
=
?+?
=
?
 
Me0.0189
52.8(ml)
(mol)e
V
n
[Xe]
4.98)(
4.98)(
AN
Xe
AN
?
?
?===  
 
ml0.81
)0.998(g/ml
)(g)181131.29(e
d
MnMn
V
)46.9(
OH
OHOHXeXe
OH
2
22
2
=
?+?
=
?+?
=
?
 
Me0.0554
18.0(ml)
(mol)e
V
n
[Xe]
)96.9(
)46.9(
H
Xe
OH
2
2
?
?
?===
O
 
 
For Xe (aq) ? Xe (AN),  
K
eq
 = [Xe (AN)] / [Xe (H
2
O)] 
      = (0.0189 ? e
(-4.98)
 M)/(0.0554 ? e
(-9.96)
 M) 
= 30.0 ? 0.04 
?G
?
4
 = - RT ln K
eq
 
= - (8.314 J mol
-1
 K
-1
)* (298 K)* (3.40 ? 0.04)  
= - (8.43 ? 0.10) kJ mol
-1 
(B-7) - (B-4):  
I
?
 (AN) + Ag (s) ? I
-
 (AN) + Ag
+
 (AN)                                                                 (B-8) 
?G
?
8
 = ?G
?
7
 - ?G
?
4
 = - (57.5 ? 3.5) ? (- (8.43 ? 0.10)) = - (49.07 ? 3.60) kJ mol
-1
 
 194
?E
o
8
 = (-49.07 kJ mol
-1
) / {(-1)* 96485 C)} = (0.508 ? 0.037) V 
For reaction (B-5), E
o
 ([Fe(Cp)
2
]
+/0
) = 0.089 V
162
 vs Ag/AgNO
3
(0.01 N)(AN) 
E
o
 ([Fe(Cp)
2
]
+/0
) vs Ag/AgNO
3
 (0.01 N) ) =  
E
o
 ([Fe(Cp)
2
]
+/0
) vs Ag/AgNO
3
 (1.0 N) )-0.059 log (0.01) 
Thus, E
o
 ([Fe(Cp)
2
]
+/0
) vs Ag/AgNO
3
(1.0 N) ) = -0.029 V 
(B-8) ? (B-5): 
Fe(Cp)
2
 (AN) + I
?
 (AN) ? [Fe(Cp)
2
]
+
 (AN) + I
-
 (AN)                                        (B-9) 
?G
?
9
 = ?G
?
8
- ?G
?
5
 ? ?E
o
9
 = 0.508 + 0.029 = (0.537 ? 0.037) V 
? E
o
 (I
?
/I
-
) = (0.537 ? 0.037) V vs ([Fe(Cp)
2
]
+/0
 
 
2. METHOD TWO 
 
I
?
 (AN) + Ag (s) ? I
-
 (AN) + Ag
+
 (AN)                                           ?G
?
8
,            (B-8) 
Ag
+
 (aq) ? Ag
+
 (AN),              ?G
?
tr
 (Ag
+
) = (16.8 ? 2) kJ mol
-1
                        (B-10) 
?
f
G
?
Ag(+)
(aq) = 77.11 kJ mol
-1
 
185
 
?G
?
tr
 (Ag
+
) = ?
f
G
?
Ag(+)
(AN) - ?
f
G
?
Ag(+)
(aq) ? 
?
f
G
?
Ag(+)
(AN) = ?G
?
tr
 (Ag
+
) + ?
f
G
?
Ag(+)
(aq) 
= ((16.8 ? 2) + 77.11) kJ mol
-1
 = (93.91 ? 2) kJ mol
-1
 
I
-
 (aq) ? I
-
 (AN)                    ?G
?
tr
 (I
-
) = (-23.2 ? 2) kJ mol
-1
                             (B-11) 
?
f
G
?
I(-)
(aq) = -51.57 kJ mol
-1
 
185
 
?G
?
tr
(I
-
) = ?
f
G
?
I(-)
(AN) - ?
f
G
?
I(-)
(aq) ? 
?
f
G
?
I(-)
(AN) = ?G
?
tr
(I
-
) + ?
f
G
?
I(-)
(aq) 
= {(-23.2 ? 2) + (-51.57)} kJ mol
-1
 
 195
= (-74.77 ? 2) kJ mol
-1
 
I
?
 (g) ? I
?
 (AN)                                                                              ?G
?
12
,          (B-12) 
Xe (g) ? Xe (AN)                                                                        ?G
?
13
,          (B-13) 
 
Me0.0189
52.8(ml)
(mol)e
V
n
[Xe]
4.98)(
4.98)(
AN
Xe
AN
?
?
?===  
 
?G
?
13
 = -R*T*ln[Xe]
AN
 
= (-)* 8.314* 298.15*ln(0.0189*e
(-4.98)
) 
= (5.05 ? 0.03) kJ mol
-1
 
??G
?
12
 
=
 ?G
?
13
 = (5.05 ? 0.03) kJ mol
-1
 
?
f
G
?
I?
 (g) = 70.25 kJ mol
-1
 
185
 
??
f
G
?
I?
 (AN) = ?G
?
12
 + ?
f
G
?
I?
 (g) 
= (5.05 ? 0.03) + 70.25 = (75.30 ? 0.03) kJ mol
-1
 
From reaction B-8,  
?G
?
8
 = ?
f
G
?
Ag(+)
(AN) + ?
f
G
?
I(-)
(AN) -?
f
G
?
I?
 (AN) 
= (93.91 ? 2) + (-74.77 ? 2) - (75.30 ? 0.03) 
= -(56.4 ? 2.8) kJ mol
-1
 
?E
o
 (I
?
/I
-
) = (?G
?
8
)/(-n*F) = (-(56.4 ? 2.8) kJ mol
-1
)/(-1*96485 C mol
-1
) 
= (0.58 ? 0.04) vs Ag/AgNO
3
(1.0 N) 
Since E
o
 ([Fe(Cp)
2
]
+/0
) vs Ag/AgNO
3
(1.0 N) )= -0.029 V 
?E
o
 (I
?
/I
-
) = (0.58 ? 0.04) - (-0.029) = (0.61 ? 0.04) V vs [Fe(Cp)
2
]
+/0
 
 196
 
 
 
 
 
 
APPENDIX C 
PRELIMINARY STUDY OF THE OXIDATION OF I
-
 BY VARIOUS OUTER-
SPHERE OXIDANTS IN ACETONITRILE 
 
I. Oxidation of iodide by Ru
III
(hfac)
3
 
 
Preparation of K[Ru
II
(hfac)
3
]. K[Ru
II
(hfac)
3
] was prepared by following standard 
procedure.
186
 2.40 mmol of ruthenium chloride was added to 50.0 mL of degassed 
EtOH/H
2
O (1:1). The solution was refluxed for 4 hours. Then 1.0 mmol of 1,1,1,5,5,5-
hexafluoro-2,4-pentanedione (hfac) was added to the solution, and it was continued 
refluxing for 2 hours. 12.0 mmol of potassium bicarbonate was added to the solution after 
it was cooled to room temperature. It was heated a little bit and a lot of bubble was 
formed. Then it was refluxed for several hours until the solution turned into pink. Collect 
the precipitate by vacuum filtration. Rinse it with 10.0 mL of toluene. Then dissolve it in 
25.0 mL of acetone, and purify it by passing through one column of silica gel (200-400 
mesh, 60 ?). Remove the acetone in the eluant by rotary evaporation. Dry it in hood for 
12 hours. Note: the solution was purged with nitrogen gas during the refluxing processes. 
Yield of K[Ru
II
(hfac)
3
]: 35%. 
1
H-NMR (400 MHz/CD
3
CN): 2.16 (s, 2H). 
Preparation of Ru
III
(hfac)
3
. Ru
III
(hfac)
3
 was prepared according to Endo?s 
report.
186
 0.13 mmol of K[Ru
II
(hfac)
3
] was added to 6.0 mL of water, and suspension was 
 197
formed. 6.0 mL of toluene and 0.30 mL of 4.0 M HCl was added to the suspension, 
respectively. Then 0.75 mL of 30% H
2
O
2
 was added dropwise, and some bubble was 
formed. The Ru
III
(hfac)
3
 was extracted with 6.0 mL of toluene and 12.0 mL of n-hexane, 
respectively. Remove the n-hexane to obtain Ru
III
(hfac)
3
 by rotary evaporation. Dry it in 
hood for 30 minutes. 
Characterization of K[Ru
II
(hfac)
3
] and Ru
III
(hfac)
3
. The UV-visible absorbance 
characteristics of the ruthenium complexes are shown in Table C-1. The spectral 
properties of K[Ru
II
(hfac)
3
] are close to that reported by Endo et al.
186
; while the 
extinction coefficient of Ru
III
(hfac)
3
 at 529 nm is about 10% higher than their report.
186
 
 
Table C-1. UV-visible absorbance and electrochemical characteristics of the Ru and Fe 
complexes in acetonitrile 
Compounds Band ?
max
, nm
?, M
-1
 
cm
-1
 
E
1/2
, mV 
a
I 289 21091 
K[Ru
II
(hfac)
3
] 
II 529 16697 
352 
I 289 11977 
Ru
III
(hfac)
3
 
II 375 7760 
355 
[Fe
III
(CH
3
Cp)
2
]
+
 I 618 418 0 
a
 E
1/2 
vs [Fe(CH
3
Cp)
2
]
+/0
, at 22.0 
o
C and ? = 0.10 M. E
1/2
 = E
f 
 
The CV (Figure C-1) of 1.0 mM K[Ru
II
(hfac)
3
] in 0.10 M Et
4
NBF
4(AN)
, with 1.0 
mM decamethylferrocene as internal reference, is quasi-reversible, with ?E
p/p
 = 76 mV, 
E
1/2
 = 352 mV vs [Fe(Cp)
2
]
+/0
. The OSWV (Figure C-2) of 1.0 mM K[Ru
II
(hfac)
3
] in 0.10 
M Et
4
NBF
4(AN)
, with 1.0 mM decamethylferrocene as internal reference, has the half-
 198
potential, E
1/2
, of 355 mV vs [Fe(Cp)
2
]
+/0
, 19 mV higher than that reported by Baird et 
al.
187
. 
 
 
Figure C-1. The CV of 1.0 mM of K[Ru
II
(hfac)
3
] and 1.0 mM Fe(Cp*)
2
 in 0.10 M 
Et
4
NBF
4
. With a platinum disc as working electrode, an Ag/AgCl
(s)
 referebce electrode, 
and a Pt wire as a counter electrode. 
 199
 
 
Figure C-2. The OSWV of 1.0 mM of K[Ru
II
(hfac)
3
] and 1.0 mM Fe(Cp*)
2
 in 0.10 M 
Et
4
NBF
4
. With a platinum disc as working electrode, an Ag/AgCl
(s)
 electrode as reference, 
and a Pt wire as a counter electrode. 
 
Kinetics. The reaction trace of oxidation of 8.0 ? 10
-5
 M Et
4
NI by 1.2 ? 10
-5
 M 
Ru
III
(hfac)
3
, in the presence of 2.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C, is 
shown in Figure C-3, with the half-life of 3.7 ms. Generally, the dead time for stopped-
flow instrument is 2 ms. If five times concentration of iodide (0.40 mM) were employed, 
the rate of the reaction would be out of limit of the stopped-flow instrument. So the 
oxidation of iodide by Ru
III
(hfac)
3
 in acetonitrile can not be performed on SF-51. 
 200
 
Figure C-3. Reaction trace for oxidation of 8.0 ? 10
-5
 M Et
4
NI by 1.2 ? 10
-5
 M  
Ru
III
(hfac)
3
. In the presence of 2.0 mM bpy, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 
II. Oxidation of iodide by [Fe
III
(Cp)
2
]
+
. 
 
Preparation of [Fe
III
(Cp)
2
]PF
6
. The synthesis of [Fe
III
(Cp)
2
]PF
6
 was referred to 
Yang et al.'s paper.
188
 2.70 mmol of Fe
II
(Cp)
2
 was added to 10.0 mL of concentrated 
H
2
SO
4
. The solution changed to blue color. Half an hour later, 150 mL of water was 
added to the blue solution. Then it was filtered by vacuum filtration, and 12.9 mmol of 
NH
4
PF
6
 was added to the filtrate. Some precipitate was formed. Filter and rinse it with 
water. Obtain the sample by drying the precipitate in vacuum disiccator for 6 hours. 
Characterization of [Fe
III
(Cp)
2
]PF
6
. The UV-visible absorbance characteristics of 
[Fe
III
(Cp)
2
]PF
6
 is shown in Table C-1, from which it has the same extinction coefficient 
as previous report.
188
 
 201
Kinetics. Firstly, copper-catalysis effect was tested by us. For the oxidation of 1 ? 
10
-2
 M Et
4
NI by 1.0 ? 10
-3
 M [Fe
III
(Cp)
2
]
+
 in acetonitrile, the following two experiments 
were performed on SF-51 instrument. Without the addition of Cu
2+
, the half-life of the 
reaction is 11.0 s; with the addition of 5.0 ?M Cu
2+
, the half-life decreases by around 23 
times, as shown in Table C-2. It indicates that copper cation is a good catalyst for the 
oxidation of iodide by [Fe
III
(Cp)
2
]
+
 in acetonitrile. To inhibit its copper-catalysis effect, 
5.0 mM bpy and 5.0 mM 2,9-dimethyl-1,10-phenanthroline (neocuproine) was added to 
the reactants, respectively (for neocuproine, 5 ? 10
-2
 M, not 1 ? 10
-2
 M, of iodide was 
added). Neither of them can suppress the copper catalysis effectively: the addition of bpy 
does not change the half-life of the reaction; while the rate of the reaction increases with 
the addition of neocuproine. All of the results are shown in Table C-2. These abnormal 
results are ascribed to the air sensitive property of [Fe
III
(Cp)
2
]PF
6
 in acetonitrile. 
Therefore, [Fe
III
(Cp)
2
]PF
6
 is not a good candidate for the oxidation of iodide in 
acetonitrile, either. 
 
Table C-2. Effect of Cu
2+
, bpy, and neocuproine on the oxidation of iodide by 1.0 ? 10
-3
 
M [Fe
III
(Cp)
2
]
+
 in acetonitrile 
[I
-
]
0
, mM [Cu
2+
], ?M [bpy], mM t
1/2
, s 
10 0.0 0.0 11.0 
10 5.0 0.0 0.464 
10 0.0 5.0 12.5 
50 0.0 0.0 0.87 
50 0.0 0.0
a 
0.25 
50 5.0 0.0
a 
0.27 
a
 5.0 mM 2,9-dimethyl-1,10-phenanthroline (neocuproine) 
 202
 
 
 
 
 
 
APPENDIX D  
COPPER-CATALYZED OXIDATION OF IODIDE BY [Fe(CH
3
Cp)
2
]
+
 IN 
CH
3
CN/H
2
O (99/1, v/v) CO-SOLVENT 
 
Introducation 
The redox potentials of the copper(II)-copper(I) and copper(I)-copper(0) are very 
sensitive to solvent: the redox potentials of Cu
2+/+
 and Cu
+/0
 are 0.153 V and 0.521 V vs 
NHE in aqueous media, while that of Cu
2+/+
 and Cu
+/0
 are 1.181 V and -0.011 V vs NHE 
in acetonitrile.
189
 This solvent-dependent electrochemical protery of cuprous cation 
necessitates that it has totally different thermodynamic characterics in these two media: it 
disproportionates completely to copper(II) and metallic copper in aqueous solution, while 
the stronger solvation of copper(I) makes it more favorable than copper(II) in acetonitrile. 
Cu(I), as a d
10
 system, tends to adopt a tetrahedral geometry. In anhydrous 
acetonitrile, Cu(I) is believed to cooridinated to four molecules of acetonitrile. In the 
presence of excess iodide, acetonitrile ligand can be replaced by iodide to form 
Cu(CH
3
CN)
3
(I) or [Cu(CH
3
CN)
2
(I)
2
]
-
. The stability constants of copper(I) iodide in 
acetonitrile was determined by Ahrland et al. using potentiometric and calorimetric 
measurements, with 1.3 ? 10
5
 M
-1
 for ?
1
 and 9.3 ? 10
5
 M
-2
 for ?
2
.
190
 
Very recently, we found that copper cation is a good catalyst for the oxidation of 
iodide by a series of outer-sphere transition metal complexes in acetonitrile,
164,191
 totally 
 203
different from the general kinetic feature of this redox reaction in aqueous media. The 
above copper catalytic behavior may arise from the unsual high redox potential of Cu
2+/+
 
and copper(I)-iodide chemistry in acetonitrile. In this research, the copper-catalysis of the 
oxidation of iodide by [Fe
III
(CH
3
Cp)
2
]
+
 is described. 
 
Experimental section 
Reagent and solution. Acetonitrile (Fisher, Certified A. C. S grade), n-hexane 
(Fisher), methanol (Fisher), diethyl ether (Fisher), hydrochloric acid (Fisher), ferric 
nitrate (Fisher), sodium hexafluorophosphate (Aldrich), 2,9-dimethyl-1,10-
phenanthroline (Aldrich), silver nitrate (Fisher). Tetrakis(acetonitrile)copper(I) 
hexafluorophosphate (Strem Chemicals) was recrystallized from acetonitrile/diethyl 
ether.
192
 Copper (I) iodide (Strem Chemicals) was purified according to Kauffman?s 
method.
193
 Due to their highly air-sensitive property, the dry and purified 
tetrakis(acetonitrile)copper(I) hexafluorophosphate and copper(I) iodide were then 
transferred to vacuum desiccator immediately. The concentrations of Cu(I) are 
standardized by UV-vis spectra of the mixture of the Cu(I) with the excess of 2,9-
dimethyl-1,10-phenanthroline in acetonitrile.
194
 1,1?-dimethylferrocene (Aldrich) was 
recrystallized from ethanol. Et
4
NI (Aldrich) was recrystallized from water and dried 
under vacuum at 100 
o
C for 12 hours. The concentration of Et
4
NI in acetonitrile/H
2
O 
(99/1, v/v) was standardized by titration with standard aqueous AgNO
3
, using Eosin as 
indicator.
152
 Et
4
NBF
4
 (Aldrich) was recrystallized three times from a mixture of methanol 
and n-hexane (4:1) and dried under vacuum at 96.0 
o
C.
195
 Sample of [Fe
III
(CH
3
Cp)
2
]PF
6
 
is from our previous experiments. (refer to Chapter 3) 
 204
Distilled deionized water was obtained from a Barnstead NANO pure infinity 
ultrapure water system. Solutions of [Fe
III
(CH
3
Cp)
2
]PF
6
 and iodide in CH
3
CN/H
2
O were 
prepared just prior to use, and kept in the dark to prevent any photochemical change. 
Both Fe(III) and Et
4
NI were prepared maintaining the appropriate concentrations of the 
other reagents, purged with argon gas, and transferred via gastight glass syringe.  
Methods. The instrumental conditions are referred to that described in Chapter 3. 
Reactions were monitored at fixed wavelengths (When the initial concentration of Fe(III) 
is at 0.25 mM, the reaction was monitored at 650 nm; when the initial concentration of 
Fe(III) is less than 50 ?M, the reaction is monitored at 363 nm), and the rate constants 
were obtained by fitting the data with OLIS-supplied first-order functions. A nonlinear-
least squares computer program was used to fit the overall rate law to the values of k
obs
. 
The water content in the product?s solution was determined by AF7 Coulometric Karl 
Fischer Titrator. 
 
Results 
Effect of small amount of water on the kinetics. The copper-catalyzed 
reaction was performed under nominally anhydrous condition. The experimental 
conditions are as follows: [[Fe
III
(CH
3
Cp)
2
]
+
]
0
 = 0.25 mM, [Et
4
NI]
0
 = 8.0 mM, [CuI]
0
 = 
5.0 ?M. It was observed that the rate of the reaction is irreproducible, with the half-life 
ranging from 0.44 to 0.50 s. This irreproducibility is ascribed to the water sensitivity of 
the standard potential of Cu
2+/+
 in acetonitrile.
189
 The standard potential of Cu
2+/+
 in 
anhydrous acetonitrile is 1.181 V vs NHE; while in CH
3
CN/H
2
O (95/5, v/v) co-solvent, 
the standard potential of Cu
2+/+
 decreases to 0.873 V vs NHE.
189
 Therefore, the different 
 205
content of water in each shot contributes to the above irreproducible kinetic results. Then 
a series of experiments was performed with the addition of very small amounts of water 
to the above reaction system. Table D-1 shows that, even with the addition of 0.50% (by 
volume) of water, reproducible results were obtained, and the rate of the reaction is as 
fast as that under nominally anhydrous condition. With the addition of 1.0% water, the 
reaction runs slower, with the half-life increase by 20%. The half-life increases by a 
factor of 2 with the addition of 2.5% of water. Based on the above results, all of the 
following experiments are performed in CH
3
CN/H
2
O (99/1, v/v) co-solvent. 
 
Table D-1. Water effect on the rate of the reaction between [Fe
III
(CH
3
Cp)
2
]
+
 and I
-
 in 
CH
3
CN/H
2
O co-solvent, in the presence of 5.0 ?M CuI. [Fe
III
(CH
3
Cp)
2
+
]
0
 = 0.25 mM, 
[Et
4
NI]
0
 = 8.0 mM. (Reactions are monitored at 650 nm) 
CH
3
CN/H
2
O, t
1/2
, s 
100/0 0.44-0.50 
99.5/0.5 0.42 
99/1 0.50 
97.5/2.5 0.80 
 
The inhibition by Fe
II
(CH
3
Cp)
2 
on the CuI-catalyzed oxidation of iodide. For 
the CuI-catalyzed oxidation of Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), a 
series of experiments was performed by adding various concentrations of Fe
II
(CH
3
Cp)
2
 
(Fe
II
(CH
3
Cp)
2
, one of the products in the reaction, see product identification section) to 
the reactants. The initial concentrations of the reactants are as follows: [Et
4
NI]
0
 = 5.0 mM, 
[[Fe
III
(CH
3
Cp)
2
]
+
]
0
 = 10.0 ?M, [CuI]
0
 = 10.0 ?M, [Fe
II
(CH
3
Cp)
2
]
0
 = 0.10 to 2.40 mM, ? 
 206
= 0.10 M (Et
4
NBF
4
). The data shown in Table D-2 indicate that the rate of the reaction 
decreases with the increase of Fe
II
(CH
3
Cp)
2
. It follows perfectly pseudo-first-order 
kinetics even though 0.20 mM of Fe
II
(CH
3
Cp) was added, i.e., [Fe(II)]
0
/[Fe(III)]
0
 = 20. 
 
Table D-2. Effect of Fe
II
(CH
3
Cp)
2
 on the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O (99/1). In the presence of 10.0 ?M CuI, at ? = 0.10 M and 25.0 
o
C. 
[Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M, [Et
4
NI]
0
 = 5.0 mM. 
 [Fe(II)]
0
, mM k
obs
, s
-1 
2.40 0.0173 
1.80 0.0197 
1.20 0.0228 
1.00 0.0250 
0.80 0.0312 
0.60 0.0355 
0.40 0.0415 
0.20 0.0600 
 
 
Comparison CuI with [Cu(NCCH
3
)
4
]PF
6
. Cu
+
 is well known to combine with I
-
 
to form CuI and CuI
2
-
 in acetonitrile.
190,196
 To avoid the inevitable consumption of iodide 
by Cu
+
, CuI, rather than [Cu(NCCH
3
)
4
]PF
6
, is firstly selected as catalytic reagent in the 
reaction. For the oxidation of 5.0 mM Et
4
NI by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O 
(99/1, v/v), with 0.60 mM Fe(II) and ionic strength of 0.10 M, 7.5, 10, and 100 ?M of 
 207
crude CuI were added to the reactants, respectively. The experiments were repeated three 
times for each concentration of CuI. The data in Table D-3 demonstrate that the rate of 
the reaction has good reproducibility at 7.5 ?M of CuI, while it is poorly reproducible at 
100 ?M of CuI. The irreproducible kinetic behavior is probably due to the poor stability 
of solid CuI in the air. Then 100 ?M crude and purified CuI were separately added to the 
following reaction system: [Et
4
NI]
0
 = 5.0 mM, [[Fe
III
(CH
3
Cp)
2
]
+
]
0
 = 10.0 ?M, [CuI]
0
 = 
100.0 ?M, [Fe
II
(CH
3
Cp)
2
]
0
 = 0.40 mM, ? = 0.10 M (Et
4
NBF
4
). With the use of purified 
CuI, the pseudo-first-order rate constant (0.0367 s
-1
) is much smaller than that of crude 
CuI. Then the above kinetic experiments were repeated with the addition of another fresh, 
purified CuI solution. However, the rate constant increases by 25%, as shown in Table D-
4. It implies the purified CuI solid sample decomposes quickly even though it is exposed 
to the air for a few minutes. Due to the poor stability of CuI in the air, then 
[Cu(NCCH
3
)
4
]PF
6
 was selected in our kinetic studies. For the oxidation of 5.0 mM Et
4
NI 
by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), with 0.40 mM Fe(II) and ionic 
strength of 0.10 M, the rate constant is nearly the same for both crude and purified 100 
?M [Cu(NCCH
3
)
4
]PF
6
, i.e., 0.0350 s
-1
, as shown in Table D-4. It indicates that the solid 
state of [Cu(NCCH
3
)
4
]PF
6
 is more stable than that of CuI in the air. Therefore, purified 
[Cu(NCCH
3
)
4
]PF
6
 will be used for all of later kinetic experiments. 
 208
Table D-3. Effect of CuI on the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, 
v/v). In the presence of 0.60 mM Fe
II
(CH
3
Cp)
2
, at ? = 0.10 M and 25.0 
o
C. 
[Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M, [Et
4
NI]
0
 = 5.0 mM. 
 [CuI]
0
, ?M 
k
obs
, s
-1
 
7.5 0.0292 
7.5 0.0272 
7.5 0.0282 
10.0 0.0355 
10.0 0.0304 
10.0 0.0371 
100 0.0490 
100 0.0361 
100 0.0519 
 
Table D-4. Comparison of CuI and [Cu(acn)
4
]PF
6
 for the oxidation of I
-
 by 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O(99/1). In the presence of 0.40 mM Fe
II
(CH
3
Cp)
2
 and 
100 ?M Cu(I), at ? = 0.10 M and 25.0 
o
C. [Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M, [Et
4
NI]
0
 = 5.0 mM. 
 [CuI]
0
, ?M [Cu(acn)
4
+
]
0
, ?M
k
obs
, s
-1 
100 0.0 0.0770
a 
100 0.0 0.0367
b 
100 0.0 0.0457
c 
0.0 100 0.0359
d 
0.0 100 0.0345
e
 
a
 crude CuI; 
b
 firstly prepared fresh, purified CuI solution; 
c
 secondly fresh, purified CuI 
solution; 
d
 crude [Cu(NC
3
HC)
4
]PF
6
; 
e
 purified [Cu(NC
3
HC)
4
]PF
6
. 
 
Reexamination of the effect of small amount of water on the kinetics. 
Does [Cu(NCCH
3
)
4
]PF
6
 play the same kinetic role as CuI in water/acetonitrile co-solvent? 
 209
For the oxidation of 8.0 mM of Et
4
NI by 10 ?M of [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN, with 5.0 
?M [Cu(NCCH
3
)
4
]PF
6
 and ionic strength of 0.008 M, it was initially performed under 
nominally anhydrous condition. It shows that the reaction is perfectly reproducible, with 
the half-life of 0.44 s. Then 1.0 % and 2.5% (volume) of water was deliberately added to 
the above reactants, respectively. The half-lives shown in Table D-5 indicate that the rate 
of the reaction decreases by a factor of two even with the addition of 1% of water, totally 
different from that of CuI. In those water-effect experiments, the initial concentration of 
[Fe
III
(CH
3
Cp)
2
]
+
 is 0.25 mM, not 10 ?M, and the non-purified CuI was employed. So 
these two experimental results are not comparable. To counteract the varying 
concentration of water in nominally anhydrous acetonitrile, 1.0 % of water, i.e., 
CH
3
CN/H
2
O (99/1, v/v), was added to acetonitrile in the kinetic studies. 
 
Table D-5. Water effect on the copper-catalyzed oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O co-solvent, in the presence of 5.0 ?M [Cu(acn)
4
]PF
6
. At ? = 0.008 M and 
25.0 
o
C, [Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M, [Et
4
NI]
0
 = 8.0 mM. 
CH
3
CN/H
2
O, t
1/2
, s 
100/0 0.10 
99/1 0.21 
97.5/2.5 0.44 
 
 210
The inhibition by Fe
II
(CH
3
Cp)
2 
on the [Cu(NCCH
3
)
4
]
+
-catalyzed oxidation of 
iodide. For the oxidation of 5.0 mM Et
4
NI by10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O 
(99/1, v/v), with 10.0 ?M [Cu(NCCH
3
)
4
]PF
6
 and ionic strength of 0.10 M, various 
concentrations of Fe
II
(CH
3
Cp)
2
 (Fe
II
(CH
3
Cp)
2
, one of the products in the reaction, see 
product identification section) were added to the reaction system. It follows perfectly 
pseudo-first-order kinetic behavior even though 0.10 mM Fe(II) was added, i.e. 
[Fe(II)]/[Fe(III)] = 10. All of the data corresponding to Fe(II) effect are shown in Table 
D-6. Plot of k
obs
 vs [Fe(II)] shown in Figure D-1 indicates that the rate constant decreases 
with the increase of Fe(II). 
 
0.00 0.25 0.50 0.75 1.00 1.25 1.50
0.00
0.01
0.02
0.03
0.04
0.05
0.06
0.07
0.08
0.09
[Fe(II)], mM
k
ob
s
, s
-1
 
Figure D-1. Plot of k
obs
 vs [Fe(II)] for the oxidation of Et
4
NI by  
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v). with 10.0 ?M [Cu(acn)
4
]PF
6
, at 
? = 0.10 M and 25.0 
o
C. [Et
4
NI]
0
 = 5.0 mM, [Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M 
 211
Table D-6. Effect of [Fe
II
(CH
3
Cp)
2
] on the oxidation of 5.0 mM Et
4
NI by 10 ?M 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1). In the presence of 10.0 ?M [Cu(NCCH
3
)
4
]PF
6
,  
at ? = 0.10 M and 25.0 
o
C. 
[Fe(II)], mM k
obs
, s
-1 
1.20 0.0160 
0.60 0.0263 
0.20 0.0534 
0.10 0.0832 
 
 
 
Figure D-2. 
1
H-NMR spectrum of product for the oxidation of 8.0 mM Et
4
NI  
by 2.0 mM [Fe
III
(CH
3
Cp)
2
]
+
 in CD
3
CN. In the presence of 10.0 ?M CuI. 
 212
Water content in the products? solution. The concentration of water in the products? 
solution was determined by Karl Fischer Titration.All of the results are shown in Table 
D-7, which demonstrates that the concentration of water from Karl Fischer Titration is 
very close to theoretical value. 
 
Table D-7. The water content in the products for the oxidation of 5.0 mM I
-
 by 
[Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v). At ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. 
 
[H
2
O], M 
[Fe(III)]
0
, ?M 
[Fe(II)]
0
, 
mM 
[Cu(acn)
4
+
]
0
, 
?M 
theory titration 
10.0 0.60 2.50 0.556 0.570 
10.0 0.60 5.00 0.556 0.545 
10.0 0.60 10.0 0.556 0.536 
40.0 0.60 10.0 0.556 0.550 
40.0 0.60 40.0 0.556 0.536 
40.0 0.60 60.0 0.556 0.518 
 
 
Identification of product and yield of Fe(II). Under argon protection, 8.0 mM 
Et
4
NI and 20 ?M CuI was mixed with 4.0 mM [Fe
III
(CH
3
Cp)
2
]
+
 equally, then the mixture 
was kept purging with argon for 10 minutes. The 
1
H-NMR spectrum of the product 
shown in Figure D-2 has one very broad peak at 3.99 ppm, indicates that one of products 
in the reaction is Fe
II
(CH
3
Cp)
2
, and implies that some paramagnetic [Fe
III
(CH
3
Cp)
2
]PF
6
 
still remains in the solution. The yield of Fe
II
(CH
3
Cp)
2
 (around 90.2%) was calculated 
from the integrated peak intensity of the spectrum, which confirms the incomplete 
conversion of [Fe
III
(CH
3
Cp)
2
]PF
6
. In the later kinetic experiments, the absorbance at 363 
 213
nm increases with the time of the reaction, which implies that I
3
-
 is one of the products. 
Therefore, the products of the copper-catalyzed reaction are Fe
II
(CH
3
Cp)
2
 and I
3
-
, the 
same as the direct oxidation of iodide in acetonitrile. 
Kinetics. Figure D-3 shows the kinetic trace of the oxidation of 5.0 ? 10 
-3 
M Et
4
NI 
by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
, in the presence of 5.0 ?M Cu(I) and 0.60 mM Fe(II), at ? = 
0.10 M and 25.0 
o
C. It follows perfectly pseudo-first-order kinetics. For the oxidation of 
5.0 mM Et
4
NI by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), with various 
concentrations of [Cu(NCCH
3
)
4
]PF
6
, at ionic strength of 0.10 M and 25.0 
o
C, 0.10, 0.20, 
0.60 and 1.20 mM Fe(II) were added to the above reactants, respectively. The results 
shown in Figure D-5 demonstrate that the rate constants increase with the increase of the 
concentration of Cu(I). Then the rate constants are saturated when the concentration of 
Cu(I) is greater than 10 ?M. At the same concentration of Cu(I), the rate constants 
decrease with the increase of [Fe(II)]. It is of importance to remember that the initial 
concentration of [Fe
III
(CH
3
Cp)
2
]
+
 in the above experiments is exactly at 10 ?M. If the 
initial concentration of [Fe
III
(CH
3
Cp)
2
]
+
 changes, does the reaction still follow such 
copper-saturated feature? Firstly, 40 ?M [Fe
III
(CH
3
Cp)
2
]
+
 was added to such reactants: 
[Et
4
NI]
0
 = 5.0 mM; [Fe(II)]
0
 = 0.60 mM; [Cu(I)]
0
 = 10 ?M; ? = 0.10 M (Et
4
NBF
4
). The 
reaction does not follow first-order kinetics, and the half-life decreases by half as that of 
10 ?M [Fe
III
(CH
3
Cp)
2
]
+
, which demonstrates that the initialconcentrations of [Fe(III)] 
and [Fe(II)] play a key role in this copper-catalyzed reaction, and that the inhibition by 
Fe
II
(CH
3
Cp)
2
 is obviously very important. Then, for the oxidation of 5.0 ? 10
-3
 M Et
4
NI 
by 40 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), in the presence of 2.40 mM 
 214
Fe
II
(CH
3
Cp)
2
, i.e., [Fe(II)]
0
/[Fe(III)]
0
 = 60, at ionic strength of 0.10 M and 25.0 
o
C, 
various concentrations of Cu(I) were added. The rate constants in Table D-8 shows that 
the reaction keeps constant when the concentration of Cu(I) is over 40 ?M. Therefore, the 
copper-saturation is a general feature in such copper-catalyzed reaction. 
 
 
Figure D-3. Kinetic trace of the oxidation of 5.0 mM Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
. In the 
presence of 5.0 ?M [Cu(AN)
4
]PF
6
 and 0.60 mM Fe(II), at ? = 0.10 M and 25.0 
o
C. 
[Fe
III
(CH
3
Cp)
2
+
]
0
 = 10.0 ?M, CH
3
CN/H
2
O = 99/1, v/v. (wavelength: 363 nm) 
 
 215
Table D-8. The rate constants for the oxidation of 5.0 mM Et
4
NI by [Fe
III
(CH
3
Cp)
2
]
+
 in  
CH
3
CN/H
2
O (99/1, v/v), in the presence of various concentrations of [Cu(NCCH
3
)
4
]PF
6
 
and Fe
II
(CH
3
Cp)
2
. At ? = 0.10 M and 25.0 
o
C.  
[Fe(III)]
0
, ?M [Fe(II)]
0
, mM [Fe(II)]
0
/[Fe(III)]
0
[Cu(I)]
0
, ?M k
obs
, s
-1
 
10 0.10 10 0.625 0.0428 
10 0.10 10 1.25 0.0472 
10 0.10 10 2.50 0.0633 
10 0.10 10 5.00 0.0758 
10 0.10 10 10.0 0.0836
 
10 0.10 10 20.0 0.0872
 
10 0.20 20 0.625 0.0244 
10 0.20 20 1.25 0.0310 
10 0.20 20 2.50 0.0419 
10 0.20 20 5.00 0.0472 
10 0.20 20 10.0 0.0534
 
10 0.20 20 20.0 0.0539
 
10 0.60 60 0.625 0.00916 
10 0.60 60 1.25 0.0122 
10 0.60 60 2.50 0.0181 
10 0.60 60 5.00 0.0222 
10 0.60 60 10.0 0.0263
 
10 0.60 60 20.0 0.0270
 
10 0.60 60 60.0 0.0278 
10 0.60 60 100 0.0293
 
10 1.20 120 0.625 0.00910 
10 1.20 120 1.25 0.0101 
10 1.20 120 2.50 0.0117 
10 1.20 120 5.00 0.0132 
10 1.20 120 10.0 0.0160 
10 1.20 120 20.0 0.0166 
40 0.60 15 10 0.0590
a 
40 2.40 60 10 0.0238 
40 2.40 60 20 0.0280 
40 2.40 60 60 0.0304 
a
 not perfect pseudo-first order 
 216
For the direct oxidation of 5.0 ? 10 
-3 
M iodide by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in 
acetonitrile, at ? = 0.10 M and 25.0 
o
C, it was observed that the pseudo-first-order rate 
constant is 0.00208 s
-1
.
191
 When 1.0 mM Fe(II) was deliberately added to the above 
reactants, the rate constant decreases only by 9%. For the copper-catalyzed oxidation of 
5.0 ? 10
-3
 M Et
4
NI by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v), in the presence 
of 0.10, 0.20, 0.60 and 1.20 mM Fe(II), at ionic strength of 0.10 M and 25.0 
o
C, the rate 
constants due to the non-copper-catalyzed contribution is assumed as 0.0019 s
-1
. Plot of 
(k
obs
-0.0019)
-1
 vs [Cu(I)]
-1
 is shown in Figure D-5. The intercept and slope in Figure D-5 
increase with the increase of Fe(II), although it is not dramatic in the case of 1.20 mM 
Fe(II). 
 
0 5 10 15 20 25
0.000
0.025
0.050
0.075
0.100
0.60 mM Fe(II)
1.20 mM Fe(II)
0.20 mM Fe(II)
0.10 mM Fe(II)
[[Cu(acn)
4
]
+
], ?M
k
obs
, s
-1
 
Figure D-4. Plot of k
obs
 vs [Cu(acn)
4
+
] for the oxidation of 5.0 mM Et
4
NI 
by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v). With 0.10, 0.20, 
0.60, and 1.20 mM Fe
II
(CH
3
Cp)
2
, at ? = 0.10 M and 25.0 
o
C. 
 
 217
0 250000 500000 750000 1000000
0
50
100
150
0.60 mM Fe(II)
1.20 mM Fe(II)
0.20 mM Fe(II)
0.10 mM Fe(II)
1/[Cu(I)], M
-1
1/
(
k
obs
-
0
.0
0
1
9
)
, s
 
Figure D-5. Plot of 1/(k
obs
-0.0019) vs 1/[Cu(I)] for the oxidation of 5.0 mM Et
4
NI  
by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1). With 0.10, 0.20, 0.60, and 1.20 mM 
Fe
II
(CH
3
Cp)
2
, at ? = 0.10 M and 25.0 
o
C. 
 
0.0000 0.0005 0.0010 0.0015
0
25
50
75
100
125
2.5 ?M Cu(I)
5.0 ?M Cu(I)
10 ?M Cu(I)
[Fe(II)], M
1/
(
k
obs
-
 0.
0019)
,
 s
 
Figure D-6. Plot of (k
obs
-0.0019)
-1
 vs [Fe(II)] for the oxidation of 5.0 mM 
Et
4
NI by 10.0 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, v/v). With 
2.5, 5.0 and 10.0 ?M [Cu(acn)
4
]PF
6
, at ? = 0.10 M and 25.0 
o
C. 
 218
For the oxidation of 5.0 ? 10 
-3 
M Et
4
NI by 10 ?M [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O 
(99/1, v/v), the inhibition by Fe
II
(CH
3
Cp)
2
 was studied by adding various concentrations 
of [Cu(NCCH
3
)
4
]PF
6
. A plot of (k
obs
-0.0019)
-1
 vs [Fe(II)] is shown in Figure D-6. The 
slope decreases with the increases of Cu(I), while the intercept is nearly the same, i.e., 
8.20 s.  
Based upon the above results, at constant concentration of iodide (5.0 ? 10 
-3 
M), 
the rate of the copper-catalyzed oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in CH
3
CN/H
2
O (99/1, 
v/v) is expressed as follows: 
 
  termcatalysisnon
)[Fe(II)])(([Cu(I)]
]Cu(I)][Fe(III)[
dt
]d[I
Rate
321
3
?+
++
==
?
kkk
      (D-1) 
 
where 0019.0
)[Fe(II)])(([Cu(I)]
]Cu(I)[
321
obs
+
++
=
kkk
k                                   (D-2) 
 
The copper content in the reactants was determined from Atomic Absorption 
Spectra. It was observed that most of copper comes from the recrystallized Et
4
NI, and 
there is approximately 0.12 ?M of Cu in 5.0 mM Et
4
NI. So 0.12 ?M of Cu was added in 
the following fitting. 
 
0019.0
)c)]a)(b[Fe(II([Cu(I)]
]Cu(I)[
obs
+
++
=k                                                  (D-3) 
 
 219
Then the rate constants in Table D-8 were fitted by Equation D-3, using DataFit 8.1 
software package. The calculated rate constants are consistent with that of experimental 
data, except those of 0.625 ?M of [Cu(NCCH
3
)
4
]PF
6
. Due to the uncontrollably varying 
concentration of copper in Et
4
NI, CH
3
CN, water, etc, it is very diffcult to obtain 
reproducible kinetic results at the concentration of Cu(I) less than 1.0 ?M. The values of 
k
1
, k
2
 and k
3
 were obtained from the above fitting, with k
1
 =a = (1.46 ? 0.34) ? 10
-6
 M, k
2
 
= b = (5.35 ? 0.64) ? 10
4
 M
-1
 s, and k
3
 = c = (5.43 ? 1.39) s. 
Unfortunately, we have difficulty in obtaining reproducible results in our kinetic 
experiments, especially at the concentration of Cu(I) no less than 20 ?M, as shown in 
Table D-9. Currently, no reasonable mechanism could be proposed. 
 220
Table D-9. Irreproducible kinetic data for the oxidation of I
-
 by [Fe
III
(CH
3
Cp)
2
]
+
 in 
CH
3
CN/H
2
O (99/1). In the presence of 0.60 mM Fe
II
(CH
3
Cp)
2
 and various concentrations 
of [Cu(NC
3
HC)
4
]PF
6
, at ? = 0.10 M (Et
4
NBF
4
) and 25.0 
o
C. [Fe
III
(CH
3
Cp)
2
+
]
0
 = 10 ?M, 
[Et
4
NI]
0
 = 5.0 mM. 
 [[Cu(acn)
4
]
+
]
0
, ?M k
obs
, s
-1
 Date 
5.0 0.0222 03/01/06 
5.0 0.0252 05/21/06 
5.0 0.0263 05/22/06 
5.0 0.0245 05/31/06 
5.0 0.0237 07/11/06
 
  
20.0 0.0270 03/01/06 
20.0 0.0336 05/11/06 
20.0 0.0342 05/26/06 
20.0 0.0283 05/31/06 
20.0 0.0302 07/11/06 
   
40.0 0.0394 05/26/06
 
40.0 0.0408 05/29/06
 
40.0 0.0382 05/31/06 
40.0 0.0349 07/11/06
 
  
60.0 0.0278 03/01/06 
60.0 0.0460 05/11/06 
60.0 0.0410 05/26/06 
60.0 0.0372 05/29/06