Labile and inert metal ions - Kinetic effects
Water exchange rate constants (s-1) for selected metal centers
Li+ Na+K+
Mg2+
Cr3+ Co3+
Fe3+
Al3+
Ru2+
10-6
10-4
Ti3+
V3+
Ni2+
V2+
Pt2+
Ca2+ Sr2+
Co2+ Fe2+Mn2+
Pd2+
10-2
100
102
Cu2+Cr2+
Cd2+ Hg2+
Zn2+
104
106
108
1010
Approximate half-lives for exchange of water molecules
from the first coordination sphere of metal ions at 25 oC
Metal ion
t1/2 , sec
Metal ion
t1/2 , sec
Metal ion
t1/2 , sec
Li+
2 x 10-9
V2+
9 x 10-3
Sn2+
< 7 x 10-5
Na+
1 x 10-9
Cr2+
7 x 10-10
Hg2+
2 x 10-9
K+
7 x 10-10
Mn2+
3 x 10-8
Al3+
0.7
Mg2+
1 x 10-6
Fe2+
2 x 10-7
Fe3+
4 x 10-3
Ca2+
2 x 10-9
Co2+
2 x 10-7
Cr3+
3 x 105
Ba2+
3 x 10-10
Ni2+
2 x 10-5
Co3+
7 x 105
Cu2+
7 x 10-10
Zn2+
3 x 10-8
Relative Stability of 3d Transition Metal Complexes
The Irving-Williams Series.
The stability order of complexes formed by divalent
3d transition metal ions.
Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+ > Zn2+
M2+ + L ↔ ML2+ (K1)
NH 2
10
en
en
NH 2
gly
8
NH 2
log K 1
gly
ox
6
O
_
O
O
O
O
O
O
_
O
_
ox
mal
4
_
mal
2
Mn
dn
LFSE (o)
_
O
Fe
Mn2+
d5
0
Co
Fe2+
d6
2/5
Ni
Co2+
d7
4/5
Cu
Ni2+
d8
6/5
Zn
O
Cu2+
d9
3/5
Zn 2+
d10
0
Ligand field stabilization energy (LFSE)
M2+(g) + nH2O
[M(H2O)6]2+
Hhydration
Jahn-Teller Effect
Spontaneous loss of degeneracy of eg and t2g orbitals
for certain dn configurations
Octahedral
Tetragonal
Some metal ions (e.g. Cu(II), d9 and Cr(II), high-spin d4) attain enhanced electronic
stability when they adopt a tetragonally distorted Oh geometry rather than a regular Oh
geometry. They therefore undergo a spontaneous tetragonal distortion (Jahn-Teller effect).
The net stabilization of the eg electrons for Cu(II), is shown above.
Jahn-Teller effect in crystalline CuCl2 lattices
Cl
2.9 Ao
Cl
Cl
2.3 Ao
Cu
2.3 Ao
Cl
Cl
2.9 Ao
Cl
Electronic spectrum of Ti3+ (d1)
Dynamic Jahn-Teller effect in electronic
excited state of d1 ion
Redox Potentials of Metal Complexes
A redox potential reflects the thermodynamic driving force for reduction.
Ox + e
Fe3+ + e
Red Eo
(Reduction potential)
Fe2+
It is related to the free energy change and the redox equilibrium constant for
the reduction process
G =  nEo F = - 2.3 RT logK
The redox potential of a metal ion couple (Mnn+/M(n-1)+) represents the relative stability
of the metal when in its oxidized and reduced states.
The redox potential for a metal ion couple will be dependent on the nature of
the ligands coordinated to the metal.
Comparison of redox potentials for a metal ion in different ligand environments provides
information on factors influencing the stability of metal centers.
The effect of ligand structure on the reduction potential
(Eored) of a metal couple
•
Ligands the stabilize the higher oxidized state lower Eo (inhibit reduction)
•
Ligands that stabilize the lower reduced state increase Eo (promote reduction)
•
Ligands that destabilize the oxidized state raise Eo (promote reduction)
•
Ligands that destabilize the reduced form decrease Eo (inhibit reduction)
•
Hard (electronegative) ligands stabilize the higher oxidation state
•
Soft ligands stabilize the lower oxidation state
•
Negatively charged ligands stabilize the higher oxidation state
Fe(phen)33+
+ e
Fe(H2O)63+
+ e
Fe(H2O)62+
Eo = 0.77 V
Fe(CN)63 +
e
Fe(CN)64
Eo = 0.36 V
Heme(Fe3+) + e
Heme(Fe2+)
Eo = 0.17 V
Fe(III)cyt-c + e-
Fe(II)cyt-c
Eo = 0.126 V
Fe(phen)32+
Eo = 1.14 V
• Soft 1,10-phenanthroline stabilizes Fe in the softer lower Fe(II) state
- i.e. it provides greater driving force for reduction of Fe(III) to Fe(II)
• Hard oxygen in H2O favors the harder Fe(III) state. - resulting in a
lower driving force for reduction of Fe(III) to Fe(II)
• Negatively charged CN- favors the higher Fe(III) oxidation state
(hard - hard interaction) - i.e. it provides a lower driving force for
reduction.
Latimer Diagrams
Eo (V)
Fe3+ + e-
Fe2+
0.771
Fe3+ + 3e-
Fe
-0.040
Fe2+ + 2e-
Fe
-0.44
Cu2+ + e-
Cu+
0.15
Cu2+ + 2e-
Cu
0.34
Cu+ + e-
Cu
0.52
Fe3+
0.771
Fe2+
-0.44
Fe
-0.040
Cu2+
0.15
Cu+
0.34
0.52
Cu
Changes in free energy are additive, but Eo values are not.
If
since
ΔGo(3) = ΔGo(1) + ΔGo(2),
ΔGo = − nEoF,
n3 (Eo)3F = n1(Eo)1F + n2(Eo)2F,
and hence
(Eo)3 = n1(Eo)1 + n2(Eo)2
n3
Dependence of Reduction Potential on pH
O2 + 4 H+ + 4 e-
2 H2O
Ε  Eo 
0.0591
log( Q)
n
Ε  Eo 
0.0591
1
log(
)
4
pO 2 [H  ] 4
Ε  1.23 
Eo = 1.23 V (1.0 M H+)
0.0591
1
log(  4 )
4
[H ]
Ε  1.23  (
0.0591
1
)4 log(  )
4
[H ]
Ε  1.23  0.0591pH
E = 0.82 V (pH 7)
2 H+ + 2 eΕ  Eo 
Eo = 0.00 V (1.0 M H+)
0.0591
log( Q)
n
Ε  0.00 
Ε
H2
0.0591
1
log(  2 )
2
[H ]
0.0591
1
(2) log(  )
2
[H ]
Ε  0.0591pH
E = -0.413 V (pH 7)