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Labile and inert metal ions - Kinetic effects
Water exchange rate constants (s-1) for selected metal centers
Cr3+ Co3+
Li+ Na+K+
Ca2+ Sr2+Mg2+
Al3+ Ti3+Fe3+ V3+
Cd2+ Hg2+Zn2+Pd2+Pt2+
Ru2+ Cu2+Cr2+Ni2+ Mn2+Co2+ Fe2+
V2+
10-6 10-4 10-2 100 102104 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)
2
4
6
8
10
Mn Fe Co Ni Cu Zn
log
K1
en
gly
ox
mal
OO
OO
_
_
mal
NH2
NH2
en
NH2
OO_
gly
O
O
O
O _
_ox
Mn2+ Fe2+ Co2+ Ni2+ Cu2+ Zn 2+
dn d5 d6 d7 d8 d9 d10
LFSE (o) 0 2/5 4/5 6/5 3/5 0
Ligand field stabilization energy (LFSE)
M2+(g) + nH2O [M(H2O)6]2+ Hhydration
Spontaneous loss of degeneracy of eg and t2g orbitals
for certain dn configurations
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
Octahedral Tetragonal
Cu
Cl
ClCl
Cl
Cl
Cl
2.3 Ao
2.9 Ao
2.9 Ao
2.3 Ao
Jahn-Teller effect in crystalline CuCl2 lattices
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 Red Eo (Reduction potential)
Fe3+ + e Fe2+
It is related to the free energy change and the redox equilibrium constant forthe reduction process
G = nEo F = - 2.3 RT logK
The redox potential of a metal ion couple (Mnn+/M(n-1)+) represents the relative stabilityof the metal when in its oxidized and reduced states.
The redox potential for a metal ion couple will be dependent on the nature ofthe ligands coordinated to the metal.
Comparison of redox potentials for a metal ion in different ligand environments providesinformation on factors influencing the stability of metal centers.
The effect of ligand structure on the reduction potential (Eo
red) 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(phen)3
2+ Eo = 1.14 V
Fe(H2O)63+ + e Fe(H2O)6
2+ Eo = 0.77 V
Fe(CN)63 + e Fe(CN)6
4 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
• 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.
Fe3+ + e- Fe2+
Eo (V)
0.771
Fe3+ + 3e- Fe -0.040
Fe2+ + 2e- Fe -0.44
Fe3+ Fe2+Fe
0.771
-0.040
-0.44
Cu2+ + e- Cu+ 0.15
Cu2+ + 2e- Cu 0.34
Cu+ + e- Cu 0.52
Cu2+ Cu+Cu
0.15
0.34
0.52
Latimer Diagrams
Changes in free energy are additive, but Eo values are not.
If ΔGo(3) = ΔGo
(1) + ΔGo(2),
since ΔGo = − nEoF,
n3 (Eo)3F = n1(Eo)1F + n2(Eo)2F,
and hence
(Eo)3 = n1(Eo)1 + n2(Eo)2
n3
Fe3+ + e- Fe2+
Eo (V)
0.771
Fe3+ + 3e- Fe -0.040
Fe2+ + 2e- Fe -0.44
Fe3+ Fe2+Fe
0.771
-0.040
-0.44
Cu2+ + e- Cu+ 0.15
Cu2+ + 2e- Cu 0.34
Cu+ + e- Cu 0.52
Cu2+ Cu+Cu
0.15
0.34
0.52
MnO4- + H+ + e- 0.90HMnO4
-
HMnO4 + 3H+ + 3e- MnO2 2.10
MnO2 + 2H+ + e- Mn3+ + H2O 0.95
Mn3+ + e- Mn2+1.54
Mn2++ 2e- Mn -1.19
MnO4- HMnO4
- MnO2 Mn3+ Mn2+ Mn0.90 2.10 0.95 1.54 -1.19
1.69 1.23
1.51
MnO4- + 8H+ + 5e- 1.51Mn2++ 4 H2O
O2 + 4 H+ + 4 e- 2 H2O
)Qlog(0591.0o
nEΕ
)]H[
1log(
4
0591.04
O
o
2
p
EΕ
)]H[
1log(
4
0591.023.1
4Ε
)]H[
1log(4)
4
0591.0(23.1 Ε
HΕ p0591.023.1
Dependence of Reduction Potential on pH
E = 0.82 V (pH 7)
Eo = 1.23 V (1.0 M H+)
2 H+ + 2 e- H2 Eo = 0.00 V (1.0
M H+)
)Qlog(0591.0o
nEΕ
)]H[
1log(
2
0591.000.0
2Ε
)]H[
1log()2(
2
0591.0Ε
HΕ p0591.0
E = -0.413 V (pH 7)