Analytical Profile of Coleus Forskohlii | Forskolin .pptx
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1. Coordination Chemistry
InCh3111
By Dr. Madhu Thomas
There are Different theoretical approaches were used for the electronic
structure of coordination complexes. They are…
1. Valence bond theory
2. Crystal field theory
3. Ligand field theory
4. Molecular orbital Theory
2. Coordination Chemistry
• Coordination compounds:- composed of a metal atom or
ion and one or more ligands (atoms, ions, or molecules)
that formally donate electrons to the metal.
• The name coordination compound comes from the
coordinate bond, which historically was considered to form
by the donation of a pair of electrons from one atom to
another.
3. Coordination Chemistry
Coordination compounds:- composed of a metal atom or ion and one
or more ligands (atoms, ions, or molecules) - donate electrons to the
metal .
The name coordination compound - Coordinate bond, - formed by the
donation of a pair of electrons – Crab’s Claw
4. Valance Bond theory
• The theory was put forward by Linus Pauling in 1930
• It uses the ideas orbital hybridization.
• For example in octahedral complexes, d2sp3 hybrids of the metal
orbitals are required.
5. Sc Ti V Cr Mn Fe Co Ni Cu Zn
Y Zr Nb Mo Tc Ru Rh Pd Ag Cd
La Hf Ta W Re Os Ir Pt Au Hg
IIIB IVB VB VIB VIIB IB IIB
VIIIB
d-Block Transition Elements
Most have partially occupied d subshells in
common oxidation states
7. Electronic Configurations
Fe [Ar] 3d64s2
Co [Ar] 3d74s2
Ni [Ar] 3d84s2
Cu [Ar]3d104s1
Zn [Ar]3d104s2
Element Configuration
[Ar] = 1s22s22p63s23p6
8. eg. Co(III) in octahedral coordination
environment.
• There are six ‘d’ electrons are there for Co(III).
• The are either in three paired condition (low spin) or one paired
and four unpaired condition(high spin).
• In the low-spin case d2sp3 hybridisation is possible.
• In the high-spin case sp3d2 hybridisation is possible by the usage of
4d outer orbital.
10. Eg. Fe(III) in Octahedral Coordination
environment
• The 3d orbital may have either one or five unpaired electrons.
• In complexes with one unpaired electron, the ligand electrons force
the metal ‘d’ electrons to pair up and leave two 3d orbitals available
for hybridization and bonding.
• In complexes with five unpaired electrons, the ligands do not bond
strongly enough to force pairing of the 3d electrons.
• Pauling proposed that the 4d orbitals could be used for bonding in
such cases
11. Crystal field theory
• Crystal field theory was developed in the 1930s.
• crystal field theory was used to describe the electronic structure of
metal ions in crystals, where they are surrounded by oxide ions or
other anions that create an electrostatic field with symmetry
dependent on the crystal structure.
• The energies of the d orbitals of the metal ions are split by the
electrostatic field
13. ligands approach along x, y, z axes
(-) Ligands attracted to (+)
metal ion; provides stability
Octahedral Crystal Field
d orbital e-’s repelled by (–)
ligands; increases d orbital
potential energy
+
-
- -
-
-
-
Crystal Field Theory
14. Crystal Field Theory
• When the d orbitals of a metal ion are placed in an octahedral field
of ligand electron pairs, or any electrons in them are repelled by the
field.
• As a result, the dx2-y2 and dz2 orbitals, which are directed at the
surrounding ligands, are raised in energy.
• The dxy, dxz and dyz orbitals, which are directed between the
surrounding ions, are relatively unaffected by the field
• The resulting energy difference is identified as ∆o (o denotes
Octahedral).
15. _ _
_ _ _
dz2
dyz
dxz
dxy
dx2- y2
_ _ _ _ _
isolated
metal ion
d-orbitals
metal ion in octahedral
complex
E
octahedral crystal field
d orbital energy levels
Crystal Field Theory
16. Molecular orbital Theory
• In octahedral complexes, the molecular orbitals can be described as
resulting from a combination of a central metal atom accepting a
pair of electrons from each of six donor ligands.
17. Molecular orbital Theory
• The dX2-y2 and dz2 orbitals can form bonding orbitals with the ligand
orbitals, but dxy, dxz and dyz orbitals cannot form bonding orbitals.
• Bonding interactions are possible with the s (weak, but uniformly
with all the ligands) and the p orbitals of the metal, with one pair of
ligand orbitals interacting with each p orbital.
• The six ligand donor orbitals (p orbitals or hybrid orbitals with the
same symmetry) match the symmetries of the 4s, 4px, 4py, 4pz,
3dz2, and 3dx2-y2 metal orbitals.
• The six bonding orbitals are filled by electrons donated by the
ligands.
• The metal T2g orbitals (dxy,dxz and dyz) do not have appropriate
symmetry to interact with the ligands and are nonbonding.
18. Ligand Field Theory
• It is an extension of CFT(Crystal Field Theory)
• In presence of octahedral ligand field ( either strong or weak ligand
), the d orbitals split into two sets of orbitals (t2g and eg).
• The energy difference between these two sets of orbitals (t2g and eg)
is called ∆o (o for octahedral).
• Ligands whose orbitals interact strongly with the metal orbitals are
called strong-field ligands.
• Ligands with small interactions are called weak-field ligands.
20. Ligand Field Theory
• Strong ligand field means large ∆o, having low spin condition.
• Weak ligand field means small ∆o, having high spin condition.
31. Preparation of Coordination Compounds
• Coordination compounds:- composed of a metal atom or
ion and one or more ligands (atoms, ions, or molecules)
that formally donate electrons to the metal.
• The name coordination compound comes from the
coordinate bond, which historically was considered to form
by the donation of a pair of electrons from one atom to
another.
32. • Ligands
• classified according to the number of donor atoms
• Examples
• monodentate = 1
• bidentate = 2
• tetradentate = 4
• hexadentate = 6
• polydentate = 2 or more donor atoms
33. • Ligands
• classified according to the number of donor atoms
• Examples
• monodentate = 1
• bidentate = 2
• tetradentate = 4
• hexadentate = 6
• polydentate = 2 or more donor atoms
chelating agents
52. Common Geometries of Complexes
Coordination Number Geometry
6
octahedral
Examples: [Co(CN)6]3-, [Fe(en)3]3+
53. Formation and Stability of coordination
compounds
• The most fundamental reaction in a complex can undergo
is ligand substitution:-
Y + M-X M-Y + X
this class of reaction includes complex formation reactions.
• Example:- the replacement of a water ligand by Cl-
[Co(OH2)6]2+ + Cl- [Co(OH2)5Cl] + + H2O
54. Formation constants
• Formation constant expresses the strength of a ligand
relative to the strength of the solvent molecules (usually
H2O) as a ligand; stepwise formation constant is the
formation constant for each individual solvent replacement
in the synthesis of the complex; an overall formation
constant is the product of the stepwise formation constants.
55. [Fe(OH2)6]3+ +SCN- = [Fe(SCN)(OH2)5]2+ + H2O
kf= [Fe(SCN)(OH2)5
2+ ]/ [Fe(OH2)6
3+] [SCN]-
• The value of kf indicates the strength of binding of the ligand
relative to water:-
If kf is large , the incoming ligand binds more strongly than the
solvent (water)
If kf is small, the incoming ligand binds more weakly than
water
56. Consider a solution containing metal ions M and monodentate
ligands L, the system equilibrium can be described the following
equations.
57. Another way of expressing the equilibrium reactions as follows
58. In both the cases there are only N independent equilibrium these Ki’s
(stepwise formation constant)and βi’s (overall formation constant)an be
related by the following equation
59. Formation and Stability of coordination compounds
LABILE AND INERT COMPLEXES
Consider the following reactions
[Cu(OH2)6]2+ + 4NH3 [Cu(NH3)4 (OH2)2]2++ 4 H2O
Blue Intense Blue
[Fe(OH2)6]3+ + SCN- [Fe (SCN)(OH2)s]2++ H2O
Very Pale Violet Red
60. LABILE COMPLEXES
• These reactions, and others like them, are very fast and form species
that can undergo a variety of reactions that are also very fast. These
types of compounds are called labile Complexes
61. INERT COMPLEXES
• A complex that react more slowly are called inert or (a term used less
often).
• An inert compound is not inert in the usual sense that no reaction
can take place; it is simply slower to react.
• Eg.
[Co(NH3)6]3+ + 6(H3O)+ [Co(H2O)6]3+ + 6(NH4)+
62. Mechanism of Ligand Substitution Reactions
• Two types 1) Dissociative Mechanism
2) Associative Mechanism
• In dissociative mechanism the intermediate has a lower coordination
number than the starting complex
MLxX →MLx + X
Intermediate Leaving group
MLx + Y → MLxY
Entering group
63. Associative Ligand substitution Reaction
• In Associative Ligand substitution mechanism the intermediate has a
higher coordination number than the starting complex.
MLxX + Y →MLxXY
entering group Intermediate
MLxXY →MLxY + X
leaving group
64. Kinetics and Reaction Mechanism
• Dissociative Mechanism
• In a dissociative reaction mechanism, loss of a ligand to form an
intermediate with a lower coordination number is followed by
addition of a new ligand to the intermediate.
k1
ML5X ML5 + X
k-1
ML5 + Y → ML5Y
k2
65. Dissociative Mechanism
• The stationary-state (or steady-state) hypothesis assumes the rate
of change of ML5 be zero during the reaction.
dML5 /dt = k1[ML5X] – k-1[ML5][X]- k2[ML5][Y] = 0
Solving for [ML5]
[ML5] = k1[ML5X]/k-1[X] + k2[Y]…………(1)
The rate law for formation of the product is
d[ML5Y]/dt = k2[ML5][Y]………..(2)
Substituting (1) in (2)
d[ML5Y]/dt = k2k1[ML5X][Y]/k-1[X] + k2[Y]
66. Associative Mechanism
• In an associative reaction, the first step, forming an intermediate
with an increased coordination number, is the rate-determining
step. It is followed by a faster reaction in which the leaving ligand is
lost
k1
ML5X + Y ML5XY
k-1
k2
ML5XY → ML5Y + X dML5Y/dt = k[ML5X][Y]
67. Electron-Transfer Reactions(Redox
Reactions)
Redox Reactions of transition metal complexes:- involve the transfer
of an electron from one species to another (from one complex to
another).
Two types of Redox Reactions
1.Inner Sphere Redox Reactions
2.Outer Sphere Redox Reactions
68. 1.Inner sphere Redox reactions
When two molecules may be connected by a common ligand through
which the electron is transferred. in which case the reaction is called
a bridging or inner-sphere reaction.
2.Outer Sphere Redox Reactions
The electron exchange may occur between two separate coordination
spheres in a non-bridging condition is called or outer-sphere reaction.
69. The Rate of reactions for electron transfer depends
on the following factors
1.The rate of substitution in the coordination sphere
of the reactants.
2.The nature of the ligands.
3.Match of energy levels of the two reactants.
4.The solvation of the two reactants.
70. Inner-sphere mechanism
Conditions
One complex must be labile and another should be inert.
Then bridging complex will be formed
A common ligand linkage make bond between two metal
atoms
72. Inner Sphere Mechanism
The reaction proceed in three steps:-
1) A substitution reaction that leaves the oxidant & reductant linked
by the bridging ligand
2) The transfer of the electron takes place
3) Separation of the products
73.
74. Outer Sphere Mechanism
• When both reactants in a redox reaction are kinetically inert,
electron transfer must take place by a tunnelling or outersphere
mechanism
75. Conditions for Outer sphere mechanism:-
*Both complexes should be inert to substitution
*Both complexes should be in the same spin state
*The M-L bond distances in both complexes must be
comparable
*It should require a little activation energy to bring both
complexes identical.
*The electron to be transferred should present in t2g set of
orbitals.
*The electron to be transferred reach the surface of complex
easily.
78. Paramagnetism
• Paramagnetism occurs because of the presence of unpaired
electrons in the metal ions.
• Magnetic moment of Paramagnetic substances are usually obtained
along with contribution from the orbital motion of the electron also.
• Magnetic data is interpreted in terms of the number of unpaired
electrons in the metal ion.
79. Diamagnetism
• Spin and orbital moment cancel each other in diamagnetic
substance
• No net magnetic moment
• Diamagnetic substances are repelled in an external magnetic field.
• The magnetic susceptibility of diamagnetic substances are negative
and is independent of field strength and temperature.
80. Ferromagnetism
• It occurs in substances where individual paramagnetic atoms or ions
are close together and each one is strongly influenced by the
orientation of the magnetic moment of its neighbours.
• In ferromagnetic substances all the spins of the electrons in the
substance tend to align in the same direction.
• This enormously enhances the susceptibility of the substances
compared to the individual moments behaved independently.
• It is generally observed in transition metal ion clusters.
81. Antiferromagnetism
• It also occurs in substances where individual paramagnetic atoms or
ions are close together and each one is strongly influenced by the
orientation of the magnetic moment of its neighbours.
• When the neighbouring paramagnetic centers interacts such as to
favour opposite orientations of their magnetic moments causing
partial cancellation.
• Anti ferromagnetic substances have magnetic susceptibilities less
than those expected for independent magnetic ions.
82. Spin Orbit Coupling
• When several electrons are present in a subshell, the overall effect of
the individual orbital angular momenta l is given by the sum of angular
quantum number L the effect of individual spins ms give by the spin
quantum number S.
• In atoms, the magnetic effects of L and S may couple giving a new
quantum number J called the total angular momentum quantum
number.
• It is obtained by the vectorial combination of L and S.
• This coupling of the resulting spin and orbital quantum numbers is
called Russel-Saunders coupling
84. _ _ _
_ _
dyz
dxz
dxy
dz2 dx2- y2
_ _ _ _ _
isolated
metal ion
d-orbitals
metal ion in
tetrahedral complex
E
d-orbital energy level
diagram
only high spin
86. dyz
dxz
dxy
dz2
dx2- y2
_ _ _ _ _
isolated
metal ion
d-orbitals
metal ion in square
planar complex
E
d-orbital energy level
diagram
__
__
__
__
__
only low spin
87. Crystal field stabilization energy(CFSE)
• The crystal field stabilization energy (CFSE) is the energy by
which the complex is stabilized relative to the free metal
atom where there is no splitting of d orbitals.
• It can be calculated by:
CFSE= {(neg X 3/5) – (nt2g X 2/5)} Δ0
Where, neg = number of electrons in eg orbitals
nt2g = number of electrons in t2g orbitals.
A pairing energy term should also added along with this
88. Calculate the CFSE for d5 configuration in octahedral environment for both low spin
and high spin cases
• Low Spin
CFSE= {(neg X 3/5) – (nt2g X 2/5)} Δ0
= {(0 X 3/5) – (5 X 2/5)} Δ0
=-2Δ0 +2P
High Spin
CFSE= {(neg X 3/5) – (nt2g X 2/5)} Δ0
= {(2 X 3/5) – (3 X 2/5)} Δ0
= 6/5Δ0 - 6/5Δ0 = 0
89. The Chelate Effect
• Complexes containing one or more five or six membered chelate
rings is more stable(having higher formation constant) than a
complex that is as similar as possible but lacks some or all of the
chelate rings.
• Consider the case of [Ni(en)3]2+ and [Ni(NH3)6]2+
• In the case of [Ni(en)3]2+ a five membered chelate ring is formed in
the following way
91. • While in the case of [Ni(NH3)6]2+ no such chelate ring formation
happens.
• The complex with three five membered chelate rings is much more
stable than the monomeric species discussed above as evident from
the thermodynamic stability constants.
92. Trans Effect
Consider the following reaction,
[PtLX3] + Y → [PtLX2Y] + X
Any one of the three labile ligands X can be replaced by the entering
ligand Y.
The ligand X that is replaced can be either Cis or Trans to L giving to
Cis or trans orientation of Y with respect to L
The proportion of Cis and Trans products varies appreciably with
nature of the ligand L.
93. Trans Effect
• Ligands L that strongly favour substitution to give trans product in
the reaction are said to be strong trans directors.
• A fairly extensive series of ligands L may be arranged in order with
respect to their tendency to be strong trans directors. They are..
H2O,OH-,NH3,PY,(NC5H5)<Cl-,Br-<SCN-,I-,NO2-,C6H5-
The above discussed effect is known as Trans Effect
94. Spectochemical Series
• In an octahedral field of ligand, the energy difference ∆o between t2g
and eg orbital depends on the following factors
• The nature of the ligands
• The charge on the metal ion
• Whether the metal is in the first, second or third row of transition
elements
95. Spectrochemical Series
• Ligands which causes only a small degree of splitting are called weak
field ligands
• Ligands which causes large splitting are called strong field ligands.
• The common ligands can be arranged in the order of strength so as
to spit the d-orbitals. That series is called the Spectrochemical
Series.
• The order is
• I- < Br- < Cl- < NO3
- < F- < OH- < H2O and oxalate anion< EDTA < NH3
and Pyridine < en < bipy < o-phenanthroline < NO2
-< CN-
96. Jahn - Teller Distortion
• Jahn – Teller Theorem
• Any nonlinear molecule in a degenerate electronic state will be
unstable and that will undergo a kind of distortion which will lower
the symmetry molecule and reduce the degeneracy.
• Consider Cu2+ in octahedral set of ligands
98. Jahn - Teller Distortion….cont…
• The electronic configuration is t2g
6eg
3
• According to Jahn-Teller theorem such complex should distort
• The distortion takes in the form of an elongation along the z-axis
• Distortion from Oh to D4h symmetry results in the stabilization of the
molecule.
• eg Orbital also will also split into lower dz2 level and higher dx2-y2
level.
99. IUPAC Nomenclature of Coordination
Compounds:
RULES
• The cation is named before the anion
• When naming a complex:
• Ligands are named first in the alphabetical order
• Metal atom/ion is named last and the oxidation state are given by Roman
Numeral in parentheses.
• Use no spaces in complex name
100. IUPAC Nomenclature: Rules
• The names of anionic ligands end with the suffix –o
• -ide suffix changed to –o
• -ite suffix changed to –ito
• -ate suffix changed to -ato
103. IUPAC Nomenclature: Rules
• Neutral ligands are referred to by the usual name for the molecule
• Example
• ethylenediamine
• Exceptions
• water, H2O = aqua
• ammonia, NH3 = ammine
• carbon monoxide, CO = carbonyl
104. IUPAC Nomenclature: Rules
• The prefixes di,tri,tetra,penta and hexa indicate the number of
ligands of that type. When the name of the ligand includes a
number, eg. Dipyridyl or ethylenediamine then bis,tris and tetrakis
are used instead of di,tri and tetra and the groups I placed in the
brackets
106. Nomenclature: IUPAC Rules
Transition
Metal
Name if in Cationic
Complex
Name if in Anionic Complex
Sc Scandium Scandate
Ti titanium titanate
V vanadium vanadate
Cr chromium chromate
Mn manganese manganate
Fe iron ferrate
Co cobalt cobaltate
Ni nickel nickelate
Cu Copper cuprate
Zn Zinc zincate
107. HOME WORK
• Write the formula of the following complexes
1) Hexamminecobalt(III)chloride
2)Chloropentamminecobalt(III)ion
3)Sulphatotetramminecobalt(II)nitrate
• Write the name of the following complexes
1)[Co(NH3)(NO2)3]
2)[Co(NH3)3NO2.Cl.CN]
3)[Cr(en)3]Cl3
108. Electronic Spectra of Transition Metal Complexes
• A characteristic feature of many d-block metal complexes is their
colours, which arise because they absorb light in the visible region.
• Absorptions arise from transitions between electronic energy levels.
• The following Transitions are possible in Transition Metal
Complexes.
1)Transitions between metal-centred orbitals possessing d character
(d–d) transitions.
2)Transitions between metal- and ligand-centred Molecular orbitals
which transfer charge from metal to ligand or ligand to metal
109. Charge Transfer Transition
• Charge transfer (CT) gives rise to intense absorptions.
• There are two types of Charge Transfer Transitions
1) MLCT – Metal to Ligand Charge Transfer Transition
2)LMCT - Ligand to Metal Charge transfer transition
110. d-d Transitions
Transitions between metal-centred orbitals possessing d character
(d–d) transitions.
In electronic Spectra the absorption between 200-400nm range shows
UV absorption and the absorption between 400-800nm shows visible
absorption(d-d transition).
Electronic spectra gives an idea about the rotational and vibrational
level of the molecule
111. • Absorption bands are described in terms of λmax corresponding to the
absorption maximum Amax
• λmax, is usually given in nm,
• The molar extinction coefficient or molar absorptivity ϵmax of an
absorption is given by the equation
ϵmax = Amax /c × l ( in dm3 mol-1 cm-1)
where c is the concentration of the solution and l is the path length
(in cm) of the spectrometer cell.
112. Selection Rules – Spin Selection Rules
Electronic Transitions obey the following selection rules.
Spin Selection Rule: ∆S = 0
Singlet to singlet or triplet to triplet transition is possible.
But a change in spin multiplicity is forbidden.
113. Selection Rules - Laporte Selection Rule
Allowed transitions: g ↔ u
Forbidden transitions : g ↔ g u ↔ u
This leads to the Selection Rule
∆l = ±1
The allowed transitions are s→p, p→d and d→f
The forbidden transitions are s → s, p → p, d → d, f → f ,
s → d, p → f etc.
114. If these selection rules are strictly followed, why do
many d-block metal complexes exhibit d–d bands in their
electronic spectra?
• A spin-forbidden transition becomes ‘allowed’ if, for example, a
singlet state mixes to some extent with a triplet state. This is
possible by spin–orbit coupling.
• But for first row metals, the degree of mixing is small and so bands
associated with ‘spin-forbidden’ transitions are very weak.
115. Vibronic Coupling
• Spin-allowed d–d transitions remain Laporte-forbidden and their
observation is explained by a mechanism called ‘vibronic coupling’.
• An octahedral complex possesses a centre of symmetry, but
molecular vibrations result in its temporary loss.
• When the molecule does not possess a centre of symmetry, mixing
of d and p orbitals can occur.
• d–d transition involving an orbital of mixed p-d character is
relatively weak
116. Vibronic Coupling-Tetrahedral Complexes
• In a molecule which is non-centrosymmetric (e.g. tetrahedral), p–d
mixing can occur to a greater extent.
• The probability of d–d transition is greater than in a
centrosymmetric complex.
• That is why tetrahedral complexes are intensely coloured than the
octahedral complexes.
117. Electronic Transitions in Tetrahedral and Octahedral
Complexes
• The electronic spectrum of Ti3+(d1) in an octahedral field arises from
a transition from the 2T2g to 2Eg term; the energy of the transition
depends on the field strength of the ligands in the octahedral Ti(III)
complex.
• For the d9 configuration (e.g. Cu2+) in an octahedral field, the ground
state of the free ion is again split into 2T2g and 2Eg terms, but, in
contrast to the d1 ion , the 2Eg term is lower than the 2T2g term.
• Thus for a d9 ion in an octahedral field, the splitting diagram is an
inversion of that for the octahedral d1 ion.
118. Energy level diagram for a d1 ion in an octahedral field
2T2g
0.6∆0
2D
0.4∆o
2Eg
119. • This relationship can be represented by an Orgel Diagram where the
right-hand side describes the octahedral d1 case and the left-hand
side, the octahedral d9 ion.
• Just as there is a relationship between the d1 and d9 configurations,
there is a similar relationship between the d4 and d6 configurations.
• we can relate the four configurations in an octahedral field as
follows.
• the Orgel Diagrams for octahedral d1 and d6 ions are the same, as
are the diagrams for octahedral d4 and d9
120. Orgel Diagram
• The Orgel Diagrams for octahedral d1 and d6 ions are the same, as
are the diagrams for octahedral d4 and d9
• The Orgel Diagram for a d1 or d9 ion is inverted by going from an
octahedral to tetrahedral field.
122. Orgel Diagram for d2, d3, d7 and d8 ions (high spin) in
octahedral coordination environment.
123. Orgal Diagram
• Three absorptions are observed in the electronic spectra of d2, d3, d7
and d8 octahedral and tetrahedral complexes
• Transitions are possible from one excited state to another, but their
probability is so low that they can be ignored.
124. The Nephelauxetic Effect
• In metal complexes, sharing of electrons between metal and ligand
happens.
• Pairing energies are lower in complexes than in gaseous M n+ ions.
• It shows that the interelectronic repulsion is less in complexes and
that the effective size of the metal orbitals has increased.
• This is the nephelauxetic effect.
• Nephelauxetic means (electron) ‘cloud expanding’.
125. • For complexes with a common metal ion, it is found that the
Nephelauxetic Effect of ligands varies according to a series
independent of metal ion:
F- < H2O < NH3 < en < [ox]-2 < [NCS-]< Cl- < [CN]- < Br- < I-
• A Nephelauxetic series for metal ions (independent of ligands) is as
follows
• Mn(II) < Ni(II) ≈ Co(II) < Mo(II) < Re(IV) < Fe(III) < Ir(III) < Co(III) <
Mn(IV)
