Isomerism

Isomers
Isomers have the same molecular formula, but their atoms are
arranged either in a different order (structural isomers) or spatial
arrangement (stereoisomers).

Because of the complicated formulae of many coordination

compounds, the variety of bond types and the number of
shapes possible, many different types of isomerism occur.
Werner's classification into polymerization, ionization,
hydrate, linkage, coordination, coordination position,

geometric and optical isomerism is still generally accepted.

3
1. Structural Isomerism
Coordination isomerism: When both positive and
negative ions are complex, isomerism may be caused by the
interchange of ligands between the two complex ions, for
example: [Co(NH3)6[[Cr(CN)6] and [Cr(NH3)6][(Co(CN)6].
Intermediate types between these extremes are also
possible.
Example
[Ni(C2H4)3][Co(SCN)4] & [Ni(SCN)4][Co(C2H4)3]
[Cr(NH3)5SO4]Br & [Cr(NH3)5Br]SO4
 Ionization isomerism: Coordination compounds which
have the same composition or molecular formula but gives
different ions in solution are called ionization isomers.
(exchange of anions between the coordination sphere and ionization
sphere.)
[Co(NH3)5Br]SO4 is red-violet [Co(NH3)5SO4]Br is dark-red
gives a precipitate with BaCl2 Vs Form a precipitate with Ag+ ions,
confirming the presence of SO42- but does not react with Ba2+

Note that the sulphate ion occupies

only one coordination position even
though it has two negative charges

Other examples of ionization isomerism are [Pt(NH3)4Cl2]Br2 and

[Pt(NH3)4Br2]Cl2,[Co(en)2NO2Cl]SCN, [Co(en)3NO2SCN]Cl and
[Co(en)2Cl.SCN]NO2
 Hydrate isomerism: exchange of water and another
ligand between inner and outer coordination sphere .
Three isomers of CrCl3.6H2O are known.

From conductivity measurements and quantitative

precipitation of the ionized chlorine, they have been given
the following formulae:
[Cr(H2O)4Cl2] Cl.2H2O dark green crystal (one ionic chlorine)

→ (dissolve H2O) [Cr(H2O)5Cl]Cl2.H2O green-blue (two ionic

chlorines)

→ (heat) [Cr(H2O)6]Cl3 violet (three ionic chlorines)

6
 Linkage Isomers : Same complex ion structure but point of attachment
of at least one of the ligands (donor atoms) differs.
In the case of the NO - ion,
2 either a
nitrogen or an oxygen atom may act as
donor, hence the possibility of
isomerism. Two different complexes
[Co(NH3)5NO2]Cl3 have been prepared,
each containing the NO - 2group in the
complex ion. One is red and is easily
decomposed by acids to give nitrous
acid, while the other is yellow and is
stable to acids. This behaviour is
[Co(NH3)5(ONO)]Cl2 analogous to the behaviour of organic
[Co(NH3)5(NO2)]Cl2
Penta ammine nitro cobalt(III) Penta ammine nitrito cobalt(III) nitrites R-ONO and nitro compounds R-
chloride
chloride NO2. The two materials are represented
The thiocyanate ion, SCN-, is also known to below. This type of isomerism occurs in
form a variety of linkage isomers, penta ambidentate ligands like : CO., NO -, 2
SCN-, CN-, S2O 32-.
ammine cobalt(III) compounds,
[Co(NH3)5NCS]2+ is nitrogen-bonded. The
corresponding penta cyano cobalt(III)
compound, however, is S-bonded.
[Co(NH3)5NO2]Cl2

[Co(NH3)5NO2]2+

[Co(NH3)4ONO]Cl2

[Co(NH3)5ONO]2+
 Polymerization isomerism; This is not really true
isomerism since it occurs between compounds having the
same empirical formula, but different molecular weights.
Thus [Pt(NH3)2Cl2], [Pt(NH3)4][PtCl4],
[Pt(NH3)4][Pt(NH3)Cl3]2 and [Pt(NH3)3Cl]2[PtCl4] all have
the same empirical formula.
Polymerization isomerism may be due to a different
number of nuclei in the complex, e.g.:
 2. Stereoisomerism
“Stereoisomers have the same connectivity's but different spatial
arrangements.”
In geometric isomers, the ligands have different spatial
arrangements about the metal ion.
Optical isomers are compounds with non-superimposable mirror images.
A. Diastereoisomers (Geometric isomer)
Diastereoisomers have different chemical and physical properties, such
as color, melting point, polarity, solubilit chemical reactivity, …… .etc.

B. Enantiomers (Optical isomer)

Most of the properties of two optical isomers are the same, such as
solubility, melting point, boiling point, color, chemical reactivity
10
 Geometrical Isomerism (cis-trans) Diastereoisomers:
Atoms or groups of atoms can assume different positions
around a rigid ring or bond.
 Cis – same side (next to each other)
 Trans – opposite sides (across from each other)
Geometrical (cis-trans) Isomerism for a Square Planar Compound

[Pt(NH3)2Cl2]

Cis isomer Trans isomer

The most common type of geometrical isomerism involves cis and trans isomers in
square planar and octahedral complexes. If the complex MX2Y2 is tetrahedral, only
one isomer exists because all of the positions in a tetrahedron are equivalent. If the
complex MX2Y2 is square planar, cis and trans isomers are possible.
• Geometrical (cis-trans) Isomerism for an Octahedral Complex Ion
Cis isomer Trans isomer

[Co(NH3)4Cl2]+

For an octahedral complex, all six positions are equivalent, so only one
compound having the formula MX5Y exists. For an octahedral complex
having the formula MX4Y2 there will be two isomers. Exemple the two
isomers of [Co(NH3)4Cl2]+ .
• If an octahedral complex has the formula MX3Y3, there are two
possible isomers. In an octahedron, the positions are assigned
numbers so the locations of ligands in the structure can be
identified. The usual numbering system for ligands in an
octahedral complex is
• The two isomers of [Co(NH3)3Cl3] have the following
structures:
• In the facial isomer (Fac), the three chloride
ions are located on the corners of one of the
triangular faces of the octahedron.
• In the meridinal i somer (Mer), the three
chloride ions are located around an edge
(meridian) of the octahedron.
• Geometrical isomers are possible for complexes
having a square-based pyramid structure. For
example, the structures for a complex having the
structure MLX2Y2 show that cis and trans
arrangements are possible for the ligands in the base.
Optical isomers : Optical isomerism is a form of isomerism
where by the different 2 isomers are the same in every way except
being non-super imposable mirror images(*) of each other.
Other stereoisomers, called optical isomers or enantiomers, are mirror
images of each other. only includes isomers that are optically active (rotate
plane-polarized light) and mirror images.
At one time it was thought that this sort of isomerism
was associated only with carbon compounds, but it
exists in inorganic molecules as well.
If a molecule is asymmetric, it cannot be superimposed on
its mirror image. The two forms have the type of symmetry
shown by the left and right hands and are called an
enantiomorphic pair. The two forms are optical isomers and
are called either dextro or laevo (d or l) depending on the
direction they rotate the plane of polarized light.
Isomers of I and II of [Co(en)3]3+ are Mirror Images
That Cannot be Superimposed

The Optical Isomers of [Co(en)2Cl2]

The two structures are non superimposable mirror images.

They are like a right hand and a left hand.
 Both four-coordinate and six-coordinate complexes
exhibit chirality. Chiral molecules have either no symmetry
elements (other than identity), or only a Cn axis.
Tetrahedral complexes can be chiral in the same way
that organic compounds are: they may have four different
ligands. They may also have unsymmetrical chelating
ligands.
Enantiomers
isomers that are non
superimposable mirror
images pairs of stereoisomers
that are chiral.
In almost all cases, a chiral
molecule lacks a plane of
symmetry.
 Optical Isomers

The upper isomer is

right handed, and the
lower one is left handed.

Octahedral complexes
with two chelating ligands
and two non-chelating ligands
can also be optically active.
 Optical isomerism is common in octahedral complexes involving bidentate groups. For
example, [Co(en)2Cl2]+ shows cis and trans forms (geometric isomerism), and the cis form is
optically active and exists in d and l forms (optical isomerism) making a total of three isomers.

Trans-dichlorobis(ethylenediamine)
cobalt(iii) ion

Optical activity occurs

also in poly nuclear
complexes, such as:

Enantiomorphic pair d and l

dichlorobis(ethylenediamine)
cobalt(iii) ion

The Optical Isomers

of [Co(en)2Cl2]
The dichlorobis(ethylenediamine)cobalt(II)
ion can exist in two geometrical isomers.
For the trans isomer, there is a plane of
symmetry that bisects the cobalt ion and
the ethylenediamine ligands, leaving one
Cl on either side of the plane. However, the
cis isomer has no plane of symmetry so
two optical isomers exist.
This is also the case for [Co(en) 3] 3+ as is
illustrated in Figure

Light behaves as though it consists of waves

vibrating in all directions around the direction
of wave propagation. For polarized light, the
propagation can be regarded as a vector, which
can be resolved into two circular vectors. If
there is no rotation of the plane, it is expected
that motion along each vector is equivalent so
that each vector traverses an equal distance
around the circle as shown in Figure .
 Isomer problem
• Co(III) and ethylenediamine react to form several
products. Cis [CoCl2(en)2]+ is violet, and the trans
isomer is green. The reaction also forms a yellow
product, [Co(en)3]3+. Determine the number of isomers
of each of the products. Label any enantiomers with the
proper prefix (∆ or Λ).
The yellow product is [Co(en)3]+3. It exists as an enantiomeric pair.
 Optical Isomers
The right-handed isomer
requires going clockwise to
get from the upper triangle
to the lower one. The prefix
for this isomer is ∆.

The left-handed isomer

requires going counter
clockwise to get from the
upper triangle to the lower
one. The prefix for this
isomer is Λ.
21
 Effective Atomic Number (EAN) Rule
This rule is given by English Chemist Nevil V. Sidgwick. Effective
atomic number (EAN) is the total number of electrons in metal
atom/ion (atomic number) plus the electrons gained from ligands.
This EAN is the atomic number of a noble gas. Therefore, EAN
decides stability of coordination compound. If a coordination
compound follow EAN rule, than it is stable one. EAN= Atomic
number of metal atom/ion + number of e- donated by ligands or 2 x
number of ligands (as each ligand can donate two electrons to metal
atom/ion).
• Effective atomic number (EAN), number that
represents the total number of electrons surrounding the nucleus
of a metal atom in a metal complex. It is composed of the metal
atom’s electrons and the bonding electrons from the surrounding
electron-donating atoms and molecules.
 Effective Atomic Number (EAN) Rule
• Effective atomic number (EAN) is the total
number of electrons in metal atom/ion (atomic
number) plus the electrons
gained from ligands.
• This EAN is the atomic number of a noble gas.
(36, 54 or 86)
• Therefore, EAN decides stability of
coordination compound.
• If a coordination compound follow EAN rule,
then it is stable one.
 Thus the effective atomic number of the cobalt atom in the complex
[Co(NH3)6]3+ is 36, the sum of the number of electrons in the
trivalent cobalt ion (24) and the number of bonding electrons from
six surrounding ammonia molecules, each ligands such as (NH3,
H2O, CO, CN-, Cl- etc) contributes an one electron pair (2 etctrons).
– For Fe(CO)5
• Fe26 =[Ar], 3d64s2
• Electrons from Fe=[Ar], 3d6,4s2 = 26 e- Electrons
from 5CO = 5x 2 = 10 e-
• Effective atomic number (EAN) = 36
– For Ni(CO)4
• Ni28 =[Ar], 3d8,4s2
• Electrons from Ni=[Ar], 3d8,4s2 = 28 e- Electrons
from 4CO = 4x 2 = 8 e-
• Effective atomic number (EAN) = 36
• For [Fe(CN)6]4-
• Atomic number of Fe=26; Atomic number of Fe2+=24; there are six
ligands hence electrons donated 6 ligands = 6 x 2
• EAN = 24 + (6 x 2) = 36 (atomic number of Krypton;

• For [Ag(NH3)4]+
• Atomic number of Ag=47; Atomic number of Ag+=46; thereare
four ligands, hence electrons donated by 4 ligands = 4 x 2
• EAN = 46 + (4 x 2) = 54 ( atomic number of Xenon; Xe)
• = 54 (atomic number of Xenon; Xe)

• For [V(CO)6]-
• Atomic number of V=23; Atomic number of V- = 24; there are six
ligands hence electrons donated 6 ligands = 6 x 2
• EAN = 24 + (6 x 2) = 36 (atomic number of Krypton; Kr)
• For [Co(CO)4]2+
• Co27 =[Ar], 3d7 ,4s 2 Co +2 =[Ar], 3d7,4s0 Electrons from Co +2
=[Ar], 3d7 ,4s0 = 25 e- Electrons from 4CO = 4x 2 = 8 e-
• Effective atomic number (EAN) = 33
• For [Cr(NH3)6]3+
• Cr24 =[Ar], 3d5 ,4s1 Cr3+ =21 =[Ar], 3d3,4s
• Electrons from Cr3+ =[Ar], 3d3,4s = 21 e- Electrons from
6NH3 = 6x 2 = 12 e-
• Effective atomic number (EAN) =33

• For [Mn(CN)4]2-
• Atomic number of Mn=25; Atomic number of Mn2+ = 23; there
are four ligands hence electrons donated 4 ligands = 4 x 2
• EAN = 23 + (4 x 2) = 31
1) [Co(NH3)6]3+
E.A.N. = Z – n + (2 X C.N.)
Z = 27, n = 3 and C.N. = (2 X 6) = 12 Thus, E.A.N.
= 27 – 3 + 12= 36
Atomic Number of Krypton (Kr) is 36, i.e.
compound obeys E.A.N. rule, thus it is stable.

2) [Fe(CN)6]4-
E.A.N. = Z – n + (2 X C.N.)
Z = 26, n = 2 and C.N. = (2 X 6) = 12 Thus,
E.A.N. = 26 – 2 + 12 = 36
Atomic Number of Krypton (Kr) is 36, i.e.
compound obeys E.A.N. rule, thus it is stable.
1) [Ni(NH3)6]Cl2
E.A.N. = Z – n + (2 X C.N.)
Z = 28, n = 2 and C.N. = (2 X 6) = 12
Thus, E.A.N. = 28 – 2 + 12= 38
Atomic Number of Krypton (Kr) is 36, i.e. compound does
not obeys E.A.N. rule, and thus it is unstable.

2) [Pt(NH3)6]4+
E.A.N. = Z – n + (2 X C.N.)
Z = 78, n = 4 and C.N. = (2 X 6) = 12
Thus, E.A.N. = 78 – 4 + 12 = 86
Atomic Number of Radon (Ra) is 86, i.e. compound obeys
E.A.N. rule, thus it is stable.
 VALENCE BOND THEORY (VBT) OF
TRANSITION METAL COMPLEXES
• Valence bond theory was given by Pauling and Slater in 1935.
According to this theory:
• In coordination compounds, the ligands form covalent-coordinate
bonds to the metal atom/ ion. The central metal atom/ ion provide
vacant orbitals (s, p and/or d atomic orbitals) equal to its
coordination number. These vacant orbitals
• hybridize and form the same number of new hybridized orbitals
(atomic orbitals overlap) of equal energy.
• Ligands can donate at least one lone pair (in σ orbital) of electrons
to the empty hybrid orbitals of the central metal atom/ ion.
• ligand with filled σ orbital then overlap with the empty hybrid
orbital of central metal atom/ ion.
• This theory helps in predicting the shape, stability and
calculating magnetic moment (magnetic property: µ=√n(n+2)
of the metal complexes.
 The Valence-Bond Approach to Bonding in Complexes

• The idea that atoms form covalent bonds by sharing pairs of electrons was
first proposed by G. N. Lewis . The last major step in the evolution of this
theory was the suggestion by Linus Pauling that atomic orbitals mix to
form hybrid orbitals, such as the sp, sp2, sp3, dsp3, and d 2sp 3 orbitals.
• It is easy to apply the valence-bond theory to some coordination
complexes, such as the[ Co(NH3)6 ]3+ ion. We start with the electron
configuration of the transition- metal ion.
• Co3+: [Ar] 3d 6
• We then look at the valence-shell orbitals and note that the 4s and
4p orbitals are empty.
• Co3+: [Ar] 3d6 4s0 4p0
• Concentrating the 3d electrons in the dxy, dxz, and dyz orbitals in this
subshell gives the following electron configuration.
• The 3dx2-y2, 3dz2, 4s, 4p x, 4py and 4pz orbitals are then mixed
to form a set of empty d 2sp3 orbitals that point toward the
corners of an octahedron. Each of these orbitals can accept a
pair of nonbonding electrons from a neutral NH3 molecule to
form a complex in which the cobalt atom has a filled shell of
valence electrons.
•

• Practice :
• Use valence-bond theory to explain why Fe2+ ions form the
Fe(CN)64- complex ion.
 Important Hybrid orbitals

Exercises
What kind of complex will Ni(2+) form ?
(a) With weak ligands like Cl(-)
(b) With strong ligands like CN(-)
• [NiCl4]2- : Electronic configuration of Ni is 3d84s2
and hence, electronic configuration of Ni+2 is 3d8.
• In this complex, the ligand is weak and no pairing of
electrons will occur. There are two unpaired
electrons in d orbitals of Ni2+. Hence, it is a
paramagnetic in nature. The central metal ion; Ni2+ is
sp3 hybridized and the geometry is tetrahedral.
•

•
 paramagnetic unpaired
[NiCl4]2- electrons
Consider [NiCl4]2+:

Ni(II) -
[NiCl4]2–
sp3
High-spin complex

Ni2+

28Ni

3d 4s 4p
“Strong” ligands as CN(-) cause single
electrons in d-orbitals to combine Similar to the Carbon sp3
together ! (forming “low-spin” in CH4
compounds)
Consider [Ni(CN)4]2-:
diamagnetic paired
[Ni(CN)4]2- electrons

Strong ligands Ni(II) -

cause electron [Ni(CN)4]2–
pairing dsp2

Ni2
valence
+

Ni2
+Ni

3d 4s 4p

Low-spin complex

Ni0 = 4 s2 3d8 Ni2+ = 4s0 3d8

[Ni(CO)4] : Ni = 3d8 4s2 . In this complex, Ni is in 0
oxidation state. The ligand is strong and hence, pairing of
electrons will occur. There are no unpaired electrons in d
orbitals of Ni2+. Hence, it is a diamagnetic in nature. The
central metal atom Ni is sp3 hybridized and the geometry
is tetrahedral.
[MnCl4]2- : Electronic configuration of Mn (atomic
number 25) is 3d5 4s2, and hence, the electronic
configuration of Mn+2 is 3d5. In this complex, again
the ligand is weak and no pairing of electrons will occur.
There are five unpaired electrons in d orbitals of Mn2+.
Hence, it is a paramagnetic in nature. The central metal
ion; Mn2+ is sp3 hybridized and the geometry is
tetrahedral.
 Explanationof labiality andinertness accordingto
VBT
• VBTclassifiesoctahedral complexesinto two types.
• Inner orbital complexes–d2sp3
• Outer orbital complex–sp3d2
• The two d- orbitals involved in thehybridization are the egsetof orbitals.

Outer orbitalcomplexes
• The complexes having sp3d2hybridization are called outer
orbital complexes.
• In terms of VBT these bonds are weaker.
• They are generally labile.
• Mn(II), Fe(II),Fe(III),Co(II),Ni(II),Cu(II) and Cr(II) arelabile.
 Inner orbitalcomplexes
• These complexes generally haved2sp3hybridization.
• The hybrid orbitals are filled withthe ligand electrons.
• The t2g orbitals of metal accommodate the d electrons of the
metal.
• If the t2glevels are leftvacant then the complexcanassociate
with an incoming ligand and the complex islabile
• If all the t2g levelsare occupiedthen the complexbecomes
inert.
• At first glance, complexes such as the Ni(NH3)62+ ion seem hard to explain
with the valence-bond theory. We start, as always, by writing the
configuration of the transition-metal ion. Ni2+: [Ar] 3d 8
• This configuration creates a problem, because there are eight electrons in the
3d orbitals. Even if we invest the energy necessary to pair the 3d electrons,
we can't find two empty 3d orbitals to use to form a set of d 2sp3 hybrids.

• There is a way around this problem. The five 4d orbitals on nickel are empty,
so we can form a set of empty sp3d 2 hybrid orbitals by mixing the 4d x2 -y 2,
4dz2, 4s, 4px, 4py and 4pz orbitals. These hybrid orbitals then accept pairs of
nonbonding electrons from six ammonia molecules to form a complex ion.

• The valence-bond theory therefore formally distinguishes between "inner-

shell" complexes, which use 3d, 4s and 4p orbitals to form a set of d
2sp3 hybrids, and "outer-shell" complexes, which use 4s, 4p and 4d orbitals to
form sp3d 2 hybrid orbitals.
•
• [Mn(CN)6]4- : Manganese (atomic no. 25) has 3d54s2 as
valence shell configuration, with manganese in +2
oxidation state; Mn2+ having 3d5 configuration. In this
complex, the CN- ligand is a strong field ligand. Hence,
pairing of electrons takes place. However, there is one
unpaired electron in d orbitals of Mn2+, it is a paramagnetic
in nature. The central metal ion Mn2+ provides inner d
orbitals and is d2sp3 hybridized (inner shell orbital
complex). Therefore, the complex, [Mn(CN)6]4- has an
octahedral structure:
[Fe(CN)6]3- : Iron (atomic no. 26) has 3d6 4s2 as valence
shell configuration, with iron in +3 oxidation state; Fe+3 having
3d5 configuration. In this complex, the ligand CN- is a strong
field ligand, hence, pairing of electrons will takes place. There
will be one unpaired electron in d orbital of Fe3+. Hence, it will
be paramagnetic in nature. The central metal ion, Fe3+,
providesinner d orbitals and thus, is d2sp3 hybridized (inner
shell orbital complex) and octahedral in shape.
[Fe(H2O)6]3+ : Iron (atomic no. 26) has 3d6 4s2 as
valence shell configuration, with iron in +3 oxidation state;
Fe+3 having 3d5 configuration. In this complex, the ligand
water is a weak field ligand hence; pairing of electrons does not
take place. There are five unpaired electrons in d orbitals of
Fe3+. Hence, it is paramagnetic in nature. The central metal ion
Fe3+ provides outer d orbitals and is thus, sp3d2 hybridized
(outer shell orbital complex) and octahedral in shape.
• Limitations of VBT:
• Cannot explain colour of complexes.
• Cannot explain why magnetic moments of some
metal complexes are temperature dependent.
• Cannot explain the structure of Cu2+ complexes.
  Bonding In Transition Metal Complexes
There are three modern theories of metal to ligand bonding in complexes.

1. Crystal field theory. In this the attraction between the central metal and
ligands in the complex is considered to be purely electrostatic, either
ion-ion attraction between positive and negative ions, or ion-dipole
attractions if the ligand is a neutral molecule.

2. Ligand field theory. This is a development of the electrostatic approach

which allows for some covalent interaction between the orbitals on the
metal and the ligand. There are three types of interaction possible: σ
overlap of orbitals, π overlap of orbitals, or dπ-pπ bonding (back-
bonding) due to π overlap of full d orbitals on the metal with empty p
orbitals on the ligand.
3. Molecular orbital theory. The bonds are considered to be mainly
covalent. The ligands provide electron pairs which occupy σ and π
bonding, antibonding, and sometimes non-bonding molecular orbitals in
the complex.
• Crystal Field theory
• Crystal field theory (CFT) describes the breaking of orbital
degeneracy in transition metal complexes due to the presence
of ligands. CFT qualitatively explains the strength of the
metal-ligand bonds. Based on the strength of the metal-ligand
bonds, the energy of the system (metal –ligand system) is
altered (changed). This may lead to a change in magnetic
properties as well as color. This theory was developed by
Hans Bethe and John Hasbrouck van Vleck.

• When examining a single transition metal ion, the five d-

orbitals have the same energy. When ligands approach the
metal ion, some experiences are observed from the d- orbital
electrons and based on the geometric structure of the
molecule. Since ligands approach from different directions,
not all d-orbitals interact directly. These interactions,
however, create a splitting due to the electrostatic
environment.
• The transition metal (central metal atom) is regarded
as a positive ion and is surrounded by negative or
neutral ligands which have a lone pair of electrons. If the
ligand is neutral like NH3, the negative end (δ-) of the
dipole (two poles one is positive and other is negative)
in the molecule is directed towards the metal atom.

•In order to understand CFT, one must understand the

description of the lobes:
dxy: lobes lie in-between the x and the y axes.
dxz: lobes lie in-between the x and the z axes.
dyz: lobes lie in-between the y and the z axes.
 dx2-y2: lobes lie on the x and y axes.
dz2: there are two lobes on the z axes and there is a donut
shape ring that lies on the xy plane around the other two
lobes.

•The electrons on the central metal atom are under repulsive

forces from the ligands electrons; hence, they occupy the d
orbital furthest ( far as much as possible) from the direction of
the ligands.
 1. Crystal Field Theory
•. Assumptions In the crystal field theory the following assumptions are made;

1. Ligands are treated as point charges.

2. There is no interaction between metal orbitals and ligand orbitals.
3. The d orbitals on the metal which were all the same energy that is degenerate in
a free atom, have their degeneracy destroyed by the ligands when a complex is
formed.
 Ligands: negative point charges
Metal-ligand bonding: entirely ionic: The transition metal which forms
the central atom in the complex is regarded as a positive ion, and is
surrounded by negative ligands or neutral molecules which have a lone
pair of electrons.
If the ligand is a neutral molecule such as NH3, the negative end of
the dipole in the molecule is directed towards the metal ion. The
electrons on the central metal are under repulsive forces from the
ligands, hence they will occupy the d orbitals furthest away from the
direction of approach of ligands.
  Octahedral Complexes
In an octahedral complex, the metal is at the center of the
octahedron, and the ligands are at the six corners.
The directions x, y and z point to the corners of the octahedron (Fig.
8.6).
The eg orbitals (dx2-y2 and dz2) are directed along the axes and the
teg orbitals {dxy, dxz, dyz point in between the axes x, y and z.
It follows that the approach of six ligands along the x, y, z, -x, -y
and -z directions will increase the energy of the dx2-y2 and dz2
orbitals which point along the axes much more than the dxy, dxz, dyz
orbitals which point in between the axes.
 dz2 and dx2-y 2 point their lobes
directly at the point-charge ligands.
 dxz, dyz and dxy point their lobes
between the point charges.
51
Thus under the influence of an octahedral ligand field the d orbitals split
into two groups of different energy (Fig. 8.7). Rather than referring to the
energy level of an isolated metal atom, the weighted mean of these two
sets of perturbed orbitals is taken as the zero.
The difference in energy between the two d levels is given either of the
symbols ∆O or 10 Dq. It follows that the orbitals are +0.6 ∆O above the
average level, and the teg orbitals are -0.4 ∆O, below the average (Fig. 8.8).

 The dxz, dyz and dxy orbitals are at a lower energy in the
octahedral complex than are the dz2 and dx2-y 2 orbitals.

Splitting of d-Orbital Energies

by an Octahedral Field of Ligands

54
 Strong Field and Low Spin
• The splitting of d orbital energies explains the color and magnetism of
complex ions.
• If the splitting produced by the ligands is very large, a situation called
strong field case, the electrons will pair in the low energy t2g orbitals.
• The strong field case is also called low spin case.
• This gives a diamagnetic complex in which all electrons are pairs.
• △0>P

Weak Field and High Spin

If the splitting produced by the ligands is small, the electrons
will occupy all five orbitals before pairing occurs called weak
field case.
The weak field case is also called high spin case. In this case, the
complex is paramagnetic. △0< P
weak field case strong field case
with paramagnetic with diamagnetic
• The magnitude of ∆o depends on the
following factors.
1- Nature of ligand
• Studies have shown that keeping the same central metal
atom, on varying the ligands attached to it, the magnitude of
∆o is different in each case. Ligands which cause only a
small degree of splitting are termed weak field ligands and
those which cause a large splitting are called strong field
ligands. The common ligands are arranged in order of their
strength and the series is called the spectrochemical series.

Weak field I- < Br- < S2- < SCN-< F- <Cl- < NO3- <
OH-< H2O ≈ C2O42- < EDTA4-<NH3≈pyridine <
ethylenediamine < dipyridyl < o-phenanthroline< NO2- <
CN- < CO. Strong field.
 2-Position of the metal in the periodic table
• The magnitude of ∆o increases on descending down the
group of transition metals in periodic table. Depending
upon the d electron configuration of the metal ion, and
under specific environment around it, its complex
acquires stability.
3-Charge on the metal atom
• The magnitude of ∆o increases as the charge on central
metal atom increases and in general M3+ complexes
have double the value than that of M2+ complexes.
•
 ΔE = E(eg) + E(t2g)

Crystal Field Stabilization Energy

(CFSE):
 In Octahedral field, configuration is: t 2gx egy
E eg

CFSE =ΔO(- 2/5 x +3/5y) +nP

x: numbers of electrons in the t2g

y: numbers of electrons in the eg

t2g
57
 Example 1: Cr3+ : d 3
[Cr(NH3)6]3+

CFSE =ΔO(- 2/5 * 3 +3/5*0) +0*P = -6/5 ΔO

Example 2: Mn3+ : d 4
[MnF6]-3 [Mn(CN)6]-3

configuration is: t2g3 eg1 configuration is: t2g4 eg0

CFSE =ΔO(- 2/5*3 +3/5*1)+0*P= -3/5 ΔO
CFSE =ΔO(- 2/5*4 +3/5*0)+1*P= -8/5ΔO+P
Weak field =High spin
Strong field =Low spin
The magnitude of ∆O depends on three factors:
1. The nature of the ligands,
2. The charge on the metal ion,
3. Whether the metal is in the first, second or third row of transition
elements.
By examining the spectra of a whole series of complexes, it has been
found that the position of the absorption band and hence the value of ∆O
for any given transition metal ion, varies depending on the ligands
attached (Table 8.5).
[CrIIICl6]3- 13640 cm-1 163 kJ/mol
[CrIII(H2O)6]3+ 17830 213
[CrIII(NH3)6]3+ 21680 259
[CrIII(CN)6]3- 26280 314
Ligands which cause only a small degree of splitting are
termed weak ligands, and those which cause a large splitting
are called strong ligands.
The common ligands can be arranged in order of strength, and this series is
called the spectrochemical series
weak ligands
I- < Br- < CI- < NO3 - < F- < OH- < H2 O and oxalate < EDTA< NH 3 and
pyridine < ethylenediamine < dipyridyl < o-phenanthroline < NO 2- < CN-
strong ligands
The splitting produced by strong CN- ligands is about double that from
weak ligands. The magnitude of ∆O increases as the charge on the metal
ion increases, and M3+ complexes have roughly double the values for
complexes (Table 8.6).
The value of ∆O also increases on descending a group of transition
elements (Table 8.7).

[Co(NH3)6]3+ 24800 cm-1

[Rh(NH3)6]3+ 34000

[Ir(NH3)6]3+ 41000
 Spectrochemical Series
• I- < Br- < SCN- ~Cl- < F- < OH- ~ ONO- < C2O42- < H2O< NCS- <
EDTA4- < NH3 ~ pyr ~ en < phen < CN- ~ CO
• Mn2+ < Ni2+ < Co2+ < Fe2+ < V 2+ < Fe3+ < Co3+ < Mn3+ < Mo3+ <
Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+
OH
N

pyr: pyridine phen: phenol

 Photophysical Properties
Larger  The Spectrochemical Series Smaller 
CO, CN- > phen > NO2- > en > NH3 > NCS- > H 2O > F- >
RCO2 - > OH- > Cl- > Br- >I-

Increasing  73
 Crystal field stabilization energy
(CFSE) in Tetrahedral complexes
• In tetrahedral molecular geometry, a central atom is
located at the center of four substituent atoms, which
form the corners of a tetrahedron. The bond angles
are approximately 109.5° when all four substituent's
are the same.
• d-Orbitals splitting for tetrahedral complex. A cube,
an octahedron, and a tetrahedron are related
geometrically. Octahedral coordination results when
ligands are placed in the centers of cube faces.
Tetrahedral coordination results when ligands are
placed on alternate corners of a cube.
Tetrahedral Splitting

dxy dyz dxz

M

dx2-y2 dz2

Tetrahedral
Why Tetrahedral complexes are favoured ?
1. where the ligands are large and bulky and could cause crowding in an
octahedral complex

2. when the ligands are weak, so the loss in CFSE is less important
3. where the central metal has a low oxidation state, since this reduces the
magnitude of ∆

4. where the electronic configuration of the central metal is d0, d5 and d10
since there is no CFSE or d1 or d6 where the loss in CFSE is small.

Many transition metal chlorides, bromides and iodides form tetrahedral

structures.
 Crystal field stabilization energy
(CFSE) in Square planar complexes:
• A square planar arrangement of ligands can be formally
derived from an octahedral array by removal of two trans
ligands.
• If we remove the ligands lying along the z axis, then the dz2
orbital is greatly stabilized and the energies of the dxz and dyz
orbitals are also lowered.
•The square planar are formed by d8 ions with strong field ligands,
e.g.[Ni(CN)4]2-.
The CFSE is large for heavier elements and in highly charged species,
and as a consequence all complexes of Pt2+, Pd2+and Au3+ are square
planar, whether with strong or weak field.
Square planar is also common for complexes of the d9 ion, Cu2+
e.g.[Cu(py)4]2+.
For this only one electron has to be placed in the high energy
dx2-y2 orbital.
Other Geometries

66
 Magnetic Properties
Larger  The Spectrochemical Series Smaller 
CO, CN- > phen > NO - > en > NH > NCS- > H O > F- > RCO - > OH > Cl- > Br- > I-
2 3 2 2 -

High Spin Low Spin

Paramagnetic- Diamagnetic-
unpaired electrons. all electrons paired.
 The Color of Complexes
• Very commonly for the first transition series, the
energy corresponds to that of visible light, so that
d-d transitions are the cause of the delicate
colours of so many of the complexes.
The size of the energy gap ∆O between the teg and eg levels can be
measured easily by recording the uv-visible spectrum of the
complex.
Consider a complex like [Ti(H2O)6]3+. The Ti3+ ion has one d
electron. In the complex this will occupy the orbital with the lowest
energy that is one of the teg orbitals (Fig. 8.9a).
The complex absorbs light of the correct wavelength (energy) to
promote the electron from the teg level to the eg level (Fig. 8.9b). The
electronic spectrum for [Ti(H2O)6]3+ is given in Fig. 8.10.
The steep part of the curve from 27 000 to 30 000 cm-1 is due to
charge transfer and occurs in the uv region. The d-d transition is the
single broad peak with a maximum at 20 300 cm-1.
The d-d transition is the single broad peak with a maximum at 20
300 cm-1. Since 1 kJ mol-1 = 83.7 cm-1 the value of for [Ti(H2O)6]3+
is 243 kJ mol-1.
 Colour of complexes

E=h*ν

Depends on the energy difference between the lower and higher

metal d-orbital levels !
Visible light is absorbed and pushes electrons up => The higher
Δ, the more “blue” is the light absorbed
We observe the appearance of a shoulder in the case of
[Ti(H2O)6]3+. Perfectly octahedral [Ti(H2O)6]3+ should give only
one d-d Transition. However, distortion occurs to eliminate the
degeneracy of the system. If a complex distorts from regular
octahedral geometry, the t2g and eg levels are split, the consequence
of which is the appearance of a shoulder as shown in the figure
right.

Aqueous [Ti(H2O)6]3+
 Jahn-Teller Distortion
Jahn-Teller theorem:
“there cannot be unequal occupation of orbitals with identical energy”
Molecules will distort to eliminate the degeneracy!

Distortion

d9
d3

1 u.e. 1 u.e.
equal occupation unequal occupation
 Jahn-Teller Distortion

2.45 Å
2.00 Å
dx2-y2
eg
E dz2
dxy
t2g
dxz dyz

[Cu(H2O)6]2+
• Tetragonal distortion of octahedral
complexes (Jahn-Teller distortion)
• The stereochemistry or shape of transition metal complexes
is determined by the tendency of electron pairs to occupy
positions as far away from each other. If the d electrons are
symmetrically arranged with respect to an octahedral field,
they will repel all the six ligands equally and a regular
octahedral geometry is achieved. The various symmetrical
electronic arrangements are observed in d1 (strong and weak
field), d3 (strong and weak field), d5 (weak field), d6 (strong
field), d8 (weak field) and d10 (strong and weak field).
• In all other cases (d4, d7, and d9) the electronic
arrangements are unsymmetrical. As the t2g orbitals point
in between the ligands, the asymmetric filling of these
orbitals has little effect on the shape of the complex.
However, the eg orbitals point directly towards the ligands
and as a result the asymmetric filling of these orbitals
leads to more repulsion of some ligands compared to
others.

The effect of a tetragonal distortion of an octahedral

crystal field on the energies of d orbitals:
• This leads to a significant deviation from octahedral
geometry. The reason for this is that, in an octahedral environment
when eg orbitals, dx2-y2 and dz2 are unsymmertically filled (for
example t2g3 eg1, t2g6 eg1 and t2g6 eg3 ) their degeneracy is
removed and the two orbitals become unequal in energy. Out of
the two orbitals, the dz2 orbital has its two lobes pointing towards
z axeswhile the dx2-y2 has four lobes pointing towards x, y, -x
and –y axes. In order to minimize the repulsion with the ligands
the single eg electron will occupy the dz orbital. Therefore two
ligands approaching along the z and -z axes will experience
greater repulsion than the four ligands along the other four
axes. This situation results in tetragonal distortion or strictly
speaking tetragonal elongation of octahedral complexes with
four short and two long bonds. This effect is called Jahn-Teller
effect or Jahn-Teller distortion.
Examples of this distortion are Cu(II) complexes (e.g.
[Cu(NO2)6]4-), where Cu2+ has t2g6 eg3 configuration,
show tetragonal distorted octahedral structures. Here also
to minimize repulsion with the dz2 orbital and one
occupies the dx2-y2 orbital.
• In case where the orbital contains extra electron, then the
elongation will occur along the x and y axes and therefore,
ligands approach more closely along the z axis. Thus there
will be four long bonds and two short bonds. This means
compression of octahedron along the z-axis and is called
tetragonal compression. The tetragonal elongation is much
more common from this discussion, when there is unequal
filling of dz2 result the energy of dx2-y2 orbital increases
relative to the other orbital. If ligand field is strong enough, the
difference in energy between these two orbitals becomes
larger than the energy needed to pair up electrons. Therefore
pairing up of electrons in dz orbital takes place leaving the
other orbital empty. Now four ligands along x, y, -x and –y
axes can approach metal atom without any difficulty but
ligands approaching along z and –z axes meet very strong
repulsive forces.
 Crystal Field Theory
Merits of crystal field theory:
1) Can be used to predict the most favorable geometry for the complex.
2) Can account for why some complexes are tetrahedral and others
square planar.
3) Use full in interpreting magnetic properties.
4) The colors of many transition metal complexes can be rationalized.

Limitations of crystal field theory:

1) Becomes less accurate as delocalization increases (more covalent
character).
2) Point charge does not accurately represent complexes.
3) Does not account for pi bonding interactions.
4) Does not account for the relative strengths of the ligands.
Labileand inertcomplexeson thebasis of CFT
• According to CFT the ligand field splits the d-orbitals.
• This splitting leads to a decrease in energyof the systemwhose
magnitude depends onthe number of d electrons present.

• if the CFSE value increasesby association or dissociation of a

ligand then the complex is labile.

• On the other hand, it is inert when there isa loss inCFSE

value.
Factorsaffecting lability ofcomplexes
•Chargeof thecentralion: Highlychargedionsform
complexeswhich react slowly i.e.inert

•Radiiof the ion: the reactivitydecreaseswith

decreasing ionicradii.

•Charge to radius ratio: if all the factors are similar, the ion
with largestz/r valuereacts with the least rate.

•Geometry of the complex:Generallyfourcoordinated

complexesare more labile
 Properties of the metal ion
• Charge and size
• Natural order (or) Irving –William order of stability
• Class a and Class b metals
• Electronegativity of the metalion
 Effect of ionic radius
Chargeonthe
Complexion Ionic radii(Aₒ) Value of ᵝ Stability
ion

[BeII(OH)] + +2 0.31 107

[MgII(OH)] + +2 0.65 120

[CaII(OH)] + +2 0.99 30

[BaII(OH)] + +2 1.35 4
 Effect of charge
Charge onthe Ionic radii(Aₒ)
Complexion Value of log ᵝ stability
ion

[FeIII(CN)6]3- +3 31.0
Almost
[FeIII(CN)6]4- +2 same 8.3

CoIIIcomplex +3 high
Almost
same
CoII complex +2 low
 Irving – William order ofstability

• Stabilities of the high spin complexes of the 3d metals

from Mn2+to Zn2+with acommon ligand isusually
Mn2+< Fe2+<Co2+<Ni2+<Cu2+>Zn2+
• This is attributed to the CFSE values ofthe
complexes and called natural order of stability.
• There is adiscrepancy with Cu which is due to Jahn–
Teller distortion
Crystal Field Stabilization Energy (CFSE) of
d0 to d10 M(II) ions:
CFSE as a function of no of d-
electrons

1.4
1.2
CFSE in multiples of

Ni2+
1 double-
0.8 humped
curve
0.6
0.4
0.2
0
0 3 4 5 6 7 8 9 10 11

no of d-electrons
Log K1(EDTA) of d0 to d10 M(II) ions:
log K1(EDTA) as a function of no of d-
electrons
= CFSE
20
double-
18
logK1(EDTA.)

humped
curve
16
Zn2+
14

12 Mn2+ rising baseline

due to ionic
10 Ca2+ contraction
0 1 2 3 4 5 6 7 8 9 10 11
no of d-electrons
 Log K1(en) of d0 to d10 M(II) ions:

log K1(en) as a function of no of d-

electrons
= CFSE

12
10 double-
humped
logK1(en).

8 curve

6
Zn2+
4
rising baseline
Ca2+ Mn 2+
2 due to ionic
contraction
0
0 1 2 3 4 5 6 7 8 9 10 11
no of d-electrons
 Class a and Class b metals
• ChattandAhrlandclassifiedmetalsintothree types.
• Class a, Class b and border line.
• Class a :H,alkaliand alkalineearthmetals,Sc ➞Cr, Al ➞Cl, Zn ➞Br
,In, Sn ,Sb ,I,lathanidesand actinides
Class b:Rh,Pd,Ag ,Ir,Pt ,Au andHg
Borderline:Mn ➞ Cu,Tl ➞ Po,Mo ,Te,Ru,W, Re,Os andCd.
Class a metals form more stablecomplexeswith ligands in
which coordination atoms are from second period. ( N, O,F)
Classb metals form morestable complexes with ligands
having third period elements as ligating atoms. (P, S,Cl)
Class b metals are having capacity to form pi bonds with the ligand atoms.
The expansion is possibleonly from the third period donor atoms.
Border line metals do not showanynoticeabletrend.
• The bond between metal and ligand atomis, to some
extent due to the donation of electron pair to the
metal.

• If the metal is having a tendency attract the electron

pair (Higher electronegativity) then more stable
complexes are formed.
 Properties of ligand

• Size and charge

• Basic character

Chelateeffect

• Sizeof the chelate ring

Steric effect
 Size and chargeof the ligand
• To some extent we can say that if theligand is smaller in size
and bearing highercharge, it will form more stable complexes.
• ForexampleusuallyF- forms more stable complexesthat Cl-

• In the case of neutral mono dentate ligands, high

dipole moment and smallsize favour more stable
complexes.
 Basiccharacter of ligands

• If the ligand is more basicthen it will donate

the electron pair moreeasily.
• So with increased basiccharacter more
stable complexes canbe expected.
• Usually the ligands which bind strongly with H+
form more stable complexes.
• Thisis observedfor IA, IIA, 3d, 4f and5f
elements
What are the implications of the following results?
NiCl2 + 6H2O  [Ni(H2O)6]+2

[Ni(H2O)6]+2 + 6NH3  [Ni(NH3)6]2+ + 6H2O log  =8.6

[Ni(H2O)6]+2 + 3 NH2CH2CH2NH2 (en) log  =18.3

[Ni(en)3]2+ + 6H2O

[Ni(NH3)6]2+ + 3 NH2CH2CH2NH2 (en) log  =9.7

[Ni(en)3]2+ + 6NH3
 Complex Formation: Major
Factors
[Ni(H2O)6] + 6NH3 [Ni(NH3)6]2+ + 6H2O

NH3 is a stronger (better) ligand than H2O

 O NH3 >  O H2O
 [Ni(NH3)6]2+ is more stable

 G = H - TS (H -ve, S 0)

 G for the reaction is negative
 Chelate Formation: Major Factors
[Ni(NH3)6]2+ + 3 NH2CH2CH2NH2 (en)

[Ni(en)3]2+ + 6NH3

en and NH3 have similar N-donor environment

but en is bidentate and chelating ligand
rxn proceeds towards right, G negative
 G = H - TS (H -ve, S ++ve)
rxn proceeds due to entropy gain
S ++ve is the major factor behind chelate
effect
 Chelate Formation: Entropy Gain

Cd2+ + 4 NH3  [Cd(NH3)4]2+ Cd2+ + 4 MeNH2  [Cd(MeNH2)4]2+

Cd2+ + 2 en  [Cd(en)2]2+

Ligands log  G H S
kJmol-1 kJmol-1 JK-1mol-1

4 NH3 7.44 -42.5 - 53.2 - 35.5

4 MeNH2 6.52 -37.2 -57.3 - 67.3

2 en 10.62 -60.7 -56.5 +13.8

 Chelate Formation: Entropy Gain
Reaction of ammonia and en with Cu2+

[Cu(H2O)6]2+ + 2NH3  [Cu(NH3)2(H2O)2]2+ + 2 H2O

Log 2  = 7.7  H = -46 kJ/mol  S = -8.4 J/K/mol

[Cu(H2O)6]2+ + en  [Cu(en)(H2O)4]2+ + 2 H2O

Log K 1 = 10.6  H = -54 kJ/mol  S = 23 J/K/mol

 Chelateeffect
•• Thestability of the complex of a metal ion with a
bidentate ligand such asen is invariably significantly
greater than the complex of the same ion with two
monodentate ligands of comparable donor ability,
i.e., for example two ammoniamolecule.
•The attainment of extra stability by formation of ring
structures, bybi orpolydentateligands which includethe
metal, is termedas chelate effect.
 Stericfactors
• when bulky groups are present near or on the ligating
atom, the steric forces come into play.
• Presence of bulkier groups near coordination
sites reduce the chances of ligand getting closer
to the metal.
• Evenwhen complex is formed, to get relieved from
the steric hindrance the bond may dissociate. This
reduces the stability of complex
 Charge Effect of Metal Ions
• As the metal ion charge increases, the
ligands are drawn closer to the metal ion
because of its increased charge density.
• As the ligands move closer, they cause greater
splitting of the d orbitals, thereby producing a
larger Δ value.
• The magnitude of Δ for a given ligand
increases as the charge on the metal ion
increases.
NH3-Co+2 (weak field) NH3-Co+3 (strong field)
Molecular Orbital theory
• Molecular Orbital theory

• The bonding of many molecules can be explained using

valence bond theory and crystal field theory.
• However, another bonding theory, the molecular orbital
theory (MOT), is necessary to explain complete and accurate
electronic structure of molecules.
• Atomic orbitals are regions of space around the nucleus of
an atom, where an electron is likely to be found.
• Atomic orbitals allow atoms to make covalent bonds.
• The most commonly orbitals are s, p, d, and f.
• On the other hand, the space around the nucleus is zero
probability of finding any electron is called Node.
• The s orbital has no angular nodes and are spherical.
• The p orbitals have a single angular node across the
nucleus and are shaped like dumbbells.

• The d and f have two and three angular nodes, respectively.

• Only two electrons will be found in any orbital space as
defined by the Pauli Exclusion Principle.
[The Pauli Exclusion Principle states that, in an atom or molecule,
no two electrons can have the same four electronic quantum
numbers.
• As an orbital can contain a maximum of only two
electrons, the two electrons must have opposing spins. This
means if one is assigned an up- spin ( +1/2), the other must
be down-spin (-1/2).]
Molecular Orbital theory

• In molecular orbital theory, the molecule as a whole unit

rather than concentrating on individual atoms.

• Molecular orbital, is a space in a molecule where the

probability of finding an electron is higher.

• Molecular orbitals have some similarities to the atomic

orbitals.

• Each molecular orbital has a specific energy-level and a

specific shape, and each molecular orbital can be occupied
by a maximum of two electrons with opposite spins.
PX PX

PX PX

PY PY

PY PY
 Free metal Octahedra Complexes ligands lone
ion orbitals ion orbitals pair orbitals

4p
4s

3d
M
9 orbitals on metal
No. of e- depende on Ligands
the metal. 6 orbitals from ligands
12e-
 Molecular Orbital Model of [CoF6]3-
Octahedral Metal Complexes
Octahedral Metal Complexes
 Ligand Field Theory
• Ligand field theory can be considered an
extension of crystal field theory such that all
levels of covalent interactions can be
incorporated into the model. Treatment of the
bonding in LFT is generally done using
molecular orbital theory.
• Ligand Field Theory uses a molecular orbital
approach. Initially, the ligands can be viewed
as having a hybrid orbital or a p orbital
pointing to ward the metal to make σ bonds.
 Ligand Field Theory
• Application of molecular orbital theory to transition metal
complexes.
• Ligands are not point charges.
• Takes into account  bonding.
• Can be used to explain spectrochemical series.
• Better than valence-bond model or crystal field theory at
explaining experimental data.
 Ligand Field Strength
Strong Field Weak Field
Stronger  donor Weak  donor
 bonding Strong Lewis base Weak Lewis base
Stronger bonding interaction Weaker bonding interaction

eg

eg
o o
t2g

t2g
Empty  acceptor Filled  donor
 bonding  acid  base
Accepts from M Donates to M
 Ligand Field Strength
Strong Field Weak Field
Stronger  donor Weak  donor
1-  Bonding only Strong Lewis base Weak Lewis base
Stronger bonding interaction Weaker bonding interaction

Larger 
 Bonding +  bonding ( acceptor)

2-  Bonding +  bonding

 Bonding +  bonding ( donor)

Smaller 
 Ligand Field Strength
eg
eg
o o
t2g
t2g
Pure  donating ligands:
: en > NH3
 donating ligands:
 : H2O > F > RCO2 > OH > Cl > Br > I
 accepting ligands:
 : CO, CN-, > phenanthroline > NO2- > NCS-

Larger  The Spectrochemical Series Smaller 

CO, CN- > phen > NO2- > en > NH3 > NCS- > H2O > F- > RCO2- > OH- > Cl- > Br- > I-
 Ligand Field Theory
The A1g group orbitals have the
same symmetry as an s orbital on
the central metal.

The T1u group orbitals have the

same symmetry as the p
orbitals on the central metal.
(T representations are triply
degenerate.)
 Ligand Field Theory
The Eg group orbitals have
the same symmetry as the
dz2 and dx2-y2 orbitals on
the central metal. (E
representations are doubly
degenerate.)
Since the ligands don’t have a
combination with t2g
symmetry, the dxy, dyz and dxy
orbitals on the metal will be
non-bonding when
considering σ bonding.
Ligand Field Theory
The molecular orbital
diagram is consistent
with the crystal field
approach.
Note that the t2g
set of orbitals is non-
bonding, and the eg
set of orbitals is
antibonding.
• lowest unoccupied
molecular orbitals
(LUMO)
Ligand Field Theory
• The electrons from
the ligands (12
electrons from 6
ligands in octahedral
complexes) will fill
the lower bonding
orbitals.
• orbitals are almost
always the highest
occupied
{ • (HOMO)
Ligand Field Theory
• The electrons from
the 4s and 3d
orbitals of the
metal (in the first
transition row)
{ will occupy the
middle portion of
the diagram.
 Molecular Orbital of Complex

• The dz2, dx2-y2, 4s, 4px, 4py and 4pz orbitals will be
involved in the MOs in the  complex ions.
• The dxz, dyz and dxy orbitals (the t2g set) of the
metal ion do not overlap with ligand orbitals.
They are called nonbonding orbitals.
• The eg* orbitals is relatively little contribution
from ligand orbitals. This lack of mixing is
caused by the large energy difference between
the ligand orbitals and the netals ion 3d orbitals.
 The Effect of Weak Field Ligands
• A ligand with a electronegative donor atom will have lone
pair orbitals of very low energy (the electrons are very
firmly bound to the ligand); these orbitals do not mix very
thoroughly with the metal ion orbitals. This will result in a
small difference between the t2g and eg * orbitals.

The Effect of Strong Field Ligands

• The strong field ligands produce larger degree of mixing
between the orbitals of ligands and metal ions
• This gives a relatively large amount of d-orbital splitting,
and low spin case results.
 Considering π Bonding
• To obtain for π bonding, a set of cartesian
coordinates is established for each of the ligands.
The direction of the σ bonds is arbitrarily set as the
y axis (or the py orbitals). The px and pz orbitals
are used in π bonding.
• Consider only the px and pz orbitals on each of the
ligands.
• This reduces to T1g + T2g + T1u + T2u. The T2g set
has the same symmetry as the dxy, dyz and dxz
orbitals on the metal. The T1u set has the same
symmetry as the px, py and pz orbitals on the metal.
 π Bonding
The main source of π bonding is between the
dxy, dyz and dxz orbitals on the metal and the d, p
or π* orbitals on the ligand.

The ligand may have empty d or π* orbitals and serve as

a π acceptor ligand, or full p or d orbitals and serve as a
π donor ligand.
 π Bonding
The empty π antibonding orbital on CO can
accept electron density from a filled d orbital
on the metal. CO is a pi acceptor ligand.

filled d empty π*
orbital orbital
 π Donor Ligands (LM)
All ligands are σ donors. Ligands with
filled p or d orbitals may also serve as pi donor
ligands. Examples of π donor ligands are I-,
Cl-, and S2-. The filled p or d orbitals on these
ions interact with the t2g set of orbitals (dxy, dyz
and dxz) on the metal to form bonding and
antibonding molecular orbitals.
π Donor Ligands (LM)

The bonding orbitals,

which are lower in energy,
are primarily filled with
electrons from the ligand,
the and antibonding
molecular orbitals are
primarily occupied by
electrons from the metal.
π Donor Ligands (LM)
The size of ∆o
decreases, since it is now
between an antibonding
t2g orbital and the eg*
orbital.
This is confirmed by
the spectrochemical
series. Weak field ligands
are also pi donor ligands.
 π Acceptor Ligands (ML)
Ligands such as CN-, NO2-
and CO have empty π
antibonding orbitals of the
proper symmetry and energy
to interact with filled d orbita ls
on the metal.
The metal uses the t2g set of the
orbitals (dxy, dyz and dxz) to
engage in pi bonding with the
ligand. The π* orbitals on the
ligand are usually higher in
energy than the d orbitals on
the metal.
 π Acceptor Ligands (ML)
The metal uses the
t2g set of orbitals (dxy,
dyz and dxz) to engage
in pi bonding with the
ligand. The π* orbitals
on the ligand are
usually higher in
energy than the d
orbitals on the metal.
 π Acceptor Ligands (ML)

The interaction
causes the energy of
the t2g bonding orbitals
to drop slightly, thus
increasing the size of
∆o.
 Summary
1. All ligands are σ donors. In general, ligand that
engage solely in σ bonding are in the middle of the
spectrochemical series. Some very strong σ donors,
such as CH3- and H- are found high in the series.
2. Ligands with filled p or d orbitals can also serve as
π donors. This results in a smaller value of ∆o.
3. Ligands with empty p, d or π* orbitals can also

serve as π acceptors. This results in a larger value

of ∆o.
I-<Br-<Cl-<F-<H2O<NH3<PPh3<CO
π donor< weak π donor<σ only< π acceptor
 • Applications of chelates:
• Chelates are useful in water softening, medicical and analytical
chemistry and different industries such as chemical and food
industry and agriculture.
• In water softening:
• Calcium (Ca2+) and magnesium (Mg2+) ions are responsible for
hardening of water. These ions on reaction with soaps precipitate
out. In the presence of chelating ligands such as polyphosphates
and polydentate amino acids, no precipitation occurs as these ions
form chelate complexes with polyphosphate and polydentate
ligands present in soap and thus, soften water.
• In food industry:
• Metal-amino acid chelates are helpful in enhancing mineral
absorption such as Ferrous bis-glycinate. Chelates are also used
to preserve fruits, fruit juices, food stuffs etc.
• In agriculture:
• Metal chelates are used as common components of fertilizers in
agriculture. Micronutrients such as manganese, iron, zinc and copper
are required for the overall health of the plants. These micronutrients
along with EDTA form chelate fertilizers. Presence of chelates
enhance uptake of micronutrients by the plants. Chelating ligands are
also used to detoxify poisonous metal such as mercury, arsenic
and lead present in polluted water.
• In medical field:
• Tetracycline and quinolone can form chelate with Fe2+, Ca2+ and Mg2+
ions and thus, these chelates can be used as suppliments of these ions.
As EDTA softens the dentin, it is used in root canal treatment as an
intracanal irritant. Chelates of gadolinium are used as contrast agents
in MRI scans. Metal poisoning can be decreased by chelation with
EDTA as toxic metals such as mercury, arsenic, lead and other
radioactive metals can be excreted without further interaction with
the body by converting them into chemically inert form (EDTA metal
complex). Chelation is also used in the treatment for autism.
• Chemical applications:
• Homogeneous catalyst such as ruthenium(II) chloride
chelated with (a bidentate phosphine) is used in Noyori
asymmetric hydrogenation and asymmetric isomerisation
for the manufacture of synthetic (–)-menthol. Bio- Rust
and Evapo-Rust are chelating agents used for the
removal of rust from iron and steel. Metal chelates are
also used in dyeing industry.

• Physiological chemistry (in human body):

• In body fluids, citric, malic and tartaric acids, the
natural chelating agents, keep the metal ions away
from precipitation. The other physiologically important
chelates are haemoglobin, vitamin B12, chrorophyll,
cytochrome and plastocyanine
 Transition Metal Complexes in
Biological Molecules
• Metal ion complexes are used in humans for the transport and
storage of oxygen, as electron-transfer agents, as catalysts, and as
drugs.
 Biological Importance of Coordination
Complexes－Hemoglobin

• The principal electron transfer molecules in the

respiratory chain are iron-containing species
called cytochromes, consisting of two main part:
an iron complex called heme and a protein.
• A metal ion coordinated to a rather complicated
planar ligand is called a porphyrin.
• The various porphyrin molecules act as
tetradentate ligands for many metal ions,
including iron, cobalt and magnesium
heme complex

Chlorophyll
 Myoglobin

• Iron plays a principal role in the transport and storage

of oxygen in mammalian blood and tissues.
• Oxygen is stored using a molecule called myoglobin,
which contains a heme complex and a protein.
• In myoglobin, the Fe+2 ion is coordinated to four
nitrogen atoms of porphyrin ring and to one nitrigen
atom of the protein chain.
• Since Fe+2 ion is normally six-coordinate, this leaves
one position open for attachment of an O2 molecule.
 Hemoglobin
The transport of O2
in the blood is
carried out by
hemoglobin, a
molecule consisting
of four myoglobin
molecules units.
Chemistry In Action: Coordination Compounds in Living Systems
Chemistry In Action: Cisplatin – The Anticancer Drug
• Porphyrins are complexes
containing a form of the
porphine molecule shown
at the right.
• Important biomolecules
like heme and chlorophyll
are porphyrins.
• Porphines (like
chlorophyll a) are
tetradentate ligands .
• Coordination Complexes in Nature and Technology
• Chlorophyll, the green pigment in plants, is a complex that contains magnesium
. This is an example of a main group element in a coordination complex. Plants
appear green because chlorophyll absorbs red and purple light; the reflected light
consequently appears green. The energy resulting from the absorption of light is
used in photosynthesis.

• . (a) Chlorophyll comes in several different forms, which all have the same basic
structure around the magnesium center. (b) Copper phthalocyanine blue, a square
planar copper complex, is present in some blue dyes.