Coordination Chemistry

For BTech.
 Coordination Chemistry
➢ Coordination complexes have been known since the beginning of modern chemistry. Early
well-known coordination complexes include dyes such as Prussian blue.
➢ Early contribution by Swedish mineralogist and chemist Christian Wilhelm Blomstrand and
Danish chemist Sophus Mads Jørgensen.
➢ It was not until 1893 that the most widely accepted version of the theory today was
published by Alfred Werner (Nobel prize in Chemistry in 1913)

A Girl with a Kitten by Jean-Baptiste Entombment of Christ painted by

Perronneau in 1743. Pieter van der Werff in 1709.
Alfred Werner (1866 -1919)
Prussian blue is used as pigment in the girl’s dress, Mary’s
veil, and in the sky.
 Werner’s Theory
Werner postulated that metal ions have two different kinds of valency:
1. Primary valency = oxidation state of the metal ion
2. Secondary valency = Coordination number

Coordination number = number of bonds formed between the metal ion

and the ligands in the complex ion. It can range from 2 to 9.
• 6 and 4 (most common)
• 2, 5 and 8 (less common)

Ligands
• The molecules or ions coordinating to the metal are
the ligands.
• They are usually anions or polar molecules.
• The must have lone pairs to interact with metal

Coordination number = 6
 Coordination Number and Coordination Geometry

Coordination number = 2 Rare for most metals; common

for d10 metal ions, especially:
Cu+, Ag+, Au+, and Hg2+

Coordination number = 3 Less common.

Encountered with d10 metal
ions e.g. Cu+ & Hg2+
trigonal planar structure

[Ag(NH3)2+]Cl or [Ag(NH3)2+]NO3
 Coordination Number # 4 and Possible Geometries

Two possible geometries – tetrahedral

and square planar

Square planar: 4-coordinate complexes of

Tetrahedral 2nd & 3rd row transition metals with d8 e-
configurations, e.g. Rh+ , Pt2+ and Pd2+, also
encountered in some Ni2+ & Cu2+ complexes.

Permanganate (MnO42-) [PtCl4]2-, [PdCl4]2-, K[AuCl4]

 Coordination Number # 5 and Possible Geometries

Square pyramidal (spy)

Trigonal bipyramidal (tbp)

Coordination Number # 6 and Possible Geometries

Most common: six ligands at vertices of an octahedron or a distorted octahedron.

Octahedral Trigonal prismatic

[Co(NH3)6]3+ ion
 Transformation of Octahedron geometry to Trigonal
Prismatic geometry

octahedron trigonal prism octahedron

octahedron trigonal prism octahedron

➢ Trigonal prismatic coordination is related to octahedral coordination as shown above.

➢ Rotation of one triangular face in an octahedron, relative to its opposite until the two are
eclipsed, gives a triganal prismatic geometry.
➢ In fact, since continuation of this rotation gives another octahedral complex the trigonal
prismatic geometry is an intermediate in isomerization reactions involving octahedral
complexes.
Coordination Numbers – 7
Higher coordination number mostly occur with the complexes of lanthanides.

The ligand atom, which is capped, is shown with blue lines

 Coordination Numbers – 8

Keeping the base fixed, twist the

top face

➢ Both the square antiprismatic and dodecahedral coordination geometries are distorted cubic
geometries.
➢ The square antiprismaric coordination geometry is just a cubic coordination geometry in which one face
has been rotated 45° relative to its opposite face.
➢ The dodecahedral geometry may be thought of as a cube in which opposing faces are folded up and
down relative to one another as shown above.
 Coordination Numbers – 9
Nine coordinate complexes typically require larger transition metals, the lanthanides, and
actinides. Coordination geometries are typically tricapped trigonal prismatic. Simple
Examples include the aqua complexes [Sc(H2O)9]3+, [Y(H2O)9]3+ , and [La(H2O)9]3+ as well as
[TcH9]2- and [ReH9]2-.

More can be found:

https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Map%3A_Inorga
nic_Chemistry_(Miessler_Fischer_Tarr)/09%3A_Coordination_Chemistry_I_-
_Structure_and_Isomers/9.04%3A_Coordination_Numbers_and_Structures
 Coordination Numbers – 12; Icosahedron

➢ Twelve coordinate [Ce(NO3)6]2-, in which the nitrate oxygens define an

icosahedral coordination geometry as shown below.

Icosahedron

It is a polyhedron with 20 faces.

 Chelating Ligands and Chelate Complexes
• When a bidentate ligand binds to a metal,

Chelate
Dentate = Teeth

Bidentate, bridging ligand

Here, pyrazine, the middle ligand is a

bidentate ligand. But it bridges two metal
centres, does not chelate.
 Polydentate Ligands

Chelating agents generally form more stable complexes than do monodentate ligands -
which is known as chelate effect.
 Metal – EDTA complex
• Ethylenediaminetetraacetate, abbreviated EDTA, has six donor atoms.
• Wraps around the central atom like an octopus.

➢ Chelate effect: metal complexes of polydentate

ligands are more stable than complexes of
chemically similar monodentate ligands.

➢ EDTA is used as ligand in complexometric

titration.
 What is Plane Polarized Light?
❖ Unpolarized Light Consists of Waves Vibrating in Many Different
Planes
❖ The Rotation of the Plane of Polarized Light by an Optically Active
Substance

Chiral molecules can rotate the

plane of polarization of light
 Enantiomers
➢ Most of the physical properties of chiral molecules are the same, boiling point,
freezing point, density, etc.
➢ One exception is the interaction of a chiral molecule with plane-polarized light. If
one enantiomer of a chiral compound is placed in a polarimeter and polarized
light is passed through it, the plane of polarization of the light will rotate.
➢ Two enantiomers rotate the plane of polarization of light equal in opposite
direction.
➢ If one enantiomer rotates the light 32° to the right, the other will rotate it 32° to
the left.

➢ Equal mixture of two enantiomers (50:50) is called racemic mixture. The rotation of plane
polarized light by a racemic mixture is zero.
Q: Does [Co(en)2Cl2]Cl exhibit geometrical isomerism?
Yes, => trans- and cis- isomers. en = ethylenediamine

Does it exhibit optical isomerism?

Trans form – No
Cis form – Yes

These two forms are identical (superimposable)

– So, only one trans-isomer possible.
• Optically inactive

Two enantiomers of the cis-form

• Optically active

Two enantiomers of cis-form – these two forms are non-superimposable

 Octahedral Chiral Complexes
➢ Octahedral complexes of bidentate chelating ligands, e.g. ethylenediamine,
oxalate, 1,10-phenanthroline etc produces a series chiral metal complexes.
➢ Chirality of such complexes are related to helical chirality. Two enantiomers are
represented as delta (Δ) and lambda (Λ).

Anti-clockwise
clockwise
 Bonding Theories in Coordination Compounds
➢ There are mainly three theories which are used to describe the nature the nature of
metal-ligand bonding in coordination compounds.

1. Valence Bond Theory (VBT): VBT was developed by Linus Pauling and Others in 1930.
2. Crystal Field Theory (CFT): CFT was proposed by Hans Bethe in 1929.
3. Ligand Field Theory (LFT) or Molecular Orbital Theory (MOT): Developed by J. H. Van
Vleck, Muliken and Hund.

Linus Pauling Hans Bethe Vleck Mulliken (Nobel

Nobel 1954,1962 Nobel 1967 Nobel 1977 1966) & Hund
 Valence Bond Theory
Valence Bond Theory was the first theory used to explain the geometry and
magnetic property of many to coordination compounds. The basic idea of the
theory is that the formation of a complex is a reaction between a Lewis base (ligand;
electron donor) and a Lewis acid (metal or metal ion; electron acceptor) with the
formation of a coordinate-covalent bond (dative bond) between the ligand and the
metal. This is based on following assumptions:

1. The central metal atom or ion provides number of vacant s, p & d orbitals equal
to its coordination number to form coordinate bond with the ligand orbitals.
2. Each ligands has at least one σ-orbital containing a lone pair of electrons.
3. The empty orbitals of the metal atom or ion undergo hybridisation to form
same number of hybrid orbitals. These hybrid orbitals overlap with the filled σ-
orbitals of the ligands to form ligand to metal coordinate σ-bond.
4. The geometry of complex ion depends on hybridisation of metal orbitals.

It is usually possible to predict the geometry of a complex from the knowledge of its
magnetic behaviour on the basis of the valence bond theory.
 Valence Bond Theory
Limitations of VBT: The VBT reigned for a period of two decades in the realm of
coordination chemistry because of its simplicity and ease in explaining structural and magnetic
properties. It could adequately explain low-spin square-planar, high-spin tetrahedral and both
low- and high-spin octahedral complexes.
But with the progress of time following shortcomings were noticed with the VBT and it is now
largely abandoned.

Disadvantages:
1. It fails to predict whether a 4-coordinate complex will be tetrahedral or square-planar and
whether an octahedral complex will be low-spin or high-spin.
2. It fails to distinguish certain geometries like tetragonal or distorted octahedral.
3. It completely neglects excited states in a complex and can not explain absorption
spectrum.
4. It doesn't have scope for quantitative calculation of bond energy and stability of
complexes.
5. It does not adequately explain the magnetic data beyond specifying the number of
unpaired electrons.
6. Too much attention has been given on metal ion while the importance of ligands is not
properly addressed.
 Crystal Field Theory
➢ This theory is based on the concept that when the negative charges of the incoming
ligands (or the negative ends of dipolar molecules like NH3 and H2O) attract the
positively charged metal ion, there is also repulsive interaction between the d-electrons
present on the metal ion and the ligands. Certain assumptions are taken while dealing
with CFT.
1. The ligands are treated as point charges.
2. The interactions between metal ion and ligand are treated as purely electrostatic, no
covalent interactions are considered. (Covalent nature is considered in Ligand Field
Theory, and Molecular Orbital Theory).
3. In isolated gaseous metal ion, all of the five d-orbitals are degenerate.
4. When a hypothetical spherical field of ligand approaches the metal ion, d-orbitals still
remain degenerate, but their energy level is raised a bit due to repulsion between the
orbitals of metal & ligand. This energy level is called Bary center.
• But in the transition metal complexes, the geometry about the metal ions are
octahedral, tetrahedral or square planar etc., the field provided by the ligands is not at
all spherically symmetrical. Therefore, the d-orbitals are unequally affected by the
ligands, and degeneracy of the d-orbitals in metal is removed and split into different
energy levels (e.g. t2g or eg).
 Shapes of the d-orbitals
• To understand CFT, it is essential to understand the description of the lobes of d-orbitals
z
y z
z y

x x x

z
z y y

x x
Orbitals affected when ligands approach a metal in an octahedral
arrangement

Ligands with their pair of electrons

approach to the metal along the X, Y and Z
axes for an octahedral complex formation

d-orbitals having lobes pointing directly at x-, y-

and z-axis (yellow colour) are repelled most by
electron-electron repulsion.
Bary center

d-orbitals not pointing directly at x-, y- and z-axis (red

colour) are stabilized to maintain the overall energy same
Repeat…………..
Crystal Field Effects on Octahedral Complexes
 Crystal Field Effects on Octahedral Complexes

Bary center

➢ The energy gap between eg and t2g is called crystal field splitting energy and it is
denoted by Δo or Δoct or 10 Dq, where Δ represent Crystal field splitting energy, "o"
in Δo is for octahedral.
➢ Because the overall energy is maintained, the energy of the three t2g orbitals are
lowered or stabilised by 0.4 Δo (4 Dq) and the energy of the two eg orbitals are
raised or repelled by 0.6Δo (6 Dq) with respect to hypothetical the spherical crystal
field or Bary Centre.
 Ti3+ => d1 system

eg t2g

Energy

1 KJmol-1 = 83.7 cm-1, eg

Δo = 20300/83.7 = 243 KJmol-1

t2g
 Crystal Field Effects on Tetrahedral Complexes
In a tetrahedral crystal field, imagine four ligands lying at
alternating corners of a cube
The dx2-y2 and dz2 orbitals on the metal ion at the center
of the cube lie between the ligands, and the dxy, dxz, and
dyz orbitals point toward the ligands. As a result, the
splitting observed in a tetrahedral crystal field is the
opposite of the splitting in an octahedral complex.

0.4Δt

Because a tetrahedral complex has fewer ligands, the

0.6Δt
magnitude of the splitting is smaller. The difference
between the energies of the t2 and e orbitals in a
tetrahedral complex (Δt) is slightly less than half of the
splitting in analogous octahedral complexes (Δo).
Δ t = 4/ 9 Δ o
 Tetrahedral Crystal Field

the splitting observed in a tetrahedral crystal field is opposite of

splitting in octahedral complex.

Δ t = 4/ 9 Δ o
➢ Tetrahedral splitting is seldom large enough to result in pairing of the electrons.
As a result, low-spin tetrahedral complexes are not common.
➢ A rare example is Cr[N(SiMe3)2]3[NO]
Factors Affecting The Magnitude of ∆

5d > 4d > 3d
 Spectrochemical Series (strength of ligand interaction)

Increasing Δ

I- < Br- < SCN- (S-bonded) < Cl- < F- < OH- < C2O42- < O2- < H2O < NCS- (N-bonded)
< py < NH3 < en < 1,10-phenanthroline < NO2- < PPh3 < CN- < CO

• Weak field ligands: H2O, F−, Cl−, OH−

• Strong field ligands: CO, CN−, NH3, PPh3

Increasing 
Different ligands on same metal give different colors
Addition of NH3 ligand to [Cu(H2O)6]2+ changes its color

[Cu(H2O)6]2+ [Cu(NH3)6]2+
 Crystal Field Stabilization Energy (CFSE)
▪ The energy difference between the distribution of electrons in a particular crystal
field and that for all electrons in the hypothetical spherical or uniform field levels
is called the crystal field stabilization energy (CFSE). [relative to their mean
energy, Bary Centre]
n = number of electrons
CFSE = {-0.4 x n(t2g) + 0.6 x n(eg)} Δ0 in the respective levels

➢ For [Ti(H2O)6]3+ ion, the CFSE is 0.4 * 243 KJmol-1 = 97 KJmol-1

➢ Hence, the complex is stable.
Metal ions with d4, d5, d6, or d7 e- configurations can be
either high spin or low spin, depending on magnitude of o

Magnitude of o
Large o = low spin complex
Smaller o = high spin complex
CFSE for d4 system:
• For high spin, CFSE = 0.6 o
• For low spin, CFSE = 1.6 o
As Energy difference increases, electron configuration changes
“Low spin”
“High spin”
Number of unpaired
electron = 4

Number of unpaired
electron = 0

Co(III) is d6
 High Spin, Low Spin – the Pairing Energy (P)
➢ In order to force an electron to pair with another an energy called Pairing energy (P) has
to be invested.
➢ This is made up of two terms: (1) the Coulombic repulsion arising out of forcing two
electrons to occupying the same orbital; (2) the loss of exchange energy that occurs as
electrons with parallel spins are forced to have anti-parallel spins.
➢ If P > Δo, the configuration t2g3eg1 is preferred. This is known as weak field or high
spin situation.
➢ If Δo > P, the configuration t2g4eg0 is adopted. This leads to strong field or low spin
situation.
➢ The nomenclature high and low spin arises from the magnetic moment
differences between the two configurations. Thus, t2g4eg0, has only two unpaired
electrons and therefore would have a magnetic moment of 2.83 BM. In contrast
the configuration t2g3eg1 would have four unpaired electrons with a magnetic
moment of 4.90 BM
o P
[Fe(H2O)6]2+ d6 9350 19150 High Spin o < P
[Fe(CN)6]4- d6 32200 19150 Low Spin o > P
 Square planar complexes are different still

octahedral Distorted octahedral, Square planar

z-elongated

• Square planar geometry can be considered by removing the z-ligands gradually away from
the metal, and finally completely removing it to infinitely apart.
 Magnetic properties: Spin only and effective
The spin-only formula (μs) applies reasonably well to metal ions from the first
row of transition metals: (units = μB,, Bohr-magnetons)

Metal ion dn configuration μs(calculated) μeff (observed)

Ca2+, Sc3+ d0 0 0
Ti3+ d1 1.73 1.7-1.8
V3+ d2 2.83 2.8-3.1
V2+, Cr3+ d3 3.87 3.7-3.9
Cr2+, Mn3+ d4 4.90 4.8-4.9
Mn2+, Fe3+ d5 5.92 5.7-6.0
Fe2+, Co3+ d6 4.90 5.0-5.6
Co2+ d7 3.87 4.3-5.2
Ni2+ d8 2.83 2.9-3.9
Cu2+ d9 1.73 1.9-2.1
Zn2+, Ga3+ d10 0 0
 Tetrahedral Octahedral

• Number of unpaired electron

is 3, same for both the cases.

➢ The difference in observed magnetic moment is due to Orbital contribution

Temperature Dependence of Magnetic Moments

Fe(Phen)2(NCS)2

Fe2+ = d6

How crystal field theory explains temperature dependence of eff !

Molecular orbital picture required to explain charge transfer spectra
Includes both  and π (back bonding) bonding

t1u

Ferric Potassium
KMnO4 Prussian Blue thiocyanate dichromate
 As

As lone pairs of electrons on ligands approach

along x, y, and z axes.
 Charge-Transfer Complexes
• Intense color can come from “charge transfer”
• Ligand electrons jump to metal orbitals

KMnO4 KCrO4 KClO4

• No d orbitals in Cl,
• Mn(VII): electron configuration d0 orbitals higher in energy
• Cr(VI): electron configuration is d0 • No charge transfer in
• Charge transfer takes place from ligand to metal (LMCT) KClO4.
• Colour intensity from charge–transfer complex is always
very high.

K2Cr2O7
 Advantages and Disadvantages of Crystal Field Theory
Advantages over Valence Bond theory
1. Explains colors of complexes
2. Explains magnetic properties of complexes (without knowing
hybridization) and temperature dependence of magnetic moments.
3. Classifies ligands as weak and strong
4. Explains anomalies in physical properties of metal complexes
5. Explains distortion in shape observed for some metal complexes

Disadvantages or drawbacks
1. Evidences for the presence of covalent bonding (orbital overlap) in metal
complexes have been disregarded.
e.g. it does not explain why CO although neutral is a very strong ligand

2. Cannot predict shape of complexes (since not based on hybridization)

3. Charge Transfer spectra not explained by CFT alone
Q. Co[(NH3)6]3+ ion is:
(a) Paramagentic
(b) Diamagnetic
(c) Ferromagnetic
(d) Ferri magnetic.

Q. In K4[Fe(CN)6] the number of unpaired electrons in iron are:

(a) 0
(b) 2
(c) 3
(d) 5.

Q. A complex compound in which the oxidation number of a metal is zero, is

(a) K4[Fe(CN)6]
(b) K3[Fe(CN)6]
(c) [Ni(CO)4]
(d) [Pt(NH3)4]Cl2
Q. The magnetic moment (spin only) of [NiCl4]2- is
(a) 1.82 BM
(b) 5.46 BM
(c) 2.82 BM
(d) 1.41 BM
Q. Among the ligands NH3, en, CN – and CO the correct order of their increasing field
strength, is
(a) CO < NH3 < en < CN–
(b) NH3 < en < CN– < CO
(c) CN– < NH3 < CO < en
(d) en < CN– < NH3 < CO
Q. The spin only magnetic moment value (in Bohr magneton units) of Cr(CO)6 is
(a) 0
(b) 2.84
(c) 4.90
(d) 5.92

Q. The spin only formula (μs) for octahedral complexes is

(a) (4S(S+1))1/2 S = spin angular momentum = n/2
(b) (4S(S+1))1/2 + (L (L+1))1/2 (n = number of unpaired electrons)
(c) (L (L+1)) 1/2

(d) L(L+1)
Q. Which of the following species will be diamagnetic?
(a) [Fe(CN)6]4-
(b) [FeF6]3+
(c) [Co(C2O4)3]4-
(d) [CoF6]3-
Q. Which of the following octahedral complexes of Co (at. no. 27) will have highest
magnitude of Δoct?
(a) [Co(CN)6]3-
(b) [Co(C2O4)3]3-
(c) [Co(H2O)6]3+
(d) [Co(NH3)6]3+
Q. The magnetic moment of [Co(NH3)6]Cl3 is
(a) 1.73
(b) 2.83
(c) 6.6
(d) Zero
Q. In the complex compound K4[Ni(CN)4] oxidation state of nickel is ?
(a) -1
(b) 0
(c) +1
(d) +2
Q. The spin only magnetic moment value (in Bohr magneton units) of Cr(CO)6 is
a) 0
b) 2.84
c) 4.90
d) 5.92

Q. Calculate the CFSE values for d3 and d8 configurations of weak field octahedral
complexes.
a) 0 Δo and -1.2 Δo
b) 1.2 Δo and -1.2 Δo
c) -1.2 Δo and -1.2 Δo
d) -1.2 and 0
Q. Calculate the CFSE values for d4 and d7 configurations of high spin tetrahedral
complexes.
a) 0 Δo and 0 Δo
b) 0.18 Δo and 0.54 Δo
c) -0.54 Δo and -0.18 Δo
d) -0.18 Δo and -0.54 Δo

Ans: to solve this, you first need to calculate in terms of tetrahedral geometry, and then convert
to octahedral values by: Δt = 4/9 Δo

d4: e2, (t2)2

2 CFSE = (-0.6 x 2) + (0.4 x 2) Δt
= - 0.4 Δt
3 = - 0.177 Δo

d7: e4, (t2)3

CFSE = (-0.6 x 4) + (0.4 x 3) Δt
= - 1.2 Δt
= - 0.533 Δo
Thank you
 Isomerism in Coordination Compounds
➢ Isomers have the same molecular formula, but their atoms are arranged either
in a different order (structural isomers), or spatial arrangement (stereoisomers).
• Some Classes of Isomers

Isomers
(same chemical composition, but
different bonding and properties)

Structural Isomers Stereo-Isomers

(different bonds) (same bonds, different
spatial arrangements)

Ionisation Hydrate Coordination Linkage

isomerism isomerism isomerism isomerism Geometric Optical
isomerism isomerism
(cis-trans)
 Structural Isomerism
1. Ionization isomerism
2. Hydrate isomerism
3. Coordination isomerism
4. Linkage isomerism

➢ Ionization isomerism:
• Same overall Composition. But the composition of the coordination sphere
varies.
• Some isomers differ in what ligands are bonded to the metal and what is
outside the coordination sphere; these are also called coordination-sphere
isomers.

• [Cr(NH3)5SO4]Br and [Cr(NH3)5Br]SO4

• Here, the counter ions (anion) are different. In the former, sulfate
coordinates to metal and bromide is counter ion. In the latter one, opposite
happens.
• Upon ionization, or upon dissolution in water, we get bromide anion in the
former case, where as, we will get sulphate anion in the latter case.
➢ Hydrate (or solvate) isomerism

• Three isomers of CrCl3(H2O)6 are

• The violet [Cr(H2O)6]Cl3,
• The green [Cr(H2O)5Cl]Cl2 ∙ H2O, and
• The (also) green [Cr(H2O)4Cl2]Cl ∙ 2H2O.

➢ Coordination isomerism
• They differ in coordination entities
• For example, [Co(NH3)6][Cr(CN)6] and [Co(CN)6][Cr(NH3)6] are coordination
isomers.

[Co(NH3)6]3+ & [Cr(CN)6]3- [Cr(NH3)6]3+ & [Co(CN)6]3-

• They have same total compositions, but they have different complex entities
(complex ions).

• Another example: [Zn(NH3)4][CuCl4] and [Cu(NH3)4][ZnCl4] compounds are also

coordination isomers.
➢ Linkage Isomerism
• Composition of the complex ion is the same, but the point of attachment of
at least one of the ligands differs.
• The donor atom of ligand varies

➢ Linkage Isomerism of NO2–

Same ligand, but the donor Nitro- Nitrito-

atom is different

➢ Linkage Isomerism of thiocyanate (SCN−) & isothiocyanate (NCS−)

• linkage isomers are violet-coloured

[(NH3)5Co(SCN)]2+ and orange-
coloured [(NH3)5Co(NCS)]2+

Violet Orange
 Stereoisomer
❖Geometrical Isomers
• Geometrical isomers are most important for square planar & octahedral
complexes.
➢ Square planar complexes:
• All the vertices of a square are equivalent, it does not matter which vertex is occupied by
ligand B in a square planar MA3B complex.
• Only one geometrical isomer is possible

• For MA2B2 square planar complexes, there are other possible arrangements.

Cis-form

Trans-form
 Trans- and Cis- isomers of Square Planar Complexes
➢ With these geometric isomers, two
chlorines and two NH3 groups are
bonded to the platinum metal, but
are clearly different.
➢ cis-Isomers have like groups on the
same side.
➢ trans-Isomers have like groups on
Cis-[Pt(NH3)2Cl2] Trans-[Pt(NH3)2Cl2] opposite sides.
Cis-Platin Trans-Platin

❖ Number of each atom is the same, the mode of bonding the same, but the arrangement
in space is different

❖ Cis-platin is used to treat various types of cancers. Where as, the trans-form does not
exhibit a comparably useful pharmacological effect.
 Isomerism in Octahedral Complexes
• Only one structure possible for octahedral complexes (if only one ligand is
different from other five): (MA5B)
• since all six vertices of an octahedron are equivalent.

➢ If two ligands in an octahedral complex are different from other four (MA4B2),
two isomers are possible: trans- and cis- isomers.
 Example: Isomerism of octahedral complex [Co(NH3)4Cl2]Cl

Cis-form Trans-form

• Same chemical composition

• Same number of groups of same types are attached to same metal.
 Isomerism in Octahedral Complexes
Replacing another A ligand by B gives an MA3B3 complex for which there are two
isomers:
• Fac: 3 ligands of each kind occupy opposite triangular faces of the octahedron
• Mer: 3 ligands of each kind lie on the meridian (cut across flat mid-plane)

Example: Cis-cis and cis-trans

geometrical isomers [Co(NH3)3Cl3]

Fac = facial
Mer = meridional
Fac- Mer-
 Stereoisomerism
❖ Optical Isomerism: Isomers have opposite effects on plane-polarized light.

❖ Optical Activity
• Exhibited by molecules that have nonsuperimposable mirror images (chiral
molecules).
• Enantiomers – isomers of nonsuperimposable mirror images.

➢ A Human Hand Exhibits a Nonsuperimposable Mirror Image

• Just as a right hand will not fit into a left glove, two enantiomers cannot be superimposed
on each other.
Thank You for Your
Attention