Coordination compounds

Co-ordination compounds play a vital role. The importance can be realized that life would not
have been possible without the existence of chlorophyll (Mg - complex) in plants and haemoglobin
(Fe- complex) in the blood of human beings, Vitamin B-12 (Cobalamin), Cis-Platin [Pt(NH3)2Cl2],
Prussian Blue Fe4[Fe(CN)6]3 etc… The study of these compounds will enlarge our understanding
of chemical bonding, physical properties such as magnetic properties of co-ordination compounds.

In all the examples we can observe the presence of metal elements which belongs to d-Block
element. Other than Mg. d-Block elements used as active metals due to availability of vacant d-
orbitals. So it accepts electron and acts as Lewis acid. And those ions or neutral molecules which
donates electrons acts as Lewis base.
Simple salt: The compounds which form only one type Cation and one type Anion on
dissolution.
Ex: KCl, NaCl, Na2CO3. etc…
Molecular/Addition compounds: When two or more compounds combine stoichiometrically
with definite ratios crystalline compounds are formed these are called molecular/addition
compound.
These are of two types 1) Double salts 2) Complex compounds or Coordination compounds.

1. Double Salts: These are the addition molecular compounds which are stable in solid state but
dissociate into constituent ions in the solution. Or those which produce one ion in common,
either it is Cation or Anion.
Ex: 1. Mohr’s salt - [FeSO4. (NH4)2SO4.6H2O] get dissociated into Fe2+, NH4+ and SO42- ions.
2. Carnalite - KCl.MgCl2.6H2O
3. Potash alum - K2Al2(SO4)3.24H2O.
2. Coordination compounds: The compounds in which metal atom or ion is bounded to ions
or neutral molecules through covalent coordinate bonds. Or Compounds which retain their
identity in solid state as well as in dissolved state.

Ex: 1.

2.

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Double salts Coordination compounds

1) Double salts exist only in solid state, in 1) It may exist in solid or aqueous state, in
aqueous solution they dissociate into their solution they do not dissociate into their
constituent ions completely constituent ions
2) These are ionic and do not contain coordinate 2) These may not be ionic but they always form
bond coordinate bond.
3) Properties of double salts are same as their 3) Properties are different from their constituents
constituents
4) Metal ions exhibit their normal valency 4) Metal ions surrounded by ions/ neutral
molecules exhibit beyond their normal valencies

Werner’s Theory of Coordination Compounds

A famous chemist Alfred Werner studied the nature and bonding of coordination compounds

He tried to explained change in colour and bonding in coordination compounds by treating above
solutions with AgNO3 solution. In first case

In second case

In third case:

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Postulates of Werner’s theory:

From conductivity measurement studies Werner Postulated theory of coordination compounds.
According to him
 Metals exhibit two types of valencies in the formation of complexes. They are primary
valency and secondary valencies.
 Primary valency correspond to Oxidation State (OS) of the metal and are satisfied by anions.
 These are ionisable and non-directional.
 Secondary valencies correspond to Coordination Number(CN) of the metal atom and are
satisfied by ligands.
 These are non-ionisable and directional. Hence, geometry is decided by these valencies.

Terms Related to Coordination Compounds

1. Complex ion or Coordination Entity
It is an electrically charged species in which central metal atom or ion is surrounded by number of
ions or neutral molecules
A) Cationic entity: the complex ion which carries positive charge. Ex: [Pt(NH 3)4]2+ .
[Co(NH3)4]2+
B) Anionic entity: the complex ion which carries negative charge. Ex: [Fe(CN)6]4-, [Cr(CN)6]3-

2. Central Atom or Ion: The atom or ion to which a fixed number of ions or molecules are
bound in a definite geometrical arrangements is termed as a central atom or ion. Transitional
metals are used as a central atom due to
 Relatively smaller size of the metal ions
 High ionic charge
 The availability of vacant d-orbital for bond formation.
Since it accepts a lone pair of electrons for the formation of coordinate bond, it is also referred as
Lewis acid. Ex: [NiCl2(H2O)4]. Ni is central metal atom.

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3. Ligands: Ligands are the atoms or molecules or ions which donate a pair of electrons to
the Central metal atom or ion and form a coordinate bond with it.
Depending upon the number of donor atoms available for coordination, the ligand may be
classified as:
3.1 Unidentate / Monodentate ligand: ligand, which has one donor site or atom,
i.e., the ligand bound to a metal ion through a single donor site. The number of such ligating
groups is called the Denticity of the ligand so, denticity =1for monodentate ligands.

Ex: H2O, NH3, I-, Br-, Cl-, OH-, etc.

3.2 Bidentate / Didentate ligand: It is the ligand. which have two donor sites or atoms,

Denticity=2
3.3 Polydentate: It is the ligand, which have several donor sites. EDTA have denticity of 6

3.4 Ambidentate ligands: A ligand which contains two donor atoms but only one of them forms
a coordinate bond at a time with central metal atom/ ion. Ex: NO2-, SCN-, CN- etc. i.e

3.5 Chelating ligands: Di or polydentate ligands cause cyclisation around the metal atom which
are known as chelate, such ligands uses two or more donor atoms to bind a single metal ion and
are known as chelating ligands.
The number of such ligating groups is called the Denticity of the ligand.

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A complex compound in which the donor atoms are attached to the metal so that the metal becomes
the part of the heterocyclic ring is called Chelate Complex.
 More the number of chelate rings, more is the stability of complex.
 The stabilization of coordination compounds due to chelation is known as chelate effect.

4. Coordination Number (CN)

The number of Unidentate ligands directly bonded to the central metal atom / ion is known as
coordination number of that metal ion / atom.
Ex: In [PtCl6]2-, Pt has coordination number 6. [Ag(NH3)2]2+, [ Zn(CN)4]2-, & [Ni(NH3)6]2+ the
coordination number of Ag, Zn & Ni are 2, 4 and 6 respectively.
In case of monodentate ligands, Coordination number = number of ligands in polydentate
ligands.
Coordination number = number of ligands * denticity
5. Coordination Sphere
The central ion and the ligands attached to it are enclosed in square bracket is known as
coordination sphere. The ionisable group written outside the bracket are known as counter ions.
Ex: The complex K4[Fe(CN)6] the coordination sphere is [Fe(CN)6]4-, and the counter ion is K+.

6. Coordination Polyhedron
The spatial arrangement of the ligands which are directly attached to the central atom or ion, is
called coordination polyhedron around the central atom or ion.
Ex: The coordination polyhedral of [Ni(CO) 4], [Co(NH3)6]3+ and [PtCl4]2- are Tetrahedral,
Octahedral and Square planar, Trigonal bipyramid respectively.

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7. Oxidation Number of Central Atom

The charge of the complex if all the ligands are removed along with the electron pairs that are
shared with the central atom, is called oxidation number of central atom.
Ex: [Cu(CN4)]3-, oxidation number of copper is +1, and represented as Cu(I).
Types of Complexes
1. Homoleptic complexes
Complexes in which the metal atom or ion is linked to only one kind of donor atoms, are called
homoleptic complexes. Ex: [Co(NH3)6]3+
2. Heteroleptic complexes
Complexes in which the metal atom or ion is linked to more than one kind of donor atoms are
called Heteroleptic complexes. Ex: [Co(NH3)4Cl2]+

IUPAC Naming of Complex Compounds

Naming is based on set of rules given by IUPAC.
1. Name of the compound is written in two parts (i) name of cation, and (ii) name of anion.
2. The cation is named first in both positively and negatively charged coordination complexes.
3. Naming the complex
a) Name the ligands first in alphabetical order before the name of central metal atom or ion.
b) while naming anionic ligands the terminal “e” is replaced by ‘0’, “-ide” as “-ido”, “-ate” as
“”-ato”, “-ite” as “-ito”. names of positive ligands end with ‘ium’ and names of neutral ligands
remains as such. But exception are there as we use aqua for H 2O, ammine for NH3, carbonyl
for CO and nitrosyl for NO

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4. Greek prefixesdi, tri, tetra,penta, hexa, etc are used to designate the number of indivdual ligand
in the complex ion, when the names of the ligands include a numerical prefix. Then the termsbis,
tris, tetrakis are used.
Ex: [NiCl2(PPh3)2] is named as dichlorobis(triphenylphosphine)nickel(II).
5. After naming ligands the name of central metal, If the complex part is cation the name of the
central metal is named with normal name. Ex: Co in a cationic complex is named as Cobalt.
Pt in a cationic complex is named as Platinum
Fe in a cationic complex is named as Iron.
If the metal atom is in anionic complex then Co is named as Cobaltate,
Pt is named as Platinate
Fe is named as ferrate.. etc ….
6. Following the name of the metal oxidation state of the metal in the complex is given as a roman
numeral in parentheses.
7. The neutral complex molecule is named similar to that of the complex cation. Ex:
(i) [Cr(NH3)3(H2O)3]Cl3 - triammine-triaquachromium (III) chloride
(ii) [Co(H2NCH2CH2NH2)3]2(SO4)3 - tris (ethane-l,2-diamine) cobalt (III) sulphate
(iii) [Ag(NH3)2] [Ag(CN)2] – diamminesilver(I) dicyanoargentate(I)
(iv) K4 [Fe(CN)6] - potassium hexacyanoferrate (II)
(v) Hg[Co(SCN)4] – Mercury tetrathiocyanato-cobaltate(III)
(vi) [CoCl2(en)2]Cl – Dichlorido-bis(ethane-1,2-diamine)cobalt (III) chloride.
(vii) [CoBr(NH3)5]SO4 - pentaammine-bromido-cobalt(III)sulphate
(viii)[Fe(NH3)6][Cr(CN)6] –Hexammine-iron(III)hexacyanidochromat(III)
(ix) [Co(SO4)(NH3)5]+ -pentaammine-sulphatoCobalt(III)ion
(x) K2[Zn(OH)4] – Potassium-tetrahydroxido-zincate(II)
(xi)[Fe(NH3)6](NO3)3 – Hexammine-Iron(III)nitrate
(xii) K3[CoF6] – Poyassiumhexafluridocobaltate(III)
(xiii) Fe4[Fe(CN)6]3 – Iron(III) hexacyanoferrate (II)

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Isomerism in Coordination Compounds

Isomers are those compounds which have the same chemical formula but different structural
arrangement of their atoms due to their structures are responsible for their different physical or
chemical properties.

A. Structural Isomerism
In this isomerism. isomers have different bonding pattern. Different types of structural isomers
are
1. Linkage isomerism: is shown by the coordination compounds having ambidentate ligands.
Ex: 1. [Co(NH3)5(NO2)]Cl and [Co(NH3)5(ONO)]Cl

2. [Mn(CO)5SCN] and [Mn(CO)5NCS]

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2. Coordination isomerism: arises due to the interchange of ligands between cationic and
anionic complexes of different metal ions present in a complex,
Ex: 1. [Pt(NH3)4] [MnCl4] and [PtCl(NH3)3] [MnCl3(NH3)]

2. [Co(NH3)6] [Cr(CN)6] and [Cr(NH3)6] [Co(CN)6]

3. Ionisation isomerism: arises due to exchange of ionisable anion with anionic ligand. Ex:

4. Solvate isomerism: This is also known as hydrate isomerism. In this isomerism, water is
taken as solvent. It has different number of water molecules in the coordination sphere and
outside it
Ex: 1. [Cr(H2O)6]Cl3 and [Cr(H2O)5Cl]Cl2. H2O

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2. [Co(NH3)4Cl2]H2O Cland [Co(NH3)4H2OCl]Cl2

B. Stereoisomerism
Stereoisomers have the same chemical formula and chemical bonds but they have different
spatial arrangement. These are of two types:
1. Geometrical isomerism Geometrical isomers are of two types i.e., cis and trans isomers.
This isomerism is common in complexes with coordination number 4 and 6.
Geometrical isomerism in complexes with coordination number - 4
1. Tetrahedral complexes do not show geometrical isomerism, because the molecule is
symmetrical in shape.
2. Square planar complexes of formula [MX2L2] (X and L are unidentate) show geometrical
isomerism. The two X ligands may be arranged adjacent to each other in a cis isomer, or opposite
to each other in a Trans isomer. Ex:

a)Square planar complex of the type [MABXL] (where A, B, X, L, are unidentate ligands) shows
three isomers, two cis and one trans. Ex: [Pt(NH3) (NO2)(NH2OH)(Py)].

Geometrical isomerism in complexes with coordination number – 6

a) Octahedral complexes of formula [MX2L4], in which the two X ligands may be oriented cis
or trans to each other, Ex: [Co(NH3)4Cl2]+.

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b) Octahedral complexes of formula [MX2L2], where X are unidentate ligands and L are bidentate
ligand form cis and trans isomers, Ex: [CoC12(en)2]

c) In octahedral complexes of formula [MX3L3], if three donor atoms of the same ligands occupy
adjacent positions at the corners of an octahedral face. it is known as facial (fac) isomer, when the
positions are around the meridian of the octahedron, it is known as meridional (mer) isomer. Ex:
[Co(NH3)3(NO2)3]

2. Optical isomerism: These are the complexes which have chiral structures. It arises when
mirror images cannot be superimposed on one another. These mirror images are called
enantiomers. The two forms are called dextro (d) and laevo (l) forms.
Tetrahedral complexes with formula [M(L) 2X2] show optical isomers [Co(L)2Cl2] and octahedral
complexes of [M(AA)X2L2] exhibit optical isomerism. Ex: Type [Co(en)(NH3)Cl2] &

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Valence Bond Theory (VBT)

This theory was proposed by L. Pauling in 1930 s.

All electrons are paired, therefore complex will be diamagnetic in nature.

Complex has unpaired electrons, therefore, it will be paramagnetic in nature with octahedral

geometry. 𝑀𝑎𝑔𝑛𝑒𝑡𝑖𝑐 𝑚𝑜𝑚𝑒𝑛𝑡 = √𝑛(𝑛 + 2)

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Where: n - is number of unpaired electrons

All electrons are paired so complex will be diamagnetic in nature. Square planar structure.
Since, complex has unpaired electrons. So it will be paramagnetic in nature. Tetrahedral
structure.

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Limitations of VBT
 This theory could not explain the quantization of the magnetic data
 Existence of inner orbital and outer orbital complex
 Change of magnetic moment with temperature and colour of complexes.

Crystal Field Theory

This theory was proposed by Bethe and Vleck.
 It assumes the central metal atom and ligands as point charges. when a complex is formed
central metal atom carries positive charge, Ligands possess negative charge
 This theory considers the interaction between central metal atom and ligand is purely
electrostatic
 When a complex is formed the central metal atom is surrounded by oppositely charged
ligands
 No hybridization takes place
 To form a bond the ligand molecule must approach towards central metal atom
 In absence of external magnetic field, the d orbital of central metal atom is degenerate but
this degeneracy breaks when ligand approaches.
 The d orbital splits into two sets: dxy, dyz, dzx & dx2-y2,dz2
Repulsive forces occur between electrons of metal and with lone pair ligands due to which energy
of electron fluctuate or changes.

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Crystal field splitting in octahedral complexes

To form octahedral complex the ligands have to approach central metal atom along the
coordination axis. During the approach the d orbitals whose lobes lie along the axis will experience
more repulsion due to this their energy will increase and the other non axial set will suffer less
repulsion. as a result, the non-axial will have less energy as compare to axial set (e g greater
than t2g). In case of octahedral complexes, energy separation is denoted by “Δ o” octahedral).

In octahedral complexes the six ligands approach the central metal ion along the axis of d x2-y2and
d z2 orbitals,The energy of eg orbitals will increase by (3/5) and t 2g will decrease by (2/5).
i.e. Energy of eg set of orbitals > energy of t 2g set of orbitals.
[The energy difference between t 2g and eg level is designated by “Δo” and is called crystal field
splitting energy.]

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a) If Δo < PE, the fourth electron enters one of the eg orbitals, ligands which produce this
effect are called as weak field ligands and form high spin complexes. (where, PE = energy
required for e- pairing in an orbital).
b) If Δo > PE, it becomes more energetically favorable for the fourth electron to occupy a t 2g
orbital with configuration t 42geog. Ligands which produce this effect are known as strong
field ligands and form low spin complexes.

Crystal field splitting in tetrahedral complex:

The ligands have to approach central metal atom in between the coordination axis. During the
approach the d orbital’s whose lobes lie along the axis will experience less repulsion due to this
their energy will increase and the other non-axial (dxy, dyz, dzx) set will suffer more repulsion. As
a result, the non-axial will have more energy as compared to axial set (t 2g greater than eg )

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Energy of ‘t 2’ set of orbitals > Energy of ‘e’ set of orbitals.

In such complexes d – orbital splitting is inverted and is smaller as compared to the octahedral
field splitting.
Orbital splitting energies are so low that pairing of electrons are not possible so these are high
spin complexes.

Color in Coordination Compounds

Most of coordination compounds are colored. It is due to electronic transitions between t 2g and eg
energy levels. The energy of an electron is increased by absorbing light energy and it moves to a
higher energy level.
Energy of a photon = Energy difference between the ground state and an excited state

h = Planck’s constant (6.63x10-34 J.sec.),

u = frequency of light,
E = energy of photon (measured with UV. or visible spectroscopy)
They absorb radiations which falls under particular color from the visible radiations and then emit
their complementary color.
Ex: Emerald absorbs yellow red color of visible light and emit green colourDue to d-d transition.

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These are Colored due to d-d transition. In the absence of ligands, crystal field splitting does not
occur and hence the substance is colorless.

Limitations of Crystal Field Theory

1. It does not consider the formation of ligand field theory of bonding in complexes.
2. It is also unable to account satisfactorily for the relative strengths of ligands e.g., it does
not explain why H2O is stronger ligand than OH-.
3. It gives no account of the partly covalent nature of metal-metal bonds.

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Ligand Field or Molecular Orbital Theory

This theory was put forward by Hund and Mullikan.
According to this theory, all the atomic orbitals of the atom participating in molecule formation
get mixed to give rise an equivalent number of new orbitals, called the molecular orbitals. The
electrons are now under the influence of all the nuclei.
Metal Carbonyls
The compounds of carbon monoxide with certain transition metals are known as metal
carbonyls
The organic compounds having metal atom directly attached to the carbon are known as

organometallic compounds. Here carbon monoxide (CO) acts as ligand.

Polynuclear metal carbonyls are also known e.g. Fe3(CO)12 Mn2(CO)10

The metal-carbon σ-bond is formed by the donation of lone pair of electron of carbonyl carbon
to vacant orbitals of the metal, and the M-C 𝜋-bond is formed by the donation of a pair of
electrons back, from a filled bonding π d-orbital of metal to the vacant anti bonding 𝜋* orbital
of carbon monoxide is called as
Backbonding or Synergic effect. Due
to synergic effect the bond between
CO and a Metal is strengthened.

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Bonding in Metal Carbonyls: 𝝈-bonded

R - Mg - X Alkyl magnesium halide commonly known as Grignard’s reagent
π - bonded
(CH3)4 Sn (Tetramethyl tin), (C2H5)2 Zn (Diethyl Zinc), n–C4H9Li (n–butyllithium)

(Zeise’s salt) Ferrocene Dibenzene chromium

Stability of Coordination Compounds

Consider the following equilibrium between undissociated complex ion and dissociated ion.

[MLn]b+ Ma+ +nLx-

The equilibrium constant Kc =

The smaller the value of Kc, the greater is the stability of complex ion and vice versa.
The reciprocal of equilibrium constant is called stability constant.

The higher the value of Ks, the more is the stability of complex ion.
Ex: Where stability constant for [Cd(CN)4]2- is 7.1 𝑋 1018 and stability constant for [Cu(CN)3]2-
is 2 𝑋 1027 . That’s why [Cu(CN)3]2- is more stable than [Cd(CN)4]2-.
The value of Ks depends on.
 Nature of central metal atom - The more the polarizing power of the central metal ion the

more is the stability of complex ion.

Thus complex of Fe3+ is more stable than Fe2+
 Nature of ligand - Since ligand is a Lewis base the more the basic character of ligand the more

is the stability of complex ion. Chelating ligands give much larger values of stability constant.

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Importance and Applications of Coordination Compounds

1. They are used in many qualitative and quantitative analysis.
a) Qualitative: In identification of Cu+2, and in separation of Ag+ from Hg+2.
b) Quantitative: In estimation of Ni+2 by using DMG (dimethyl glyoxime).
2. Hardness of water is estimated by complexometric titration of Ca+2 and Mg+2 with Na2 EDTA.
3. Purification of metals can be achieved through formation and subsequent decomposition of
their coordination compounds.

4. They have great importance in biological systems.

a) Ex: Heamoglobin: the red pigment of blood acts as is a oxygen carrier is coordination
compound of Iron(Fe),
b) Chlorophyll: the plant pigment responsible for photosynthesis is coordination compound
of Magnesium(Mg)
c) Vitamin B12: Cyanocobalamin: the anti-pernicious anemia factor, is a coordination
compound of Cobalt (Co)
5. Metal complex of Ag, Au, Cu, etc are used for electroplating of these metals
on the desired objects.
6. They are used as catalyst for many industrial processes. Ex: Wilkinson
Catalyst i.e. [RhCl(PPh3)3]. Used for hydrogenation of alkenes.
7. In medicinal chemistry, there is a growing interest of chelating therapy.
Cis Platina: used as a anticancer drug for treatment of various types of cancers.
EDTA: used in treatment of lead poisoning, to remove Lead (Pb+2) etc…
8. As an inorganic dye: Prussian blue- Fe4[Fe(CN)6]3.

--------------------- o --------------------

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