Chapter -09

CO-ORDINATION COMPOUNDS

 A coordination compound is a molecular compound contains a central metal atom or ion

(cation) surrounded by number of oppositely charged ions or neutral molecules.
 These ions or molecules are bonded to the metal atom or ion by a coordinate bond.
Co-ordinate bond
 A coordinate bond is a covalent bond (a shared pair of electrons) in which both electrons
come from the same atom

Difference b/w Double salt and Complex salt

 A salt that keeps its identity only in solid state is called a double salt. In solution they
dissociate into component ions.
Eg: Mohr’s salt [FeSO4. (NH4)2SO4.6H2O],
Carnallite [KCl.MgCl2.6H2O]
Potash alum [K2SO4.Al2(SO4)3.24H2O].
 The salt that keeps its identity both in solid and solution states is called a complex salt.
eg: potassium ferrocyanide {K4[Fe(CN)6]}, [Cu(NH3)4]SO4, K2[PtCl4], [Ni(CO)4] etc.
Terms related to co-ordination compounds
 Central meatal ion

ligand

[Co (NH3)6]Cl3 Counter ion (anion)

Co-ordination sphere Co-ordination no

Co-ordination entity:
 The central metal atom or ion and ligands form a co-ordination entity.
 For example, [CoCl3(NH3)3] is a co-ordination entity in which the cobalt ion is surrounded
by three ammonia molecules and three chloride ions.
Central atom/ion:
 In a co-ordination entity, the atom/ion to which a fixed number of ions/neutral
molecules are attached is called the central atom or ion.
 For example, the central atom/ion in the co-ordination entities: [NiCl2(H2O)4],
[CoCl(NH3)5]2+ and [Fe(CN)6]3– are Ni2+, Co3+ and Fe3+ respectively.
 These central atoms/ions are also referred to as Lewis acids, since they accept electron
pairs from ligands.
Ligands:
 A molecule, ion or group that is bonded to the metal atom or ion in a complex or
coordination compound by a coordinate bond is called ligand.
 For a species to act as ligand, it can donate at least one pair of electrons to the central
atom.
 The atom of the ligand which is directly bonded to the central atom or ion is called co-
ordinating atom or donor atom.
 Examples for ligands are Cl-, Br-, F-, I-, OH-, CN-, NC-, CNO-, NCO-, SO42-, NO3-, CNS-, H2O,
NH3, CO etc.
Types of ligands

 Based on the number of donor atoms of the ligand that binds to a metal ion or atom, the
ligands are classified as follows:
▪ Monodentate ligand
▪ Di dentate ligand
▪ Polydentate ligand

a) Monodentate or unidentate ligand:

 A ligand that binds to the central atom/ ion through a single donor atom, is said to be
unidentate ligand.
 Eg: Cl-, Br-, I-, OH-, CN-, CNO-, H2O, NH3, CO etc.

b) Bidentate (Didentate) ligands:

 A ligand that binds to the central atom through two donor atoms is called a bidentate
ligand.
 Eg: Ethane-1,2-diamine or ethylenediamine (H2NCH2CH2NH2) notated as ‘en’ and oxalate
ion (C2O42–).

c) Polydentate ligand:

 A ligand that binds to the central atom through more than two donor atoms is called
polydentate ligand.
 Eg: Triethylamine ammonia [N(CH2-CH2-NH2)3], Ethylenediamine tetraacetate ion
(EDTA4–) etc.
 Ethylenediamine tetraacetate ion (EDTA4–) is an important hexadentate ligand. It can
bind through two nitrogen and four oxygen atoms to a central metal ion.
Ambidentate ligands:

 They are unidentate ligands which contain more than one donor atoms.
 They can co-ordinate through two different atoms.
 Examples of such ligands are the NO2- , CN-, SCN–, CNO- etc.
 NO2- ion can co-ordinate either through nitrogen or through oxygen atom to the central
metal atom/ion. If the donor atom is N, it is written as NO2- and is called nitro (N) and if
it is O, it is written as ONO- and is called nitrito(O).
 Similarly, SCN– ion can co-ordinate either through sulphur atom (←SCN – thiocyanato) or
through nitrogen atom (←NCS – isothiocyanato).

Chelating Ligands:

 Di- or polydentate ligands can bind to the central atom through two or more donor
atoms and form ring complexes. Such complexes are called chelates and such types of
ligands are said to be chelating ligands.
 Complexes containing chelating ligands are more stable than those containing
unidentate ligands.
 For eg: the complex [Co(en)3]3+ is a chelate and ethane-1,2-diamine (en) is a chelating
ligand.

Co-ordination number:
 The co-ordination number (CN) of a metal ion in a complex can be defined as the total
number of ligand donor atoms to which the metal is directly bonded.
 For example, in the complex ion [PtCl6]2– the coordination number of Pt is 6,
in the complex ions, [Fe(C2O4)3]3– and [Co(en)3]3+ the co-ordination number of both Fe
and Co, is 6 because C2O42– and en (ethane-1,2- diamine) are bidentate ligands.
 The co-ordination number of the central atom/ion is determined only by the number of
sigma bonds formed by the ligand with the central atom/ion. Generally, the co-
ordination number of most of the complexes is 2, 4 or 6
 Oxidation number of central atom:
 The oxidation number of the central atom in a complex is defined as the residual charge
on it, if all the ligands are removed along with their electron pairs that are shared with
the central atom.
 The oxidation number is represented by a Roman numeral in simple brackets.
 For example, oxidation number of copper in [Cu(CN)4]3– is +1 and it is written as Cu(I).

Homoleptic and Heteroleptic complexes:

 Complexes which contain only one type of ligand are called homoleptic complexes.
Eg: [Co(NH3)6]3+, [Fe(CN)6]4- etc.
 Complexes which contain more than one type of ligands are called heteroleptic
complexes.
Eg: [Co(NH3)4Cl2]+ , [Cu(NH3)2Cl2] etc.

Nomenclature of Co-ordination Compounds

Rules:
 The cation is named first.
 Ligands arranged in the alphabetic order.
 The name anionic ligand end with ‘o’
 If two or more same ligands are present use di (2),tri(3)..etc
 In the case of ligand like en, PPh3 use bis (2),tris,tetrabis,..etc
 The oxidation number of the metal is written in roman numeral after the metal.
Note:
In the case of negative complexes, the name of the metal end in `ate’
In the case of negative complexes, use Latin name of the metal.
Fe- ferrum
Cu- cuprum
Au- aurum
Ag-argentum
Sn- stannum
Isomerism in Co-ordination Compounds
 Compounds that have the same molecular formula but different structural formula or
spatial arrangement of atoms are called isomers and the phenomenon is called
isomerism.
 Isomers differ in physical or chemical properties.
 The isomerism shown by co-ordination compounds are broadly divided into two –
▪ structural isomerism
▪ stereo isomerism.

(I) Structural Isomerism

 These are isomers which differ in the structural arrangement of ligands around the
central atom.
 They are of four types:
 ▪ Ionisation Isomerism
▪ Linkage isomerism
▪ Co-ordination Isomerism
▪ Solvate isomerism

1) Ionisation Isomerism:

 It arises due to the inter change of ions between the inside and outside of co-ordination
sphere.
 They give different types of ions in aqueous solution.
 In order to show this isomerism, the ion outside the co-ordination sphere can also act as
ligand.
 An example is [Co(NH3)5SO4]Br and [Co(NH3)5Br]SO4.

2) Linkage isomerism:

 It arises in a co-ordination compound containing ambidentate ligand, which can bind to

the central atom through more than one donor atoms.
 For example complexes containing thiocyanate ligand, SCN–, may bind either through
nitrogen to give M–NCS or through sulphur to give M–SCN.
 Another e.g. is [Co(NH3)5(ONO)]Cl2, in Co-ordination Compounds-which the nitrite ligand
is bound through oxygen (–ONO), and [Co(NH3)5(NO2)]Cl2 in which the nitrite ligand is
bound through nitrogen (–NO2).

3) Co-ordination Isomerism:

 If both anionic and cationic parts are complexes, the isomerism arises due to the
interchange of ligands between cationic and anionic entities. This type of isomerism is
called co-ordination isomerism.
 An example is [Co(NH3)6][Cr(CN)6], in which the NH3 ligands are bound to Co3+ and the
CN– ligands to Cr3+ and [Cr(NH3)6][Co(CN)6], the NH3 ligands are bound to Cr3+ and the
CN– ligands to Co3+ .

4) Solvate isomerism:

 This form of isomerism is also known as ‘hydrate isomerism’ if water is involved as the
solvent. This is like ionisation isomerism.
 Solvate isomers differ in the no. of solvent molecule which are directly bonded to the
metal ion as ligand.
 An example is [Cr(H2O)6]Cl3 (violet) and its solvate isomer [Cr(H2O)5Cl]Cl2.H2O (grey-
green).

(II) Stereoisomerism

 These are isomers which differ only in the spatial arrangement of ligands around the
central atom. They have same atom to atom bond.
 These are of two types:
▪ Geometrical isomerism
▪ Optical isomerism

i) Geometrical Isomerism:

 This type of isomerism is shown by heteroleptic complexes.

 It arises due to the different possible geometric arrangements of the ligands around the
central atom.
 It is mainly found in co-ordination complexes with co-ordination numbers 4 (square
planar complexes) and 6 (octahedral complexes).
 Geometrical isomer in which the same ligands are on the same side of the central metal
atom is called cis isomer and the isomer in which the same ligands are on the opposite
side is called trans isomer.
 Square planar complexes with formula [MX2L2] (X and L are unidentate ligands) can show
this isomerism. E.g.: [Cu (NH3)2Cl2]
 Octahedral complexes with formula [MX2L4] can also show this type of isomerism. Here
the two ligands X may be oriented cis or trans to each other.

 This type of isomerism also arises when bidentate ligands (L – L) are present in
complexes with formula [MX2(L – L)2] e.g: [Co(en)2Cl2] +

 Tetrahedral complexes do not show geometrical isomerism because in a tetrahedron all

the positions are equivalent. So, the relative positions of the ligands attached to the
central metal atom are the same with respect to each other.

Fac-mer isomerism:

 It is a type of geometrical isomerism occurs in octahedral co-ordination entities of the

type [Ma3b3].
 If similar ligands occupy three adjacent positions of an octahedral face, it is called facial
(fac) isomer.
 When the positions are around the meridian of the octahedron, it is called meridional
(mer) isomer.
 Eg. [Co(NH3)3(NO2)3].
ii) Optical Isomerism

 Optical isomers are mirror images that cannot be superimposed on one another.
 These are also called enantiomers.
 The molecules or ions that cannot be superimposed are called chiral.
 There are two forms of optical isomers –
• dextro (d)
• laevo (l)
 depending upon the direction they rotate the plane of polarised light in a polarimeter (d
rotates to the right, l to the left).
 Optical isomerism is common in octahedral complexes involving bidentate ligands.
 In a co-ordination entity of the type [PtCl2(en)2]+, only the cis-isomer shows optical
activity.
 The trans- isomer has a plane of symmetry and is optically inactive.

Theories of Co-ordination Compounds

Werner’s Co-ordination Theory

 It was Alfred Werner who first proposed a theory for co-ordination compounds.
 He studied the amino complexes of Co, Pt etc.
 Postulates of Werner’s Co-ordination Theory:
o Every metal has two types of valencies – primary (10) valency and secondary (20)
valency.
o Primary valency is ionisable, while secondary valency is non-ionisable.
o Primary valency gives the oxidation state of the metal, while secondary valency
gives the co-ordination number of the metal.
o Primary valency is always satisfied by –ve ions, while secondary valency may be
satisfied by –ve ions or neutral molecules.
o Every metal has a fixed number of secondary valency
o The primary valencies are non-directional, while the secondary valencies are
directional and give definite geometry to the complex.

Demerits:

 Werner could not explain why only certain elements form co-ordination compounds.
 he could not explain the directional nature of bonds in co-ordination compounds and
their magnetic and optical properties.

In order to explain the above properties, many theories such as Valence Bond Theory
(VBT), Crystal Field Theory (CFT), Ligand Field Theory (LFT) and Molecular Orbital Theory
(MOT) are proposed.

The Valence Bond Theory (VBT)

 This theory was put forward by Linus Pauling.

 The important postulates of this theory are:
o The M-L bond in complexes is formed by the donation of pairs of electrons by
ligands to the metals
o In co-ordination compounds, the central metal atom/ion provides some vacant
orbitals in order to accommodate the electrons donated by the ligands.
o The number of vacant orbitals formed is equal to the co-ordination number of
the metal atom
o The vacant orbitals of the metal undergo hybridisation to form a set of new
orbitals called hybrid orbitals.
o The type of hybridisation gives the shape of the compound.
 CN HYBRIDISATION SHAPE
6 d2sp3 Octahedral
6 sp3d2 Octahedral
4 dsp2 Square planar
4 sp3 Tetrahedral
5 Sp3d TBP

o Each ligand should contain at least one pair of electrons.

o The vacant hybrid orbitals of the metal overlap with the filled orbitals of the
ligands to form M-L co-ordinate bond.
o If a complex contains unpaired electron, it is paramagnetic and if it contains only
paired electron, it is diamagnetic.

Applications of VBT

 It is usually possible to predict the geometry of a complex from the knowledge of its
magnetic behaviour based on the valence bond theory.
 Example:

[Fe (CN)6]3-

Here the central atom Fe is in +3 oxidation state.

26Fe – [Ar]3d6 4s2

Fe3+ - [Ar]3d5 4s0 4P0

Fe3+ - ↑ ↑ ↑ ↑ ↑

3d5 4s0 4p0

In this complex, the co-ordination number of Fe is 6 and hence the no. of vacant orbitals
required = 6. In presence of the ligand CN-, the electrons in 3d level get paired.

↑↓ ↑↓ ↑↓

Now the two 3d orbitals, one 4s orbital and three 4p orbitals undergo d2sp3 hybridization to
form 6 new orbitals.
Six pairs of electrons, one from each CN- ions, occupy these six hybrid orbitals. Thus, the
complex has octahedral geometry and is paramagnetic because of the presence of one
unpaired electron.

In the formation of this complex, since the inner d orbitals (3d) are used for hybridization,
the complex is called an inner orbital or low spin or spin paired complex.

↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

Magnetic Properties of Co-ordination Compounds

 By knowing the magnetic moment, we can predict the geometry of complexes.

o For eg. [Mn(CN)6]3– has magnetic moment of two unpaired electrons while
[MnCl6]3– has a paramagnetic moment of four unpaired electrons.
o [Fe(CN)6]3– has magnetic moment of a single unpaired electron while [FeF6]3– has
a paramagnetic moment of five unpaired electrons.
o [CoF6]3– is paramagnetic with four unpaired electrons while [Co(C2O4)3]3– is
diamagnetic. This can be explained by valence bond theory in terms of formation
of inner orbital and outer orbital co-ordination entities. [Mn(CN)6]3–, [Fe(CN)6]3–
and [Co(C2O4)3]3– are inner orbital complexes involving d2sp3 hybridisation, the
former two complexes are paramagnetic and the latter diamagnetic.
o On the other hand, [MnCl6]3–, [FeF6]3– and [CoF6]3– are outer orbital complexes
involving sp3d2 hybridisation and are paramagnetic corresponding to four, five
and four unpaired electrons

Limitations of Valence Bond Theory

Even though the VB theory explains the formation, structures, and magnetic behaviour of
co-ordination compounds, it has the following limitations:

 It involves several assumptions.

 It does not give quantitative interpretation of magnetic data.
 It does not explain the colour exhibited by co-ordination compounds.
 It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities
of co-ordination compounds.
 It does not make exact predictions regarding the tetrahedral and square planar
structures of 4-co-ordinated complexes.
 It does not distinguish between weak and strong ligands

Crystal Field Theory

 According to CFT, a complex is a combination of a positive central metal ion and ligands
which may be negative ions or neutral molecule.
 The -ve ligands are regarded as point charges and neutral ligands are considered as
dipoles.
 The M-L bond is considered as purely ionic which is arises from the electrostatic
interaction b/w positive charged metal ion and -ve charged ligand.
 The five d orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they
are degenerate. This degeneracy is maintained if a spherically symmetrical field of
negative charges surrounds the metal atom/ion.

Crystal field splitting

 Splitting of degenerate metal `d’ orbital in the presence of ligand into sets of different
energies is called Crystal field splitting.

(a) Crystal field splitting in octahedral complexes

 in a octahedral complex, the metal ion is at the centre of the octahedral and the ligands
are at the six corners.
 The six ligands are approach the metal through the axis.
 So the d-orbital present on the axis [dx2-y2, dz2] will have higher energy (called eg orbitals)
due to the repulsion while the orbital present in b/w the axis (dxy, dyz and dxz orbitals)
will have lower energy (called t2g orbitals)
(a) Crystal field splitting in tetrahedral complex

 In tetrahedral complexes, the 4 ligands approach the metal in b/w the axis.
 So the metal d orbital presence in b/w the axis will have higher energy and the orbitals
present on the axis will have lower energy.

Filling of electrons

 For d1, d2 and d3 coordination entities, the d electrons occupy the t2g orbitals singly in
accordance with the Hund’s rule.
 For d4 ions, two possible patterns of electron distribution arise:
▪ the fourth electron could either enter the t2g level and pair with an
existing electron, or
▪ it could enter into the eg level. Here the electron distribution depends on
the relative magnitude of the crystal field splitting (Δo) and the pairing
energy (P). If Δo < P, the fourth electron enters one of the eg orbitals
giving the configuration t2g3 eg1.
 ▪ Ligands for which Δo < P are known as weak field ligands and form high
spin complexes. If Δo > P, the fourth electron occupy a t2g orbital with
configuration t2g4 eg0.
▪ Ligands for which Δo < P are known as strong field ligands and form low
spin complexes

Colour in Coordination Compounds

 Most of the complexes of transition metals are coloured. This can be explained in terms
of the crystal field theory.
 In presence of the ligands, the crystal field splitting occurs. So the electrons from lower d
level (t2g level) can excite to higher d level (eg level). For this some energy is required,
which is absorbed from the white light.
 The colour of the complex is complementary to that which is absorbed.
 Thus according to crystal field theory the colour of the coordination compounds is due
to d-d transition of the electron.
 For example, the complex [Ti(H2O)6]3+ is violet in colour. This is an octahedral complex
where the single electron (Ti3+ is a 3d1 system) in the metal d orbital is in the t2g level in
the ground state of the complex.
 The next higher state available for the electron is the empty eg level. If light
corresponding to the energy of yellow-green region is absorbed by the complex, it would
excite the electron from t2g level to the eg level. Consequently, the complex appears
violet in colour.
 In the absence of ligand, crystal field splitting does not occur and hence the substance is
colourless. For example, when [Ti(H2O)6]3+ is heated it becomes colourless. Similarly,
anhydrous CuSO4 is white, but CuSO4.5H2O is blue in colour.
 The colour of a complex depends on the strength of ligand and the nature of the field.

Limitations of Crystal Field Theory

• From the assumptions, that the ligands are point charges, it follows that anionic
ligands should exert the greatest splitting effect. But the anionic ligands are found at
the low end of the spectrochemical series.
 • It does not consider the covalent character of bonding between the ligand and the
central atom.

Stability of Coordination Compounds

 The stability of a complex in solution refers to the degree of association between the
metal ion and ligands in the state of equilibrium.
 The magnitude of the (stability or formation) equilibrium constant for the association,
quantitatively expresses the stability.
For a reaction of the type:
M + 4L → ML4, β = [ML4]/[M][L]4
 As the value of the stability constant increases, the stability of the complex also
increases. The above reaction can be considered to take place in 4 steps:
M + L → ML, K1 = [ML]/[M][L]
ML + L → ML2, K2 = [ML2]/[ML][L]
ML2 + L → ML3, K3 = [ML3]/[ML2][L]
ML3 + L → ML4, K4 = [ML4]/[ML3][L]
Where K1, K2, K3 and K4 are referred to as stepwise stability constants.
 The overall stability constant, β = K1 × K2 × K3 × K4
 The stability of a complex depends on the following factors:
o The nature of the metal ion: Greater the charge to radius of the metal ion,
greater will be the stability of the complex.
o Nature of the ligand: The greater the basic strength of the ligand, the greater will
be the stability of the complex.
o Presence of chelating ligands: increases the stability of the complex.

Application of Co-ordination Complexes

 In Qualitative & Quantitative Analysis: Co-ordination compounds find use in many

qualitative and quantitative chemical analyses.
For e.g. Ni2+ is detected and estimated by the formation of a complex with Dimethyl
Glyoxime (DMG). The brown ring test for the detection of nitrate ion is due to the
formation of the brown complex [Fe(H2O)5NO]2+. The Ca2+ and Mg2+ ions are
estimated by the formation of stable complexes with EDTA.
 In water treatment: The Hardness of water is estimated by simple titration with
Na2EDTA (sodium salt of EDTA). The Ca2+ and Mg2+ ions form stable complexes with
EDTA.
The hardness of water can be removed by the formation of a complex with calgon
(Sodium polymetaphosphate)
 In Metallurgy: Metals like silver and gold are extracted by the formation of complexes
with CN- ligands.
Gold forms the complex [Au(CN)2] - and silver forms [Ag(CN)2] - which are separated
with Zn.
Similarly, coordination compounds also find application in the refining of some
metals. For example, impure nickel is converted to [Ni(CO)4], which is decomposed
to yield pure nickel
 Biological Applications: Coordination compounds are of great importance in biological
systems.
Chlorophyll is a coordination compound of magnesium,
Haemoglobin, is a co-ordination compound of iron
Vitamin B12 (cyanocobalamine) is a co-ordination compound of cobalt.
 In Catalysis: Co-ordination compounds are used as catalysts for many industrial
processes.
For e.g. Tris(triphenylphosphine)rhodiumchloride, [(Ph3P)3RhCl] (Wilkinson catalyst),
is used for the hydrogenation of alkenes.
 In electroplating: Articles can be electroplated with silver and gold by using the solutions
of the complexes, [Ag(CN)2] - and [Au(CN)2] - respectively as electrolytes.
 In Photography: In black and white photography, the developed film is fixed by washing
with hypo solution which dissolves the undecomposed AgBr to form a complex ion,
[Ag(S2O3)2]3– .
 In medicine: Cis-platin is used for the treatment of cancer.
Excess of copper and iron in animal or plant body are removed by the chelating
ligands D–penicillamine and desferrioxime B through the formation of co-ordination
compounds.
EDTA is used in the treatment of lead poisoning.