Coordination Compounds

Werner's Theory of Coordination Compounds & Definition of Important Terms in Coordination

Compounds

• Coordination Compounds − Complex compounds in which transition metal atoms are bound to a
number of anions or neutral molecules

Postulates of Werner’s Theory of Coordination Compounds

• In coordination compounds, there are two types of linkages (valences) − primary and secondary.

• The primary valences are ionisable, and are satisfied by negative ions.

• The secondary valences are non-ionisable, and are satisfied by negative ions or neutral molecules.
The secondary valence is equal to the coordination number of a metal, and remains fixed for a
metal.

• Different coordination numbers have characteristic spatial arrangement of ions or groups bound
by the secondary linkages.

Such spatial arrangements are called coordination polyhedra. The species within the square
brackets are called coordination entities or complexes, and the ions outside the square brackets are
called counter ions.

Difference between a Double Salt and a Complex

• In water, a double salt dissociates completely to give simpler ions. Examples of double salt:
carnallite (KCl.MgCl2.6H2O), Mohr’s salt [FeSO4.(NH4)2SO4.6H2O]

• Complex ions do not dissociate further to give simpler ions; for example, [Fe(CN)6]4−, [Fe(C2O4)3]3−.

Definition of Important Terms in Coordination Compounds

• Coordination Entity

• Constitutes a central metal atom or ion bonded to a fixed number of ions or molecules

• Example: [CoCl3(NH3)3] is a coordination entity

• Central Atom or Ion

• The atom or ion to which a fixed number of ions/groups are bound in a definite geometrical
arrangement around it in a coordination entity
• Example: Ni2+ in [NiCl2(H2O)4] and Fe3+ in [Fe(CN)6]3−

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 • Ligands

• Ions or molecules bound to the central metal atom or ion in a coordination entity
• Unidentate- A ligand bound to the central metal atom through one donor atom only.
Example: Cl−Cl-, NH3.

• Didentate- A ligand bound to central metal atom/ion through two donor atoms. Example: ethylene
diamine, oxalate ion etc.

• Polydentate- A ligand bound to central metal atom/ion through multiple donor atoms.
Example: ethylenediamnietetraacetate ion.

• Ambidentate- A ligand which can bind through two different atoms. For example thiocyanate
ion.

• Coordination Number

• Number of ligand-donor atoms bonded directly to the metal

• Example: The coordination number of Pt and Ni in [PtCl6]2− and [Ni(NH3)4]2+ are 6 and 4
respectively.

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 • Coordination Sphere

• The central atom or ion and the ligands attached to it are enclosed in square brackets, which
are collectively known as the coordination sphere.

• Example: In the complex K4[Fe(CN)6], the coordination sphere is [Fe(CN)6]4−.

• Coordination Polyhedron

• The spatial arrangement of the ligand atoms which are directly attached to the central atom
or ion

• Example: Octahedral, square planar, tetrahedral

• Oxidation Number of Central Atom

• The charge central metal atom would carry if all the ligands are removed along with the
electron pairs that are shared with the central atom

• Example: Oxidation number of copper in [Cu(CN)4]3− is 1.

• Homoleptic and Heteroleptic Complexes

• Homoleptic − Complexes in which the metal is bound to only one kind of donor group.
Example: [Co(NH3)6]3+

• Heteroleptic − Complexes in which the metal is bound to more than one kind of donor groups.
For example: [Co(NH3)4Cl2]+

Nomenclature of Coordination Compounds

Formulas of Mononuclear Coordination Compounds

The following are the rules for writing the formulas.

• The central atom is listed first.

• The ligands are then listed in the alphabetical order.

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 • Polydentate ligands are also listed in the alphabetical order. In the case of abbreviated ligands, the
first letter of the abbreviation is used for determining the position of the ligands in the alphabetical
order.

• The formula of the entire coordination entity is enclosed in square brackets. Ligand abbreviations
and formulas for polyatomic ligands are enclosed in parentheses.

• There should be no space between the ligands and the metal within a coordination sphere.

• For the charged coordination entity, the charge is indicated outside the square brackets, as a right
superscript, with the number before the sign. Example: [Fe(CN)6]3−, [Cr(H2O)6]3+, etc.

• The charge of the cation(s) is balanced by the charge of the anion(s).

Naming of Mononuclear Coordination Compounds

The following rules are followed while naming coordination compounds.

• The cation is named first in both positively and negatively charged coordination entities.

• The ligands are named in alphabetical order before the name of the central atom/ion.

• Names of the anionic ligands end in −o and those of neutral and cation ligands are the same.

[Exceptions: aqua (H2O), ammine (NH3), carbonyl (CO), nitrosyl (NO)]

• To indicate the number of the individual ligands, the prefixes mono−, di−, tri−, etc., are used. If
these prefixes are present in the names of ligands, then the terms −bis, −tris, −tetrakis, etc., are
used. For example, [NiCl2(PPh3)2] is named as dichlorobis(triphenylphosphine)nickel(II).

• Oxidation state of the metal is indicated by a Roman numeral in parentheses.

• If the complex ion is cation, then the metal is named as the element. For example, Fe in a complex
cation is called iron and Pt is called platinum.

• If the complex ion is anion, then the metal is named with ‘−ate’ ending. For example, Co in a
complex anion, [Co(SCN)4]2− is called cobaltate.

• The neutral complex molecule is named as the complex cation.

IUPAC name of some coordination compounds are as follows:

[Ni(CO)4] − Tetracarbonlynickel(0)

[Co(NH3)3(H2O)3]Cl3− Triamminetriaquacobalt(III) chloride

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 [Pt(NH3)2Cl(NO2)] − Diamminechloridonitrito-N-platinum(II)

[CoCl2(en)2]Cl − Dichloridobis(ethane-1,2-diammine)cobalt(III) chloride

K2[Zn(OH)4] − Potassium tetrahydroxozincate(II)

Isomerism in Coordination Compounds

• Stereoisomers have same chemical formula and chemical bonds, but have different spatial
arrangement.

• Structural isomers have different chemical bonds.

Geometrical Isomerism

• Arises in heteroleptic complexes due to different possible geometrical arrangement of ligands

• Generally found with compounds having coordination numbers 4 and 6

• Square planar complex of formula [MX2L2] (X and L are uni-dentate) exhibits geometrical
isomerism − cis-isomer and trans-isomer.

• Example − Geometrical isomers of Pt(NH3)2Cl2

• Octahedral complex of formula [MX2L4] also exhibits geometrical isomerism.

• Example − Geometrical isomers of [Co(NH3)4)Cl2]+

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 • Octahedral complex of formula [MX2 (L − L)2] exhibits geometrical isomerism, where L−L is a
bidentate ligand. [e.g., NH2 CH2 CH2 NH2(en)]

• Example − Geometrical isomers of [CoCl2(en)2]+

• Octahedral complexes of type [Ma3b3] exhibit another type of geometrical isomerism − facial (fac)
isomer and meridional (mer) isomer.

• Example − [Co(NH3)3(NO2)3]

• Facial (fac) isomer − Three donor atoms of the same ligands occupy adjacent positions at the
corners of an octahedral face.

• Meridional (mer) isomer − When the positions of the donor atoms are around the meridian of the
octahedron

Optical Isomerism

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 • Optical isomers (also called enantiomers) − Mirror images that cannot be superimposed on one
another

• Molecules or ions that cannot be superimposed are called chiral.

• Two forms − dextro (d) and laevo (l)

• d-isomer rotates the plane of polarised light to the right and l-isomer rotates it to the left.

• Example − ‘d’ and ‘l’ isomer of [Co(en)3]3+

• → Only cis isomer shows optical activity.

Linkage Isomerism

• Arises in the compound containing ambidentate ligand (which can ligate through two different
atoms)

• Example: Thiocyanate ligand (NCS−) can bind through nitrogen to give M−NCS or through sulphur
to give M−SCN.

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 • also exhibits two forms:

• Red form − Nitrite ligand binds through oxygen (−ONO).

• Yellow form − Nitrite ligand binds through nitrogen (−NO2).

Coordination Isomerism

• Arises from the interchange of ligands between cationic and anionic entities of different metal ions
present in the complex

• Example −

Ionisation Isomerism

• Arises when the counter ion in the complex salt is itself a potential ligand and can displace a ligand,
which can then become the counter ion

• Example −

Solvate Isomerism

• Known as ‘hydrate isomerism’ when water is involved as a solvent

• Similar to ionisation isomerism

• Solvate isomers differ by whether or not a solvent molecule is directly bonded to the metal ion or
merely present as free solvent molecules in the crystal lattice.

• Example − (violet) and (grey green)

Valence Bond Theory and its Limitations & Magnetic Properties of Coordination Compounds

Valence Bond Theory

• The metal atom or ion under the influence of ligands can use its (n−1)d, ns, np or ns, np, nd orbitals
for hybridisation, to yield a set of equivalent orbitals of definite geometry such as octahedral,
tetrahedral, square planar, and so on.

Coordination number Type of hybridisation Distribution of hybrid

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 orbitals in space

4 sp3 Tetrahedral

4 dsp2 Square planar

5 sp3d Trigonal bipyramidal

6 sp3d2 Octahedral

6 d2sp3 Octahedral

• These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs
for bonding.

• Octahedral Complexes

• The hybridisation involved can be d2sp3 or sp3d2.

• Inner-orbital or low-spin or spin-paired complexes: Complexes that use inner d-orbitals in

hybridisation; for example, [Co(NH3)6]3+. The hybridisation scheme is shown in the following
diagram.

• Outer-orbital or high-spin or spin-free complexes: Complexes that use outer d-orbitals in

hybridisation; for example, [CoF6]3−. The hybridisation scheme is shown in the following diagram.

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 • Tetrahedral Complexes

• The hybridisation involved is sp3

• Example: [NiCl4]2−

• The hybridisation scheme is shown in the following diagram.

• Square planar Complexes

• The hybridisation involved is dsp2

• Example: [Ni(CN)4]2−

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 • The hybridisation scheme is shown in the following diagram.

Magnetic Properties of Coordination Compounds

• Complexes with unpaired electron(s) in the orbitals are paramagnetic.

• Complexes with no unpaired electron(s) in the orbitals (i.e., all the electrons are paired) are
diamagnetic.

• Example: [MnCl6]3−, [FeF6]3− and [CoF6]3− are paramagnetic. It can be explained on the basis of the
valence bond theory. These coordination compounds are outer-orbital complexes
with sp3d2 hybridisation, and are paramagnetic containing four, five and four unpaired electrons
respectively.

On the other hand, [Co(C2O4)3]3−is diamagnetic. According to the valence bond theory, it is an
inner-orbital complex involving d2sp3hybridisation, with no unpaired electron(s), and is
diamagnetic.

Limitations of Valence Bond Theory

• A number of assumptions are involved.

• Quantitative interpretation of magnetic data is not given.

• The exhibition of colour by coordination compounds is not explained.

• The thermodynamic or kinetic stabilities of coordination compounds are not quantitatively

interpreted.

• Whether a complex of coordination number 4 is tetrahedral or square planar cannot be exactly

predicted.

• Weak and strong ligands cannot be distinguished.

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 Crystal-Field Theory

• An electrostatic model which considers the metal−ligand bond to be ionic, which arises purely from
the electrostatic interaction between the metal ion and the ligand

• Ligands are treated as point charges in the case of anions, or dipoles in the case of neutral
molecules.

• The five d-orbitals in an isolated gaseous metal atom/ion are degenerate (i.e., have the same
energy).

• Due to the negative fields of the ligands (either anions or the negative ends of dipolar molecules),
the degeneracy of the d-orbitals is lifted, resulting in the splitting of the d-orbitals.

Crystal-Field Splitting in Octahedral Coordination Entities

Crystal-field splitting is the splitting of the degenerate energy levels due to the presence of ligands.

• The splitting of d-orbitals in an octahedral crystal-field is shown in the given figure.

• The and orbitals (i.e., eg set), which point towards the axes along the direction of the
ligand, will experience more repulsion, and will be raised in energy.

• The dxy, dyz and dxz orbitals (i.e., t2g set), which are directed in between the axes, will be lowered in
energy relative to the average energy in the spherical crystal-field.

• The energy separation is denoted by Δo (the subscript ‘o’ is for octahedral)

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 • The energy of two eg orbitals will increase by (3/5)Δo, and that of three t2g orbitals will decrease by
(2/5)Δo.

• Spectrochemical series

Ligands are arranged in a series in the increasing order of the field strength as follows:

• Ligands for which Δo (crystal-field splitting) < P (pairing energy), are called weak-field ligands, and
form high-spin complexes.

• Ligands for which Δo (crystal-field splitting) > P (pairing energy), are called strong-field ligands,
and form low-spin complexes.

Crystal-Field Splitting in Tetrahedral Coordination Entities

• The splitting of d-orbitals in a tetrahedral crystal field is shown in the given figure.

• Δt = (4/9) Δo

• The orbital-splitting energies are not sufficiently large for forcing pairing, and therefore, low-spin
configurations are rarely observed.

Colour in Coordination Compounds

• The colour of the coordination compounds is attributed to d−d transition of electrons.

• For example, the complex [Ti(H2O)6]3+ is violet in colour, which is due to d−d transition (shown in
the figure)

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 • The energy of yellow-green region is absorbed by the complex, and then, the electron gets excited
from t2g level to the eg level, i.e.,

• The relationships between the wavelength of light absorbed and the colour observed in some
coordination entities are listed in the given table.

Coordination Wavelength of light absorbed Colour of light Colour of coordination

entity (nm) absorbed entity

[CoCl(NH3)5]2+ 535 Yellow Violet

[Co(NH3)5(H2O)]3+ 500 Blue-green Red

[Co(NH3)6]3+ 475 Blue Yellow-orange

[Co(CN)6]3− 310 Ultraviolet Pale yellow

[Cu(H2O)4]2+ 600 Red Blue

[Ti(H2O)6]3+ 498 Blue-green Purple

• In the absence of ligand, crystal-field splitting does not occur; hence, the substance is colourless.
For example, the removal of water from [Ti(H2O)6]Cl3 on heating renders it colourless. However,
the complex [Ti(H2O)6]3+ is violet in colour.

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 Limitations of Crystal-Field Theory

• Crystal-field theory assumes that ligands are point charges. Hence, the anionic ligands should exert
the greatest splitting effect. However, the anionic ligands are found at the lower end of the
spectrochemical series.

• Crystal-field theory does not take into account the covalent character of bonding between ligand
and the central metal atom.

Bonding in Metal Carbonyls, Stability of Coordination Compounds & Applications of Coordination

Compounds

Bonding in Metal Carbonyls

• Homoleptic carbonyls i.e. the compounds containing carbonyl ligands only are formed by most of
the transition metals.

• Structures of some representative homoleptic metal carbonyls are given below.

• The metal−carbon bonds in metal carbonyls have both σ and π characters.

• A σ bond is formed when the carbonyl carbon donates a lone pair of electrons to the vacant orbital
of the metal.

• A π bond is formed by the donation of a pair of electrons from the filled metal d- orbital to the
vacant anti-bonding π* orbital (also known as back bonding of the carbonyl group).

• Thus, a synergic effect is created due to this metal-to-ligand bonding. This synergic effect
strengthens the bond between CO and the metal.

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 Stability of Coordination Compounds

• The stability of a complex in a solution refers to the degree of association between the two species
involved in the state of equilibrium.

• Stability can be expressed quantitatively in terms of stability constant or formation constant.

• For this reaction, the greater the value of the stability constant, the greater is the proportion of
ML4 in the solution.

• Free metal ions are usually surrounded by solvent molecules in solution, which will compete with
the ligands molecules (L) and be successively replaced by them as follows:

Where, K1, K2, …= Stepwise stability constants

• Alternatively,

Where, β1, β2, …. = Overall stability constant

The stepwise and overall stability constants are therefore related as follows:

or more generally,

• The reciprocal of the formation constant is called instability constant or dissociation constant.

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 Applications of Coordination Compounds

• In many qualitative and quantitative chemical analyses: Colour reactions are given by metal ions
with a number of ligands (especially chelating ligands), which help in detection and estimation.

• Hardness of water can be estimated by simple titration with Na2EDTA, which forms stable
complexes with Ca2+ and Mg2+ ions.

• Some important extraction processes of metals such as gold and silver make use of complex
formation.

• Coordination compounds are used as catalysts. For example, Wilkinson catalyst, [(Ph3P)3RhCl], is
used for the hydrogenation of alkenes.

• Chelate therapy is used in medicinal chemistry. For example, EDTA is used in the treatment of lead
poisoning. Some coordination compounds such as cis-platin and related compounds inhibit the
growth of tumours.

• Coordination compounds are of great importance in biological systems. For example, chlorophyll is
a coordination compound of magnesium, haemoglobin is a coordination compound of iron and
vitamin B12 is a coordination compound of cobalt.

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