COORDINATION COMPOUNDS

A coordination complex consists of a central atom or ion, which is

usually metallic and is called the coordination centre, and a surrounding array

of bound molecules or ions, that are in turn known as ligands or complexing

agents. Many metal-containing compounds, especially those of transition metals,

are coordination complexes. A coordination complex whose centre is a metal atom

is called a metal complex of d block element.

The complex ion is generally written in a square box and the ion (cation or

anion) is written outside complex ion.

eg : [Co (NH3)6] Cl3

[Complex ion] anion

eg : K4 [Fe (CN)6]

cation [Complex ion]

General formula : Ax [MLn]/[MLn]By

where : M is the central metal atom/ion

L is the ligand A is the cation B is the anion.

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WERNER’S THEORY

Alfred Werner in 1898 proposed Werner’s theory explaining the structure of

coordination compounds.

Werner’s Experiment: By mixing AgNO3 (silver nitrate) with CoCl3·6NH3, all

three chloride ions got converted to AgCl (silver chloride).

CoCl3·6NH3 + 3AgNO3 → Co.6NH3 +3AgCl

However, when AgNO3 was mixed with CoCl3·5NH3, two moles of AgCl were

formed.

CoCl3·5NH3 + 2AgNO3 → CoCl·5NH3 + 2AgCl

Further, on mixing CoCl3·4NH3 with AgNO3, one mole of AgCl was formed.

Based on this observation, the following Werner’s theory was postulated:

POSTULATES OF WERNER’S THEORY

The central metal atom in the coordination compound exhibits two types of

valency, namely, primary and secondary linkages or valencies.

Primary linkages are ionizable and are satisfied by the negative ions.

Secondary linkages are non-ionizable. These are satisfied by negative ions.

Also, the secondary valence is fixed for any metal and is equal to its

coordination number.

The ions bounded by the secondary linkages to the metal exhibit

characteristic spatial arrangements corresponding to different coordination

numbers.

DOUBLE SALT AND COORDINATION COMPOUND

Both double salts as well as complexes are formed by the combination

of two or more stable compounds in stoichiometric ratio. However,

they differ in the fact that double salts such as carnallite,

KC1.MgCl2• 6H2O.Mohr’s salt, FeSO4.(NH4)2SO46H2O, potash alum,

KAl(SO4)2.12H2O, etc. dissociate into simple ions completely when

dissolved in water.

KCl.MgCl2• 6H2O → k++ 3Cl-+Mg2+

However, complex ions such as [Fe(CN)6l4+ of M [Fe(CN)6] do not

dissociate into Fe2+ and CN- ions

[Fe (CN)6] → NO REACTION

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IMPORTANT TERMS OF COORDINATION COMPOUNDS

COORDINATION ENTITY

A chemical compound in which the central ion or atom (or the coordination centre)

is bound to a set number of atoms, molecules, or ions is called a coordination

entity.

Examples: [CoCl3(NH3)3], and [Fe(CN)6]4-.

CENTRAL ATOMS AND CENTRAL IONS

As discussed earlier, the atoms and ions to which a set number of atoms,

molecules, or ions are bound are referred to as the central atoms and the central

ions.

In coordination compounds, the central atoms or ions are typically Lewis

Acids and can, therefore, act as electron-pair acceptors

Eg: [CoCl3(NH3)3], and [Fe(CN)6]4-

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LIGAND

The ions or molecules bound to the central atom/ion in the coordination entity are

called ligands. These may be simple ions such as Cl–, small molecules such as H2O

or NH3 , larger molecules such as H2NCH2CH2NH2.

When a ligand is bound to a metal ion through a single donor atom, as

with Cl, H2O or NH3, the ligand is said to be UNIDENTATE.

When a ligand can bind through two donor atoms as in

H2NCH2CH2NH2 (ethane-1,2-diamine) or C2O42-(oxalate), the ligand is

said to be DIDENTATE

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POLYDENTATE

when several donor atoms are present in a single ligand is said to be polydentate.

CHELATE LIGAND

Chelation is a type of bonding of ions and molecules to metal ions. It involves the

formation or presence of two or more separate coordinate bonds between

a polydentate (multiple bonded) ligand and a single central atom.

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AMBIDENTATE LIGAND

Ambidentate ligand is a type of ligands which have the ability to bind to the central

atom via the atoms of two different elements.

COORDINATION NUMBER

It is important to note here that coordination number of the central

atom/ion is determined only by the number of sigma bonds formed by

the ligand with the central atom/ion. Pi bonds, if formed between the

ligand and the central atom/ion, are not counted for this purpose.

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COORDINATION SPHERE

The central atom/ion and the ligands attached to it are enclosed in

square bracket and is collectively termed as the coordination

sphere. The ionisable groups are written outside the bracket and

are called counter ions.

For example, in the complex K4[Fe(CN)6] . the coordination sphere is

[Fe(CN)6]4- and the counter ion is K+.

COORDINATION POLYHEDRON

The spatial arrangement of the ligand atoms which are directly attached to the central atom /ion defines a coordination polyhedron about the central atom. The most common coordination polyhedral are octahedral, square planar and tetrahedral. For example,

[Co(NH3)6] is octahedral, [Ni(CO)4] is tetrahedral and [PtCl4] ° is

square planar.

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OXIDATION NUMBER OF CENTRAL ATOM

The oxidation number of the central atom in a complex is defined as

the charge it would carry if all the ligands are removed along with

the electron pairs that are shared with the central atom. The oxidation

number is represented by a Roman numeral in parenthesis following

the name of the coordination entity.

For example, oxidation number of cobalt in [Co(H2O)(CN)(en)2]2+ is +3

and it is written as Co(III).

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VALENCE BOND THEORY

According to this 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 hybridization to yield a set of equivalent orbitals of definite geometry such as octahedral tetrahedral, square planar and so on.

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

It is usually possible to predict the geometry of a complex from

the knowledge of its magnetic behavior on the basis of the VBT

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FORMATION OF [CoF6]3-

FORMATION OF [Fe(CN)6]3-

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FORMATION OF [Ni(CN)4]2-

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SPECTROCHEMICAL SERIES

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CRYSTAL FIELD THEORY

Crystal field theory (CFT) describes the breaking of degeneracies of d orbitals,

due to a static electric field produced by a surrounding charge distribution (anion

neighbors). This theory has been used to describe various spectroscopies

of transition metal coordination complexes, in particular optical spectra (colors).

CFT successfully accounts for some magnetic properties,

colors, hydration enthalpies, and spinel structures of transition metal complexes,

but it does not attempt to describe bonding.

According to crystal field theory, the interaction between a transition metal

and ligands arises from the attraction between the positively charged metal cation

and the negative charge on the non-bonding electrons of the ligand. The theory is

developed by considering energy changes of the five degenerate d-orbitals upon

being surrounded by an array of point charges consisting of the ligands. As a

ligand approaches the metal ion, the electrons from the ligand will be closer to

some of the d-orbitals and farther away from others, causing a loss of degeneracy.

The electrons in the d-orbitals and those in the ligand repel each other due to

repulsion between like charges. Thus the d-electrons closer to the ligands will have

a higher energy than those further away which results in the d-orbitals splitting in

energy. This splitting is affected by the following factors:

The nature of the metal ion.

The metal's oxidation state. A higher oxidation state leads to a larger splitting

relative to the spherical field.

The arrangement of the ligands around the metal ion.

The coordination number of the metal (i.e. tetrahedral, octahedral...)

The nature of the ligands surrounding the metal ion. The stronger the effect of

the ligands then the greater the difference between the high and low

energy d groups.

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CRYSTAL FIELD SPLITTING IN OCTAHEDRAL COMPLEX

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ELECTRONIC CONFIGURATION

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COLOR OF COORDINATION COMPLEXES

The variety of color among transition metal complexes has long fascinated the

chemists. For example, aqueous solutions of [Fe(H2O)6]3+ are red, [Co(H2O)6]

2+ are

pink, [Ni(H2O)6]2+ are green, [Cu(H2O)6]

2+ are blue and [Zn(H2O)6]2+ are colorless.

Although the octahedral [Co(H2O)6]2+ are pink, those of tetrahedral [CoCl4]

2- are

blue. The green color of [Ni(H2O)6]2+ turns blue when ammonia is added to give

[Ni(NH3)6]2+. Many of these facts can be rationalized from CFT.

Why we see Color?

When a sample absorbs light, what we see is the sum of the remaining colors that

strikes our eyes. If a sample absorbs all wavelength of visible light, none reaches

our eyes from that sample, and then the sample appears black. If the sample

absorbs no visible light, it is white or colorless. When the sample absorbs a

photon of visible light, it is its complementary color we actually see. For

example, if the sample absorbed orange color, it would appear blue; blue and

orange are said to be complementary colors.

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THE COLOR WHEEL

The visible part of the electromagnetic spectrum contains light of wavelength 380-750 nm. The color wheel

below gives information on the wavelength of different color and also the complementary color.

COLOR OF COORDINATION COMPLEXES

The color of coordination complexes arises from electronic transitions between

levels whose spacing corresponds to the wavelengths available in the visible

light. In complexes, these transitions are frequently referred to as d-d transitions

because they involve the orbitals that are mainly d in character (for examples: t2g

and eg for the octahedral complexes and eg and t2g for the tetrahedral complexes).

Obviously, the colors exhibited are intimately related to the magnitude of the

spacing between these levels. Since this spacing depends on factors such as the

geometry of the complex, the nature of the ligands and the oxidation state of the

central metal atom, variation on colors can often be explained by looking

carefully at the complexes concerned.

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