1. COORDINATE COMPOUNDS

2. COORDINATE COMPOUNDS
Complex ion. An electrically charged species in which a metal atom or a simple cation is
coordinated to two or more anions or neutral molecules and which retains its identity in
solution is called a complex ion. For example, [Fe(CN)6]4– (ferrocyanide) is a complex ion.
A complex ion may have positive or negative charge on it. A positively charged complex ion is
called cationic complex ion, while the negatively charged complex ion is called anionic
complex ion.
Cationic complex ions: [Cu(NH3)4]2+, [Ag(NH3)2]+, [Ni(NH3)6]2+ etc.
Anionic complex ions: [Ag(CN)2]–, [Fe(CN)6]3–, [Fe(CN)6]4–, [CuCl4]2– etc.
Coordination compound (or complex compound). A compound in which a metal atom or ion
is coordinated to two or more anions or neutral molecules and which retains its identity in the
solid as well as in solution is called a complex compound, or coordination compound. A coordi-
nation compound (or complex compound) may be molecular (non-ionic) or ionic.
Central atom or ion. The metal atoms or ion in a complex to which two or more anions or
neutral molecules are attached is called central atom or central ion.
Ligands. The molecular or ionic species which gets attached directly to the central metal atom
or ion during the formation of a complex ion is called a ligand. The ligands are attached to the
central atom, or ion through coordinate bonds.
Ligands can attach themselves to the central atom/ion at one or more than one points. Accord-
ingly, the ligands are termed as monodentate (one point of attachment), bidentate (two points of
attachment), tridentate (three points of attachment) and so on.
Coordination number. The total number of ligands attached directly to the central atom or ion

3. in a complex is called the coordination number of that metal atom or ion. For example, in the
complex ion [Fe(CN)6]4– (ferrocyanide ion), the coordination number of Fe2+ is six.
A bidentate ligand attaches itself to the central atom/ion through two points. That is, a bidentate
ligand occupies two coordinating positions around the central atom/ion. Therefore, while calcu-
lating the coordination number of any central atom/ion, each bidentate ligand contributes two to
the coordination number. Similarly, a tridentate ligand contributes three to the coordination
number, and so on.
Coordination sphere. The molecular or ionic species attached directly to the central ion
constitute the coordination sphere of the central atom/ion. The central atom/ion with its
coordination sphere is written inside a square bracket.
The bonding between the central metal ion/atom, and the ligands in its coordination sphere is
non-ionisable.
Charge number of a complex ion. The net electrical charge on a complex ion is called its
charge number (or commonly as valency).
Coordination sphere
Charge number
[Co Cl2 (NH3)4]+ Cl–
Ionisable anion
Central ion
(Co3+)
Ligands Coordination number (4 + 2 = 6)
–
(NH3 molecules and Cl ions)
Double salt. The molecular addition compounds which exist in the solid state, but dissociate
into their constituent ions when dissolved in water (or any other solvent) are called double salts.
The solution of a double salt gives tests of its constituent ions.

4. Example: Mohr’s salt, (FeSO4.(NH4)2SO4.6H2O). This compound when dissolved in water gives
Fe2+, NH4+ and SO42– ions in the solution, which gives the tests of all these ions.
Chelate and chelating agent. A polydentate ligand which gets attached to the central metal ion
through two or more donor atoms forming one or more ring around it, is called a chelating agent
or chelating ligand. The complex ion/compound having a ring structure around its central ion is
called a chelate: (Chelate is a Greek word meaning Claw).
A chelate formed due to the reaction of ethylenediamine with Cu2+ is,
H2C – H2N: :NH2 – CH2 2+
Cu or [Cu(en)2]2+
H2C – H2N: :NH2 – CH2
Modern notation for representing coordination compounds
According to the latest IUPAC (1989) nomenclature rules, the formula of a coordination
compound or, a complex ion is written as follows:
• Write down the symbol of the metal forming the central ion/atom.
• Write the symbols/formulae of the ligands on the right of the symbol of the metal in the order:
negative ligands, neutral ligands, and then positive ligands.
In each group of ligands (negative, neutral or positive), the names of the ligands are written
alphabetically according to the first letter of the symbol/formula. These are written as a single
word (with no gap in between).
• The whole complex ion (metal and all the ligands) is enclosed in a square bracket ([ ]).
• The net electrical charge on the complex ion is shown as superscript on the square bracket on the
right hand side, e.g., [ ]z+ or [ ]z–.
• The ionisable groups are written outside the square bracket on the right.

5. Thus, a coordination compound may be represented as,
[MLn]z+ Xz–
where, M is the central metal atom or ion,
L is the ligand,
X – is the ionisable group,
z+ is the net electrical charge on the complex ion.
Ionisation isomerism. The coordination compounds which have the same molecular formula
but give different ions in the solution are called ionisation isomers. This type of isomerism
occurs due to the interchange of groups or ions between the coordination sphere of the metal ion
and the ions outside this sphere. For example, there are two distinct compounds of the formula,
CoBr(NH3)5SO4, viz.,
[CoBr(NH3)5]2+SO42– and [CoSO4(NH3)5]+Br–
pentaamminebromocobalt(III) sulphate pentaamminesulphatocobalt(III) bromide
(red-violet) (red)
Coordination isomerism. This type of isomerism is shown by the compounds which contain
complex cation, and a complex anion. Coordination isomerism is caused by the interchange of
the ligands between the complex cation and complex anion. Some common examples of the
coordination isomerism are given below.
[Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6]
hexaamminecobalt(III) hexaamminechromium(III)
hexacyanochromate(III) hexacyanocobaltate(III)
Linkage isomerism. The compounds which have the same molecular formula, but differ in the
mode of attachment of a ligand to the central atom/ion are called linkage isomerism. Linkage
isomerism is shown by the coordination compounds which contain ligands that are able to
attach themselves to the central ion through more than one donor atoms, (ambidentate ligands).

6. [CoONO(NH3)5]Cl2 and [CoNO2(NH3)5]Cl2
pentaamminenitritocobalt(III) chloride pentaamminenitrocobalt(III) chloride
(red), (decomposed by acids) (yellow-brown), (stable towards acids)
Geometrical isomerism. The isomerism due to different geometrical positions of the ligands
around the central metal ion is called geometrical isomerism.
When the similar ligands occupy the adjacent positions. This arrangement gives rise to
cis-isomer.
When the similar ligands occupy positions opposite to each other. This arrangement gives rise
to trans-isomer.
The geometrical isomerism is also called cis-trans isomerism.
In tetrahedral geometry, all four ligands are symmetrically placed relative to each other.
Therefore, cis-trans isomerism is not possible in tetrahedral coordination compounds.
The coordination compounds of the type Ma2b2, Mabcd, Ma2bc, M(ab)2 having square planar
geometry show geometrical (cis-trans) isomerism.
Square planar complexes of the types MA4, MA3X and MAX3 do not show geometrical isomerism
because in these cases all spatial arrangements of the ligands relative to each 1
other are equivalent.
6 2
In an octahedral geometry, the positions, 1-4, 2-5, 3-6 are trans positions,
M
while the positions 1-2, 1-3, 2-3, 3-4, 4-5 are cis relative to each other.
5 3
The octahedral complexes which exhibit geometrical isomerism are of the
type, Ma4b2, Ma2b4, Ma3b3, Ma4bc. 4
Optical isomerism. The compounds which have the same molecular formula but rotate the
plane of polarised light in the opposite directions are called optical isomers. The isomer which
rotates the plane of polarised light to the right is called dextro rotatory (+), and the isomer which
rotates the plane of the polarised light on the left is called laevo rotatory (–).

7. The structures of the two optical isomers of a compound are the mirror images of each other.
These mirror images do not superimpose on each other. Thus, the structures of the two optical
isomers are not superimposable on each other. This would be possible only if the molecule is
asymmetric. Thus, the essential condition for a compound to show optical activity is that the
molecule of the compound should not have a plane of symmetry.
The most well-known cases of optical isomerism occur among octahedral complexes of the
type [M(aa) 3], where ‘aa’ is a bidentate ligand, e.g., trioxalatochromate(III) ion, and
tris(ethylenediamine) cobalt(III) ion show optical isomerism.
Inner orbital and outer orbital octahedral complexes.
Type of Orbitals involved in Designation of the complexes Magnetic
hybridisation hybridisation moment
d2sp3 (n – 1) d2, ns, np3 Inner orbital complexes: Low spin Low
sp3d2 ns, np3, nd2 Outer orbital complexes: High spin High
Crystal field splitting in octahedral complexes. In an
octahedral complex, the metal ion is at the centre of the dx2 – y 2 dz2
eg
octahedron and the six ligands are at its six corners. The
octahedral field of the ligands disturbs the degeneracy
of the d-orbitals and the d-orbitals split into two groups ∆o
of different energies. The difference in energy between
the groups of d-levels is denoted by ∆o or 10 ∆q. Five d-orbitals
t2g
• Crystal field splitting. The energy gap ∆o between d xy d yz d zx
the two groups of d-orbitals is called crystal field
splitting.
• Weak and strong field ligands. The ligands which cause only a small crystal field splitting are
called weak field ligands. The ligands which cause a large crystal field splitting are called strong
field ligands.

8. Spectrochemical series. The arrangement of common ligands in the ascending order of crystal
field splitting (∆o) is called spectrochemical series.
Weak field ligands
I– < Br– < S2– < Cl– < NO3– < F– < OH– < EtOH < C2O42– < H2O < EDTA < NH3,
Py < Ethylenediamine < dipyridyl < o-phenanthroline < NO2– < CN– < CO
Strong field ligands
The weak ligands produce high-spin complexes. [Fe(H2O)6] 2+ and [CoF ]3– are high spin
6
complexes.
The strong field ligands produce low-spin complexes. [Fe(CN)6]4– and [Co(NH3)6]3+ are low
spin complexes.
Crystal field splitting in tetrahedral complexes. In a tetrahedral geometry, the central metal
ion is surrounded by four ligands along the four corners of a tetrahedron. The tetrahedral field of
the four ligands, splits the d-orbitals into two groups dxy dyz dzx
(t2g and eg) of different energies. t2g
The tetrahedral crystal field splitting (∆t) is nearly half
of the octahedral crystal field splitting (∆o). Even in
∆t
the case of strong field ligands, the tetrahedral crystal
field splitting is not very large. As a result, electron
pairing is not energetically favourable and all tetrahe- eg
dral complexes are high-spin. dx 2 – y 2 dz 2
Stability of a complex ion. The stability of a complex can be expressed in terms of the
equilibrium constant for the reaction leading to the formation of the complex ion from the
metal ion and the ligands.
Such an equilibrium constant is called the stability constant of the complex.
The stability constant for a complex [MLn]p formed by the reaction,

9. Mz+ + n Lz– l [MLn]p
where, p = z+ + n × z–
is given by
[(MLn ) p ]
K=
[M z + ][Lz − ]n
The stability of a complex depends upon the following factors.
(a) Nature of the central ion – High charge, small size favours the stability of the complex.
(b) Nature of the ligand – More basic ligand forms more stable complex.