This document provides an introduction to coordination compounds. It defines coordination compounds as complex compounds formed when a central metal ion bonds with surrounding neutral or charged ligand species via dative bonds. The central concepts discussed include: - Coordination number which represents the total number of bonds between the metal ion and ligands. - Coordination sphere which comprises the metal ion and surrounding ligands. - Ligands which contain lone pair electrons and can be monodentate, polydentate, etc. - Werner's theory which proposed that metal ions exhibit both primary and secondary valences satisfied by anions and ligands respectively. - Modern electronic theories like Sidgwick's effective atomic number rule and valence bond

1. Coordination Compounds
By
Dr. P. Amutha
Assistant Professor
Department of Chemistry
PSGR Krishnammal College for Women

2. INTRODUCTION
• Simple Salts
When an acid reacts with an alkali, a simple salt is produced
NaOH + HCl NaCl + H2O
• Addition (or) molecular compounds
Molecular (or) addition compounds are formed when solution containing
two or more salts are allowed to evaporate . It is further classified in to two types
 Double salts / Lattice Compounds
 Coordination / Complex compounds

3. Double Salts Co-ordination compound
These exist only in solid state and dissociate into
constituent species in their solution
Eg: Mohr’s salt: FeSO4.(NH4)2SO4.6H2O
They retain their identity in solid as well in
solution state
They lose their identity in dissolved state
Eg: An aqueous solution of potash alum will give
the tests for K+, Al3+, and SO4
-2
K2SO4.Al2(SO4)3.24H2O →2K++ 2Al+3 + 4SO4
-2+ 24H2O
They do not lose their identity in dissolved state
Their properties are essentially the same as those of
constituent species
Their properties are different from those of their
constituents
Eg: K4[Fe(CN)6] does not show the test of Fe2+ and
CN- ions

4. Coordination Compounds
• The central metal ion in the complex forms dative (or) coordinate
covalent bonds with the neutral/anionic/cationic species surrounding it
[Fe(CN)6]4-

5. Coordination Number
Total Number of ligands
that can coordinate to the central
metal ion (i.e.) it represents the
total number of bond formed
between the central metal ion
and donor atoms of the ligand
NH2CH2CH2NH2
–
Ethylenediamine (en)

6. Coordination Sphere
[Fe(CN)6]3- - Comprises of metal ion and ligands
Oxidation State
K3[Fe(CN)6] = 3K+ + [Fe(CN)6]3- [Co(NH3)4Cl2]Cl = [Co(NH3)4Cl2]1+ + Cl-
Oxidation state of Fe = 3 Oxidation state of Co = 3
X + 6 (-1) = -3 X +4 (0)+2 (-1) = +1
X = -3+6 X = +1+2
= 3 = 3

7. LIGANDS
• They are neutral molecule (or) ionic species which contain one or more lone
pair of electrons
Classification of Ligands Based on Number of Donor Atoms
 Monodendate / Unidendate Ligands
The ligands contain one donar atom and hence can coordinate to
the central metal atom at one position
Ex. Negative Ligands: Cl-, F-, CH3COO-, Br-, CN-, SCN-, SO4
2-
Positive Ligands: NO+
Neutral Ligands: H2O, NH3, CO, CS, NO

8.  Polydentate Ligands
The ligands having two (or) more donar site

9.  Ambidentate Ligands
Ligands which can attach
them with the central metal atom
by two different atoms are called
as ambidentate ligands.

10.  Flexidentate Ligands
Certain polydentate ligands
not uses all its donar atoms to get
coordinated to metal ion
EDTA – can also act as
pentadentate, tetradentate &
hexadentate ligand

11.  Bridging Ligands
Monodentate ligands having
more than one electron pair
can coordinate simultaneously
with two (or) more central
metal atom

12. Naming Metal Complexes
(i)Ligands
Neutral Ligands - It is named as such
Ex. NH2CH2CH2NH2 – Ethylenediamine (en)
NH2CH2CH2NHCH2CH2NH2 – diethylenetriamine (dien)
H2O – aqua
NH3 – ammine
NO – nitrosyl
CO – carbonyl

13. Naming Metal Complexes
(i)Ligands
NegativeLigands – If the anion name ends in ide, ite or ate, the final ‘e’ is
replaced by ‘o’

14. Naming Metal Complexes
(i)Ligands
Positive Ligands
It is named by giving suffix ‘ium’
Ex. NH2NH3
+ - hydrazinium
NO+ - nitrosylium

15. Naming Metal Complexes
(i)Ligands
 Two (or) more ligands of one type is present then it is indicated with Greek prefix
di, tri, tetra, penta, hexa, etc.
 If the ligand has di, tri, in its name then it is indicated with prefix
bis, tris, tetrakis, pentakis etc.
Eg: bisethylenediamine
 If there is more than one ligand type
• Listed in alphabetical order with no space between them
• Prefixes do not affect the order
diamminedichlorodifluoro
triamminedichlorofluoro

16. Naming Metal Complexes
(ii) Metal

17. Naming Metal Complexes
(ii) Metal

18. Naming Metal Complexes
• Counter ions are named as separate words and are not numbered
Anionic Complex
[Cr(en)I4]- - ethylenediamminetetraiodochromate(III) ion
[Zn(OH)4]2- - tetrahydroxozincate(II) ion
K3[Co(CN)6] - potassium hexacyanocobaltate(III)
Cationic Complex
[Ag(NH3)2]+ - diamminesilver(I) ion
[Co(NH3)4Cl2]Cl – tetraamminedichlorocobalt(III) chloride
Neutral Complex
[Ni(CO)4] - tetracarbonylnickel(0)

19. 1. [Cr(NH3)3(H2O)3]Cl3
triamminetriaquachromium(III) chloride
2. [Pt(NH3)5Cl]Br3
pentaamminechloroplatinum(IV) bromide
3. [Pt(H2NCH2CH2NH2)2Cl2]Cl2
dichlorobis(ethylenediamine)platinum(IV) chloride
4. Na2[NiCl4]
sodium tetrachloronickelate(II)
4. [Ag(NH3)2][Ag(CN)2] - diamminesilver(I) dicyanoargentate(I)

20. Werner's Theory
Alfred Werner a Swiss chemist put forward a theory to explain the
formation of complex compounds
Postulates
The central metal atom (or) ion in a coordination compound exhibits two
types of valencies - primary and secondary.
Primary valencies are ionisable and correspond to the oxidation state of
metal
Primary valencies are satisfied by negative ions & represented by dotted
lines

21. Werner's Theory
 Secondary valencies correspond to the coordination number
of the metal atom
 Secondary valency is satisfied either by negative (or) neutral
molecules & represented by solid lines
 Every metal atom has fixed number of secondary valencies
 Secondary valencies are directional (ie) they are directed
towards fixed position in space and this leads to definite
geometry.
 Every metal atom tends to satisfy both its primary and
secondary valencies. In certain cases negative ions satisfy both
types of valencies
 The postulates of Werner's coordination theory were actually
based on experimental evidence rather than theoretical

22. a) yield three moles of AgCl(s) per mole of compound
b) yield only two moles of AgCl(s) per mole of compound
c) yield only one moles of AgCl(s) per mole of compound
a)
b)
d)
c)

23. Some chloride ions must do
double duty and help satisfy
both primary valency and
secondary valency

24. Drawbacks of Werner’s Theory
It doesn't explain why only certain elements form coordination
compounds.
It does not explain the stability of the complexes
compounds have directional properties.
It does not explain the colour, and the magnetic and optical properties of
complexes.

25. Sidgwick’s Electronic Theory
 Sidgwick in 1927 extended the Lewis theory of electron pair bond
formation to explain the bonding in coordination compounds.
 He suggested that metal ion accepts electron pairs from the ligands
until it achieves the next noble gas configuration
 The total number of electrons on central atom including those gained
from ligands is called effective atomic number
 Sidgwick Theory or EAN Rule

26. Effective Atomic Number
• EAN = Z – X + Y
Where,
Z = atomic no. of the metal
X = no. of electron lost during the formation of the metal ion
from its atom
Y = no. of electrons donated by the ligands

27. Example 1
EAN of Fe (II) in [Fe(CN)6]4-
Fe (0) atom Z= 26 electrons
Fe (II) ion (Z-X) = 26-2 = 24
Electrons donated by
6 CN- = Y = 2X6 = 12
EAN = Z – X + Y
= 24 + 12 = 36 (Kr)
Example 2
EAN of Pt (IV) in [Pt(NH3)6]4+
Pt (0) atom Z = 78 electrons
Pt (IV) ion (Z-X) = 78 – 4 = 74
Electrons donated by
6 NH3 = Y = 2 X 6 = 12
EAN = Z – X + Y
= 74 + 12 = 86 (Rn)

28. Draw Backs of Sidgwick Theory
• Many complexes are stable but do
not follow EAN rule
• Theory doesn’t predict the
geometry of the complexes
• Theory doesn’t predict the
magnetic behavior of the
complexes

29. VALENCE BOND THEORY
Valence bond theory – Originally proposed by Linus Pauling in the 1930’s, using hybridization ideas
Postulates
 Central metal atom makes available empty s, p, d & f atomic orbitals equal to its coordination
number.
 These vacant atomic orbitals hybridize together to form hybrid orbitals. These hybrid orbitals have
same energy and definite geometry. Number of hybrid orbital formed is equalent to number of
atomic orbital involved in hybridisation
 The ligands have at lest one σ orbital with lone pair of electrons & each ligand donates a pair of
electron to central metal ion
 Vacant hybrid orbital of metal overlaps with the filled σ orbital of the ligands to form L→ M
coordinate bond

30. COORDINATION
NUMBER
GEOMETRY HYBRIDISATION Orbitals
2 Linear sp
3 Trigonal sp2
4 Tetrahedral
Square Planar
sp3
dsp2
5 Trigonal
bipyramidal
sp3d
6 Octahedral d2sp3 / sp3d2

31. Octahedral Complexes

32. Octahedral Complexes
i) Inner orbital complexes (d2sp3)
ii) Outer orbital complexes (sp3d2)
Inner Orbital Complexes (d2sp3)
 dz
2 & dx
2-y
2 - lie in path of the ligands.
 Atomic orbitals used for hybridization
are (n-1)d2 ns np3
 This type of hybridization takes place
in complexes with strong ligands
Outer Orbital Complexes (sp3d2)
 dz
2 & dx
2-y
2 - lie in path of the ligands.
 Atomic orbitals used for hybridization
are nd2 ns np3
 This type of hybridization takes place
in complexes with weak ligands

33. Example:
[Fe(CN)6]3- - Hexacyanoferrate(III) ion (paramagnetic with one unpaired electron)
Fe = 3d6 4s2 4p0 & Fe3+ = 3d5 4s0 µ = √n(n+2) BM = √1(1+2) = 1.73 BM
Fe3+
Fe3+ in [Fe(CN)6]3-
[Fe(CN)6]3-
Octahedral d2sp3 Geometry
↑↓ ↑↓ ↑
xx xx xx xx
↑ ↑ ↑ ↑ ↑
mix
d2sp3
↑↓ ↑↓ ↑ xx xx
3d 4s 4p
3d
3d
4s
4s
4p
4p
dx
2-y
2 dz
2
dx
2-y
2 dz
2

34. Octahedral sp3d2 Geometry
Example:
[CoF6]3- - Hexafluorocobaltate (III) ion (paramagnetic with four unpaired electron)
Co = 3d7 4s2 4p0 & Co3+ = 3d6 4s0 µ = √n(n+2) BM
Co3+
Co3+ in [CoF6]3-
[CoF6]3-
↑↓ ↑ ↑ ↑ ↑
xx xx xx xx
↑↓ ↑ ↑ ↑ ↑
mix
sp3d2
↑↓ ↑ ↑ ↑ ↑
4dz
2 4dx
2-y
2
3d 4s 4p
3d 4s 4p
3d
4s 4p
xx xx
4dz
2 4dx
2-y
2
dx
2-y
2 dz
2
dx
2-y
2 dz
2

35. Tetrahedral sp3 Geometry
Example:
[MnCl4]2- (paramagnetic with five unpaired electron)
Mn = 3d5 4s2 4p0 & Mn2+ = 3d54s0 µ = √n(n+2) BM = √5(5+2) = 5.9 BM
Mn2+
Mn2+ in [MnCl4]2-
[MnCl4]2-
↑ ↑ ↑ ↑ ↑
xx xx xx xx
↑ ↑ ↑ ↑ ↑
mix
sp3
↑ ↑ ↑ ↑ ↑
3d 4s 4p
3d 4s 4p
3d
4s 4p

36. Square planar dsp2 Geometry
Example:
[Ni(CN)4]2- (Diamagnetic)
Ni = 3d8 4s2 4p0 & Ni2+ = 3d8 4s0
Ni2+
Ni2+ in [Ni(CN)4]2-
[Ni(CN)4]2-
↑↓ ↑↓ ↑↓ ↑↓
xx xx xx
↑↓ ↑↓ ↑↓ ↑ ↑
mix
dsp2
↑↓ ↑↓ ↑↓ ↑↓ xx
3d 4s py
3d 4s 4p
3d
4s 4p
3 dx
2-y
2 px
4p

37. Example:
[Cu(NH3)4]2+
Ni = 3d9 4s2 4p0 & Cu2+ = 3d9 4s0
a) Cu2+
b) Cu2+ in [Cu(NH3)4]2+
c) Cu2+ in [Cu(NH3)4]2+
↑↓ ↑↓ ↑↓ ↑↓ ↑
3d 4s 4p
↑↓ ↑↓ ↑↓ ↑↓ ↑
3d 4s 4p pz
↑↓ ↑↓ ↑↓ ↑↓
3d 4s 4p
↑
5s
↑↓ ↑↓ ↑↓ ↑↓ ↑ xx xx xx
3d 4s 4p
xx
5d
dx
2-y
2
sp2d

38. Limitations of VBT
• Assumes all the d-orbitals are in same energy
• Octahedral, square planar and tetrahedral complexes of d1, d2, d3 & d9 have same
number of unpaired electrons and hence cannot be distinguished based on unpaired
electrons.
• The theory is unable to correlate the electronic spectra and magnetic moment of
complexes – both colour and magnetic moment of the complex is due to unpaired
electrons
• More stress is laid on central metal ion while the importance of ligand is not properly