Class 12 Chemistry Notes

Chapter 5 – Coordination Compounds (Part 1)

1. Introduction

Coordination compounds are compounds in which a central metal atom/ion is attached to surrounding ligands through coordinate (dative) bonds.

Importance

Present in biological molecules like haemoglobin and chlorophyll.
Used in medicines.
Important in electroplating.
Used as catalysts.
Useful in analytical chemistry.

2. Werner’s Theory

Alfred Werner explained the structure of coordination compounds.

Main Postulates

Primary Valency

Equal to oxidation state.
Ionisable.
Satisfied by negative ions.

Secondary Valency

Equal to coordination number.
Non-ionisable.
Satisfied by ligands.
Has fixed geometry.

Common Geometries

Coordination Number	Geometry
2	Linear
4	Tetrahedral / Square planar
6	Octahedral

3. Important Definitions

(A) Coordination Entity

The complete species enclosed inside square brackets.

Example:

[Co(NH3)6]³⁺

(B) Central Metal Atom/Ion

The atom or ion to which ligands are attached.

Examples:

Co³⁺
Fe²⁺
Ni²⁺

(C) Ligands

Atoms, ions or molecules that donate an electron pair to the metal ion.

Examples

NH₃
H₂O
Cl⁻
CN⁻

Types of Ligands

1. Unidentate

Donates one lone pair.

Examples

NH₃
H₂O
Cl⁻

2. Bidentate

Donates two lone pairs.

Examples

Oxalate (C₂O₄²⁻)
Ethylenediamine (en)

3. Polydentate

Donates more than two lone pairs.

Example

EDTA⁴⁻ (Hexadentate)

4. Ambidentate

Can attach through two different atoms.

Examples

NO₂⁻
SCN⁻

4. Coordination Number (CN)

The number of donor atoms directly attached to the metal ion.

Examples

Complex	CN
[Co(NH₃)₆]³⁺	6
[Ni(CO)₄]	4
[PtCl₆]²⁻	6

5. Coordination Sphere

Everything inside the square brackets forms the coordination sphere.

Example

K₄[Fe(CN)₆]

Coordination sphere → [Fe(CN)₆]⁴⁻
Counter ions → K⁺

6. Coordination Polyhedron

Represents the arrangement of ligands around the metal ion.

Common shapes

Octahedral
Tetrahedral
Square planar

7. Oxidation Number

Charge present on the central metal after removing all ligands.

Example

[Fe(CN)₆]⁴⁻

Fe oxidation state = +2

8. Homoleptic and Heteroleptic Complexes

Homoleptic

Only one type of ligand present.

Example

[Co(NH₃)₆]³⁺

Heteroleptic

More than one type of ligand present.

Example

[Co(NH₃)₄Cl₂]⁺

9. Difference Between Double Salt and Complex Compound

Double Salt	Complex Compound
Dissociates completely in water	Does not dissociate completely
Gives all constituent ions	Gives complex ion intact
Less stable in solution	More stable in solution

Examples

Double salt:

Mohr’s salt
Potash alum

Complex:

K₄[Fe(CN)₆]

Quick Revision

✅ Primary valency = Oxidation state

✅ Secondary valency = Coordination number

✅ Ligands donate lone pair

✅ EDTA = Hexadentate ligand

✅ CN depends on donor atoms

✅ Species inside [] = Coordination sphere

✅ One ligand type = Homoleptic

✅ Multiple ligand types = Heteroleptic

10. IUPAC Nomenclature of Coordination Compounds

Rules for Writing Formula

Write the central metal atom first.
Write ligands in alphabetical order.
Enclose the coordination entity in square brackets [ ].
Write the charge of the complex outside the brackets.
Counter ions are written outside the brackets.

Rules for Naming Coordination Compounds

Rule 1

Name the cation first, then the anion.

Rule 2

Name ligands before the metal.

Rule 3

Use prefixes to indicate the number of ligands.

Number	Prefix
1	Mono
2	Di
3	Tri
4	Tetra
5	Penta
6	Hexa

For complex ligand names use:

bis
tris
tetrakis

Rule 4

Names of common ligands

Formula	IUPAC Name
H₂O	Aqua
NH₃	Ammine
CO	Carbonyl
NO	Nitrosyl
Cl⁻	Chlorido
Br⁻	Bromido
CN⁻	Cyanido
OH⁻	Hydroxido
NO₂⁻	Nitrito
SCN⁻	Thiocyanato

Rule 5

Write oxidation state of the metal in Roman numerals.

Example

Cobalt(III)

Iron(II)

Platinum(IV)

Rule 6

If the complex ion is anionic, the metal name ends with “-ate”.

Examples

Metal	Name in Anionic Complex
Fe	Ferrate
Cu	Cuprate
Ag	Argentate
Au	Aurate
Sn	Stannate
Pb	Plumbate
Co	Cobaltate
Ni	Nickelate
Cr	Chromate

Important Examples

Formula	IUPAC Name
[Co(NH₃)₆]Cl₃	Hexaamminecobalt(III) chloride
K₃[Fe(CN)₆]	Potassium hexacyanidoferrate(III)
K₂[PtCl₄]	Potassium tetrachloridoplatinate(II)
[Pt(NH₃)₂Cl₂]	Diamminedichloridoplatinum(II)

Shortcut for Naming

Ligands → Metal → Oxidation State → Counter Ion

Example

[Co(NH₃)₅Cl]Cl₂

→ Pentaamminechloridocobalt(III) chloride

11. Isomerism in Coordination Compounds

Definition

Compounds having the same molecular formula but different arrangement of atoms or ligands are called isomers.

Types of Isomerism

A. Structural Isomerism

Different bonding arrangement.

Includes

Ionisation
Hydrate (Solvate)
Linkage
Coordination

B. Stereoisomerism

Same bonding but different spatial arrangement.

Includes

Geometrical
Optical

12. Geometrical Isomerism

Occurs due to different positions of ligands.

Cis Isomer

Similar ligands are adjacent (90°).

Example

Cl     Cl
 \     /
  Pt
 /     \
NH3   NH3

Trans Isomer

Similar ligands are opposite (180°).

Cl    NH3
 \    /
  Pt
 /    \
NH3   Cl

Important Points

Occurs in

Square planar complexes
Octahedral complexes

Does not occur in tetrahedral complexes because all ligand positions are equivalent.

Facial (fac) and Meridional (mer)

Found in octahedral complexes of the type MA₃B₃.

Facial (fac)

Three identical ligands occupy one face of the octahedron.

Meridional (mer)

Three identical ligands lie in one plane (meridian).

13. Optical Isomerism

Occurs when a compound and its mirror image cannot be superimposed.

These forms are called enantiomers.

Types

d-form (dextro)
l-form (laevo)

Conditions

Common in octahedral complexes containing bidentate ligands such as ethane-1,2-diamine (en).

14. Structural Isomerism

(A) Linkage Isomerism

Occurs due to ambidentate ligands.

Examples

NO₂⁻
SCN⁻

These ligands can attach through different donor atoms.

(B) Coordination Isomerism

Occurs when ligands exchange between the cationic and anionic complex ions.

(C) Ionisation Isomerism

Occurs when a ligand inside the coordination sphere exchanges places with the counter ion.

Example

[Co(NH₃)₅SO₄]Br
[Co(NH₃)₅Br]SO₄

These produce different ions in solution.

(D) Solvate (Hydrate) Isomerism

Occurs due to different positions of solvent molecules (usually water).

Water may be:

inside the coordination sphere, or
outside as water of crystallisation.

NCERT Exam Tips

Geometrical Isomerism

✔ Cis–Trans

✔ Fac–Mer

Optical Isomerism

✔ Mirror-image isomers

✔ Found mainly in octahedral complexes with bidentate ligands

Structural Isomerism

✔ Linkage

✔ Coordination

✔ Ionisation

✔ Solvate (Hydrate)

One-Minute Revision

Cis → Similar ligands together.
Trans → Similar ligands opposite.
Fac → Three same ligands on one face.
Mer → Three same ligands in one plane.
Optical isomers → Non-superimposable mirror images.
Linkage isomerism → Ambidentate ligand binds through different atoms.
Ionisation isomerism → Ligand and counter ion exchange.
Coordination isomerism → Ligands exchange between two complex ions.
Hydrate isomerism → Different positions of water molecules.

15. Bonding in Coordination Compounds (Valence Bond Theory – VBT)

Basic Idea

According to Valence Bond Theory (VBT):

The central metal atom/ion undergoes hybridisation.
Empty hybrid orbitals accept lone pairs from ligands.
Coordinate bonds are formed by overlap between ligand orbitals and metal hybrid orbitals.

16. Hybridisation and Geometry

Coordination Number	Hybridisation	Shape
4	sp³	Tetrahedral
4	dsp²	Square Planar
5	sp³d	Trigonal Bipyramidal
6	d²sp³	Octahedral (Inner orbital)
6	sp³d²	Octahedral (Outer orbital)

17. Octahedral Complexes

(A) Inner Orbital Complex

Hybridisation

d²sp³

Features

Uses inner (n−1)d orbitals.
Strong field ligands cause electron pairing.
Usually diamagnetic or weakly paramagnetic.
Also called Low Spin Complex.

Example

[Co(NH₃)₆]³⁺

(B) Outer Orbital Complex

Hybridisation

sp³d²

Features

Uses outer d orbitals.
Weak field ligands do not pair electrons.
Usually paramagnetic.
Also called High Spin Complex.

Example

[CoF₆]³⁻

18. Tetrahedral Complexes

Hybridisation

sp³

Properties

Always tetrahedral.
Usually paramagnetic.
Electron pairing generally does not occur.

Example

[NiCl₄]²⁻

19. Square Planar Complexes

Hybridisation

dsp²

Properties

Common for d⁸ metal ions.
Usually diamagnetic.
Strong field ligands are present.

Example

[Ni(CN)₄]²⁻

20. Magnetic Properties

Paramagnetic

Contains one or more unpaired electrons.

Examples

[NiCl₄]²⁻
[CoF₆]³⁻

Diamagnetic

All electrons are paired.

Examples

[Ni(CN)₄]²⁻
[Co(NH₃)₆]³⁺

21. Strong Field and Weak Field Ligands

Strong Field Ligands

Cause electron pairing.

Examples

CN⁻
CO
NH₃
en (ethane-1,2-diamine)

Result

Inner orbital complex
Low spin
Usually diamagnetic

Weak Field Ligands

Do not cause electron pairing.

Examples

F⁻
Cl⁻
Br⁻
I⁻

Result

Outer orbital complex
High spin
Usually paramagnetic

22. Inner Orbital vs Outer Orbital Complex

Inner Orbital	Outer Orbital
d²sp³	sp³d²
Strong ligand	Weak ligand
Low spin	High spin
Electron pairing occurs	Pairing usually absent
Often diamagnetic	Usually paramagnetic

23. Magnetic Moment

Magnetic moment depends on the number of unpaired electrons.

Formula

$\mu = \sqrt{n(n+2)} \text{ BM}$ μ=n(n+2) BM

Where:

μ = Magnetic moment
n = Number of unpaired electrons
BM = Bohr Magneton

24. Limitations of Valence Bond Theory

VBT explains geometry and magnetic behaviour but has some limitations:

Does not explain the colour of complexes.
Cannot explain why some ligands are strong and others are weak.
Does not give quantitative magnetic data.
Does not explain stability of complexes.
Cannot always predict square planar vs tetrahedral geometry correctly.

Board Important Examples

Complex	Hybridisation	Shape	Nature
[Co(NH₃)₆]³⁺	d²sp³	Octahedral	Diamagnetic
[CoF₆]³⁻	sp³d²	Octahedral	Paramagnetic
[NiCl₄]²⁻	sp³	Tetrahedral	Paramagnetic
[Ni(CN)₄]²⁻	dsp²	Square planar	Diamagnetic
[Ni(CO)₄]	sp³	Tetrahedral	Diamagnetic

NCERT-Based Questions

Q1. Why is [Ni(CN)₄]²⁻ diamagnetic?

Answer:
CN⁻ is a strong field ligand. It pairs the electrons of Ni²⁺, resulting in dsp² hybridisation. Since all electrons are paired, the complex is diamagnetic.

Q2. Why is [NiCl₄]²⁻ paramagnetic?

Answer:
Cl⁻ is a weak field ligand. It does not pair the electrons of Ni²⁺. The complex has sp³ hybridisation with two unpaired electrons, making it paramagnetic.

Q3. Why is [Co(NH₃)₆]³⁺ called an inner orbital complex?

Answer:
NH₃ is a strong field ligand. It causes electron pairing, allowing the use of inner d orbitals for d²sp³ hybridisation, so it is called an inner orbital (low-spin) complex.

One-Minute Revision

VBT → Metal hybridises and ligands donate lone pairs.
sp³ → Tetrahedral.
dsp² → Square planar.
d²sp³ → Inner orbital, low spin.
sp³d² → Outer orbital, high spin.
Strong ligands → Pair electrons.
Weak ligands → Do not pair electrons.
Paramagnetic → Unpaired electrons.
Diamagnetic → All electrons paired.
Magnetic moment depends on the number of unpaired electrons.

25. Crystal Field Theory (CFT)

Introduction

Crystal Field Theory (CFT) explains the electronic structure, colour, magnetic properties, and stability of coordination compounds.

Basic Assumptions

Ligands are treated as point charges (or point dipoles for neutral ligands).
The interaction between the metal ion and ligands is mainly electrostatic.
When ligands approach the metal ion, the five d-orbitals no longer have the same energy (degeneracy is removed).
This results in crystal field splitting.

26. Crystal Field Splitting in Octahedral Complexes

When six ligands surround a metal ion in an octahedral arrangement:

Two orbitals (dx²−y² and dz²) experience greater repulsion and move to higher energy.
Three orbitals (dxy, dxz, dyz) experience less repulsion and remain at lower energy.

Orbital Sets

Lower Energy	Higher Energy
t₂g	e<sub>g</sub>

Splitting Energy

The energy difference between these two sets is called the crystal field splitting energy (Δ₀).

27. High Spin and Low Spin Complexes

The arrangement of electrons depends on the comparison between:

Crystal field splitting energy (Δ₀)
Electron pairing energy (P)

High Spin Complex

Occurs when:

Δ₀ < Pairing Energy

Features:

Electrons occupy higher-energy orbitals before pairing.
More unpaired electrons.
Usually paramagnetic.
Formed by weak field ligands.

Examples:

F⁻
Cl⁻
Br⁻

Low Spin Complex

Occurs when:

Δ₀ > Pairing Energy

Features:

Electrons pair in lower-energy orbitals first.
Fewer unpaired electrons.
Usually diamagnetic.
Formed by strong field ligands.

Examples:

CN⁻
CO
NH₃

28. Factors Affecting Crystal Field Splitting

The value of Δ₀ depends on:

Nature of the metal ion.
Oxidation state of the metal.
Nature of the ligand.
Geometry of the complex.

Higher oxidation state generally increases Δ₀.

29. Spectrochemical Series

The spectrochemical series arranges ligands according to their ability to split d-orbitals.

Weak → Strong Field Ligands

I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO

Important Points

Ligands on the left produce small splitting.
Ligands on the right produce large splitting.

30. Colour of Coordination Compounds

Many coordination compounds are coloured because of d–d electronic transitions.

Explanation

White light contains many colours.
An electron absorbs energy equal to Δ₀ and moves from the lower-energy d-orbitals to the higher-energy d-orbitals.
The absorbed colour is removed from white light.
The remaining (complementary) colour is observed.

Colour Depends On

Metal ion.
Oxidation state.
Type of ligand.
Magnitude of Δ₀.

Why Are Some Complexes Colourless?

Complexes are colourless when:

The d-orbitals are completely empty (d⁰ configuration), or
The d-orbitals are completely filled (d¹⁰ configuration),

because no d–d transition is possible.

31. Limitations of Crystal Field Theory

Although CFT explains many properties, it has limitations:

Assumes purely electrostatic interactions.
Does not explain covalent character in bonding.
Cannot explain π-bonding between ligands and metal.
Does not account for all spectral observations.

32. Comparison: VBT vs CFT

Valence Bond Theory (VBT)	Crystal Field Theory (CFT)
Explains hybridisation	Explains d-orbital splitting
Predicts geometry	Explains colour
Explains magnetic behaviour	Explains high-spin and low-spin complexes
Does not explain colour	Explains colour and stability better

Board Important Questions

Q1. What is crystal field splitting?

Answer:
It is the separation of the five degenerate d-orbitals into two groups of different energies when ligands approach the central metal ion.

Q2. Why are coordination compounds coloured?

Answer:
They absorb specific wavelengths of visible light due to d–d electronic transitions. The complementary colour is reflected, giving the complex its observed colour.

Q3. What is the spectrochemical series?

Answer:
It is the arrangement of ligands in order of increasing crystal field splitting power, from weak field ligands to strong field ligands.

One-Minute Revision

CFT explains colour, magnetism, and stability.
In an octahedral field:
- t₂g → lower energy.
- eg → higher energy.
Δ₀ = crystal field splitting energy.
Weak ligands → High-spin complexes.
Strong ligands → Low-spin complexes.
Spectrochemical series: I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO
Colour arises due to d–d transitions.
d⁰ and d¹⁰ complexes are generally colourless.

33. Metal Carbonyls

Introduction

Metal carbonyls are coordination compounds in which carbon monoxide (CO) acts as a ligand attached to a metal atom.

General Formula

$M_x(CO)_y$ Mx(CO)y

Examples:

Ni(CO)₄
Fe(CO)₅
Cr(CO)₆

34. Bonding in Metal Carbonyls

CO is a strong field ligand.

Bonding occurs through two interactions:

1. Sigma Bond Formation

CO donates a lone pair of electrons from carbon to the empty orbital of the metal.
This forms a coordinate bond.

2. Pi Back Bonding

Filled metal d-orbitals donate electrons back to the empty antibonding orbital of CO.
This strengthens the metal–carbon bond.

35. Properties of Metal Carbonyls

Physical Properties

Many are volatile liquids or solids.
Generally stable compounds.
Carbonyls of transition metals are usually colourless or pale coloured.

Structural Features

Mononuclear Carbonyls

Contain one metal atom.

Examples:

Ni(CO)₄ → Tetrahedral
Fe(CO)₅ → Trigonal bipyramidal
Cr(CO)₆ → Octahedral

36. Importance of Coordination Compounds

Coordination compounds have many applications in chemistry, biology, and industry.

(A) Biological Importance

Haemoglobin

Contains iron coordination complex.
Helps in oxygen transport in blood.

Chlorophyll

Contains magnesium complex.
Responsible for photosynthesis in plants.

Vitamin B₁₂

Contains cobalt coordination complex.

(B) Medicinal Uses

Cisplatin

A platinum coordination compound.
Used in cancer treatment.

(C) Analytical Chemistry

Coordination compounds are used for:

Detection of metal ions.
Separation of ions.
Estimation of hardness of water.

Example:

EDTA forms stable complexes with metal ions.

(D) Metallurgy

Coordination compounds help in extraction and purification of metals.

Example:

Gold and silver extraction uses cyanide complexes.

(E) Electroplating

Complex ions are used to deposit metals smoothly on surfaces.

37. Stability of Coordination Compounds

The stability of a complex depends on:

Nature of metal ion.
Charge of metal ion.
Nature of ligand.
Coordination number.
Chelation.

38. Chelate Effect

Chelating Ligands

Ligands that attach through two or more donor atoms form ring structures.

Examples:

Ethylenediamine (en)
EDTA

Chelate Effect

Complexes containing chelating ligands are generally more stable than complexes containing similar unidentate ligands.

Example:

[Ni(en)₃]²⁺ is more stable than complexes containing simple monodentate ligands.

39. Important Formula Revision

Oxidation State Calculation

For: $[Co(NH_3)_6]Cl_3$ [Co(NH3)6]Cl3

NH₃ is neutral.

Let oxidation state of Co = x $x + 0 = +3$ x+0=+3

Co = +3

Coordination Number

Count donor atoms directly attached to metal.

Example:

[Co(en)₃]³⁺

en is bidentate.
Each en donates two pairs.
Coordination number = 6

40. Complete Chapter Quick Revision

Coordination Terms

Central metal → accepts electron pairs.
Ligand → donates electron pairs.
Coordination number → number of donor atoms.
Coordination sphere → species inside brackets.

Ligands

Unidentate → one donor atom.
Bidentate → two donor atoms.
Polydentate → many donor atoms.
Ambidentate → can attach through different atoms.

Nomenclature

Order:

Ligands → Metal → Oxidation state

Anionic complexes:

Metal name ends with -ate.

Isomerism

Structural

Ionisation
Hydrate
Linkage
Coordination

Stereoisomerism

Geometrical
Optical

VBT

sp³ → tetrahedral
dsp² → square planar
d²sp³ → inner orbital
sp³d² → outer orbital

CFT

Explains colour and magnetic properties.
d-orbitals split into different energy levels.
Strong ligands → low spin.
Weak ligands → high spin.

Magnetic Behaviour

Unpaired electrons → Paramagnetic
Paired electrons → Diamagnetic

Formula: $\mu=\sqrt{n(n+2)}$ μ=n(n+2)

Metal Carbonyls

CO acts as ligand.
Bonding involves sigma donation and pi back bonding.
Examples:
- Ni(CO)₄
- Fe(CO)₅
- Cr(CO)₆

Important Board Exam Questions

1. Explain the role of coordination compounds in biological systems.

Answer points:

Haemoglobin transports oxygen.
Chlorophyll helps photosynthesis.
Vitamin B₁₂ contains cobalt complex.

2. What is chelate effect?

Answer:
The increased stability of complexes formed by multidentate ligands due to ring formation is called the chelate effect.

3. Why is CO a strong field ligand?

Answer:
CO causes strong splitting of d-orbitals due to its ability to accept electron density through pi back bonding.