CBSE Class 12 Chemistry Revision Notes Chapter 5 Coordination Compounds
Coordination compounds are formed when a central metal atom or ion is bonded to ligands. In CBSE Class 12 Chemistry, this chapter explains Werner’s theory, ligands, nomenclature, isomerism, VBT, CFT and applications.
Coordination Compounds is an important chapter in inorganic chemistry. It explains how metal atoms or ions bind with anions or neutral molecules to form complex compounds. These compounds are also found in daily life and biological systems. Chlorophyll is a magnesium complex, haemoglobin is an iron complex and vitamin B12 is a cobalt complex.
Use these CBSE Class 12 Chemistry Revision Notes Chapter 5 to revise the 2026–27 chapter quickly. Start with Werner’s theory and important terms. Then move to nomenclature, isomerism, bonding theories, stability and applications.
Key Takeaways
- Coordination Compound: It contains a central metal atom or ion surrounded by ligands.
- Ligand: It donates an electron pair to the central metal atom or ion.
- Coordination Number: It is the number of donor atoms directly bonded to the central metal.
- Isomerism: Coordination compounds can show geometrical, optical, linkage, coordination, ionisation and solvate isomerism.
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Access 30 Minutes Class 12 Chemistry Chapter 5 Coordination Compounds Notes
Coordination Compounds becomes easier when you revise definitions first. Learn ligand, coordination number, coordination sphere, oxidation state and coordination entity.
After that, revise nomenclature rules and isomerism. Keep Valence Bond Theory, Crystal Field Theory and applications for the final round.
Coordination Compounds Class 12 Chemistry Chapter 5 Overview
Coordination compounds are compounds in which a central metal atom or ion is surrounded by a fixed number of ligands.
The ligands may be ions or neutral molecules.
Examples:
- [Co(NH₃)₆]Cl₃
- K₄[Fe(CN)₆]
- [Cu(NH₃)₄]SO₄
- [PtCl₄]²⁻
Coordination compounds are important in inorganic chemistry, biological systems, medicines, metallurgy and industry.
Double Salts and Coordination Compounds
Double salts and coordination compounds are both formed by the combination of two or more stable compounds. But they behave differently in water.
| Basis | Double Salt | Coordination Compound |
| Behaviour in water | Dissociates completely into simple ions | Complex ion remains intact |
| Identity in solution | Lost | Retained |
| Example | Mohr’s salt, potash alum | K₄[Fe(CN)₆], [Cu(NH₃)₄]SO₄ |
| Ionisation | Gives all simple ions | Gives complex ion and counter ion |
Example of double salt:
FeSO₄.(NH₄)₂SO₄.6H₂O dissociates into Fe²⁺, NH₄⁺ and SO₄²⁻ ions.
Example of coordination compound:
K₄[Fe(CN)₆] gives K⁺ and [Fe(CN)₆]⁴⁻ ions. It does not give free Fe²⁺ and CN⁻ ions.
Werner’s Theory of Coordination Compounds
Alfred Werner explained the structure of coordination compounds using primary and secondary valence.
He studied cobalt chloride-ammonia complexes and found that some chloride ions were ionisable, while some were directly attached to cobalt.
Main Postulates of Werner’s Theory
| Postulate | Explanation |
| Two types of valence | Metals show primary and secondary valence |
| Primary valence | Ionisable and satisfied by negative ions |
| Secondary valence | Non-ionisable and satisfied by neutral molecules or negative ions |
| Coordination number | Secondary valence is equal to coordination number |
| Spatial arrangement | Ligands have definite arrangement around the metal |
Example:
| Compound | Modern Formula | AgCl Formed with AgNO₃ |
| CoCl₃.6NH₃ | [Co(NH₃)₆]Cl₃ | 3 mol |
| CoCl₃.5NH₃ | [CoCl(NH₃)₅]Cl₂ | 2 mol |
| CoCl₃.4NH₃ | [CoCl₂(NH₃)₄]Cl | 1 mol |
In these compounds, the species inside square brackets is the coordination entity.
Important Terms in Coordination Compounds
Coordination Entity
A coordination entity is formed by a central metal atom or ion bonded to ligands.
Example:
[CoCl₃(NH₃)₃]
Here, cobalt is bonded to three chloride ions and three ammonia molecules.
Central Atom or Ion
The central atom or ion is the metal atom or ion to which ligands are attached.
Examples:
| Complex | Central Atom/Ion |
| [NiCl₂(H₂O)₄] | Ni²⁺ |
| [CoCl(NH₃)₅]²⁺ | Co³⁺ |
| [Fe(CN)₆]³⁻ | Fe³⁺ |
The central atom or ion acts as a Lewis acid because it accepts electron pairs.
Ligand
A ligand is an ion or molecule that donates an electron pair to the central metal atom or ion.
Examples:
- Cl⁻
- H₂O
- NH₃
- CN⁻
- C₂O₄²⁻
- en
Ligands act as Lewis bases.
Coordination Number
Coordination number is the number of ligand donor atoms directly bonded to the central metal atom or ion.
Examples:
| Complex | Coordination Number |
| [PtCl₆]²⁻ | 6 |
| [Ni(NH₃)₄]²⁺ | 4 |
| [Fe(C₂O₄)₃]³⁻ | 6 |
| [Co(en)₃]³⁺ | 6 |
Only sigma bonds from ligand donor atoms are counted.
Coordination Sphere
The central atom or ion and ligands written inside square brackets form the coordination sphere.
Example:
K₄[Fe(CN)₆]
Coordination sphere = [Fe(CN)₆]⁴⁻
Counter ion = K⁺
Coordination Polyhedron
The arrangement of ligands around the central metal atom or ion is called coordination polyhedron.
| Complex | Shape |
| [Co(NH₃)₆]³⁺ | Octahedral |
| [Ni(CO)₄] | Tetrahedral |
| [PtCl₄]²⁻ | Square planar |
Oxidation Number
Oxidation number is the charge the central atom would carry if all ligands were removed with their shared electron pairs.
Example:
In [Cu(CN)₄]³⁻:
CN⁻ has charge -1.
Let oxidation state of Cu = x
x + 4(-1) = -3
x = +1
So, copper is Cu(I).
Homoleptic and Heteroleptic Complexes
| Type | Meaning | Example |
| Homoleptic complex | Metal is attached to only one type of ligand | [Co(NH₃)₆]³⁺ |
| Heteroleptic complex | Metal is attached to more than one type of ligand | [Co(NH₃)₄Cl₂]⁺ |
Types of Ligands in Coordination Compounds
Ligands are classified on the basis of the number of donor atoms.
| Type of Ligand | Donor Atoms | Example |
| Unidentate | 1 | Cl⁻, NH₃, H₂O |
| Didentate or bidentate | 2 | en, C₂O₄²⁻ |
| Polydentate | More than 2 | EDTA⁴⁻ |
| Hexadentate | 6 | EDTA⁴⁻ |
| Ambidentate | Can bind through two different atoms | NO₂⁻, SCN⁻ |
Chelate Ligands
When a di- or polydentate ligand attaches to the same metal ion through two or more donor atoms, it forms a chelate.
Chelate complexes are usually more stable than similar complexes with unidentate ligands.
Example:
en and C₂O₄²⁻ form chelate complexes.
Ambidentate Ligands
Ambidentate ligands have two possible donor atoms, but they attach through only one donor atom at a time.
Examples:
| Ligand | Donor Atoms |
| NO₂⁻ | N or O |
| SCN⁻ | S or N |
Nomenclature of Coordination Compounds
Nomenclature gives a systematic way to write and name coordination compounds.
Rules for Writing Formulas
| Rule | Explanation |
| Central atom first | Write metal atom or ion first |
| Ligands next | Write ligands in alphabetical order |
| Square brackets | Enclose coordination entity in [ ] |
| Polyatomic ligands | Write them in parentheses |
| Charge | Write charge outside the bracket |
| Counter ions | Balance total charge |
Example:
Potassium hexacyanidoferrate(II)
Formula: K₄[Fe(CN)₆]
Rules for Naming Coordination Compounds
| Rule | Explanation |
| Cation first | Name cation before anion |
| Ligands first | Name ligands before metal |
| Alphabetical order | Ligands are named alphabetically |
| Anionic ligands | End in -o or -ido |
| Neutral ligands | Special names are used |
| Oxidation state | Written in Roman numerals |
| Anionic complex | Metal name ends in -ate |
Common Ligand Names
| Formula | Ligand Name |
| NH₃ | ammine |
| H₂O | aqua |
| CO | carbonyl |
| NO | nitrosyl |
| Cl⁻ | chlorido |
| CN⁻ | cyanido |
| OH⁻ | hydroxido |
| C₂O₄²⁻ | oxalato |
| en | ethane-1,2-diamine |
Examples of IUPAC Names
| Formula | IUPAC Name |
| [Co(NH₃)₆]Cl₃ | hexaamminecobalt(III) chloride |
| [Co(NH₃)₅Cl]Cl₂ | pentaamminechloridocobalt(III) chloride |
| K₃[Fe(CN)₆] | potassium hexacyanidoferrate(III) |
| K₄[Fe(CN)₆] | potassium hexacyanidoferrate(II) |
| [Pt(NH₃)₂Cl₂] | diamminedichloridoplatinum(II) |
| [Ni(CO)₄] | tetracarbonylnickel(0) |
Isomerism in Coordination Compounds
Isomers are compounds with the same chemical formula but different arrangement of atoms.
Coordination compounds show two main types of isomerism:
| Type | Subtypes |
| Stereoisomerism | Geometrical and optical isomerism |
| Structural isomerism | Linkage, coordination, ionisation and solvate isomerism |
Geometrical Isomerism
Geometrical isomerism arises due to different spatial arrangements of ligands around the central metal.
It is common in square planar and octahedral complexes.
Cis and Trans Isomers
For square planar [MX₂L₂]:
| Isomer | Arrangement |
| Cis | Same ligands are adjacent |
| Trans | Same ligands are opposite |
Example:
[Pt(NH₃)₂Cl₂]
It has cis and trans forms.
Octahedral complexes like [Co(NH₃)₄Cl₂]⁺ also show cis-trans isomerism.
Fac and Mer Isomers
Octahedral complexes of type [Ma₃b₃] can show fac and mer isomerism.
| Isomer | Arrangement |
| Fac | Three same ligands occupy one face |
| Mer | Three same ligands lie around a meridian |
Example:
[Co(NH₃)₃(NO₂)₃]
Optical Isomerism
Optical isomers are non-superimposable mirror images.
They are called enantiomers.
The two forms are:
| Form | Meaning |
| d-form | Rotates plane-polarised light to the right |
| l-form | Rotates plane-polarised light to the left |
Optical isomerism is common in octahedral complexes with didentate ligands.
Example:
[Co(en)₃]³⁺ shows optical isomerism.
In [PtCl₂(en)₂]²⁺, only the cis form is optically active.
Linkage Isomerism
Linkage isomerism occurs when an ambidentate ligand attaches through different donor atoms.
Examples:
| Ligand | Forms |
| NO₂⁻ | nitrito-N and nitrito-O |
| SCN⁻ | thiocyanato-S and isothiocyanato-N |
Example:
[Co(NH₃)₅(NO₂)]Cl₂ can exist in two forms.
One form has nitrite attached through nitrogen, and the other through oxygen.
Coordination Isomerism
Coordination isomerism occurs when ligands interchange between cationic and anionic complex ions.
Example:
[Co(NH₃)₆][Cr(CN)₆]
and
[Cr(NH₃)₆][Co(CN)₆]
In the first compound, NH₃ is attached to Co and CN⁻ to Cr.
In the second compound, NH₃ is attached to Cr and CN⁻ to Co.
Ionisation Isomerism
Ionisation isomerism occurs when a ligand inside the coordination sphere and a counter ion outside the coordination sphere exchange places.
Example:
[Co(NH₃)₅(SO₄)]Br
and
[Co(NH₃)₅Br]SO₄
These compounds give different ions in solution.
Solvate Isomerism
Solvate isomerism occurs when a solvent molecule is directly bonded to the metal in one isomer and present outside the coordination sphere in another.
When water is involved, it is also called hydrate isomerism.
Example:
[Cr(H₂O)₆]Cl₃
and
[Cr(H₂O)₅Cl]Cl₂.H₂O
Bonding in Coordination Compounds
Werner’s theory explained many features of coordination compounds, but it did not explain bonding, magnetic behaviour and optical properties.
Bonding in coordination compounds is explained mainly by:
- Valence Bond Theory
- Crystal Field Theory
Valence Bond Theory
According to Valence Bond Theory, the central metal atom or ion uses vacant orbitals for hybridisation.
These hybrid orbitals accept lone pairs from ligands.
Common Hybridisations and Shapes
| Coordination Number | Hybridisation | Shape |
| 4 | sp³ | Tetrahedral |
| 4 | dsp² | Square planar |
| 6 | sp³d² | Octahedral |
| 6 | d²sp³ | Octahedral |
Inner Orbital and Outer Orbital Complexes
| Complex Type | Orbital Used | Spin Type |
| Inner orbital complex | Inner d orbitals | Low spin |
| Outer orbital complex | Outer d orbitals | High spin |
Example:
[Co(NH₃)₆]³⁺ is an inner orbital, low spin, diamagnetic octahedral complex.
[CoF₆]³⁻ is an outer orbital, high spin, paramagnetic octahedral complex.
Magnetic Properties of Coordination Compounds
Magnetic behaviour depends on the number of unpaired electrons.
| Type | Meaning |
| Diamagnetic | No unpaired electrons |
| Paramagnetic | One or more unpaired electrons |
Magnetic moment is calculated using:
μ = √n(n + 2) BM
Here, n is the number of unpaired electrons.
Strong field ligands cause electron pairing and often form low spin complexes.
Weak field ligands do not cause pairing easily and often form high spin complexes.
Crystal Field Theory
Crystal Field Theory explains bonding in coordination compounds by considering electrostatic interaction between the metal ion and ligands.
Ligands are treated as point charges or dipoles.
When ligands approach a transition metal ion, the d orbitals split into different energy levels.
Crystal Field Splitting in Octahedral Complexes
In an octahedral complex, five d orbitals split into two sets:
| Set | Orbitals | Energy |
| t₂g | dxy, dyz, dzx | Lower energy |
| eg | dx²-y², dz² | Higher energy |
The energy difference between these two sets is called crystal field splitting energy.
It is represented as Δ₀.
Crystal Field Splitting in Tetrahedral Complexes
In tetrahedral complexes, the splitting is smaller than in octahedral complexes.
The order of energy is reversed.
| Set | Energy |
| e | Lower energy |
| t₂ | Higher energy |
Tetrahedral complexes are usually high spin because the splitting energy is small.
Colour in Coordination Compounds
Many coordination compounds are coloured because of d-d transitions.
When light falls on the complex, an electron absorbs energy and jumps from a lower d orbital to a higher d orbital.
The colour observed is the complementary colour of the absorbed light.
The colour depends on:
- metal ion
- oxidation state
- ligand
- geometry
- crystal field splitting energy
Stability of Coordination Compounds
The stability of a coordination compound depends on the strength of the metal-ligand bond.
A stable complex has a high formation constant.
Formation constant is the equilibrium constant for the formation of a complex ion from the metal ion and ligands.
For:
M + 4L ⇌ ML₄
Formation constant:
Kf = [ML₄] / [M][L]⁴
Higher Kf means higher stability.
Factors Affecting Stability
| Factor | Effect |
| Charge on metal ion | Higher charge usually increases stability |
| Size of metal ion | Smaller size usually increases stability |
| Nature of ligand | Strong donor ligands form more stable complexes |
| Chelation | Chelate complexes are more stable |
| Ring size | Five and six-membered chelate rings are usually stable |
Importance and Applications of Coordination Compounds
Coordination compounds are useful in biology, medicine, analysis and industry.
Biological Applications
| Compound | Metal |
| Chlorophyll | Magnesium |
| Haemoglobin | Iron |
| Vitamin B12 | Cobalt |
These compounds are essential for photosynthesis, oxygen transport and biological functions.
Analytical Applications
Coordination compounds are used in qualitative and quantitative analysis.
Examples:
- EDTA is used in complexometric titrations.
- Dimethylglyoxime is used to detect Ni²⁺.
- Potassium ferrocyanide is used in detection reactions.
Industrial Applications
Coordination compounds are used in:
- electroplating
- textile dyeing
- metallurgy
- catalysis
- extraction of metals
Example:
Silver and gold are extracted using cyanide complexes.
Medicinal Applications
Some coordination compounds are used in medicine.
Example:
Cisplatin is used as an anticancer drug.
Quick Revision Table for Coordination Compounds
| Concept | Quick Point |
| Coordination compound | Central metal attached to ligands |
| Ligand | Electron pair donor |
| Central metal ion | Electron pair acceptor |
| Coordination number | Number of donor atoms attached to metal |
| Coordination sphere | Metal and ligands inside square brackets |
| Counter ion | Ion outside square brackets |
| Homoleptic complex | One type of ligand |
| Heteroleptic complex | More than one type of ligand |
| Werner’s primary valence | Ionisable |
| Werner’s secondary valence | Non-ionisable |
| Geometrical isomerism | Different spatial arrangement |
| Optical isomerism | Non-superimposable mirror images |
| Linkage isomerism | Ambidentate ligand binds differently |
| VBT | Explains hybridisation and geometry |
| CFT | Explains d-orbital splitting |
| Chelate effect | Chelate complexes are more stable |
Important Terms in Coordination Compounds
| Term | Meaning |
| Coordination entity | Central metal atom or ion with ligands |
| Central atom/ion | Metal atom or ion accepting electron pairs |
| Ligand | Ion or molecule donating electron pairs |
| Unidentate ligand | Ligand with one donor atom |
| Didentate ligand | Ligand with two donor atoms |
| Polydentate ligand | Ligand with many donor atoms |
| Ambidentate ligand | Ligand with two possible donor atoms |
| Chelate ligand | Ligand forming ring-like structure with metal |
| Coordination number | Number of donor atoms bonded to metal |
| Coordination sphere | Complex species inside square brackets |
| Counter ion | Ion outside coordination sphere |
| Oxidation number | Charge on metal after removing ligands |
| Homoleptic complex | Complex with one type of ligand |
| Heteroleptic complex | Complex with different types of ligands |
| Crystal field splitting | Splitting of d orbitals in ligand field |
| Formation constant | Measure of complex stability |
Useful Links for Class 12 Chemistry
| Section | Useful Links |
| Syllabus | CBSE Class 12 Chemistry Syllabus |
| Revision Notes | CBSE Class 12 Chemistry Revision Notes |
| Chemistry Notes | CBSE Class 12 Chemistry Revision Notes Chapter 1 |
| NCERT Solutions | NCERT Solutions for Class 12 Chemistry |
| Sample Papers | CBSE Sample Papers for Class 12 Chemistry |
| Important Questions | Important Questions Class 12 Chemistry |
| NCERT Books | NCERT Books for Class 12 Chemistry |
| Class 12 Support | CBSE Class 12 Syllabus |
FAQs (Frequently Asked Questions)
A coordination compound contains a central metal atom or ion bonded to ligands. The ligands donate electron pairs to the metal. Example: K₄[Fe(CN)₆] contains the complex ion [Fe(CN)₆]⁴⁻.
Coordination number is the number of donor atoms directly attached to the metal. Oxidation number is the charge the metal would have after removing all ligands with their shared electron pairs.
Many coordination compounds are coloured due to d-d transitions. Electrons absorb visible light and move from lower energy d orbitals to higher energy d orbitals. The colour seen is the complementary colour.
Chelate complexes are more stable because polydentate ligands bind to the metal through more than one donor atom. This forms ring structures and gives extra stability to the complex.
The most important topics are Werner’s theory, important terms, ligand types, nomenclature, isomerism, Valence Bond Theory, Crystal Field Theory, magnetic behaviour, stability and applications.