Coordination Compounds Class 12 Notes (2026-27) — CBSE
Class 12 Chemistry Chapter 5 notes: Werner's theory, ligands, nomenclature, isomerism, VBT, CFT, spectrochemical series and applications.
Coordination Compounds — Class 12 Chemistry Notes
Chapter Snapshot
A coordination compound contains a central metal atom or ion bonded to a fixed number of ions or neutral molecules called ligands via coordinate bonds. This chapter covers Werner's theory, IUPAC nomenclature, types of isomerism, and bonding theories — Valence Bond Theory (VBT) and Crystal Field Theory (CFT) — that explain colour, magnetism, and geometry.
Board relevance: ~8 marks in the Inorganic Chemistry unit. Expect a nomenclature question, a type of isomerism, and a CFT-based explanation of colour or magnetic behaviour. Learn the rules systematically.
Werner's Theory
Alfred Werner (Nobel Prize, 1913) proposed:
1. Metals exhibit two types of valency: primary (ionisable) valency, which is satisfied by anions and is called the oxidation state; and secondary (non-ionisable) valency, which is satisfied by neutral molecules or ions called ligands and defines the coordination number.
2. Ligands are arranged around the central metal in a fixed geometry — most often octahedral (CN = 6), square planar (CN = 4), or tetrahedral (CN = 4).
Example: In [Co(NH₃)₆]Cl₃, Co³⁺ has primary valency 3 (satisfied by 3Cl⁻) and secondary valency 6 (satisfied by 6NH₃ ligands).
Key Concepts & Definitions
Coordination entity: the central metal ion/atom plus its ligands, written inside square brackets.
Ligand: an ion or molecule that donates a lone pair to the central metal. Classified by the number of donor atoms:
- Monodentate: one donor atom (Cl⁻, H₂O, NH₃, CN⁻, NO₂⁻).
- Bidentate: two donor atoms — e.g. ethylenediamine (en), oxalate (C₂O₄²⁻).
- Polydentate: three or more — e.g. EDTA⁴⁻ (hexadentate), DPTA.
Chelation: when a polydentate ligand forms a ring with the metal, it is called a chelate; the ring is the chelate ring. Chelate complexes are more stable than analogous monodentate complexes (chelate effect).
Coordination number (CN): number of ligand donor atoms directly bonded to the central metal (not the number of ligands). En contributes 2, EDTA contributes 6.
Coordination sphere: the central atom plus its ligands enclosed in square brackets; it is treated as a single unit and does not ionise.
Counter ions: the ions outside the square brackets that balance the charge.
IUPAC Nomenclature Rules
1. Cation is named before anion (as in simple salts).
2. Within the coordination entity: ligands first (alphabetically), then the metal.
3. Anionic ligands: add -o (e.g. chloro, nitro, oxalato). Neutral ligands: keep name (except aqua for H₂O, ammine for NH₃, carbonyl for CO, nitrosyl for NO).
4. Number of ligands: di, tri, tetra… (or bis, tris, tetrakis for bidentate/ambidentate ligands to avoid confusion).
5. Oxidation state of metal in Roman numerals in parentheses: e.g. cobalt(III), iron(II).
6. Anionic complexes: metal name ends in -ate (ferrate, cuprate, plumbate, argentate).
Examples:
- [Cr(NH₃)₄Cl₂]Cl → tetraamminedichloridochromium(III) chloride
- K₂[PtCl₆] → potassium hexachloridoplatinate(IV)
- [Co(en)₂Cl₂]⁺ → dichloridobis(ethylenediamine)cobalt(III)
Types of Isomerism
Structural isomers (same formula, different connectivity):
1. Linkage isomerism: an ambidentate ligand (like NO₂⁻/ONO⁻ or SCN⁻/NCS⁻) bonds through different atoms. Example: [Co(NH₃)₅(NO₂)]²⁺ vs [Co(NH₃)₅(ONO)]²⁺.
2. Ionisation isomerism: ligand and counter ion swap positions. Example: [Co(NH₃)₅Br]SO₄ vs [Co(NH₃)₅SO₄]Br.
3. Coordination isomerism: in a salt where both cation and anion are complex ions, ligands are distributed differently. Example: [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].
4. Solvate (hydrate) isomerism: water is inside or outside the coordination sphere. Example: [Cr(H₂O)₆]Cl₃ (violet) vs [Cr(H₂O)₅Cl]Cl₂·H₂O (grey-green).
Stereo isomers (same connectivity, different arrangement in space):
1. Geometric (cis-trans) isomerism: found in square planar and octahedral complexes.
- Square planar MA₂B₂: cis has like ligands on adjacent corners; trans on opposite corners.
- Octahedral MA₄B₂: cis has the two B ligands at 90°; trans at 180°.
2. Optical isomerism: the complex is non-superimposable on its mirror image (chiral). The two forms are called enantiomers (d and l). Common example: [Co(en)₃]³⁺ which exists as d and l forms.
Bonding Theories
Valence Bond Theory (VBT)
The metal ion provides empty orbitals; ligand lone pairs occupy these orbitals forming coordinate covalent bonds. The type of hybridisation determines geometry and magnetic character:
Geometry CN Hybridisation Example
Octahedral (inner sphere) 6 d²sp³ [Co(NH₃)₆]³⁺
Octahedral (outer sphere) 6 sp³d² [CoF₆]³⁻
Square planar 4 dsp² [Ni(CN)₄]²⁻, [PtCl₄]²⁻
Tetrahedral 4 sp³ [NiCl₄]²⁻
Inner sphere (low spin): d orbitals rearrange to accommodate ligands — means there are unpaired electrons pushed together, fewer remain. Strong field ligands produce inner-sphere complexes → diamagnetic or less paramagnetic.
Outer sphere (high spin): d electrons are not disturbed, all available electrons remain unpaired → more paramagnetic. Weak field ligands produce outer-sphere complexes.
Limitation of VBT: cannot explain colour of complexes or the different stabilities of cis and trans isomers.
Crystal Field Theory (CFT)
Treats ligands as point negative charges (or dipoles) that create an electrostatic field around the metal. This field splits the five degenerate d orbitals into sets of different energy.
Octahedral field: d orbitals split into two sets:
- eg (d{x²-y²} and d{z²}): point directly at ligands → higher energy (raised by 6Dq).
- t₂g (d{xy}, d{yz}, d{xz}): point between ligands → lower energy (lowered by 4Dq).
The energy gap is Δo (crystal field splitting energy). Its magnitude depends on the ligand (the spectrochemical series):
I⁻ < Br⁻ < S²⁻ < SCN⁻ < Cl⁻ < NO₃⁻ < F⁻ < OH⁻ < ox < H₂O < EDTA < NH₃ < en < CN⁻ < CO
Ligands on the left are weak field (small Δo, high-spin complexes); on the right are strong field (large Δo, low-spin complexes).
Colour explained: when Δo falls in the visible range, a photon is absorbed to promote an electron from t₂g to eg. The complementary colour is seen. If Δo is 0 (no d electrons, or all d orbitals full), the compound is colourless.
Tetrahedral field: only 4 ligands → smaller splitting (Δt ≈ 4/9 Δo) and the sets are inverted (e below t₂). Tetrahedral complexes are almost always high spin.
Stability of Complexes
Stability is expressed as the formation (stability) constant Kf:
M + nL ⇌ [MLₙ] Kf = [[MLₙ]] / [M][L]ⁿ
A larger Kf means greater stability. Factors that increase stability:
- High charge density on the metal ion (high charge, small radius).
- Ligands matching the metal's size and charge.
- Chelate effect: polydentate ligands form thermodynamically more stable complexes because of the positive entropy contribution when multiple monodentate ligands are displaced.
Applications
1. Metallurgy: cyanide complexes extract gold and silver from ores (hydrometallurgy).
2. Analytical chemistry: EDTA complexes quantify metal ions by complexometric titration.
3. Medicine: [Pt(NH₃)₂Cl₂] (cisplatin) — an anticancer drug; cis form is active, trans is not. EDTA treats lead and mercury poisoning (chelation therapy).
4. Photography: sodium thiosulphate dissolves unexposed AgBr as [Ag(S₂O₃)₂]³⁻.
5. Biological systems: haemoglobin (Fe²⁺ in porphyrin ring), chlorophyll (Mg²⁺), vitamin B₁₂ (Co³⁺).
Important Reactions
Reaction Equation
Complex formation (general) M^n+ + x L → [MLx]^(n−x)
Ag in photography (fixing) AgBr + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaBr
Gold extraction 4Au + 8NaCN + O₂ + 2H₂O → 4Na[Au(CN)₂] + 4NaOH
Cisplatin formation [Pt(NH₃)₄]²⁺ + 2Cl⁻ → [Pt(NH₃)₂Cl₂] + 2NH₃
Important Question Patterns
1. Nomenclature (2–3 marks): name a given formula or write the formula for a name; follow the rules step by step.
2. Isomerism (2–3 marks): identify the type; draw/describe cis-trans or enantiomers; explain linkage isomers.
3. C
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