Coordination Compounds Class 12 Notes - CBSE Chemistry Chapter 9

Table of Contents

• Key Concepts - Werner's theory, ligands, nomenclature, isomerism, VBT, CFT, stability, uses
• Weightage in Board & Entrance Exams
• Important Definitions
• Solved Examples
• Important Questions for Board Exams
• Quick Revision Points

Key Concepts

1. What Is a Coordination Compound?

Take a bottle of copper sulphate and add ammonia. The pale blue solution turns deep royal blue because a new species, [Cu(NH₃)₃]²⁺, has formed. That coloured ion is a coordination compound - a compound in which a central metal atom or ion is bonded to a fixed number of surrounding ions or molecules by coordinate (dative) bonds.

These molecules or ions do not lose their identity even in solution. That is the key difference from a double salt. A double salt such as Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O, breaks up completely into all its ions in water (Fe²⁺, NH₄⁺, SO₄²⁻). A complex salt such as K₄[Fe(CN)₆] gives K⁺ and the intact complex ion [Fe(CN)₆]⁴⁻, which does not give the test for Fe²⁺ or CN⁻ ions.

Coordination compounds are everywhere: the haemoglobin that carries oxygen in your blood, the chlorophyll that lets plants make food, and the medicines and catalysts used in industry are all complexes.

2. Werner's Coordination Theory

Alfred Werner (Nobel Prize, 1913) was the first to explain the bonding and shapes of these compounds. His main ideas are still used today.

• Metals show two kinds of valency. The primary valency is ionisable and is satisfied by negative ions - it is basically the oxidation state of the metal. The secondary valency is non-ionisable and is satisfied by ligands - it is the coordination number.
• The primary valency is non-directional, but the secondary valency is directional: it fixes the geometry (shape) of the complex.
• Every metal tries to satisfy both its primary and its secondary valencies.

For example, in CoCl₃·6NH₃ (that is, [Co(NH₃)₆]Cl₃), cobalt has a primary valency of 3 (satisfied by three Cl⁻) and a secondary valency of 6 (satisfied by six NH₃ ligands). All three chloride ions are ionisable, so adding AgNO₃ precipitates all three as AgCl. In CoCl₃·5NH₃ ([Co(NH₃)₅Cl]Cl₂), only two chlorides are ionisable, so only two are precipitated - exactly as Werner predicted.

3. Important Terms You Must Know

• Central atom / ion: the metal atom or ion (usually a transition metal) to which the ligands are bonded, for example Fe²⁺ in [Fe(CN)₆]⁴⁻. It behaves as a Lewis acid (electron-pair acceptor).
• Ligand: the ion or molecule bound to the central atom. It has at least one lone pair to donate, so it is a Lewis base (for example NH₃, H₂O, Cl⁻, CN⁻).
• Donor atom: the exact atom of the ligand that donates the lone pair (the N in NH₃, the C in CN⁻).
• Coordination number (CN): the number of donor atoms directly bonded to the central metal - that is, the number of coordinate bonds. In [PtCl₆]²⁻ the CN is 6; in [Ni(CN)₄]²⁻ it is 4. Note: for a bidentate ligand each donor atom counts separately, so [Co(en)₃]³⁺ has CN 6, not 3.
• Coordination sphere: the central metal plus its ligands, written together inside square brackets. Everything inside the bracket is non-ionisable; the counter ions written outside are ionisable.
• Coordination polyhedron: the spatial shape formed by the ligands around the metal (octahedral, tetrahedral, square planar, and so on).
• Oxidation number of the metal: the charge the central atom would carry if all ligands were removed with their lone pairs. In K₄[Fe(CN)₆], iron is +2.

4. Ligands and Their Denticity

Ligands are classified by how many donor atoms they use to grip the metal - this is called denticity.

• Monodentate: one donor atom (Cl⁻, H₂O, NH₃, CN⁻, CO).
• Bidentate: two donor atoms, for example ethane-1,2-diamine (en, two N atoms) and oxalate (ox, two O atoms).
• Polydentate: several donor atoms. EDTA⁴⁻ is hexadentate - it grips the metal with two N and four O atoms.
• Chelate ligand: a bi- or polydentate ligand that forms a ring with the metal (a chelate). Chelate rings make a complex much more stable - the chelate effect.
• Ambidentate ligand: a monodentate ligand that can bind through either of two different donor atoms, for example NO₂⁻ (through N as nitro, or through O as nitrito) and SCN⁻ (through S as thiocyanato, or through N as isothiocyanato).

Ligands are also classified by charge: neutral (H₂O, NH₃, CO), anionic (Cl⁻, CN⁻, OH⁻), or cationic (NO⁺, nitrosonium).

5. Coordination Number and Shape

The coordination number decides the shape of the complex. The three shapes you must know for the board and NEET are shown below.

• CN 2 → linear, for example [Ag(NH₃)₂]⁺ and [CuCl₂]⁻.
• CN 4 → tetrahedral or square planar. Tetrahedral: [NiCl₄]²⁻, [Ni(CO)₄]. Square planar: [Ni(CN)₄]²⁻, [Pt(NH₃)₂Cl₂]. Square planar is common for d⁸ metals such as Ni²⁺, Pd²⁺, Pt²⁺.
• CN 6 → octahedral, the most common of all, for example [Co(NH₃)₆]³⁺ and [Fe(CN)₆]⁴⁻.

6. IUPAC Nomenclature of Coordination Compounds

Naming complexes looks scary but follows a short set of rules. Learn the rules once and you can name any complex.

• Cation is named first, then the anion - just like any ionic compound.
• Within the complex ion, ligands are named first (in alphabetical order), then the metal.
• Anionic ligands end in ‑o (chlorido, cyanido, hydroxido, oxalato); neutral ligands keep their name, with special names aqua (H₂O), ammine (NH₃), carbonyl (CO) and nitrosyl (NO).
• The number of each ligand is shown by di-, tri-, tetra- for simple ligands, and by bis-, tris-, tetrakis- for ligands that already contain a prefix (like ethane-1,2-diamine → bis(ethane-1,2-diamine)).
• The oxidation state of the metal is written in Roman numerals in brackets right after its name.
• If the complex ion is an anion, the metal name ends in ‑ate (ferrate, cuprate, cobaltate, argentate, plumbate).

7. Isomerism in Coordination Compounds

Two complexes can have the same formula yet a different arrangement of atoms - they are isomers. Isomerism is split into structural isomerism (different bonds) and stereoisomerism (same bonds, different arrangement in space).

Geometrical isomerism is shown by square planar MA₂B₂ complexes and by octahedral complexes. In the cis form the like ligands are next to each other; in the trans form they are opposite.

This is not just theory. The cis isomer of [Pt(NH₃)₂Cl₂], called cisplatin, is a powerful anti-cancer drug, while the trans isomer is medically useless. The shape decides the biology.

Optical isomerism appears when a complex has no plane of symmetry, so it exists as two mirror-image forms that rotate plane-polarised light in opposite directions. Complexes with bidentate ligands such as [Co(en)₃]³⁺ and cis-[Co(en)₂Cl₂]⁺ are classic optical isomers.

8. Valence Bond Theory (VBT) and Hybridisation

Linus Pauling's Valence Bond Theory explains bonding as follows: the ligand lone pairs are donated into empty hybrid orbitals of the metal ion. The type of hybridisation fixes the geometry.

For octahedral complexes there are two possibilities:

• Inner-orbital (low-spin) complex: a strong-field ligand forces the d electrons to pair up, freeing inner (n−1)d orbitals. Hybridisation is d²sp³. Fewer unpaired electrons → usually diamagnetic or weakly paramagnetic. Example: [Co(NH₃)₆]³⁺ is d²sp³, diamagnetic.
• Outer-orbital (high-spin) complex: a weak-field ligand does not pair the electrons, so the metal uses outer nd orbitals. Hybridisation is sp³d². More unpaired electrons → more paramagnetic. Example: [CoF₆]³⁻ is sp³d², paramagnetic (4 unpaired electrons).

Similarly, for CN 4: [Ni(CN)₄]²⁻ is dsp² (square planar, diamagnetic), while [NiCl₄]²⁻ is sp³ (tetrahedral, paramagnetic). Magnetic moment µ = √[n(n+2)] BM, where n is the number of unpaired electrons, lets you predict or check these structures.

VBT has limits: it explains geometry and magnetism but does not explain the colour of complexes or give a quantitative order of ligand strength. For that we need Crystal Field Theory.

9. Crystal Field Theory (CFT)

Crystal Field Theory treats the metal-ligand bond as purely electrostatic (ionic). The ligands are seen as point negative charges that repel the metal's d electrons. In a free metal ion the five d orbitals are degenerate (equal energy), but when ligands approach, the d orbitals that point towards the ligands are raised in energy more than those that point between them. The d orbitals therefore split.

In an octahedral field the d orbitals split into a lower set of three (t2g: dₓₙ, dₙₔ, dₓₔ type) and an upper set of two (eg: dₓ²−ₙ², dₔ² type). The energy gap is the crystal field splitting energy, Δₒ. The t2g set is lowered by 0.4Δₒ and the eg set is raised by 0.6Δₒ relative to the average (barycentre). In a tetrahedral field the order is reversed (e lower, t2 higher) and the split is smaller, Δₜ ≈ (4/9)Δₒ, so tetrahedral complexes are almost always high spin.

The size of Δₒ depends on the ligand. Arranging ligands by increasing field strength gives the spectrochemical series:

I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO

• Weak-field ligands (I⁻, Br⁻, Cl⁻, F⁻, H₂O) give a small Δₒ. If Δₒ is less than the electron pairing energy P, electrons stay unpaired and spread into eg → high-spin complex.
• Strong-field ligands (CN⁻, CO, NH₃, en) give a large Δₒ. If Δₒ is greater than P, electrons pair up in the lower t2g set before filling eg → low-spin complex.

Origin of colour: most transition-metal complexes are coloured because of a d-d transition. When white light falls on the complex, an electron in the t2g set absorbs the exact energy Δₒ and jumps to the eg set. The colour we see is the complementary colour of the light absorbed. This is why [Ti(H₂O)₆]³⁺ is purple. Complexes of ions with no d electrons (Sc³⁺, d⁰) or a full d subshell (Zn²⁺, d¹⁰) cannot undergo a d-d transition and are colourless.

10. Stability of Coordination Compounds

The stability of a complex in solution is measured by its stability constant (formation constant), K₌. A large K₌ means the complex forms readily and does not fall apart. Stability rises with:

• a higher charge and smaller size of the central metal ion (greater charge density);
• a higher basicity (donor strength) of the ligand - CN⁻ and NH₃ form more stable complexes than Cl⁻ or H₂O;
• the chelate effect - complexes with chelating (bidentate/polydentate) ligands such as en and EDTA are far more stable than those with the same number of monodentate ligands.

11. Importance and Applications

• In biology: haemoglobin (an iron-porphyrin complex) carries oxygen in blood; chlorophyll (a magnesium complex) drives photosynthesis; vitamin B₁₂ is a cobalt complex.
• In medicine: cisplatin and related platinum complexes treat cancer; EDTA is used to treat lead poisoning by chelating the Pb²⁺ ions.
• In metallurgy: silver and gold are extracted as their cyanido complexes [Ag(CN)₂]⁻ and [Au(CN)₂]⁻; nickel is purified through [Ni(CO)₄] in the Mond process.
• In analysis: EDTA is used to estimate the hardness of water (Ca²⁺ and Mg²⁺); many metal ions are detected as coloured complexes.
• In industry: Wilkinson's catalyst and many other complexes speed up important reactions; complexes are used in electroplating, dyes and photography.

Weightage in Board & Entrance Exams

Important Definitions

• Coordination compound: a compound in which a central metal is bonded to ligands by coordinate bonds and which retains its identity in solution.
• Ligand: an ion or molecule with a lone pair that binds to the central metal (a Lewis base).
• Coordination number: the number of donor atoms directly bonded to the central metal.
• Chelate: a ring complex formed when a bidentate or polydentate ligand binds the same metal.
• Ambidentate ligand: a monodentate ligand that can attach through either of two different donor atoms (e.g. NO₂⁻, SCN⁻).
• Crystal field splitting energy (Δ): the energy gap between the split d-orbital sets in a complex.
• Spectrochemical series: ligands arranged in order of increasing crystal field strength.
• Stability constant (K₌): the equilibrium constant for the formation of a complex from the metal ion and its ligands.

Solved Examples

Example 1

Q. Write the IUPAC name of K₃[Fe(CN)₆] and give the oxidation state and coordination number of iron.

A. The complex ion is an anion, so iron becomes ferrate. Six cyanido ligands and three K⁺ outside mean the ion is [Fe(CN)₆]³⁻, so iron is +3. Name: Potassium hexacyanidoferrate(III). Coordination number = 6.

Example 2

Q. [Co(NH₃)₆]³⁺ is diamagnetic while [CoF₆]³⁻ is paramagnetic. Explain using VBT.

A. Co³⁺ is d⁶. NH₃ is a strong-field ligand, so it pairs the electrons, freeing two inner 3d orbitals → d²sp³ (inner orbital), no unpaired electrons → diamagnetic. F⁻ is a weak-field ligand, so it cannot pair them; the metal uses outer 4d orbitals → sp³d² (outer orbital) with 4 unpaired electrons → paramagnetic.

Example 3

Q. Calculate the spin-only magnetic moment of [NiCl₄]²⁻.

A. Ni²⁺ is d⁸. Cl⁻ is a weak-field ligand giving a tetrahedral sp³ complex with 2 unpaired electrons. µ = √[n(n+2)] = √[2(2+2)] = √8 = 2.83 BM.

Example 4

Q. How many ionisable chloride ions are present in [Co(NH₃)₅Cl]Cl₂, and how many moles of AgCl will 1 mole of it give with excess AgNO₃?

A. Only the two Cl⁻ outside the square bracket are ionisable; the one inside is bonded to cobalt. So it gives 2 moles of AgCl.

Example 5

Q. Why is [Ti(H₂O)₆]³⁺ coloured but [Sc(H₂O)₆]³⁺ colourless?

A. Ti³⁺ is d¹: its single d electron can absorb light and undergo a t2g → eg (d-d) transition, so the complex is purple. Sc³⁺ is d⁰: it has no d electron to promote, so no d-d transition and no colour.

Example 6

Q. Give one example each of linkage, ionisation and coordination isomerism.

A. Linkage: [Co(NH₃)₅(NO₂)]²⁺ vs [Co(NH₃)₅(ONO)]²⁺. Ionisation: [Co(NH₃)₅SO₄]Br vs [Co(NH₃)₅Br]SO₄. Coordination: [Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆].

Important Questions for Board Exams

1-Mark Questions (VSA)

• What is a ligand? Give one example of a bidentate ligand.
• Define coordination number. State the CN of cobalt in [Co(en)₃]³⁺.
• What is an ambidentate ligand? Give one example.
• Why is [Zn(NH₃)₄]²⁺ colourless?

2–3-Mark Questions (SA)

• State the postulates of Werner's coordination theory.
• Write the IUPAC names of [Cr(H₂O)₆]Cl₃ and [Pt(NH₃)₂Cl₂].
• Explain the difference between a double salt and a complex salt with one example each.
• Using VBT, explain why [Ni(CN)₄]²⁻ is square planar and diamagnetic while [NiCl₄]²⁻ is tetrahedral and paramagnetic.

5-Mark Questions (LA)

• Discuss crystal field splitting in octahedral complexes. Explain high-spin and low-spin complexes using the spectrochemical series.
• Describe the types of isomerism shown by coordination compounds with one example of each.
• What is the crystal field theory explanation for the colour of transition-metal complexes?

Quick Revision Points

• Complex salts keep their identity in solution (K₄[Fe(CN)₆]); double salts break up fully (Mohr's salt).
• Werner: primary valency = oxidation state (ionisable); secondary valency = coordination number (directional, fixes shape).
• Coordination number counts donor atoms; a bidentate ligand contributes 2 (so [Co(en)₃]³⁺ has CN 6).
• Naming: ligands alphabetically then metal; anionic complex → metal name ends in ‑ate; oxidation state in Roman numerals.
• Structural isomers: ionisation, hydrate, linkage, coordination. Stereoisomers: geometrical (cis/trans) and optical (d/l).
• VBT: d²sp³ = inner orbital / low spin; sp³d² = outer orbital / high spin; dsp² = square planar; sp³ = tetrahedral.
• µ = √[n(n+2)] BM gives the spin-only magnetic moment from the number of unpaired electrons.
• CFT: octahedral splits d into lower t2g and upper eg by Δₒ (t2g −0.4Δₒ, eg +0.6Δₒ); tetrahedral is inverted, Δₜ ≈ 4/9 Δₒ.
• Spectrochemical series (weak → strong): I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < en < CN⁻ < CO. Strong field → low spin.
• Colour comes from a d-d transition; d⁰ and d¹⁰ ions are colourless.
• Chelate complexes (en, EDTA) are extra stable - the chelate effect. Haemoglobin (Fe) and chlorophyll (Mg) are natural complexes.