Coordination Compounds — NEET Notes

Werner's theory — the breakthrough that built the field

Before 1893 no one knew what to make of compounds like CoCl₃·6NH₃. Cobalt's normal valence is three; ammonia is already saturated. Why should they combine at all, and in fixed proportions? Alfred Werner, a Swiss chemist still in his twenties, solved the puzzle and won the 1913 Nobel Prize for it. He observed that adding excess AgNO₃ to four cobalt-ammonia complexes precipitated different amounts of AgCl — three moles from CoCl₃·6NH₃, two from CoCl₃·5NH₃, one from CoCl₃·4NH₃. He concluded that some chlorides were free to ionise (precipitable) while others were locked inside a stable inner sphere (not precipitable). Conductivity measurements confirmed the same partition.

Werner proposed that every metal exercises two kinds of valence. The primary valence is the metal's oxidation number, satisfied by negative ions in the outer sphere and ionisable in solution. The secondary valence is the coordination number, satisfied by neutral molecules or anions directly bound to the metal, locked in a fixed geometry. The species inside the square brackets is the coordination entity; the ions outside are counter ions.

"Metals possess two distinct types of valence — primary, ionisable; secondary, directional. The secondary valence is fixed for a metal and determines the spatial arrangement of its complex."

Alfred Werner · Coordination theory, 1898

Werner went further. He argued that the secondary valences point in definite directions in space — what we now call coordination polyhedra. The most common are octahedral (CN 6), tetrahedral and square planar (both CN 4). The four cobalt-ammonia complexes resolve cleanly under this scheme:

Key terms — the vocabulary you cannot skip

Every line of this chapter depends on six terms. NEET tests them by demanding you identify them in unfamiliar complexes — so master the definitions cold.

Denticity — how many teeth does a ligand bite with?

The number of donor atoms a ligand uses simultaneously to bind the same metal is its denticity. A ligand that uses two or more donor atoms to lock onto a single metal forms a ring; such ligands are called chelating ligands, and the rings they form are chelate rings. Chelating ligands produce dramatically more stable complexes than equivalent monodentate ligands — this is the chelate effect, and NEET 2023 tested it directly.

IUPAC nomenclature — the four-step formula

The IUPAC rules look fiddly until you realise they reduce to four moves applied in a fixed order. Treat naming a complex like reading a recipe: cation first, anion second; inside the bracket name ligands alphabetically and then the metal; show metal's oxidation state in Roman numerals; close with the suffix -ate if the whole complex is an anion.

Ligand naming has its own micro-rules. Anionic ligands take an -o suffix: chloride → chlorido, hydroxide → hydroxido, cyanide → cyanido, sulphate → sulphato, oxalate → oxalato. Neutral ligands keep their own names with a few exceptions: H₂O is aqua, NH₃ is ammine (two m's — not the same as amine), CO is carbonyl, NO is nitrosyl. Note the alphabetical ordering ignores prefixes like di- and tri- but respects bis-/tris- in compound names.

Worked example: [Pt(NH₃)₂Cl(NO₂)]. Ligands alphabetically — ammine (×2), chlorido, nitrito-N. Sum of charges: 2(0) + (−1) + (−1) = −2; complex is neutral, so Pt = +2. Name: diamminechloridonitrito-N-platinum(II). Worked example two — K₃[Al(C₂O₄)₃]: three oxalato ligands, K is +1 each so complex is −3, ligand charges 3(−2) = −6, so Al = +3. Name: potassium trioxalatoaluminate(III). NEET 2023 used this exact complex.

Isomerism — when formula does not equal structure

Two complexes with the same molecular formula but different arrangement of atoms are isomers. NEET treats this topic as a guarantee — at least one question every year, usually testing whether you can spot the type of isomerism from a formula. Isomerism splits into two broad families: structural (different connectivity) and stereo (same connectivity, different spatial arrangement).

Structural isomerism — four sub-types

Stereoisomerism — geometry and chirality

Two complexes with identical connectivity but different spatial arrangement are stereoisomers. Two flavours appear in NEET: geometrical (cis-trans) and optical (d-l).

Geometrical (cis-trans) isomerism

This appears in square planar and octahedral complexes when the ligand arrangement allows non-equivalent positions. In a square planar complex [Ma₂b₂] the two a-ligands can be adjacent (90° apart, cis) or opposite (180° apart, trans) — for example, the famous anticancer drug cisplatin is cis-[Pt(NH₃)₂Cl₂], biologically active; its trans isomer is not. Octahedral complexes show cis-trans in [Ma₄b₂] and [M(AA)₂b₂] patterns; NEET 2018 asked about [CoCl₂(en)₂], which has exactly this kind of cis-trans pair.

Optical (d-l) isomerism

Octahedral complexes can be chiral when they lack any plane of symmetry. Classic case: [Co(en)₃]³⁺, where three bidentate ethylenediamine ligands wrap around the cobalt as a three-bladed propeller. The two propellers cannot be superimposed — they are mirror images (enantiomers), rotating plane-polarised light in opposite directions (d for dextrorotatory, l for levorotatory). The cis isomer of [CoCl₂(en)₂]⁺ is also chiral; the trans isomer has a plane of symmetry and is not.

Valence Bond Theory — hybridisation, geometry, magnetism

Linus Pauling adapted his hybridisation framework to coordination compounds in the 1930s. The metal provides empty hybrid orbitals; each ligand donates an electron pair into one of those orbitals to form a coordinate covalent bond. The geometry of the complex is dictated by the hybridisation, and the number of unpaired electrons left in non-hybrid d orbitals determines its magnetic behaviour. Four hybridisations dominate NEET-level questions:

Magnetic moment is the experimental fingerprint of VBT's predictions. For n unpaired electrons, the spin-only formula is μ = √n(n+2) BM. With one unpaired electron, μ = 1.73 BM; two electrons, 2.83 BM; three, 3.87 BM; four, 4.90 BM; five, 5.92 BM. NEET 2021 asked you to match four iron(II/III) complexes to their magnetic moments — and the entire question collapses once you know which ligand is strong-field (CN⁻ pairs electrons) versus weak-field (H₂O does not).

Crystal Field Theory — splitting, spin state, colour

VBT explains geometry and magnetism but stays silent on colour. Crystal Field Theory (Bethe, 1929; van Vleck, 1932) fills that gap. CFT treats the metal-ligand interaction as purely electrostatic: ligands are point negative charges that approach the metal d orbitals and raise the energy of those d orbitals which point directly at them. The originally five-fold degenerate d orbitals split into two sets, separated by an energy gap Δ that is the crystal field splitting energy.

Octahedral field — t₂g and eg

In an octahedral field, six ligands sit along the ±x, ±y, ±z axes. They point directly at the d_x²-y² and d_z² orbitals, raising their energy. The other three orbitals (d_xy, d_yz, d_xz) point between the axes and are lowered. The lower set of three is labelled t₂g and the upper set of two is eg. The gap between them is Δ₀ (sometimes written 10 Dq). Relative to the centre of gravity (barycentre), t₂g is stabilised by 0.4 Δ₀ each and eg destabilised by 0.6 Δ₀ each — this gives Crystal Field Stabilisation Energy (CFSE).

Tetrahedral field — e and t₂

In a tetrahedral field the four ligands approach between the axes; the splitting pattern is inverted (e lower, t₂ upper) and the gap is much smaller. The mathematical result is Δ_t = (4/9) Δ₀ for the same metal and ligand. Because Δ_t is almost always smaller than the electron-pairing energy, tetrahedral complexes are nearly always high-spin; low-spin tetrahedral complexes are extraordinarily rare.

Strong-field vs weak-field ligands — the spectrochemical series

The magnitude of Δ₀ depends on both metal and ligand. Empirically, ligands can be ordered by the size of the splitting they produce — the spectrochemical series:

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

Ligands on the left (halides, water) are weak-field; they produce small Δ₀, electrons stay unpaired, giving high-spin complexes. Ligands on the right (CN⁻, CO) are strong-field; they produce large Δ₀ that exceeds the pairing energy, forcing electrons to pair in t₂g and giving low-spin complexes. NEET 2020 asked the order SCN⁻ < F⁻ < C₂O₄²⁻ < CN⁻, and NEET 2017 asked the corresponding wavelength order (Δ₀ ∝ 1/λ, so stronger ligands absorb at shorter wavelengths).

Why complexes are coloured

The splitting Δ₀ corresponds to energies that fall in the visible region of the spectrum. A d-electron can absorb a photon and jump from t₂g to eg — this is a d-d transition. The complex absorbs that particular wavelength and transmits its complement. [Ti(H₂O)₆]³⁺ absorbs green and looks purple; [Cu(H₂O)₄]²⁺ absorbs red and looks blue. d⁰ ions (Sc³⁺, Ti⁴⁺) and d¹⁰ ions (Zn²⁺, Cu⁺) are colourless because no d-d transition is possible.

Metal carbonyls — the synergic bond

Compounds with carbon monoxide as the only ligand are called metal carbonyls. They span a remarkable range — from the colourless gas Ni(CO)₄ to deep-red Fe(CO)₅, dark-green Cr(CO)₆, and yellow-orange Mn₂(CO)₁₀. They obey the eighteen-electron rule and are central to organometallic chemistry, catalysis, and the Mond process for nickel refining.

The metal-carbonyl bond is unusual in that it has two reinforcing components — a synergic pair:

1. σ-donation: the lone pair on the carbon atom of CO donates into an empty σ-symmetry orbital on the metal. Carbon is the donor, not oxygen.
2. π-back-donation: a filled d orbital of the metal donates electron density into the empty π* antibonding orbital of CO. The metal is now the donor.

The two donations reinforce each other — each one increases the electron density that drives the other — hence the term synergic bonding. The net effect is a strong M–C bond and a weakened C–O bond. The more negative the metal centre, the more it can back-donate, and the longer the C–O bond becomes. NEET 2016 tested this directly: in the series Ni(CO)₄, [Mn(CO)₆]⁺, [Co(CO)₄]⁻, [Fe(CO)₄]²⁻, the most negative complex [Fe(CO)₄]²⁻ has the longest C–O bond.

M ← C≡O (σ donation)  +  M → C≡O (π* back-donation)

Synergic bond — both donations strengthen the M–C bond and weaken the C–O bond

Geometries follow the eighteen-electron rule: Ni(CO)₄ is tetrahedral (Ni in zero oxidation state, d¹⁰, sp³, diamagnetic — NEET 2018), Fe(CO)₅ is trigonal bipyramidal (mononuclear — NEET 2018), Cr(CO)₆ is octahedral. Polynuclear carbonyls like Mn₂(CO)₁₀ feature direct M–M bonds in addition to terminal and bridging CO ligands.

Applications — biology, industry, medicine

NCERT's emphasis on applications is not decorative. Two PYQ-relevant facts: most enzymes that handle redox chemistry use a coordination complex at their active site, and EDTA is used as the standard titrant for water-hardness analysis.