Coordination CompoundsNEET Chemistry · Class 12 · NCERT Chapter 5

You can expect 3 to 5 questions from this chapter in NEET every year, making it one of the highest-weightage chapters in Class 12 Chemistry. The topics tested most are IUPAC naming (especially getting the ligand names and order right), identifying types of isomerism, determining hybridisation (d2sp3 or sp3d2), and crystal field theory (high spin vs low spin, colour). Practise IUPAC naming with real examples until it becomes automatic.

A chelate is a complex formed when a polydentate ligand (one with two or more donor atoms) binds to a single metal ion through multiple points simultaneously, forming a ring structure. For example, ethylenediamine (en) has two NH2 groups that can both bond to the metal, forming a 5-membered ring. Chelating complexes are more stable than those with an equivalent number of monodentate ligands — this extra stability is called the chelate effect. The reason is entropy: replacing two monodentate ligands (e.g., 2NH3) with one bidentate ligand (en) increases the number of free particles in solution, which increases entropy (ΔS > 0) and makes the reaction more spontaneous (more negative ΔG).

Follow these steps in order: (1) Name the cation before the anion (just like for ionic compounds). (2) Within the complex ion, name the ligands first in alphabetical order (ignore multiplying prefixes like di, tri when alphabetising). (3) Then name the central metal atom, followed by its oxidation state in Roman numerals in brackets — for example, chromium(III). (4) For anionic complexes, add the suffix -ate to the metal name, e.g., cobaltate, ferrate, platinate. (5) Common ligand name changes: F- = fluoro, Cl- = chlorido, CN- = cyanido, NH3 = ammine, H2O = aqua, CO = carbonyl, NO = nitrosyl, OH- = hydroxido, en = ethylenediamine. Example: [Co(NH3)4Cl2]+ is tetraamminedichloridocobalt(III) ion.

Geometric isomers (cis-trans isomers) differ in the spatial arrangement of ligands around the metal. In a square planar complex like [PtCl2(NH3)2], cis has two Cl on the same side, trans has them on opposite sides. In octahedral complexes, cis/trans isomerism occurs for MA4B2 type. Geometric isomers are not mirror images of each other and cannot be interconverted without breaking bonds. Optical isomers (enantiomers) are non-superimposable mirror images of each other. They rotate plane-polarised light in opposite directions. Octahedral complexes with 3 bidentate ligands like [Co(en)3]3+ or cis-[Co(en)2Cl2]+ show optical isomerism. A mixture of equal amounts of both optical isomers is called a racemic mixture.

In an isolated transition metal ion, all five d-orbitals are degenerate (same energy). When ligands approach the metal in an octahedral arrangement, they repel the d-electrons. The d-orbitals split into two sets: a lower-energy set t2g (dxy, dxz, dyz, pointing between the axes) and a higher-energy set eg (dx2-y2, dz2, pointing along the axes directly at the ligands). The energy gap between these sets is called crystal field splitting energy, Δo. When white light hits the complex, electrons absorb photons with exactly the energy Δo and jump from t2g to eg. The colour you see is the complementary colour. For example, if violet light is absorbed, the complex appears yellow-green.

In octahedral complexes, the d-electrons can be arranged in two ways depending on the balance between Δo (the crystal field splitting energy) and P (the pairing energy, the energy cost of putting two electrons in the same orbital). If Δo is small (weak field ligands like F-, Cl-, H2O), it is easier for electrons to go to eg than to pair up. This gives a high-spin complex (more unpaired electrons, more paramagnetic). If Δo is large (strong field ligands like CN-, CO, en, NH3), it is cheaper to pair up in t2g than to promote to eg. This gives a low-spin complex (fewer unpaired electrons, less paramagnetic or diamagnetic). The spectrochemical series orders ligands from weak to strong field: I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < CN- < CO.

The sum of the charges on the central metal and all ligands must equal the overall charge of the complex. Steps: (1) Note the total charge of the complex (positive, negative, or neutral). (2) Assign known charges to all ligands (e.g., Cl- = -1, CN- = -1, NH3 = 0, H2O = 0, en = 0, NO2- = -1, ox2- = -2). (3) Let the metal oxidation state = x, and solve: x + sum of ligand charges = complex charge. Example: [Fe(CN)6]4- — total charge = -4; six CN- each = -1, total ligand charge = -6; so x + (-6) = -4; x = +2. Fe is +2.

The EAN rule (proposed by Sidgwick) states that stable metal complexes tend to have the metal surrounded by enough electrons from ligands to reach the electron count of the nearest noble gas. EAN = atomic number of metal - electrons lost (oxidation state) + electrons donated by all ligands. Each ligand donates 2 electrons (one lone pair). Example: [Fe(CO)5] — Fe is Z=26, no charge lost (Fe0), five CO each donate 2 electrons = 10. EAN = 26 - 0 + 10 = 36 (same as Kr). For [Ni(CO)4] — Ni Z=28, Ni0, four CO donate 8 electrons: EAN = 28 + 8 = 36 (Kr). The EAN rule works well for metal carbonyls but has exceptions for other complexes.