Coordination Compounds

Inorganic Chemistry Weightage: 2-3 Questions (8-12 Marks) JEE Unit 12

“Welcome back! You already know the NCERT basics of double salts vs coordination compounds and simple naming. This guide stays on the JEE Main syllabus line: Werner's theory, ligands, coordination number and denticity, IUPAC nomenclature, isomerism, Valence Bond Theory and Crystal Field Theory (d-orbital splitting, magnetic properties). Almost every question here reduces to just THREE moves — identify the metal's oxidation state and d-electron count, decide whether the ligand set is weak-field or strong-field, then read off the geometry, hybridisation, magnetic moment and colour from that one decision. Learn to make that one call correctly and the rest becomes arithmetic.”

— SCORECHEM ACADEMIC TEAM

1. Double Salts, Complexes and Key Terminology

Double salt vs coordination compound: a double salt (e.g. carnallite KCl.MgCl₂.6H₂O, Mohr's salt FeSO₄.(NH₄)₂SO₄.6H₂O, potash alum K₂SO₄.Al₂(SO₄)₃.24H₂O) loses its identity in water, dissociating completely into its simple constituent ions. A coordination compound (e.g. K₄[Fe(CN)₆]) retains its complex ion's identity in solution — the metal-ligand bonds inside the coordination sphere do NOT break apart.
Coordination entity/sphere: the metal atom/ion + its directly bonded ligands, written inside square brackets, e.g. [Cu(NH₃)₄]²⁺ in [Cu(NH₃)₄]SO₄. Counter ions sit outside the bracket and ARE freely dissociable/ionisable.
Central atom/ion: a Lewis ACID (electron-pair acceptor). Ligands: Lewis BASES (electron-pair donors) — each donor atom supplies a lone pair to form a coordinate (dative) bond.
Coordination number (C.N.): the number of σ-bonds (ligand donor atoms directly bonded to the metal) — count σ-bonds ONLY, never π-bonds, even for a polydentate ligand contributing more than one donor atom.
Coordination polyhedron: the spatial arrangement of the donor atoms around the metal (e.g. octahedral for C.N. 6, tetrahedral or square planar for C.N. 4).
Homoleptic vs heteroleptic: homoleptic = all ligands identical (e.g. [Co(NH₃)₆]³⁺); heteroleptic = more than one kind of ligand (e.g. [Co(NH₃)₄Cl₂]⁺).

⚠️ JEE Trap: A double salt and a coordination compound can have the SAME empirical formula pattern. The test is behaviour in solution, not the formula's appearance — does it dissociate all the way to simple ions (double salt), or does a complex ion survive intact (coordination compound)?

2. Ligands, Denticity and the Chelate Effect

Denticity = number of donor atoms from ONE ligand bonded to the SAME metal. Monodentate (Cl^-, NH₃, H₂O, CO), bidentate (ethane-1,2-diamine "en", oxalate C₂O₄^2-), polydentate (EDTA^4-, normally hexadentate).
Chelate ligand: a poly/bidentate ligand that binds the same metal at 2+ points, forming a ring (a "chelate"). Chelate effect: chelate complexes are markedly MORE stable than the equivalent number of monodentate complexes — e.g. K_f for [Ni(H₂O)₆]²⁺ + 6NH₃ ⇌ [Ni(NH₃)₆]²⁺ is ~10⁸, but with 3 "en" (chelating, same total donor atoms) it jumps to ~10¹⁸. 5- and 6-membered chelate rings are the most stable.
Flexidentate ligand: a ligand that does NOT always use its maximum possible denticity — EDTA^4- is normally hexadentate but can bind as penta- or tetradentate depending on the metal and complex; SO₄^2- can be mono- or bidentate.
Ambidentate ligand: has TWO different possible donor atoms, but binds through only ONE at a time in any given complex — NO₂^- (binds via N = nitrito-N, or via O = nitrito-O/ONO), SCN^- (via S = thiocyanato, or via N = isothiocyanato/NCS), CN^- (via C, normal, or via N = isocyanide). Binding through a different donor atom of the SAME ambidentate ligand gives linkage isomers (Section 7).

⚠️ JEE Trap: "Bidentate" and "ambidentate" are opposite ideas — don't confuse them. A bidentate ligand donates through TWO atoms AT ONCE (e.g. en, ox^2-). An ambidentate ligand has two possible donor atoms but uses only ONE at a time (e.g. SCN^-, NO₂^-).

3. IUPAC Nomenclature

Order: cation named first, then anion (as in any ionic compound). Within the complex ion: ligands are named in ALPHABETICAL order (ignoring multiplying prefixes like di-, tri-), followed by the metal name.
Multiplying prefixes: simple ligands use di-, tri-, tetra-... (dichlorido). Ligands with their OWN complex/substituted names (en, PPh₃, a numerical prefix already in the name) use bis-, tris-, tetrakis-... with the ligand name in parentheses, e.g. bis(ethane-1,2-diamine).
Anionic complex: metal name takes the suffix -ate (ferrate, cobaltate, cuprate for Fe/Co/Cu; most others just add -ate, e.g. manganate, chromate, zincate). Neutral or cationic complexes use the normal element name.
Oxidation state: the metal's oxidation number is written in Roman numerals in parentheses immediately after the metal name, e.g. cobalt(III).
Bridging ligands: prefixed with μ- (or μ₂-, μ₃- for higher bridging), and if more than one bridging ligand of the same kind, di-μ-, etc.
Common ligand names to know: Cl^- = chlorido, CN^- = cyanido, OH^- = hydroxido, O^2- = oxido, H₂O = aqua, NH₃ = ammine, CO = carbonyl, NO = nitrosyl, NO₂^-(via N) = nitrito-N, ONO^-(via O) = nitrito-O.
Worked examples: [CoCl₂(en)₂]⁺ = dichloridobis(ethane-1,2-diamine)cobalt(III) ion. K₃[Fe(CN)₆] = potassium hexacyanidoferrate(III).

4. Werner's Theory and the EAN Rule

Fig. 1: Coordination number stays 6 throughout — only the split between PRIMARY (ionisable) and SECONDARY (non-ionisable, bound) valency changes as NH3 is replaced by Cl.

Werner's postulates: metals show two types of valency — primary valency (ionisable, satisfied by negative ions, non-directional, equals oxidation state) and secondary valency (non-ionisable, satisfied by ligands occupying a FIXED spatial position, equals coordination number). The classic CoCl₃.nNH₃ series (above) shows coordination number staying fixed at 6 while the ionisable-Cl count (moles of AgCl precipitated per mole of complex) falls as NH₃ is replaced by Cl inside the coordination sphere.
Reading a Werner-series question: moles of AgCl formed = number of Cl^- OUTSIDE the coordination sphere (free/ionisable). The remaining Cl^- (total 3, minus ionisable) sit INSIDE the bracket, coordinated to the metal, and don't react with AgNO₃.
EAN (Effective Atomic Number) rule: EAN = Z (atomic number of metal) − oxidation state + 2 × (coordination number). Many stable complexes have EAN equal to the NEAREST noble gas (36, 54 or 86) — e.g. [Co(NH₃)₆]³⁺: 27−3+12=36. Metal carbonyls (Ni(CO)₄, Fe(CO)₅, Cr(CO)₆) obey the EAN rule rigorously. It is a guideline, not a law: [Fe(CN)₆]^3- (EAN=35), [Co(CN)₆]^4- (EAN=37) and [Ni(NH₃)₆]²⁺ (EAN=38) are all perfectly stable despite NOT hitting a noble-gas EAN.

⚠️ JEE Trap: A Werner-series numerical almost always hides TWO unknowns — the number of AgCl moles tells you ionisable Cl^-, then charge balance (using the REMAINING coordinated Cl^- and a fixed coordination number, usually 6) gives you the metal's oxidation state. Solve for ionisable ligands FIRST from the precipitate data, then use charge balance for the oxidation state.

5. Valence Bond Theory

The core VBT picture: the metal ion hybridises a set of empty orbitals to accept a lone pair from each ligand donor atom, forming coordinate covalent bonds. The hybridisation used depends on how many, and which, orbitals are needed for the observed geometry.
Inner vs outer orbital complex: if a STRONG field ligand causes the metal's own (n−1)d electrons to pair up, the resulting empty inner (n−1)d orbitals get used → inner orbital complex (e.g. d²sp³, low-spin, more stable/kinetically inert). If a WEAK field ligand does NOT force pairing, the metal must reach further out to empty outer nd orbitals → outer orbital complex (e.g. sp³d², high-spin, more labile).
Magnetic behaviour from VBT: count unpaired electrons remaining AFTER the hybridisation-driven pairing (if any), then μ = √[n(n+2)] BM.
Pt(II)/Pd(II)/Ni(II)-with-strong-ligand: almost always square planar, dsp², diamagnetic (d⁸, strong field pairs the two unpaired electrons, leaving one d-orbital free for the hybrid set). Ni(II) octahedral complexes MUST be outer-orbital, sp³d², paramagnetic — d⁸ can never give a fully-paired octahedral (d²sp³ is impossible for d⁸, since only 2 inner d-orbitals would need to be free but d⁸ can free at most... none without an odd geometry), so any octahedral Ni(II) complex is automatically high-spin outer-orbital.

Fig. 2: Outer-orbital (ns,np,nd) = weak-field ligand, high-spin. Inner-orbital ((n-1)d,ns,np) = strong-field ligand, low-spin, more stable.

⚠️ JEE Trap: VBT tells you geometry and spin state but CANNOT explain colour. VBT has no concept of split d-orbital energy levels, so it cannot account for d-d transitions. That is exactly why Crystal Field Theory (Section 6) is needed alongside it.

6. Crystal Field Theory, Colour and Stability

Fig. 3: Ligands left of H2O are weak field (favour high-spin); ligands from NH3 onward are strong field (favour low-spin) in octahedral complexes.

Octahedral splitting: the metal's 5 degenerate d-orbitals split into a lower t_2g set (d_xy, d_xz, d_yz; energy −0.4Δ_o each) and a higher e_g set (d_z², d_x²-y²; energy +0.6Δ_o each), because e_g orbitals point directly AT the approaching ligands (more repulsion) while t_2g orbitals point BETWEEN them.
CFSE (Crystal Field Stabilisation Energy) = (−0.4 × n_t2g + 0.6 × n_eg)Δ_o. Bigger magnitude = more stabilised complex.
High spin vs low spin (only matters for d⁴–d⁷ in octahedral fields): if Δ_o < pairing energy P (weak field ligand), electrons spread out to BOTH sets first (Hund's rule wins) → high spin. If Δ_o > P (strong field ligand), electrons fill t_2g completely before entering e_g → low spin. For d¹–d³ and d⁸–d¹⁰, there is only ONE possible filling — no high/low spin distinction exists.
Tetrahedral splitting is INVERTED and smaller: e (lower) and t₂ (higher) sets, Δ_t = (4/9)Δ_o for the same metal/ligands — always too small to overcome pairing energy, so tetrahedral complexes are ALWAYS high-spin. Square planar splitting is the most spread out: Δ_sp ≈ 1.3Δ_o.
Colour: arises from a d-d transition — an electron absorbing a specific visible wavelength to jump from t_2g to e_g (or vice versa). The absorbed colour and the OBSERVED colour are complementary (e.g. a complex absorbing red light appears green). d⁰ and d¹⁰ are always colourless (no electron to promote, or no vacancy to promote into). High-spin d⁵ (e.g. Mn²⁺) is only very faintly coloured because the transition is spin-forbidden.

Fig. 6: Colour needs BOTH a d-electron to promote AND a vacant d-orbital to promote it into — only d1-d9 (partially filled) configurations qualify.

Stability constant: β_n = K₁×K₂×...×K_n for successive ligand addition. Irving-Williams order of M²⁺ stability for 3d metals: Mn^II < Fe^II < Co^II < Ni^II < Cu^II > Zn^II (peaks at Cu, then falls at Zn).

⚠️ JEE Trap: Co³⁺ (d⁶) is the classic exception — it forms low-spin complexes with almost EVERY ligand, even weak-field H₂O, not just strong-field ones. Only an unusually weak ligand like F^- ([CoF₆]^3-, paramagnetic, sp³d²) forces Co³⁺ into the high-spin state. Don't apply the generic "H₂O = weak field = high spin" rule blindly to Co³⁺.

7. Isomerism

Fig. 4: Structural isomers differ in WHAT is bonded to the metal; stereoisomers differ only in the 3-D ARRANGEMENT of identical ligand sets.

Ionisation isomerism: a ligand and a counter ion swap roles — [Co(NH₃)₅SO₄]Br vs [Co(NH₃)₅Br]SO₄. Each gives a DIFFERENT precipitate test (AgNO₃ only precipitates the free/ionic halide; BaCl₂ only precipitates the free/ionic sulfate).
Hydrate/solvate isomerism: water (or another solvent) swaps between being coordinated and being lattice/free — CrCl₃.6H₂O exists as three real forms: [Cr(H₂O)₆]Cl₃ (violet, all 6 H₂O coordinated, all 3 Cl^- ionic), [Cr(H₂O)₅Cl]Cl₂.H₂O (blue-green, 1:2 electrolyte), [Cr(H₂O)₄Cl₂]Cl.2H₂O (dark green, 1:1 electrolyte).
Linkage isomerism: an AMBIDENTATE ligand binds through a different donor atom — [Co(NO₂)(NH₃)₅]Cl₂ (nitrito-N) vs [Co(ONO)(NH₃)₅]Cl₂ (nitrito-O).
Coordination isomerism: in a salt with BOTH cation and anion being complex ions, ligands redistribute between the two metals — [Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆].
Geometrical isomerism (needs C.N. 4 square planar, or C.N. 6 octahedral — NEVER tetrahedral):
- Square planar [Ma₂b₂]: 2 isomers (cis, trans). [Ma₂bc]: 2 isomers. [Mabcd]: 3 isomers. [M(AB)₂] (unsymmetrical bidentate): 2 isomers. [M(AA)₂]: 0 (only 1 arrangement).
- Octahedral [Ma₄b₂]: 2 (cis, trans). [Ma₄bc]: 2. [Ma₃b₃]: 2 (fac, 3 identical ligands on one triangular face; mer, meridional, 3 in one plane). [M(AA)₂a₂]: 2 (cis, trans). [Mabcdef] (6 different ligands): 15.
Optical isomerism: requires the complex to have NO plane of symmetry and NO centre of symmetry (i.e. it's chiral, non-superimposable on its mirror image). [M(AA)₃]: both enantiomers chiral/active. [M(AA)₂a₂]: ONLY the cis isomer is optically active (no symmetry plane); the trans isomer has a mirror plane, so it is inactive. Square planar complexes are almost never optically active (very rare exceptions exist). Tetrahedral [Mabcd] (4 different groups) CAN be optically active, even though it shows no geometrical isomerism.

Fig. 5: Square-planar [Ma2b2] gives cis/trans (no optical activity); octahedral [M(AA)2a2] gives cis (chiral) and trans (achiral, has a mirror plane).

⚠️ JEE Trap: "[M(AA)₂a₂] has 3 geometrical isomers" is a very common wrong answer. It has exactly 2 (cis, trans) — students often confuse it with the 3-isomer count that only applies to a square-planar [Mabcd]-type complex with 4 different unidentate ligands.

8. How JEE Frames Coordination-Compound Questions

"Given atomic number Z, find n and μ": always work Z → ground-state configuration → ion's configuration (remove ns electrons before (n−1)d) → decide weak/strong field & geometry → count unpaired electrons → μ=√[n(n+2)]. Never skip straight to a memorised μ value.
Werner-series/AgNO₃ numericals: moles of AgCl = ionisable Cl^- only; use that plus charge balance (coordination number usually fixed at 6) to back out the metal's oxidation state and the ligand split.
Statement/assertion-reason pairs on isomer counts: re-derive the isomer count from the general formula type ([Ma₂b₂], [M(AA)₂a₂], [Ma₃b₃], etc.) rather than recalling a number — the trap is almost always an off-by-one count.
Mixture + selective precipitation questions (ionisation isomers): identify which ion is COORDINATED (inert to that specific precipitating reagent) and which is the free counter ion (reacts) in EACH isomer separately before doing any mole arithmetic.
Hybridisation/geometry/magnetic-moment multi-statement questions: for each species, independently fix oxidation state → d-count → ligand field strength → spin state → hybridisation → μ, rather than trusting a pattern across the list.
Colour/wavelength ordering questions: rank complexes by ligand field strength (spectrochemical series) to get the CFSE/Δ_o order; wavelength absorbed is INVERSELY related to that (stronger field → larger Δ_o → higher energy → SHORTER wavelength absorbed).

⚠️ JEE Trap: When a question gives you an atomic number instead of just naming the ion, it wants YOU to derive the configuration from scratch. Treat every "Z = ..." clue as the required first step, not optional decoration.

9. Quick Sheet and Checklist

Idea	Rule
Coordination number	Counts σ-bonds (donor atoms) only, never π-bonds
Ambidentate vs bidentate	Ambidentate = 2 possible donors, uses ONE at a time. Bidentate = 2 donors used AT ONCE
CFSE (octahedral)	(-0.4 n_t2g + 0.6 n_eg)Δo
High spin vs low spin	Only possible for d4-d7 octahedral; weak field=high spin, strong field=low spin
Tetrahedral	Always high spin (Δt = 4/9 Δo, too small); NEVER geometrical isomerism
Colourless configs	d0, d10 only
[M(AA)2a2] isomers	2 total: cis (optically active) + trans (inactive)
[Ma3b3] isomers	2 total: fac + mer
Co3+ exception	Low-spin with almost all ligands, even H2O; only F- forces high-spin
Ni(II) octahedral	Always outer-orbital, sp3d2, paramagnetic (d8 can't be inner-orbital octahedral)

Before the exam, check you can:

Distinguish a double salt from a coordination compound using solution behaviour, not just the formula.
Name a given complex by IUPAC rules (ligand order, bis/tris vs di/tri, -ate suffix, Roman-numeral oxidation state) and reverse the process.
Reconstruct any CoCl₃.nNH₃-style Werner series numerical from AgCl data and charge balance.
Predict hybridisation, geometry and magnetic moment for a given ion + ligand set using both VBT and CFT.
Rank a set of complexes by CFSE/Δ_o and correctly invert that into a wavelength-absorbed order.
Count geometrical and optical isomers correctly for square planar and octahedral general formula types ([Ma2b2], [Ma3b3], [M(AA)2a2], [Mabcd], etc.).
State the Co³⁺ low-spin exception and the Ni(II)-octahedral-is-always-outer-orbital rule.