Coordination Chemistry FAQs
NEET & JEE | 25 Most Asked Questions
These Coordination Chemistry FAQs cover the most asked NEET and JEE questions about crystal field theory, magnetic moment, hybridisation, colour of transition elements, CFSE, strong and weak ligands, geometrical isomerism, and coordination compounds.
01 Why are transition metals coloured? ▶
Transition metals have partially filled d-orbitals. Ligands split these d-orbitals into two energy levels (crystal field splitting). When light falls on the complex, electrons absorb a specific wavelength to jump between these levels, and the complementary colour is seen. The colour observed = complementary colour of the absorbed light.
02 Why is Zn²⁺ colourless? ▶
Zn²⁺ has a d¹⁰ configuration — all d-orbitals are completely filled. Since no d–d electronic transition is possible, no visible light is absorbed, and the ion appears colourless.
03 Why is Cu²⁺ blue? ▶
Cu²⁺ has a d⁹ configuration. When hydrated ([Cu(H₂O)₆]²⁺), water causes crystal field splitting. The electron absorbs red/orange light (~600–700 nm) for d–d transition, and the complementary blue colour is transmitted, making the solution appear blue.
04 Why is KMnO₄ purple? ▶
KMnO₄ has Mn in +7 state with d⁰ configuration — no d–d transition is possible. Its intense purple colour arises from charge transfer (CT) transitions, where an electron transfers from O²⁻ to Mn⁷⁺. CT transitions are much more intense than d–d transitions.
05 Why are d⁰ and d¹⁰ ions colourless? ▶
d⁰ ions (e.g. Ti⁴⁺, Sc³⁺) have no electrons to excite; d¹⁰ ions (e.g. Zn²⁺, Cu⁺) have no empty d-orbitals to receive them. In both cases, d–d electronic transitions are impossible, so no visible light is absorbed and the complex is colourless.
06 Strong ligand vs weak ligand ▶
Strong ligands (e.g. CN⁻, CO, en) cause large crystal field splitting (Δ), forcing electrons to pair up → low spin, diamagnetic complexes. Weak ligands (e.g. Cl⁻, F⁻, H₂O) cause small Δ, leaving electrons unpaired → high spin, paramagnetic complexes. Strong field ligands appear at the top of the spectrochemical series.
07 What is the spectrochemical series? ▶
It is the arrangement of ligands in increasing order of crystal field splitting (Δ):
I⁻ < Br⁻ < S²⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO. Ligands on the right cause large splitting (strong field); on the left, small splitting (weak field).
08 How to find hybridisation in complexes? ▶
Count the coordination number (CN): CN = 6 → sp³d² (outer) or d²sp³ (inner); CN = 4 → sp³ (tetrahedral) or dsp² (square planar); CN = 2 → sp (linear). If the ligand is strong and forces pairing using inner d-orbitals → d²sp³; if weak and uses outer d-orbitals → sp³d². Square planar is almost always dsp².
09 Inner orbital vs outer orbital complexes ▶
Inner orbital complexes use (n–1)d orbitals for hybridisation (d²sp³) — formed with strong ligands, low spin, less paramagnetic or diamagnetic. Outer orbital complexes use nd orbitals (sp³d²) — formed with weak ligands, high spin, more paramagnetic. Example: [Fe(CN)₆]⁴⁻ is inner orbital; [FeF₆]³⁻ is outer orbital.
10 Why is [Ni(CN)₄]²⁻ square planar? ▶
Ni²⁺ has d⁸ configuration. CN⁻ is a strong field ligand that forces the two unpaired electrons to pair up, vacating one d-orbital. This allows dsp² hybridisation → square planar geometry. The complex is diamagnetic with 0 unpaired electrons.
11 Why is [NiCl₄]²⁻ tetrahedral? ▶
Cl⁻ is a weak field ligand; it cannot force electron pairing in Ni²⁺ (d⁸). The electrons remain unpaired, so inner d-orbital is unavailable for dsp² hybridisation. Instead, sp³ hybridisation occurs → tetrahedral geometry. The complex is paramagnetic with 2 unpaired electrons.
12 High spin vs low spin complexes ▶
High spin: weak field ligands → small Δ → electrons occupy all d-orbitals before pairing (follow Hund's rule) → more unpaired electrons → paramagnetic. Low spin: strong field ligands → large Δ → electrons pair in lower energy orbitals first → fewer/no unpaired electrons → diamagnetic or less paramagnetic.
13 Why is CO a strong ligand? ▶
CO is a strong ligand due to synergic bonding (σ-donation + π-back bonding). It donates a lone pair to the metal (σ-bond) and simultaneously accepts electron density from filled metal d-orbitals into its π* antibonding orbital. This back-bonding stabilises the complex and increases Δ, making CO the strongest known ligand.
14 What is the chelate effect? ▶
When a polydentate ligand (e.g. en, EDTA) forms a ring by coordinating at multiple sites, the resulting complex is unusually stable. This extra stability compared to analogous monodentate complexes is the chelate effect. It arises mainly from entropy: replacing several monodentate ligands with one chelating ligand releases more free particles, increasing ΔS and thus stability.
15 What is an ambidentate ligand? ▶
A ligand that can coordinate to the metal through two different donor atoms is called ambidentate. Example: NO₂⁻ can bond via N (nitro, –NO₂) or via O (nitrito, –ONO); SCN⁻ can bond via S (thiocyanato) or via N (isothiocyanato). These give rise to linkage isomerism.
16 Why is EDTA hexadentate? ▶
EDTA (ethylenediaminetetraacetic acid) has 6 donor atoms: 2 nitrogen atoms (from the two –NH groups) and 4 oxygen atoms (from the four –COOH groups after deprotonation). All 6 can donate lone pairs to the metal simultaneously, making EDTA a hexadentate ligand that forms 5 chelate rings with the metal ion.
17 Optical isomerism in coordination compounds ▶
Optical isomers (enantiomers) are non-superimposable mirror images of each other. They arise when a complex lacks a plane of symmetry. Most common in octahedral complexes with bidentate ligands: [Co(en)₃]³⁺ shows optical isomerism (Δ and Λ forms). Square planar and tetrahedral complexes rarely show it. Key rule: if a complex has a plane of symmetry, it is optically inactive (meso form).
18 Geometrical isomerism trick ▶
For square planar MA₂B₂: cis (same groups adjacent) and trans (same groups opposite) isomers exist. For octahedral MA₄B₂: cis (B's at 90°) and trans (B's at 180°). Quick trick: geometrical isomerism exists only when the complex has at least 2 different types of ligands AND they can be arranged differently. MA₆, MA₅B types show no geometrical isomerism.
19 Magnetic moment shortcut ▶
Use the spin-only formula: μ = √(n(n+2)) BM, where n = number of unpaired electrons. Quick values: n=1 → 1.73; n=2 → 2.83; n=3 → 3.87; n=4 → 4.90; n=5 → 5.92 BM. First find the oxidation state of the metal, write its d-electron configuration, apply strong/weak ligand rule, count unpaired electrons, then apply the formula.
20 Why are tetrahedral complexes always high spin? ▶
Crystal field splitting in tetrahedral complexes (Δt) is only about 4/9 of the splitting in octahedral complexes (Δo) for the same ligand and metal. This small Δt is almost always less than the pairing energy, so electrons prefer to remain unpaired in higher energy orbitals → high spin. No tetrahedral complex is known to be low spin.
21 CFSE shortcut ▶
Crystal Field Stabilisation Energy (CFSE) = energy gained when d-electrons occupy lower energy t₂g orbitals rather than eg. Each electron in t₂g = –0.4Δo; each in eg = +0.6Δo. CFSE is maximum for d³ and d⁸ (low spin) configurations. d⁰, d⁵ (high spin), and d¹⁰ have CFSE = 0, meaning no extra stabilisation from the crystal field.
22 Which complexes show Jahn-Teller distortion? ▶
Jahn-Teller distortion occurs when the eg orbitals (dx²–y² and dz²) are unequally occupied. Most pronounced in d⁹ (e.g. Cu²⁺) and high-spin d⁴ configurations, where the asymmetric eg occupancy causes elongation or compression along the z-axis. It is weak (sometimes seen) in t₂g asymmetry but strong and always observed for unequal eg occupancy.
23 Why is Fe³⁺ more stable than Fe²⁺? ▶
Fe³⁺ has a d⁵ configuration (half-filled d-orbitals) with all 5 electrons unpaired — this is an extra-stable arrangement due to symmetry and exchange energy. Fe²⁺ has d⁶ with one paired electron, which is less stable. This extra stability of the half-filled d⁵ makes Fe³⁺ more stable, despite its higher charge state.
24 Difference between double salt and complex salt ▶
Double salts (e.g. Mohr's salt, alum) dissociate completely in water to give all their constituent ions — they lose their identity in solution. Complex salts (e.g. K₄[Fe(CN)₆]) dissociate to give a complex ion ([Fe(CN)₆]⁴⁻) that retains its identity in solution and does not give free CN⁻ or Fe²⁺ ions. The complex ion is the key difference.
25 Werner theory — simple explanation ▶
Werner proposed two types of valencies for metals in complexes: (1) Primary valency (ionisable, shown by counter ions outside the coordination sphere, satisfies oxidation state) and (2) Secondary valency (non-ionisable, satisfied by ligands inside the coordination sphere, determines geometry). Example: In [Co(NH₃)₆]Cl₃ — primary valency = 3 (three Cl⁻), secondary valency = 6 (six NH₃ inside).
