Coordination CompoundsNEET Chemistry · Class 12 · NCERT Chapter 5

Introduction to Coordination Compounds

A coordination compound is a compound in which a central metal atom (or ion) is bonded to a fixed number of ions or neutral molecules called ligands. These bonds are coordinate (dative) bonds — the ligand donates a lone pair of electrons to the metal.

Coordination compounds are widespread: haemoglobin (Fe-porphyrin), chlorophyll (Mg-porphyrin), vitamin B12 (Co complex), cisplatin (Pt anticancer drug), EDTA in food preservation, and cyanide extraction of gold are all coordination chemistry in action.

The study of coordination compounds was systematised by Alfred Werner (1913 Nobel Prize in Chemistry). His key insight was the idea of primary and secondary valence, which is the foundation of modern coordination chemistry.

Terminology and Types of Ligands

Key Terms

• Central metal atom/ion: the metal at the centre of the complex that accepts electron pairs. It is typically a transition metal.
• Ligand: a molecule or ion that donates one or more lone pairs of electrons to the central metal. Ligands must have at least one lone pair available for donation.
• Coordination number (CN): the total number of ligand atoms directly bonded to the central metal. Common coordination numbers: 4 (square planar or tetrahedral) and 6 (octahedral).
• Coordination sphere: the central metal + all directly bonded ligands, written inside square brackets, e.g., [Co(NH3)6]3+.
• Chelate: a complex in which a polydentate ligand forms one or more rings by bonding to the metal through two or more donor atoms.

Types of Ligands

Chelate effect: Polydentate ligands form more stable complexes than equivalent monodentate ligands. This extra stability arises because replacing multiple monodentate ligands with one polydentate ligand increases the number of free particles in solution (increases entropy). The more donor atoms a ligand has, the more stable the complex.

IUPAC Nomenclature

Rules for Naming Coordination Compounds

1. Name the cation first, then the anion (same as for ionic compounds).
2. Within the complex ion, name the ligands first in alphabetical order, then the central metal.
3. Alphabetise ligands by their names, ignoring multiplying prefixes (di, tri, tetra, bis, tris).
4. Use multiplying prefixes: di, tri, tetra for simple ligands; use bis, tris, tetrakis for ligands whose names already contain a number or for polydentate ligands (to avoid ambiguity).
5. Indicate the oxidation state of the central metal in Roman numerals in round brackets immediately after the metal name: e.g., cobalt(III).
6. For anionic complexes, add the suffix -ate to the metal name (using Latin-based names for some metals):
   • Fe → ferrate; Co → cobaltate; Cr → chromate; Cu → cuprate; Au → aurate; Ag → argentate; Pt → platinate
   • Most others use English names + -ate: nickelate, manganate, etc.

Common Ligand Names

Worked Examples

• [Co(NH3)6]3+ = hexaamminecobalt(III) ion
• [CoCl2(NH3)4]+ = tetraamminedichloridocobalt(III) ion (ammine before chlorido alphabetically: a before c; note tetra + ammine, di + chlorido)
• K3[Fe(CN)6] = potassium hexacyanidoferrate(III)
• K2[PtCl4] = potassium tetrachloridoplatinate(II)
• [Cu(en)2]2+ = bis(ethylenediamine)copper(II) ion (bis because en already has a number)
• [Fe(CO)5] = pentacarbonyliron(0)

Isomerism in Coordination Compounds

Structural Isomers (same molecular formula, different bonding)

1. Ionisation Isomerism

The ligand inside the coordination sphere and the counter ion outside are swapped. They give different ions in solution. Test with AgNO3 (precipitates Cl-) or BaCl2 (precipitates SO42-).

• [Co(NH3)5Cl]SO4 (gives SO42- in solution) vs [Co(NH3)5(SO4)]Cl (gives Cl- in solution)

2. Hydrate (Solvate) Isomerism

Water molecules can be inside the coordination sphere (bonded as aqua ligands) or outside (as water of crystallisation). The three hydrate isomers of CrCl3·6H2O:

• [Cr(H2O)6]Cl3 — violet; all 3 Cl- are free in solution
• [Cr(H2O)5Cl]Cl2·H2O — blue-green; 2 Cl- free
• [Cr(H2O)4Cl2]Cl·2H2O — dark green; 1 Cl- free

3. Linkage Isomerism

Ambidentate ligands can bond through different atoms. NO2- can bond through N (nitro) or O (nitrito).

• [Co(NH3)5(NO2)]Cl2 (nitro, M-N) vs [Co(NH3)5(ONO)]Cl2 (nitrito, M-O)

4. Coordination Isomerism

When both the cation and anion are complex ions, the distribution of ligands between the two complexes can vary.

• [Co(NH3)6][Cr(CN)6] vs [Cr(NH3)6][Co(CN)6]

Stereoisomers (same bonding, different spatial arrangement)

5. Geometric (Cis-Trans) Isomerism

Geometric isomers differ in the relative positions of identical ligands. Not possible for regular tetrahedral complexes (all positions are equivalent).

Square planar MA2B2 type (like [PtCl2(NH3)2]):

• Cis: two identical ligands at 90° (adjacent corners)
• Trans: two identical ligands at 180° (opposite corners)

Octahedral MA4B2 type (like [Co(NH3)4Cl2]+):

• Cis: two Cl are adjacent (90° apart)
• Trans: two Cl are opposite (180° apart)

Octahedral MA3B3 type (like [Co(NH3)3Cl3]):

• Fac (facial): three identical ligands on one triangular face. All A-M-A angles = 90°.
• Mer (meridional): three identical ligands span a meridian. One A-M-A angle = 180°.

6. Optical Isomerism

Optical isomers are non-superimposable mirror images (enantiomers). They rotate plane-polarised light in opposite directions (one is +/dextro, the other is -/laevo). An equal mixture of both is called a racemic mixture (optically inactive).

Conditions for optical isomerism in octahedral complexes:

• [M(en)3]n+ type (e.g., [Co(en)3]3+) — always optically active, no plane of symmetry
• cis-[M(en)2A2]n+ (e.g., cis-[Co(en)2Cl2]+) — optically active (cis isomer)
• trans-[M(en)2A2]n+ — optically inactive (has a plane of symmetry)

Bonding Theories: Werner, VBT, and CFT

Werner's Theory (1893)

Werner proposed that metals have two types of valence:

• Primary valence (ionisable valence): the oxidation state of the metal. Satisfied by negative ions (counter ions outside the coordination sphere).
• Secondary valence (non-ionisable valence): the coordination number. Satisfied by ligands directly bonded to the metal (inside the coordination sphere).

For CoCl3·6NH3 = [Co(NH3)6]Cl3: Co has primary valence 3 (oxidation state, satisfied by 3 Cl-) and secondary valence 6 (coordination number, satisfied by 6 NH3). Werner's theory was confirmed by conductivity measurements (number of ions in solution) and by precipitation of Cl- with AgNO3.

Valence Bond Theory (VBT)

VBT describes the bonding in terms of hybridisation of the metal's orbitals to accept lone pairs from ligands.

Inner orbital complex (d2sp3): strong field ligands (CN-, CO, en, NH3) force electron pairing in lower d-orbitals, freeing inner d-orbitals for hybridisation. Result: fewer unpaired electrons, lower (or zero) paramagnetism. Also called low-spin complex.

Outer orbital complex (sp3d2): weak field ligands (F-, Cl-, H2O, OH-) do not force pairing — electrons spread out, using outer d-orbitals for hybridisation. Result: more unpaired electrons, higher paramagnetism. Also called high-spin complex.

Example: [Co(NH3)6]3+ vs [CoF6]3-

• Co3+ = [Ar] 3d6. With NH3 (strong field): electrons pair up in 3d (3d6 → t2g6 eg0), freeing 3d3 and 4s, 4p for d2sp3. Result: diamagnetic (0 unpaired).
• Co3+ with F- (weak field): electrons spread across 3d5 and 4s, 4p, 4d for sp3d2. Result: 4 unpaired electrons, paramagnetic.

Crystal Field Theory (CFT)

CFT treats the metal-ligand interaction as an electrostatic interaction between the metal d-electrons and the negative charges (or dipoles) of the ligands. The key result is the splitting of the five degenerate d-orbitals into two sets:

Octahedral field:

• eg (higher energy, +0.6 Δo): dx2-y2 and dz2 — point along the x, y, z axes directly toward the ligands (maximum repulsion)
• t2g (lower energy, -0.4 Δo): dxy, dxz, dyz — point between the axes (less repulsion)

High Spin vs Low Spin

When filling d-electrons in an octahedral field, there is a competition between:

• Δo (crystal field splitting energy) — the energy cost of promoting an electron to eg
• P (pairing energy) — the energy cost of pairing two electrons in the same orbital

If Δo > P (strong field ligands): electrons pair up in t2g before going to eg. Low spin complex (fewer unpaired electrons).
If Δo < P (weak field ligands): electrons go to eg before pairing in t2g. High spin complex (more unpaired electrons).

Spectrochemical Series (weak field to strong field):

I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < CN- < CO

CFSE Calculation

CFSE = (number of t2g electrons × -0.4 Δo) + (number of eg electrons × +0.6 Δo)

Example: High spin d6 in octahedral field (t2g4 eg2): CFSE = 4(-0.4) + 2(+0.6) = -1.6 + 1.2 = -0.4 Δo

Example: Low spin d6 in octahedral field (t2g6 eg0): CFSE = 6(-0.4) + 0 = -2.4 Δo

Tetrahedral Crystal Field

In a tetrahedral field, the splitting is inverted: the e-set (dx2-y2, dz2) is now lower in energy and the t2-set (dxy, dxz, dyz) is higher. Also, Δt = (4/9) Δo. Because Δt is so small, almost all tetrahedral complexes are high spin.

Applications and Stability of Coordination Compounds

Biological Importance

• Haemoglobin: Fe2+ coordinated to a porphyrin ring (haem group) and a protein chain. Carries O2 in the blood. CO binds to the same site 200× more strongly than O2, causing CO poisoning.
• Chlorophyll: Mg2+ at the centre of a porphyrin ring. Absorbs red and blue light; reflects green. Essential for photosynthesis.
• Vitamin B12 (Cobalamin): Co3+ in a corrin ring. Essential cofactor for DNA synthesis and nerve function.

Industrial Applications

• Gold and silver extraction (cyanide process):
  
  $$4\text{Au}+8\text{NaCN}+{\text{O}}_2+2{\text{H}}_2\text{O}\to 4\text{Na}[{\text{Au(CN)}}_2]+4\text{NaOH}$$
• Nickel refining (Mond process): Ni reacts with CO at 50-60°C to form [Ni(CO)4], a volatile liquid. It decomposes at 230°C to give pure Ni.
• Electroplating: metals are plated from coordination compound solutions (e.g., gold plating from [Au(CN)2]-).

Medicinal Uses

• Cisplatin (cis-[PtCl2(NH3)2]): anticancer drug. Crosslinks guanine bases on the same DNA strand, blocking replication.
• EDTA as antidote: Ca-EDTA chelates heavy metals (Pb, Hg) and removes them from the body.

Stability of Complexes

Stability is measured by the stability constant (formation constant, Kf). The larger Kf, the more stable the complex. Key factors:

• Chelate effect: polydentate ligands give higher stability due to entropy increase.
• Nature of metal: higher charge density on metal = stronger bonds. Small, highly charged metal ions form more stable complexes.
• Nature of ligand: strong field ligands (CN-, CO) form very stable complexes (large Δo, more CFSE).

Worked NEET Problems

Summary Cheat Sheet

Key Terms Quick Reference