Coordination Compounds Notes for NEET, JEE Main 2027: Key Points, PDF

Coordination Compounds is Unit 5 of NCERT Class 12 Chemistry (Part I) and Unit 12 of both the NEET UG and JEE Main syllabi. These notes cover Werner’s theory, ligands and denticity, IUPAC nomenclature, structural and stereo isomerism, valence bond theory, crystal field theory, colour, magnetic properties, metal carbonyls and applications, with a free notes PDF.

Coordination Compounds Notes for NEET and JEE Main

Section-wise key points from NCERT Unit 5, comparison tables, high-yield facts and common mistakes for NEET and JEE Main 2027 revision.

Coordination Compounds at a Glance

The topic list follows the NEET UG 2026 syllabus released by the National Medical Commission (NMC) on 22 December 2025 and NTA’s JEE Main 2026 syllabus, which stay the reference until the 2027 syllabi are out. The full NEET list is on the NEET syllabus page.

What is Werner’s theory of coordination compounds?

Werner’s theory (1898) says metals in coordination compounds show two kinds of valence: an ionisable primary valence and a non-ionisable secondary valence that equals the coordination number.

• Alfred Werner (1866-1919) was a Swiss chemist. He became the first Swiss chemist to win the Nobel Prize, in 1913, for this work.
• Primary valences are normally ionisable and are satisfied by negative ions (today: the oxidation state).
• Secondary valences are non-ionisable and are satisfied by neutral molecules or negative ions. The secondary valence equals the coordination number and is fixed for a metal.
• Groups bound by secondary linkages have characteristic spatial arrangements, now called coordination polyhedra. Octahedral, tetrahedral and square planar shapes are the most common.
• Species inside the square bracket form the coordination entity; ions outside are counter ions.

Double salt or complex?

• Double salts such as carnallite (KCl·MgCl2·6H2O), Mohr’s salt (FeSO4·(NH4)2SO4·6H2O) and potash alum (KAl(SO4)2·12H2O) dissociate completely into simple ions in water.
• Complex ions such as [Fe(CN)6]4– in K4[Fe(CN)6] do not dissociate into Fe2+ and CN– ions.

Key Terms: Ligand, Denticity and Coordination Number

A ligand is an ion or molecule bound to the central atom or ion; the central atom acts as a Lewis acid.

How are coordination compounds named?

Names follow IUPAC additive nomenclature: ligands are named in alphabetical order before the metal, and the metal’s oxidation state follows in Roman numerals.

Writing Formulas

• Central atom first, then ligands in alphabetical order, whatever their charge.
• The whole coordination entity goes in square brackets; polyatomic ligands and ligand abbreviations go in parentheses.
• No space between the ligands and the metal. The charge of an entity written alone goes outside the bracket, e.g. [Co(CN)6]3–.

Writing Names

• The cation is named first, whether the complex part is the cation or the anion.
• Anionic ligands end in -o (chlorido, cyanido, hydroxido, oxalato). Neutral ligands keep their names except aqua (H2O), ammine (NH3), carbonyl (CO) and nitrosyl (NO).
• Prefixes di, tri, tetra give the number of each ligand. If the ligand name already has a numerical prefix, use bis, tris, tetrakis with the ligand in parentheses: [NiCl2(PPh3)2] is dichloridobis(triphenylphosphine)nickel(II).
• In a complex anion the metal name ends in -ate (cobaltate); some use Latin names, e.g. ferrate for Fe.
• A neutral complex is named like a complex cation.

What types of isomerism do coordination compounds show?

Coordination compounds show stereoisomerism (geometrical and optical) and structural isomerism (linkage, coordination, ionisation and solvate).

• Tetrahedral complexes do not show geometrical isomerism, because all ligand positions are equivalent relative to each other.
• Optical isomerism is common in octahedral complexes with didentate ligands.
• Optical isomers rotate plane-polarised light in opposite directions: dextro (d) to the right, laevo (l) to the left.

What does valence bond theory explain?

Valence bond theory says the metal uses hybridised (n–1)d, ns, np or ns, np, nd orbitals to accept ligand electron pairs, and the hybridisation decides the shape.

Magnetic Properties

• Magnetic susceptibility gives the number of unpaired electrons and hence the likely structure.
• Ti3+ (d1), V3+ (d2) and Cr3+ (d3) leave two d orbitals free for octahedral hybridisation, so complexes behave like the free ions.
• d4 to d6 ions complicate things: [Mn(CN)6]3– has two unpaired electrons but [MnCl6]3– has four; [Fe(CN)6]3– has one but [FeF6]3– has five; [CoF6]3– has four but [Co(C2O4)3]3– is diamagnetic.
• The cyanide and oxalate complexes above are inner orbital (d2sp3); the chloride and fluoride complexes are outer orbital (sp3d2).
• NCERT example: spin-only magnetic moment 5.9 BM for [MnBr4]2– means five unpaired electrons, so it is tetrahedral, not square planar.

Limitations of Valence Bond Theory

• It involves many assumptions and gives no quantitative interpretation of magnetic data.
• It does not explain colour, or the thermodynamic and kinetic stabilities of complexes.
• It cannot predict exactly whether a 4-coordinate complex is tetrahedral or square planar, and does not distinguish weak and strong ligands.

Crystal Field Theory: Splitting, Spin and Colour

Crystal field theory is an electrostatic model: ligands act as point charges or point dipoles and lift the degeneracy of the metal’s five d orbitals.

Octahedral and Tetrahedral Splitting

• In an octahedral field the d orbitals split into a lower t2g set and a higher eg set. The eg orbitals rise by (3/5)Δo and the t2g orbitals fall by (2/5)Δo.
• Δo depends on the field produced by the ligand and the charge on the metal ion.
• If Δo < P (pairing energy): weak field ligand, high spin complex, for example t2g3 eg1 for d4.
• If Δo > P: strong field ligand, low spin complex, for example t2g4 eg0 for d4.
• In a tetrahedral field the splitting is inverted and smaller: Δt = (4/9)Δo, so low spin tetrahedral complexes are rarely observed.
• The ‘g’ subscript is used only for octahedral and square planar complexes, which have a centre of symmetry.

Spectrochemical series (increasing field strength): I– < Br– < SCN– < Cl– < S2– < F– < OH– < C2O4 2– < H2O < NCS– < edta4– < NH3 < en < CN– < CO.

Colour

• Colour comes from a d-d transition: the complex absorbs some visible light and shows the complementary colour.
• [Ti(H2O)6]3+ absorbs blue-green light (about 498 nm) and appears violet.
• Without ligands there is no splitting and no colour: heating [Ti(H2O)6]Cl3 to remove water makes it colourless; anhydrous CuSO4 is white but CuSO4·5H2O is blue.
• Adding en to [Ni(H2O)6]2+ (green) changes the colour step by step: [Ni(H2O)4(en)]2+ pale blue, [Ni(H2O)2(en)2]2+ blue/purple, [Ni(en)3]2+ violet.
• Ruby is Al2O3 with about 0.5-1% Cr3+ ions; emerald has Cr3+ in octahedral sites of beryl. Both colours come from d-d transitions.

Limitations of Crystal Field Theory

• Treating ligands as point charges predicts that anionic ligands should split most, yet they lie at the low end of the spectrochemical series.
• It ignores the covalent character of metal-ligand bonding.

Bonding in Metal Carbonyls

Metal carbonyls are held by a synergic bond: a sigma bond from carbon to the metal plus a pi back bond from a filled metal d orbital into the empty antibonding pi* orbital of CO.

• The metal-carbon bond has both sigma and pi character; the synergic effect strengthens the bond between CO and the metal.

Why are coordination compounds important?

Coordination compounds are used in chemical analysis, metallurgy, biological systems, industry and medicine.

Most-asked Facts for NEET and JEE Main: Coordination Compounds

Revise these one-line facts from the NCERT chapter just before the exam.

• Primary valence = ionisable (oxidation state); secondary valence = coordination number, non-ionisable.
• CoCl3·6NH3, CoCl3·5NH3 and CoCl3·4NH3 give 3, 2 and 1 mol AgCl respectively.
• EDTA4– is hexadentate; en and oxalate are didentate.
• NO2– and SCN– are ambidentate and cause linkage isomerism.
• Coordination number counts only sigma bonds from donor atoms.
• Ligand names: aqua (H2O), ammine (NH3), carbonyl (CO), nitrosyl (NO); anionic ligands end in -o.
• Tetrahedral complexes show no geometrical isomerism.
• fac/mer isomers occur in octahedral [Ma3b3], e.g. [Co(NH3)3(NO2)3].
• Only cis-[PtCl2(en)2]2+ is optically active.
• [Co(NH3)6]3+: d2sp3, diamagnetic; [CoF6]3–: sp3d2, paramagnetic (4 unpaired).
• [NiCl4]2–: sp3, paramagnetic; [Ni(CN)4]2–: dsp2, square planar, diamagnetic.
• Δt = (4/9)Δo; tetrahedral complexes are almost always high spin.
• Δo > P gives low spin (strong field); Δo < P gives high spin (weak field).
• Spectrochemical series ends: I– weakest, CO strongest.
• [Ti(H2O)6]3+ is violet; anhydrous CuSO4 is white, CuSO4·5H2O blue.
• Synergic bonding in metal carbonyls: sigma donation plus pi back bonding.
• Wilkinson catalyst [(Ph3P)3RhCl] hydrogenates alkenes; cis-platin is an anticancer complex.
• Vitamin B12 contains cobalt, chlorophyll magnesium, haemoglobin iron.

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this chapter.

• Counting pi bonds when finding the coordination number. Only sigma bonds from donor atoms count.
• Counting ligands instead of donor atoms: [Co(en)3]3+ has three ligands but coordination number 6.
• Ordering ligands by charge. Both formulas and names list ligands in alphabetical order, whatever their charge.
• Writing ‘amine’ for NH3 as a ligand. The ligand name is ammine, with two m’s.
• Forgetting the -ate ending for a metal in a complex anion (ferrate, cobaltate, zincate).
• Expecting geometrical isomers for tetrahedral complexes.
• Assuming a strong field ligand always gives low spin in tetrahedral complexes. Δt is too small, so low spin tetrahedral complexes are rare.
• Saying the complex shows the colour it absorbs. It shows the complementary colour.
• Taking the anionic ligands as the strongest field ligands. Halides lie at the weak end of the spectrochemical series.
• Confusing double salts with complexes. Mohr’s salt and potash alum dissociate fully in water; K4[Fe(CN)6] gives no free Fe2+ or CN– ions.

Coordination Compounds Notes PDF and Related Pages

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• Some Basic Concepts of Chemistry
• Structure of Atom
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• Organic Chemistry: Some Basic Principles and Techniques
• Electrochemistry
• Amines
• Solutions
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• Coordination Compounds
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