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Class 12 Chemistry Notes Chapter 9 (Coordination compounds) – Chemistry-I Book

Philoid

13 Jul 2025 — 10 min read

Chemistry-I
Alright class, let's begin our detailed study of Coordination Compounds. This chapter is crucial not only for your board exams but also forms a significant part of the syllabus for various government competitive examinations. Pay close attention to the definitions, theories, and examples.

Chapter 9: Coordination Compounds - Detailed Notes

1. Introduction

Coordination Compounds: These are compounds in which a central metal atom or ion is linked to a number of ions or neutral molecules (ligands) by coordinate bonds (dative bonds). These compounds retain their identity in the solid state as well as when dissolved in water or any other solvent.
- Example: K4[Fe(CN)6] dissolves in water to give K+ ions and the complex ion [Fe(CN)6]4-, but it does not give Fe2+ or CN- ions.
Double Salts vs. Coordination Compounds:
- Double Salts: Exist only in the solid state and dissociate completely into their constituent ions when dissolved in water.
  - Example: Mohr's salt, FeSO4·(NH4)2SO4·6H2O, dissociates into Fe2+, NH4+, and SO42- ions in water.
- Coordination Compounds: Retain their complex ion identity in solution.

2. Werner's Theory of Coordination Compounds (Key Postulates)

Metals exhibit two types of valencies:
- Primary Valency (Ionisation Valency): Corresponds to the oxidation state of the central metal ion. It is satisfied by negative ions and is ionisable. Represented by dotted lines.
- Secondary Valency (Coordination Number): Corresponds to the coordination number of the central metal ion. It is satisfied by neutral molecules or negative ions (ligands). It is non-ionisable and directional, determining the geometry of the complex. Represented by solid lines.
Every metal tends to satisfy both its primary and secondary valencies.
Secondary valencies are directed towards fixed positions in space, leading to definite geometries (e.g., tetrahedral, square planar, octahedral).
- Example: In [Co(NH3)6]Cl3:
  - Central metal: Co
  - Primary valency: 3 (satisfied by 3 Cl- ions) - Oxidation state of Co is +3.
  - Secondary valency: 6 (satisfied by 6 NH3 molecules) - Coordination number is 6. Geometry is octahedral.

3. Important Terminology

Coordination Entity: The central metal atom/ion bonded to a fixed number of ligands. Enclosed in square brackets [ ].
- Examples: [Co(NH3)6]3+, [Ni(CO)4], [Fe(CN)6]4-
Central Atom/Ion: The metal atom or ion to which the ligands are attached. It acts as a Lewis acid (electron pair acceptor).
Ligands: The ions or molecules bound to the central atom/ion in the coordination entity. They act as Lewis bases (electron pair donors).
- Types of Ligands:
  - Unidentate: Ligands with only one donor atom. Examples: H2O, NH3, Cl-, CN-, CO.
  - Bidentate: Ligands with two donor atoms. Examples: Ethylenediamine (en, H2NCH2CH2NH2), Oxalate ion (ox, C2O42-).
  - Polydentate: Ligands with several donor atoms. Example: Ethylenediaminetetraacetate ion (EDTA4-), which is hexadentate.
  - Chelating Ligands: Di- or polydentate ligands that bind to a central metal ion through two or more donor atoms, forming a ring structure (chelate). Chelation increases the stability of the complex (Chelate Effect).
  - Ambidentate Ligands: Unidentate ligands that can coordinate through more than one site. Examples: NO2- (can coordinate through N as nitro or O as nitrito), SCN- (can coordinate through S as thiocyanato or N as isothiocyanato).
Coordination Number (CN): The total number of coordinate bonds formed by the central metal atom/ion with the ligands. It is determined only by the number of sigma bonds formed by the ligand with the central metal atom/ion.
- Examples: In [Co(NH3)6]3+, CN = 6. In [Ni(CO)4], CN = 4. In [Fe(C2O4)3]3-, CN = 3 × 2 = 6 (since oxalate is bidentate).
Coordination Sphere: The central atom/ion and the ligands attached to it, enclosed in square brackets [ ]. The part outside the bracket is the counter ion/ionisation sphere.
Oxidation Number (ON): The charge the central atom would carry if all the ligands were removed along with the electron pairs shared with the central atom. Calculated by considering the charge on the complex ion and the known charges of the ligands.
- Example: In [Fe(CN)6]4-, let ON of Fe be x. Then, x + 6(-1) = -4 => x = +2.
Homoleptic Complexes: Complexes in which the metal is bound to only one kind of donor group. Example: [Co(NH3)6]3+.
Heteroleptic Complexes: Complexes in which the metal is bound to more than one kind of donor group. Example: [Co(NH3)4Cl2]+.

4. Nomenclature of Coordination Compounds (IUPAC Rules)

Order of Naming Ions: Cation is named first, followed by the anion (like in simple salts).
Naming the Coordination Sphere:
- Ligands: Named first, in alphabetical order, before the central metal atom/ion.
  - Anionic ligands end in '-o' (e.g., Cl- chloro, CN- cyanido, SO42- sulfato, NO2- nitrito-N or nitro, ONO- nitrito-O or nitrito). Note: Older convention used '-o' for all anions (e.g., Cl- chlorido), but NCERT often uses chloro, cyano etc. Be aware of both, but follow NCERT/exam guidelines. Modern IUPAC prefers 'chlorido', 'cyanido'. Let's stick to NCERT's common usage for simplicity here: chloro, cyano, nitrito-N, nitrito-O.
  - Neutral ligands retain their usual names (e.g., ethylenediamine), with exceptions: H2O (aqua), NH3 (ammine), CO (carbonyl), NO (nitrosyl).
  - Cationic ligands (rare) end in '-ium' (e.g., NH2-NH3+ hydrazinium).
- Prefixes: Use di-, tri-, tetra-, etc., to indicate the number of individual ligands. For ligands whose names already contain a numerical prefix (like ethylenediamine), use bis-, tris-, tetrakis-, etc., and enclose the ligand name in parentheses.
- Metal:
  - If the complex is a cation or neutral, the metal is named same as the element (e.g., Cobalt, Nickel, Platinum).
  - If the complex is an anion, the metal name ends with the suffix '-ate' (e.g., Fe: ferrate, Co: cobaltate, Cu: cuprate, Ag: argentate, Au: aurate, Sn: stannate, Pb: plumbate).
- Oxidation State: The oxidation state of the central metal is written in Roman numerals in parentheses immediately after the name of the metal (without any space).
Bridging Ligands: Indicated by the prefix 'μ-' before the name of the ligand.

Examples:
- [Co(NH3)6]Cl3 : Hexaamminecobalt(III) chloride
- K4[Fe(CN)6] : Potassium hexacyanidoferrate(II)
- [Cr(H2O)4Cl2]Cl : Tetraaquadichloridochromium(III) chloride
- [Ni(CO)4] : Tetracarbonylnickel(0)
- [Pt(NH3)2Cl(NO2)] : Diamminechloridonitrito-N-platinum(II) (assuming NO2 coordination)
- K3[Al(C2O4)3] : Potassium trioxalatoaluminate(III)
- [Co(en)3]2(SO4)3 : Tris(ethylenediamine)cobalt(III) sulfate

5. Isomerism in Coordination Compounds

Isomers are compounds having the same chemical formula but different arrangements of atoms.

A. Structural Isomerism: Arises due to difference in the structure (connectivity) of ligands.
- 1. Ionisation Isomerism: Exchange of ions between the coordination sphere and the ionisation sphere. Example: [Co(NH3)5Br]SO4 (gives SO42- test) and [Co(NH3)5SO4]Br (gives Br- test).
- 2. Hydrate Isomerism (Solvate Isomerism): Difference in the number of solvent molecules (usually water) inside or outside the coordination sphere. Example: [Cr(H2O)6]Cl3 (violet), [Cr(H2O)5Cl]Cl2·H2O (blue-green), [Cr(H2O)4Cl2]Cl·2H2O (dark green).
- 3. Linkage Isomerism: Arises when an ambidentate ligand coordinates through different donor atoms. Example: [Co(NH3)5(NO2)]Cl2 (nitro, yellow) and [Co(NH3)5(ONO)]Cl2 (nitrito, red).
- 4. Coordination Isomerism: Exchange of ligands between cationic and anionic coordination spheres of different metal ions present in a complex salt. Example: [Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6].
B. Stereoisomerism: Same chemical formula and connectivity, but different spatial arrangement of ligands.
- 1. Geometrical Isomerism (cis-trans isomerism): Arises due to different relative positions of ligands around the central metal atom.
  - Common in square planar ([Ma2b2], [Mabcd]) and octahedral ([Ma4b2], [Ma3b3], [M(AA)2b2]) complexes (M=metal, a,b=unidentate ligands, AA=bidentate ligand).
  - cis: Similar groups on the same side. trans: Similar groups on opposite sides.
  - Example (Square Planar): cis-[Pt(NH3)2Cl2] and trans-[Pt(NH3)2Cl2].
  - Example (Octahedral): cis-[Co(NH3)4Cl2]+ and trans-[Co(NH3)4Cl2]+.
  - Octahedral complexes of type [Ma3b3] show facial (fac) and meridional (mer) isomerism.
- 2. Optical Isomerism: Arises when a molecule and its mirror image are non-superimposable (chiral). They rotate the plane of polarised light in opposite directions (dextrorotatory 'd' or '+', laevorotatory 'l' or '-').
  - Common in octahedral complexes with bidentate ligands, e.g., [Co(en)3]3+, cis-[Co(en)2Cl2]+.
  - Tetrahedral complexes with four different ligands [Mabcd] can show optical isomerism, but it's rare.
  - Square planar complexes generally do not show optical isomerism due to the presence of a plane of symmetry.

6. Bonding in Coordination Compounds

A. Valence Bond Theory (VBT) - Linus Pauling
- Postulates:
  1. The central metal atom/ion makes available a number of empty orbitals equal to its coordination number for the formation of coordinate bonds with ligands.
1. These empty orbitals hybridise to form a set of equivalent hybrid orbitals with definite geometry (e.g., sp-Linear, sp3-Tetrahedral, dsp2-Square Planar, sp3d2/d2sp3-Octahedral).
2. Ligands donate electron pairs to these empty hybrid orbitals, forming coordinate bonds.
3. Inner d-orbitals ((n-1)d) or outer d-orbitals (nd) can be used in hybridisation.
  - Inner Orbital Complex (Low Spin Complex): Uses (n-1)d orbitals (e.g., d2sp3). Often formed with strong field ligands. Pairing of electrons occurs.
  - Outer Orbital Complex (High Spin Complex): Uses nd orbitals (e.g., sp3d2). Often formed with weak field ligands. Pairing of electrons follows Hund's rule.
- Magnetic Properties: Predicts magnetic behaviour based on unpaired electrons.
  - Paramagnetic: Contains unpaired electrons.
  - Diamagnetic: Contains no unpaired electrons.
  - Magnetic Moment (μ) = √n(n+2) Bohr Magnetons (BM), where n = number of unpaired electrons.
- Examples:
  - [Co(NH3)6]3+ (Co3+: 3d6): NH3 strong field -> pairing -> 3d6 becomes t2g6 eg0 -> d2sp3 hybridisation -> Octahedral, Diamagnetic.
  - [CoF6]3- (Co3+: 3d6): F- weak field -> no pairing -> 3d6 remains t2g4 eg2 -> sp3d2 hybridisation -> Octahedral, Paramagnetic (n=4).
  - [Ni(CN)4]2- (Ni2+: 3d8): CN- strong field -> pairing in 3d -> dsp2 hybridisation -> Square Planar, Diamagnetic.
  - [NiCl4]2- (Ni2+: 3d8): Cl- weak field -> no pairing -> sp3 hybridisation -> Tetrahedral, Paramagnetic (n=2).
- Limitations of VBT: Cannot explain colour, quantitative magnetic data, relative stabilities, distinction between weak/strong ligands, or detailed spectral properties.
B. Crystal Field Theory (CFT)
- Assumptions: Interaction between metal ion and ligands is purely electrostatic. Ligands are treated as point charges or dipoles.
- Core Idea: In a free gaseous metal ion, all five d-orbitals are degenerate (have same energy). When ligands approach, this degeneracy is lost due to repulsion between ligand electrons/negative poles and the electrons in the metal d-orbitals. The pattern of splitting depends on the geometry of the complex.
- Crystal Field Splitting in Octahedral Complexes:
  - The d-orbitals split into two sets:
    - t2g set (lower energy): dxy, dyz, dzx (orbitals lying between the axes). Energy decreases by -0.4 Δo or -4 Dq.
    - eg set (higher energy): dx2-y2, dz2 (orbitals lying along the axes, experience more repulsion). Energy increases by +0.6 Δo or +6 Dq.
  - Δo (Delta Octahedral): Crystal Field Splitting Energy in octahedral field. It is the energy difference between t2g and eg sets.
- Crystal Field Splitting in Tetrahedral Complexes:
  - Splitting is inverted compared to octahedral. Ligands approach between the axes.
  - e set (lower energy): dx2-y2, dz2. Energy decreases by -0.6 Δt.
  - t2 set (higher energy): dxy, dyz, dzx. Energy increases by +0.4 Δt.
  - Δt (Delta Tetrahedral): Crystal Field Splitting Energy in tetrahedral field.
  - Relationship: Δt ≈ (4/9) Δo. Splitting is smaller in tetrahedral fields. Pairing is less common.
- Factors Affecting Δo:
  - Nature of Ligands: Ligands arranged in order of increasing field strength (increasing Δo) constitute the Spectrochemical Series:
    I- < Br- < SCN- < Cl- < S2- < F- < OH- < C2O42- < H2O < NCS- < EDTA4- < NH3 < en < CN- < CO
    (Weak field ligands <---------------------> Strong field ligands)
  - Oxidation State of Metal Ion: Higher the oxidation state, greater the Δo.
  - Nature of Metal Ion: Δo increases down a group (3d < 4d < 5d).
- High Spin vs. Low Spin (Octahedral):
  - If Δo < P (Pairing Energy): Electrons prefer to occupy eg orbitals before pairing in t2g (Weak field ligand -> High Spin Complex). Configuration follows Hund's rule.
  - If Δo > P: Electrons prefer to pair up in t2g orbitals before occupying eg (Strong field ligand -> Low Spin Complex).
- Colour of Coordination Compounds: Due to d-d transitions. When light falls on the complex, an electron from a lower energy d-orbital (e.g., t2g) absorbs energy corresponding to a specific colour/wavelength and gets excited to a higher energy d-orbital (e.g., eg). The transmitted light appears as the complementary colour. Complexes with d0 or d10 configuration are usually colourless.
- Magnetic Properties: Explained by the number of unpaired electrons in the split d-orbitals.
- Limitations of CFT: Assumes purely ionic bonding (ignores covalent character), fails to explain the relative strengths of ligands fully (e.g., why neutral NH3 is stronger than anionic OH-), doesn't consider ligand orbitals.

7. Stability of Coordination Compounds

Stability Constant (Formation Constant, K): Equilibrium constant for the formation of the complex in solution.
M + 4L ⇌ ML4 ; K = [ML4] / ([M][L]^4)
Higher the value of K, greater the stability of the complex.
Stepwise Stability Constants (k1, k2, ...) and Overall Stability Constant (βn = k1 × k2 × ... × kn).
Factors Affecting Stability:
- Charge on the central metal ion: Higher charge leads to greater stability.
- Size of the central metal ion: Smaller size (higher charge density) leads to greater stability.
- Nature of the Ligand:
  - Basic strength: More basic ligands generally form more stable complexes.
  - Chelate Effect: Complexes formed by chelating ligands are significantly more stable than those formed by analogous unidentate ligands. This is largely due to favourable entropy changes.

8. Applications of Coordination Compounds

Analytical Chemistry: Qualitative detection (e.g., Ni2+ with DMG, Fe3+ with KSCN) and quantitative estimation (e.g., EDTA titrations for Ca2+, Mg2+ hardness of water).
Metallurgy: Extraction of metals (e.g., Ag and Au using NaCN - MacArthur-Forrest cyanide process), purification of metals (e.g., Mond process for Ni using CO).
Biological Systems: Vital roles, e.g., Haemoglobin (Fe complex for O2 transport), Chlorophyll (Mg complex for photosynthesis), Vitamin B12 (Co complex).
Medicine: Cisplatin [Pt(NH3)2Cl2] used in cancer treatment, EDTA used for lead poisoning treatment.
Catalysis: Many industrial processes use coordination compounds as catalysts (e.g., Wilkinson's catalyst [(Ph3P)3RhCl] for hydrogenation).
Photography: Fixing involves dissolving unreacted AgBr as a soluble complex [Ag(S2O3)2]3- using hypo solution (Na2S2O3).

Multiple Choice Questions (MCQs)

The oxidation state and coordination number of Cobalt in K3[Co(NO2)6] are respectively:
a) +3, 6
b) +2, 6
c) +3, 4
d) +2, 4
Which of the following ligands is ambidentate?
a) NH3
b) en
c) SCN-
d) C2O42-
The IUPAC name for [Cr(H2O)4Cl2]Cl is:
a) Tetraaquadichloridochromium(III) chloride
b) Dichloridotetraaquachromium(III) chloride
c) Tetraaquadichloridochromate(III) chloride
d) Chloride tetraaquadichloridochromium(III)
The complex [Co(NH3)5Br]SO4 and [Co(NH3)5SO4]Br exhibit which type of isomerism?
a) Linkage isomerism
b) Coordination isomerism
c) Ionisation isomerism
d) Hydrate isomerism
According to VBT, the hybridisation and shape of [Ni(CN)4]2- are:
a) sp3, tetrahedral
b) dsp2, square planar
c) sp3d2, octahedral
d) d2sp3, octahedral
Which of the following complexes is expected to be colourless?
a) [Ti(H2O)6]3+
b) [Cu(NH3)4]2+
c) [Zn(NH3)4]2+
d) [Fe(H2O)6]2+
The increasing order of crystal field splitting strength (Δo) for the ligands Cl-, H2O, NH3, CN- is:
a) Cl- < H2O < NH3 < CN-
b) CN- < NH3 < H2O < Cl-
c) H2O < Cl- < NH3 < CN-
d) NH3 < CN- < H2O < Cl-
The complex [CoF6]3- is paramagnetic while [Co(NH3)6]3+ is diamagnetic because:
a) F- is a strong field ligand, NH3 is a weak field ligand.
b) F- causes less splitting (Δo < P), NH3 causes more splitting (Δo > P).
c) Cobalt is in +2 state in [CoF6]3- and +3 in [Co(NH3)6]3+.
d) [CoF6]3- is tetrahedral, [Co(NH3)6]3+ is octahedral.
The chelate effect refers to the enhanced stability of complexes containing:
a) Ambidentate ligands
b) Unidentate ligands
c) Polydentate ligands forming rings
d) Bridging ligands
Cisplatin, an anticancer drug, has the formula:
a) trans-[Pt(NH3)2Cl2]
b) [Pt(en)Cl2]
c) cis-[Pt(NH3)2Cl2]
d) K[PtCl3(NH3)]

Answer Key for MCQs:

a (+3, 6)
c (SCN-)
a (Tetraaquadichloridochromium(III) chloride)
c (Ionisation isomerism)
b (dsp2, square planar)
c ([Zn(NH3)4]2+ - Zn2+ is d10)
a (Cl- < H2O < NH3 < CN- - Spectrochemical series)
b (F- weak field, high spin; NH3 strong field, low spin)
c (Polydentate ligands forming rings)
c (cis-[Pt(NH3)2Cl2])

Study these notes thoroughly, focusing on understanding the concepts rather than just memorizing. Practice naming compounds and predicting geometry/magnetic properties. Good luck with your preparation!

Class 12 Chemistry Notes Chapter 9 (Coordination compounds) – Chemistry-I Book

Philoid

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