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Class 12 Chemistry Notes Chapter 9 (Coordination Compounds) – Examplar Problems Book

Philoid

Philoid

— 12 min read

Examplar Problems
Alright class, let's dive deep into Coordination Compounds. This chapter is crucial not just for your board exams but also forms a significant part of the syllabus for various competitive government exams. We'll be focusing on the core concepts, especially those highlighted in the NCERT Exemplar, to build a strong foundation.

Chapter 9: Coordination Compounds - Detailed Notes

1. Introduction & Basic Terminology

  • Coordination Compounds: These are compounds containing a central metal atom or ion bonded to a fixed number of surrounding molecules or ions, called ligands, through coordinate covalent bonds (dative bonds). They retain their identity in solution.
    • Example: K₄[Fe(CN)₆], [Cu(NH₃)₄]SO₄
  • Coordination Entity: The central metal atom/ion bonded to a fixed number of ligands. Enclosed in square brackets [].
    • Example: [Fe(CN)₆]⁴⁻, [Cu(NH₃)₄]²⁺
  • Central Metal Atom/Ion: The acceptor atom/ion (usually a transition metal) to which ligands are attached. It acts as a Lewis acid.
  • Ligands: The molecules or ions bonded to the central metal atom/ion. They donate at least one pair of electrons and act as Lewis bases.
    • Types of Ligands based on Charge:
      • Anionic: Cl⁻ (chlorido), CN⁻ (cyanido), SO₄²⁻ (sulfato), C₂O₄²⁻ (oxalato)
      • Neutral: H₂O (aqua), NH₃ (ammine), CO (carbonyl), NO (nitrosyl), en (ethane-1,2-diamine)
      • Cationic: Rare, e.g., NO⁺ (nitrosonium)
    • Types of Ligands based on Denticity (Number of donor sites):
      • Unidentate/Monodentate: Donate one electron pair (e.g., Cl⁻, H₂O, NH₃, CN⁻).
      • Didentate/Bidentate: Donate two electron pairs (e.g., ethane-1,2-diamine (en), oxalate ion (ox)).
      • Polydentate: Donate more than two electron pairs (e.g., EDTA⁴⁻ - ethylenediaminetetraacetate ion, which is hexadentate).
      • Chelating Ligands: Di- or polydentate ligands that bind to a single central metal ion forming a ring structure (chelate ring). This leads to increased stability (Chelate Effect).
      • Ambidentate Ligands: Unidentate ligands that can coordinate through more than one donor atom. Example: NO₂⁻ (can bind through N as 'nitro' or O as 'nitrito'), SCN⁻ (can bind 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's essentially the number of ligand donor atoms directly attached to the metal.
    • Example: In [Fe(CN)₆]⁴⁻, CN of Fe is 6. In [Cu(NH₃)₄]²⁺, CN of Cu is 4. In [Co(en)₃]³⁺, CN of Co is 6 (since 'en' is bidentate).
  • Coordination Sphere: The central metal ion and the ligands directly attached to it, enclosed in square brackets [].
  • Counter Ions: The ionisable ions outside the coordination sphere.
    • Example: K⁺ in K₄[Fe(CN)₆], SO₄²⁻ in [Cu(NH₃)₄]SO₄.
  • Oxidation Number (State) of Central Metal: The charge the central metal atom would carry if all the ligands were removed along with the electron pairs shared with the central atom. Calculated by considering the charges of ligands and the overall charge of the coordination entity.
    • Example: In [Fe(CN)₆]⁴⁻: Let oxidation state of Fe be x. x + 6(-1) = -4 => x = +2.
    • Example: In [Cu(NH₃)₄]SO₄: The complex ion is [Cu(NH₃)₄]²⁺. Let oxidation state of Cu be y. y + 4(0) = +2 => y = +2.
  • Homoleptic Complexes: Complexes where the metal is bound to only one kind of donor group. Example: [Co(NH₃)₆]³⁺.
  • Heteroleptic Complexes: Complexes where the metal is bound to more than one kind of donor group. Example: [Co(NH₃)₄Cl₂]⁺.

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

  • Metals exhibit two types of valencies:

    • Primary Valency: Corresponds to the oxidation state of the metal ion. It is ionisable and satisfied by negative ions. Represented by dotted lines.
    • Secondary Valency: Corresponds to the coordination number of the metal ion. It is non-ionisable and satisfied by ligands (neutral molecules or negative ions). It determines 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 geometry (e.g., tetrahedral, square planar, octahedral).

    • Example: CoCl₃·6NH₃ is formulated as [Co(NH₃)₆]Cl₃.
      • Primary valency of Co = 3 (satisfied by 3 Cl⁻ ions).
      • Secondary valency of Co = 6 (satisfied by 6 NH₃ molecules). Geometry is octahedral.
      • Total ions in solution = 1 [Co(NH₃)₆]³⁺ + 3 Cl⁻ = 4 ions.

3. Nomenclature of Coordination Compounds (IUPAC Rules)

  1. Order of Naming Ions: Cation is named first, then the anion (like in simple salts, e.g., NaCl).

  2. Naming the Coordination Sphere:

    • Ligands: Named first in alphabetical order, before the central metal atom/ion. Ignore prefixes (di-, tri-, etc.) for alphabetisation.
    • Ligand Names:
      • Anionic ligands end in '-o' (e.g., Cl⁻: chlorido, CN⁻: cyanido, SO₄²⁻: sulfato, NO₂⁻: nitrito-N or nitro, ONO⁻: nitrito-O).
      • Neutral ligands retain their usual names (e.g., C₂H₅N: pyridine). Exceptions: H₂O (aqua), NH₃ (ammine), CO (carbonyl), NO (nitrosyl).
      • Cationic ligands end in '-ium' (rare).
    • Prefixes for Ligand Numbers: Use di-, tri-, tetra-, etc., for simple ligands. Use bis-, tris-, tetrakis-, etc., for complex ligands (whose names already contain a numerical prefix, like ethylenediamine).
    • Metal Name:
      • If the complex ion is a cation or neutral, the metal is named as the element (e.g., Cobalt, Platinum, Nickel).
      • If the complex ion is an anion, the metal name ends in '-ate' (e.g., Fe: ferrate, Co: cobaltate, Cu: cuprate, Ag: argentate, Au: aurate, Pt: platinate).
    • Oxidation State: The oxidation state of the central metal is written in Roman numerals in parentheses immediately after the metal name (e.g., (II), (III), (0)).
    • Bridging Ligands: Indicated by the prefix 'μ-' before the ligand name. If there are two or more bridging groups of the same kind, use di-μ-, tri-μ-, etc.

  3. Formula Writing: Metal written first, followed by ligands in alphabetical order (based on symbol). The formula of the entire coordination entity is enclosed in square brackets. Polyatomic ligands are enclosed in parentheses.

    • Examples:
      • [Cr(NH₃)₃(H₂O)₃]Cl₃ : Triamminetriaquachromium(III) chloride
      • [Co(en)₃]₂(SO₄)₃ : Tris(ethane-1,2-diamine)cobalt(III) sulfate
      • K₄[Fe(CN)₆] : Potassium hexacyanidoferrate(II)
      • [Pt(NH₃)₂Cl(NO₂)] : Diamminechloridonitrito-N-platinum(II) (assuming NO₂ binding)
      • [Ni(CO)₄] : Tetracarbonylnickel(0)

4. Isomerism in Coordination Compounds

Isomers have the same chemical formula but different arrangements of atoms.

  • A. Structural Isomerism: Different connectivity/bonding.

    • Ionisation Isomerism: Exchange of counter ion with a ligand inside the coordination sphere. They give different ions in solution.
      • Example: [Co(NH₃)₅SO₄]Br (gives Br⁻ test) and [Co(NH₃)₅Br]SO₄ (gives SO₄²⁻ test).
    • Hydrate (Solvate) Isomerism: Difference in whether water (or solvent) molecules are directly bonded to the metal (as ligand) or present as free solvent molecules.
      • Example: [Cr(H₂O)₆]Cl₃ (violet) and [Cr(H₂O)₅Cl]Cl₂·H₂O (blue-green).
    • Linkage Isomerism: Arises when an ambidentate ligand coordinates to the metal through different donor atoms.
      • Example: [Co(NH₃)₅(NO₂)]Cl₂ (nitro, N-bonded, yellow) and [Co(NH₃)₅(ONO)]Cl₂ (nitrito, O-bonded, red).
    • Coordination Isomerism: Exchange of ligands between cationic and anionic coordination spheres in salts containing complex cations and anions.
      • Example: [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].

  • B. Stereoisomerism: Same connectivity but different spatial arrangement.

    • Geometrical Isomerism (cis-trans): Arises due to different relative positions of ligands around the central metal.
      • Square Planar Complexes:
        • MA₂B₂ type: cis (identical ligands adjacent), trans (identical ligands opposite). Example: cis-[Pt(NH₃)₂Cl₂], trans-[Pt(NH₃)₂Cl₂].
        • MABCD type: Can show 3 isomers.
        • Tetrahedral complexes do not show geometrical isomerism (all positions are equivalent).
      • Octahedral Complexes:
        • MA₄B₂ type: cis and trans. Example: cis-[Co(NH₃)₄Cl₂]⁺, trans-[Co(NH₃)₄Cl₂]⁺.
        • MA₃B₃ type: facial (fac) (3 identical ligands occupy corners of one octahedral face) and meridional (mer) (3 identical ligands occupy positions around the 'meridian' of the octahedron). Example: fac-[Co(NH₃)₃(NO₂)₃], mer-[Co(NH₃)₃(NO₂)₃].
        • M(AA)₂B₂ type (AA = symmetrical bidentate ligand): cis and trans. Example: cis-[Co(en)₂Cl₂]⁺, trans-[Co(en)₂Cl₂]⁺.
    • Optical Isomerism (enantiomerism): Arises when a molecule and its mirror image are non-superimposable (chiral). These isomers (enantiomers) rotate the plane of polarised light in opposite directions.
      • Common in octahedral complexes with bidentate ligands.
      • Examples: cis-[Co(en)₂Cl₂]⁺ exists as a pair of enantiomers. trans-[Co(en)₂Cl₂]⁺ is achiral (has a plane of symmetry) and optically inactive.
      • [Co(en)₃]³⁺, [Cr(ox)₃]³⁻ are chiral and exist as enantiomeric pairs.
      • Tetrahedral complexes of type M(AB)₂ (where AB is unsymmetrical bidentate ligand) can show optical isomerism.

5. Bonding in Coordination Compounds

  • A. Valence Bond Theory (VBT)

    • Focuses on hybridisation of metal orbitals and coordinate bond formation.
    • Ligands donate electron pairs to vacant hybrid orbitals of the central metal ion.
    • The number and type of hybrid orbitals determine the geometry.
    • Predicts magnetic properties based on unpaired electrons in the metal ion after hybridisation and bond formation.
    • Common Hybridisations and Geometries:
      • CN 4: sp³ (Tetrahedral), dsp² (Square Planar)
      • CN 5: sp³d (Trigonal Bipyramidal)
      • CN 6: sp³d² (Octahedral, outer orbital complex, high spin), d²sp³ (Octahedral, inner orbital complex, low spin)
    • Inner vs Outer Orbital Complexes (Octahedral):
      • Inner Orbital (Low Spin): Uses inner (n-1)d orbitals for hybridisation (d²sp³). Formed typically with strong field ligands which cause pairing of electrons. Often diamagnetic or less paramagnetic.
      • Outer Orbital (High Spin): Uses outer nd orbitals for hybridisation (sp³d²). Formed typically with weak field ligands which do not cause pairing. Usually paramagnetic.
    • Examples:
      • [Co(NH₃)₆]³⁺ (Co³⁺: 3d⁶): NH₃ is strong field -> pairing -> 3d⁶ becomes t₂g⁶ -> uses inner d orbitals -> d²sp³ -> Octahedral, Diamagnetic.
      • [CoF₆]³⁻ (Co³⁺: 3d⁶): F⁻ is weak field -> no pairing -> 3d⁶ remains t₂g⁴ eg² -> uses outer d orbitals -> sp³d² -> Octahedral, Paramagnetic (4 unpaired e⁻).
      • [Ni(CN)₄]²⁻ (Ni²⁺: 3d⁸): CN⁻ is strong field -> pairing -> one 3d orbital becomes vacant -> dsp² -> Square Planar, Diamagnetic.
      • [NiCl₄]²⁻ (Ni²⁺: 3d⁸): Cl⁻ is weak field -> no pairing -> uses outer orbitals -> sp³ -> Tetrahedral, Paramagnetic (2 unpaired e⁻).
    • Limitations of VBT: Cannot explain colour, quantitative magnetic moments, spectrochemical series (why some ligands are strong/weak), thermodynamic stability differences.

  • B. Crystal Field Theory (CFT)

    • Assumes interaction between metal ion and ligands is purely electrostatic. Ligands are treated as point charges or dipoles.
    • Focuses on the effect of the ligand field on the degeneracy of metal d-orbitals.
    • Crystal Field Splitting: In the presence of ligands, the degeneracy of the d-orbitals is lifted.
      • Octahedral Field: d-orbitals split into two sets:
        • t₂g set (lower energy): dxy, dyz, dzx (orbitals between the axes, experience less repulsion). Lowered by -0.4 Δo.
        • eg set (higher energy): dx²-y², dz² (orbitals along the axes, experience more repulsion). Raised by +0.6 Δo.
        • Δo: Crystal Field Splitting Energy in octahedral field.
      • Tetrahedral Field: Splitting is inverted compared to octahedral.
        • e set (lower energy): dx²-y², dz² (experience less repulsion). Lowered by -0.6 Δt.
        • t₂ set (higher energy): dxy, dyz, dzx (experience more repulsion). Raised by +0.4 Δt.
        • Δt: Crystal Field Splitting Energy in tetrahedral field.
        • Relation: Δt ≈ (4/9) Δo. Splitting is smaller, so tetrahedral complexes are usually high spin.
      • Square Planar Field: More complex splitting derived from octahedral by removing ligands along z-axis. Energy order (approx): dxy < dyz, dzx < dz² < dx²-y². Large energy gap after dxy.
    • Spectrochemical Series: Arrangement of ligands in order of increasing crystal field splitting (Δo):
      I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO
      (Weak field ligands ← → Strong field ligands)
    • Factors Affecting Δo: Metal ion (oxidation state, position in d-block), Nature of ligand. Higher oxidation state and 3rd row > 2nd row > 1st row transition metals lead to larger Δo.
    • High Spin vs Low Spin (Octahedral): Depends on the relative magnitude of Δo and Pairing Energy (P - energy required to pair electrons in the same orbital).
      • If Δo < P (Weak field ligand): Electrons prefer to occupy higher energy eg orbitals before pairing in t₂g. High Spin configuration.
      • If Δo > P (Strong field ligand): Electrons prefer to pair up in lower energy t₂g orbitals. Low Spin configuration.
      • Applies to d⁴, d⁵, d⁶, d⁷ configurations in octahedral fields.
    • Colour of Coordination Compounds: Due to d-d transitions. Absorption of light energy promotes an electron from a lower energy d-orbital (e.g., t₂g) to a higher energy d-orbital (e.g., eg). The colour observed is the complementary colour of the light absorbed. The energy gap (Δo) determines the frequency (colour) of light absorbed. Complexes with d⁰ or d¹⁰ configuration are usually colourless.
    • Magnetic Properties: Explained by the number of unpaired electrons predicted by CFT filling based on Δo vs P. Magnetic moment (μ) ≈ √[n(n+2)] BM (Bohr Magneton), where n = number of unpaired electrons.
    • Limitations of CFT: Assumes purely electrostatic interactions (ignores covalent character of metal-ligand bond), fails to explain the relative strengths of ligands fully (e.g., why neutral CO is stronger than anionic F⁻).

6. Stability of Coordination Compounds

  • Stability Constant (Formation Constant, K) / Overall Stability Constant (β): Equilibrium constant for the formation of the complex in solution. Higher the value, greater the stability.
    • M + L ⇌ ML ; K₁ = [ML]/[M][L]
    • ML + L ⇌ ML₂ ; K₂ = [ML₂]/[ML][L] ... etc.
    • Overall: M + nL ⇌ MLn ; βn = K₁ × K₂ × ... × Kn = [MLn]/[M][L]ⁿ
  • Factors Affecting Stability:
    • Nature of Central Metal Ion:
      • Higher charge density (higher charge, smaller size) leads to greater stability. Stability generally follows Irving-Williams series for M²⁺ ions: Mn²⁺ < Fe²⁺ < Co²⁺ < Ni²⁺ < Cu²⁺ > Zn²⁺.
    • Nature of Ligand:
      • Higher basicity of ligand generally leads to greater stability.
      • Chelate Effect: Complexes formed by chelating ligands (di- or polydentate) are significantly more stable than similar complexes with unidentate ligands. This is primarily due to a favorable entropy change upon chelation (increase in the number of free particles).
    • Steric Factors: Bulky ligands can decrease stability due to steric hindrance.

7. 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 Nickel using CO).
  • Biological Systems: Vital roles in plants and animals.
    • Haemoglobin: Iron porphyrin complex, oxygen transport.
    • Chlorophyll: Magnesium porphyrin complex, photosynthesis.
    • Vitamin B₁₂: Cobalt complex (cyanocobalamin).
    • Various enzymes contain metal ions (e.g., Carboxypeptidase A contains Zn²⁺).
  • Medicine:
    • Cisplatin [cis-Pt(NH₃)₂Cl₂]: Anti-cancer drug.
    • EDTA: Used to treat lead poisoning (forms stable lead-EDTA complex which is excreted).
  • Catalysis: Many industrial processes use coordination compounds as catalysts (e.g., Wilkinson's catalyst [(Ph₃P)₃RhCl] for hydrogenation of alkenes).
  • Photography: Fixing involves formation of a soluble complex [Ag(S₂O₃)₂]³⁻ with hypo (Na₂S₂O₃).
  • Electroplating: Stable complex solutions provide smooth metal deposition.

Multiple Choice Questions (MCQs)

  1. The IUPAC name for K₃[Al(C₂O₄)₃] is:
    a) Potassium trioxalatoaluminate(III)
    b) Potassium tris(oxalato)aluminium(III)
    c) Potassium aluminium(III) oxalate
    d) Potassium trioxalatoaluminate(II)

  2. According to Werner's theory, the primary valency and secondary valency for the complex [Co(NH₃)₅Cl]Cl₂ are respectively:
    a) 3 and 5
    b) 3 and 6
    c) 2 and 6
    d) 6 and 3

  3. Which of the following is an ambidentate ligand?
    a) NH₃
    b) en
    c) SCN⁻
    d) Cl⁻

  4. The complexes [Co(NH₃)₅(NO₂)]Cl₂ and [Co(NH₃)₅(ONO)]Cl₂ exhibit which type of isomerism?
    a) Ionisation isomerism
    b) Coordination isomerism
    c) Linkage isomerism
    d) Geometrical isomerism

  5. The hybridisation, shape, and magnetic character of [Fe(CN)₆]⁴⁻ are: (Atomic no. Fe = 26)
    a) sp³d², Octahedral, Paramagnetic
    b) d²sp³, Octahedral, Diamagnetic
    c) dsp², Square Planar, Diamagnetic
    d) sp³, Tetrahedral, Paramagnetic

  6. In the crystal field splitting for an octahedral complex, the d-orbitals which are raised in energy are:
    a) dxy, dyz, dzx
    b) dx²-y², dz²
    c) dxy, dx²-y²
    d) dz², dyz

  7. Which of the following complexes is expected to be a high spin complex? (Consider common ligand strengths)
    a) [Co(NH₃)₆]³⁺
    b) [Fe(CN)₆]⁴⁻
    c) [CoF₆]³⁻
    d) [Ni(CN)₄]²⁻

  8. The colour of coordination compounds is attributed to:
    a) Ligand to metal charge transfer
    b) Metal to ligand charge transfer
    c) d-d electronic transitions
    d) s-p electronic transitions

  9. The stability of [Cu(en)₂]²⁺ is significantly greater than [Cu(NH₃)₄]²⁺ due to:
    a) Higher charge on 'en' ligand
    b) Chelate effect
    c) Higher coordination number in the 'en' complex
    d) Pi-bonding ability of 'en'

  10. The central metal ion present in Chlorophyll is:
    a) Iron
    b) Cobalt
    c) Magnesium
    d) Manganese

Answer Key & Explanations

  1. a) Potassium trioxalatoaluminate(III)
    • Explanation: K⁺ is cation. [Al(C₂O₄)₃]³⁻ is anion. Ligand is C₂O₄²⁻ (oxalato). Three oxalato ligands = trioxalato. Metal in anionic complex = aluminate. Oxidation state: x + 3(-2) = -3 => x = +3.
  2. b) 3 and 6
    • Explanation: The formula is [Co(NH₃)₅Cl]Cl₂. The ions outside [] satisfy primary valency (2 Cl⁻) and the charge on the complex ion (+2) is balanced by the metal's oxidation state (+3) and the charge of the ligand inside (-1 for Cl⁻). So, primary valency = Oxidation state = +3. Ligands inside [] satisfy secondary valency (5 NH₃ + 1 Cl⁻ = 6 ligands). Secondary valency = Coordination Number = 6.
  3. c) SCN⁻
    • Explanation: SCN⁻ can coordinate either through Sulfur (thiocyanato) or Nitrogen (isothiocyanato). NH₃ and Cl⁻ are unidentate. 'en' is bidentate but not ambidentate.
  4. c) Linkage isomerism
    • Explanation: The complexes differ only in the atom through which the ambidentate ligand NO₂⁻/ONO⁻ is attached to the Cobalt ion (N vs O).
  5. b) d²sp³, Octahedral, Diamagnetic
    • Explanation: Fe is +2 (3d⁶). CN⁻ is a strong field ligand. It causes pairing of 3d electrons. The configuration becomes t₂g⁶ eg⁰. Two inner d-orbitals (3d), one 4s, and three 4p orbitals hybridise to form d²sp³ hybrid orbitals. Geometry is octahedral. Since all electrons are paired, it is diamagnetic.
  6. b) dx²-y², dz²
    • Explanation: In octahedral splitting, the eg orbitals (dx²-y², dz²), which point directly towards the ligands along the axes, experience greater repulsion and are raised in energy. The t₂g orbitals (dxy, dyz, dzx) point between the axes and are lowered in energy.
  7. c) [CoF₆]³⁻
    • Explanation: High spin complexes are typically formed with weak field ligands. F⁻ is a weak field ligand. NH₃ and CN⁻ are strong field ligands, likely forming low spin complexes with Co³⁺ and Fe²⁺ respectively. [Ni(CN)₄]²⁻ is square planar (low spin). Co³⁺ is 3d⁶. With weak field F⁻, Δo < P, configuration is t₂g⁴ eg², hence high spin.
  8. c) d-d electronic transitions
    • Explanation: Absorption of visible light promotes electrons from lower energy d-orbitals to higher energy d-orbitals within the same d-subshell. The energy difference corresponds to the energy of visible light.
  9. b) Chelate effect
    • Explanation: 'en' (ethane-1,2-diamine) is a bidentate ligand, forming chelate rings with Cu²⁺. Chelation leads to a significant increase in thermodynamic stability compared to complexes with analogous monodentate ligands like NH₃.
  10. c) Magnesium
    • Explanation: Chlorophyll, the pigment essential for photosynthesis in plants, is a coordination complex containing Magnesium (Mg²⁺) as the central metal ion in a porphyrin ring structure.

Study these notes thoroughly. Pay close attention to the examples and the reasoning behind bonding theories and isomerism. Good luck with your preparation!

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