Chemistry Chapter 5 Notes of Coordination Compounds CBSE Class 12

Coordination compounds are a major area of modern inorganic and bio-inorganic chemistry, featuring metal atoms or ions bound to ligands by sharing electrons. Alfred Werner pioneered the field, proposing primary (ionisable) and secondary (non-ionisable, equal to coordination number) valences for the metal, and suggesting definite spatial arrangements for ligands. Key terms include coordination entity, central atom/ion, ligand (classified by denticity and donor atoms), coordination number, coordination sphere, and coordination polyhedron. Bonding is explained by theories like Valence Bond Theory (VBT) and Crystal Field Theory (CFT), addressing structure, magnetic properties, and colour. Isomerism (geometric and structural) is common due to different atomic arrangements. Coordination compounds have wide-ranging applications in biological systems, industry, analysis, and medicine.

1. Introduction and Importance of Coordination Compounds

• Coordination compounds are also called complex compounds.
• In these compounds, metal atoms are bound to anions or neutral molecules by sharing electrons.
• They are an important and challenging area of modern inorganic chemistry.
• New concepts of chemical bonding and molecular structure from studying these compounds have provided insights into their function in biological systems.
• Examples in biology include chlorophyll (magnesium), haemoglobin (iron), and vitamin B12 (cobalt).
• They are the backbone of modern inorganic and bio–inorganic chemistry and the chemical industry.
• Applications include metallurgical processes (e.g., extraction/purification of silver, gold, nickel), industrial catalysts , analytical reagents (e.g., EDTA, DMG), electroplating , textile dyeing , medicinal chemistry (e.g., chelate therapy, anti-tumour agents like cis-platin), and photography .

2. Werner’s Theory of Coordination Compounds

• Alfred Werner (1866-1919), a Swiss chemist, was the first to formulate ideas about the structures of coordination compounds.
• He prepared and characterised many compounds, studying their behaviour using simple experimental techniques.
• Werner proposed the concept of primary valence and secondary valence for a metal ion.
• He received the Nobel Prize in 1913 for his work on the linkage of atoms and coordination theory.

2.1 Postulates of Werner's Theory

1. In coordination compounds, metals show two types of linkages (valences): primary and secondary .
2. The primary valences are normally ionisable and are satisfied by negative ions.
3. The secondary valences are non-ionisable . These are satisfied by neutral molecules or negative ions.
4. The secondary valence is equal to the coordination number and is fixed for a metal.
5. The ions/groups bound by secondary linkages have characteristic spatial arrangements corresponding to different coordination numbers. In modern formulations, these are called coordination polyhedra .
6. Octahedral, tetrahedral, and square planar geometrical shapes are more common in coordination compounds of transition metals.

2.2 Experimental Evidence (Cobalt(III) Chloride with Ammonia)

• Werner studied a series of compounds of cobalt(III) chloride with ammonia.
• He found that adding excess silver nitrate solution precipitated some chloride ions as AgCl, but some remained in solution.
• Observations:
• 1 mol CoCl₃.6NH₃ (Yellow) gave 3 mol AgCl .
• 1 mol CoCl₃.5NH₃ (Purple) gave 2 mol AgCl .
• 1 mol CoCl₃.4NH₃ (Green) gave 1 mol AgCl .
• 1 mol CoCl₃.4NH₃ (Violet) gave 1 mol AgCl .
• These results, along with conductivity measurements, suggested:
• Six groups (Cl⁻ or NH₃ or both) were bonded to the cobalt ion during the reaction.
• The compounds could be formulated with a central entity that doesn't dissociate.
• The groups bound directly to the metal ion were termed secondary valence , which was six in these examples.
• The compounds were formulated as shown in Table 5.1:
• [Co(NH₃)₆]³⁺ 3Cl⁻ (1:3 electrolyte)
• [CoCl(NH₃)₅]²⁺ 2Cl⁻ (1:2 electrolyte)
• [CoCl₂(NH₃)₄]⁺ Cl⁻ (1:1 electrolyte)
• The fact that CoCl₃.4NH₃ exists as distinct green and violet compounds with identical empirical formula but different properties indicates the existence of isomers .

3. Definitions of Some Important Terms

• Coordination Entity: A central metal atom or ion bonded to a fixed number of ions or molecules. Examples: [CoCl₃(NH₃)₃], [Ni(CO)₄], [Fe(CN)₆]⁴⁻. The species within the square bracket are coordination entities.
• Central Atom/Ion: The atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement. Examples: Co³⁺, Ni²⁺, Fe³⁺. Central atoms/ions are also called Lewis acids .
• Ligands: The ions or molecules bound to the central atom/ion in the coordination entity. These can be simple ions (Cl⁻), small molecules (H₂O, NH₃), larger molecules (ethane-1,2-diamine), or even macromolecules (proteins).
• Unidentate: Ligand binds through a single donor atom (e.g., Cl⁻, H₂O, NH₃).
• Didentate: Ligand binds through two donor atoms (e.g., ethane-1,2-diamine, oxalate (C₂O₄²⁻)).
• Polydentate: Ligand has several donor atoms (e.g., N(CH₂CH₂NH₂)₃). EDTA⁴⁻ is an important hexadentate ligand.
• Chelate Ligand: A di- or polydentate ligand that uses two or more donor atoms simultaneously to bind a single metal ion. Complexes formed are chelate complexes and are generally more stable than those with unidentate ligands.
• Ambidentate Ligand: A ligand with two different donor atoms, either of which can ligate in the complex (e.g., NO₂⁻ binds via N or O; SCN⁻ binds via S or N).
• Coordination Number (CN): The number of ligand donor atoms to which the metal is directly bonded. Determined by the number of sigma bonds formed. Examples: 6 in [PtCl₆]²⁻, 4 in [Ni(NH₃)₄]²⁺, 6 in [Fe(C₂O₄)₃]³⁻ (oxalate is didentate).
• Coordination Sphere: The central atom/ion and the ligands attached to it, enclosed in square brackets. It is collectively termed the coordination sphere. Example: [Fe(CN)₆]⁴⁻ in K₄[Fe(CN)₆].
• Counter Ions: The ionisable groups written outside the square bracket. Example: K⁺ in K₄[Fe(CN)₆].
• Coordination Polyhedron: The spatial arrangement of the ligand atoms directly attached to the central atom. Common shapes are octahedral, square planar, and tetrahedral . Examples: [Co(NH₃)₆]³⁺ is octahedral, [Ni(CO)₄] is tetrahedral, [PtCl₄]²⁻ is square planar.
• Oxidation Number: The charge the central atom would carry if all ligands are removed with their shared electron pairs. Represented by a Roman numeral in parenthesis after the entity name. Example: Cu(I) in [Cu(CN)₄]³⁻.
• Homoleptic Complex: A metal bound to only one kind of donor groups (e.g., [Co(NH₃)₆]³⁺).
• Heteroleptic Complex: A metal bound to more than one kind of donor groups (e.g., [Co(NH₃)₄Cl₂]⁺).

3.1 Double Salt vs. Complex

4. Nomenclature of Coordination Compounds

• Based on IUPAC recommendations for unambiguous naming, especially for isomers.

4.1 Writing Formulas

1. The central atom is listed first.
2. Ligands are listed in alphabetical order (regardless of charge).
3. Polydentate ligands are also listed alphabetically; for abbreviated ligands, use the first letter of the abbreviation.
4. The entire coordination entity formula is in square brackets .
5. Formulas of polyatomic ligands or ligand abbreviations are in parentheses within the square brackets.
6. No space between ligands and the metal within the coordination sphere.
7. For a charged entity without counter ion, indicate charge outside brackets as a right superscript (number then sign, e.g., ³⁻).
8. Cation charge balances anion charge in a complete compound formula.

4.2 Naming Mononuclear Coordination Compounds

1. The cation is named first , followed by the anion. This applies whether the coordination entity is the cation or anion.
2. Ligands are named in alphabetical order before the central atom/ion name (reverse of formula writing).
3. Anionic ligands end in –o (e.g., chlorido for Cl⁻), neutral and cationic ligands use their usual names, with specific exceptions: aqua (H₂O), ammine (NH₃), carbonyl (CO), nitrosyl (NO).
4. Use prefixes mono, di, tri, etc. to indicate the number of individual ligands. If the ligand name already includes a numerical prefix (like ethylenediamine), use bis, tris, tetrakis , etc., and enclose the ligand name in parentheses.
5. The oxidation state of the metal is indicated by a Roman numeral in parenthesis after the metal name.
6. If the complex ion is a cation or a neutral molecule, the metal is named same as the element (e.g., cobalt for Co).
7. If the complex ion is an anion , the metal name ends with the suffix –ate (e.g., cobaltate for Co). Some metals use Latin names in anionic complexes (e.g., ferrate for Fe).

Examples:

• [Cr(NH₃)₃(H₂O)₃]Cl₃: triamminetriaquachromium(III) chloride (Complex cation, neutral ligands, Cr is +3)
• [Co(H₂NCH₂CH₂NH₂)₃]₂(SO₄)₃: tris(ethane-1,2–diamine)cobalt(III) sulphate (Complex cation, en is neutral, 3 sulphates mean Co complex is +3, Co is +3)
• [Ag(NH₃)₂][Ag(CN)₂]: diamminesilver(I)dicyanidoargentate(I) (First entity is cation, second is anion. Metal name differs)

5. Isomerism in Coordination Compounds

• Isomers have the same chemical formula but different arrangement of atoms, leading to different properties.
• Two principal types: Stereoisomerism and Structural isomerism .
• Stereoisomers have the same bonds but different spatial arrangements.
• Structural isomers have different bonds.

5.1 Stereoisomerism

• Geometrical isomerism: Arises in heteroleptic complexes due to different geometric arrangements of ligands.
• Common in coordination numbers 4 (square planar) and 6 (octahedral).
• Square Planar [MX₂L₂]: X ligands can be adjacent ( cis ) or opposite ( trans ). (Fig. 5.2 in source)
• Square Planar [MABXL]: Shows three isomers (two cis, one trans).
• Octahedral [MX₂L₄]: X ligands can be adjacent ( cis ) or opposite ( trans ). (Fig. 5.3 in source)
• Octahedral [MX₂(L–L)₂] (L-L is didentate): Shows cis/trans isomers (Fig. 5.4 in source).
• Octahedral [Ma₃b₃]:
• Facial (fac): Three donor atoms of the same ligand occupy adjacent positions at the corners of an octahedral face.
• Meridional (mer): Three donor atoms of the same ligand are around the meridian of the octahedron. (Fig. 5.5 in source)
• Geometric isomerism is not possible in tetrahedral complexes with different unidentate ligands because relative positions are the same.
• Optical isomerism: Occurs when mirror images are non-superimposable ( enantiomers ). Such molecules/ions are chiral . Forms are dextro (d) and laevo (l) based on rotation of plane-polarised light.
• Common in octahedral complexes involving didentate ligands (Fig. 5.6 in source).
• For [PtCl₂(en)₂]²⁺, only the cis-isomer shows optical activity (Fig. 5.7 in source).

5.2 Structural isomerism

• Linkage isomerism: Arises with ambidentate ligands that can bind through different atoms. Examples: NCS⁻ (binds via N or S), NO₂⁻ (binds via N or O, giving yellow -NO₂ or red -ONO forms).
• Coordination isomerism: Interchange of ligands between cationic and anionic coordination entities of different metal ions in a complex. Example: [Co(NH₃)₆][Cr(CN)₆] vs [Cr(NH₃)₆][Co(CN)₆].
• Ionisation isomerism: The counter ion is a potential ligand and displaces a ligand inside the coordination sphere, which becomes the counter ion. Example: [Co(NH₃)₅(SO₄)]Br vs [Co(NH₃)₅Br]SO₄.
• Solvate isomerism: Similar to ionisation isomerism, specifically when water is involved as the solvent ( hydrate isomerism ). Differs by whether the solvent molecule is bonded to the metal or is free in the crystal lattice. Example: [Cr(H₂O)₆]Cl₃ (violet) vs [Cr(H₂O)₅Cl]Cl₂.H₂O (grey-green).

6. Bonding in Coordination Compounds

• Werner's theory couldn't explain why only certain elements form complexes, bond directionality, or magnetic/optical properties.
• Modern theories include Valence Bond Theory (VBT), Crystal Field Theory (CFT), Ligand Field Theory (LFT), and Molecular Orbital Theory (MOT). VBT and CFT are discussed.

6.1 Valence Bond Theory (VBT)

• Metal atom/ion uses (n-1)d, ns, np or ns, np, nd orbitals for hybridisation under ligand influence.
• Hybridised orbitals overlap with ligand orbitals donating electron pairs.
• Hybridisation type dictates geometry (e.g., sp³ for tetrahedral, dsp² for square planar, sp³d² or d²sp³ for octahedral).
• Magnetic behaviour can often predict geometry.
• Diamagnetic: All electrons are paired; attracted weakly by magnetic field.
• Paramagnetic: Contains unpaired electrons; attracted strongly by magnetic field. Magnetic moment relates to number of unpaired electrons.
• Inner Orbital / Low Spin / Spin Paired Complex: Uses inner d orbitals ((n-1)d) for hybridisation (e.g., d²sp³ for octahedral). This occurs when ligands cause electron pairing in d orbitals.
• Outer Orbital / High Spin / Spin Free Complex: Uses outer d orbitals (nd) for hybridisation (e.g., sp³d² for octahedral). This occurs when ligands do not force electron pairing in d orbitals.
• VBT explains paramagnetic [NiCl₄]²⁻ (d⁸, sp³ hybridisation, 2 unpaired electrons) and diamagnetic [Ni(CN)₄]²⁻ (d⁸, dsp² hybridisation, 0 unpaired electrons).
• Explains strongly paramagnetic [Fe(H₂O)₆]³⁺ (d⁵, H₂O is weak field, sp³d² outer orbital, 5 unpaired electrons) vs. weakly paramagnetic [Fe(CN)₆]³⁻ (d⁵, CN⁻ is strong field, d²sp³ inner orbital, 1 unpaired electron).
• Limitations of VBT:
• Involves assumptions.
• Doesn't give quantitative magnetic data interpretation.
• Doesn't explain colour.
• Doesn't give quantitative stability interpretation.
• Doesn't make exact predictions for 4-coordinate geometry (tetrahedral vs square planar).
• Doesn't distinguish weak vs strong ligands.

6.2 Crystal Field Theory (CFT)

• An electrostatic model .
• Metal-ligand bond is considered ionic , arising from electrostatic interactions.
• Ligands are treated as point charges (anions) or point dipoles (neutral molecules).
• Degeneracy of metal d orbitals is lifted (split) by the asymmetrical negative field of ligands. The splitting pattern depends on the crystal field geometry.

Crystal Field Splitting in Octahedral Entities (D₀):

• Six ligands surround the metal.
• Repulsion between metal d electrons and ligand charges/electrons.
• d orbitals pointing towards ligands (dₓ²-ᵧ², dz²) experience more repulsion and are raised in energy (e g set) .
• d orbitals pointing between ligands (dxy, dyz, dxz) experience less repulsion and are lowered in energy (t₂ g set) .
• Energy separation between t₂ g and e g is D₀ .
• e g energy increases by (3/5)D₀, t₂ g energy decreases by (2/5)D₀ relative to the average energy in a spherical field. (Fig. 5.8 in source)
• D₀ depends on ligand field strength and metal ion charge.

Spectrochemical Series:

Arrangement of ligands in order of increasing field strength (and thus increasing D₀).

• I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < edta⁴⁻ < NH₃ < en < CN⁻ < CO.

Weak Field vs. Strong Field Ligands:

• Electron distribution in d orbitals (especially d⁴-d⁷) depends on D₀ vs. Pairing Energy (P). P is energy required to pair electrons in an orbital.
• If D₀ < P : Electron enters the higher energy e g level before pairing in t₂ g . Weak field ligands cause small splitting (D₀) and form high spin complexes (maximum unpaired electrons). Configuration example for d⁴: t₂ g ³ e g ¹.
• If D₀ > P : Electron pairs up in the lower energy t₂ g level to avoid entering the e g level. Strong field ligands cause large splitting (D₀) and form low spin complexes (minimum unpaired electrons). Configuration example for d⁴: t₂ g ⁴ e g ⁰.

Crystal Field Splitting in Tetrahedral Entities (Dₜ):

• d orbital splitting is inverted compared to octahedral.
• Two orbitals (dₓ²-ᵧ², dz²) are lowered, three orbitals (dxy, dyz, dxz) are raised. (Fig. 5.9 in source)
• Splitting energy Dₜ is smaller: D ₜ = (4/9) D ₀ for the same metal, ligands, and distances.
• Splitting energies are generally not large enough to force pairing, so low spin configurations are rarely observed in tetrahedral complexes.
• The 'g' subscript is not used as tetrahedral complexes lack a center of symmetry.

Colour in Coordination Compounds:

• Colour arises from the absorption of visible light . The observed colour is complementary to the absorbed colour. (Table 5.3 in source)
• CFT attributes colour to d-d transitions .
• When light is absorbed, an electron from a lower energy d orbital (e.g., t₂ g ) is excited to a higher energy d orbital (e.g., e g ). The energy of absorbed light equals the crystal field splitting energy (D₀ or Dₜ).
• If there is no crystal field splitting (absence of ligands) or no d electrons/vacant d orbitals for transition, the substance is colourless (e.g., anhydrous CuSO₄, Ti³⁺ complexes without water).
• The ligand influences colour because it affects the magnitude of D₀ (spectrochemical series). Changing ligands changes the splitting energy, changing the wavelength of light absorbed, and thus changing the observed colour (e.g., Ni²⁺ complexes with H₂O vs en).

Limitations of CFT:

• Assumes ligands are point charges, which is incorrect as anionic ligands are low in the spectrochemical series.
• Does not account for the covalent character of bonding between the metal and ligand.

7. Bonding in Metal Carbonyls

• Compounds containing only carbonyl (CO) ligands.
• Structures: Ni(CO)₄ (tetrahedral), Fe(CO)₅ (trigonal bipyramidal), Cr(CO)₆ (octahedral). Some are multinuclear with metal-metal bonds (e.g., Mn₂(CO)₁₀, Co₂(CO)₈).
• Metal-carbon bond has both sigma (σ) and pi (π) character .
• M–C σ bond: Formed by donation of a lone pair of electrons from the carbonyl carbon into a vacant metal orbital. (Ligand to metal donation).
• M–C π bond: Formed by donation of a pair of electrons from a filled metal d orbital into the vacant antibonding π* orbital of carbon monoxide. (Metal to ligand back donation).
• This synergic effect (σ donation and π back-bonding) strengthens the bond between CO and the metal, providing stability. (Fig. 5.14 in source)

8. Importance and Applications Review

• Analytical Chemistry: Used for detection and estimation of metal ions via colour formation with ligands (e.g., EDTA, DMG). Hardness of water estimated using EDTA to complex Ca²⁺ and Mg²⁺.
• Metallurgy: Used in metal extraction (e.g., gold with cyanide) and purification (e.g., nickel carbonyl).
• Biological Systems: Essential components like chlorophyll (photosynthesis, Mg), haemoglobin (oxygen transport, Fe), Vitamin B₁₂ (Co). Enzymes like carboxypeptidase A and carbonic anhydrase contain coordinated metal ions.
• Industrial Catalysts: Examples include Wilkinson catalyst [(Ph₃P)₃RhCl] for alkene hydrogenation.
• Electroplating: Provides smoother, more even coatings from complex solutions ([Ag(CN)₂]⁻, [Au(CN)₂]⁻) than simple ions.
• Photography: Undecomposed AgBr dissolved by hypo solution (Na₂S₂O₃) forming the complex ion [Ag(S₂O₃)₂]³⁻.
• Medicinal Chemistry: Chelate therapy removes toxic metals using chelating ligands (e.g., D-penicillamine, desferrioxime B for excess Cu/Fe; EDTA for lead poisoning). Some platinum compounds (e.g., cis-platin ) inhibit tumour growth.

Frequently Asked Questions (FAQ)

1. What is the key difference between a double salt and a complex?
• Double salts dissociate completely into their constituent ions in water.
• Complexes, however, contain complex ions that maintain their identity in solution and do not dissociate into simple ions.
• This difference is demonstrated by K₄[Fe(CN)₆], where the [Fe(CN)₆]⁴⁻ ion does not break down into Fe²⁺ and CN⁻ ions.
2. How does Crystal Field Theory explain the colour of coordination compounds?
• CFT postulates that ligand presence causes the metal d orbitals to split into different energy levels (e.g., t₂ g and e g in octahedral complexes).
• Colour arises when the complex absorbs light energy from the visible spectrum to excite an electron from a lower d orbital level to a higher one (d-d transition).
• The observed colour is the complementary colour of the light absorbed.
3. Briefly explain the concept of synergic bonding in metal carbonyls.
• Synergic bonding involves two components: a σ bond from the CO ligand to the metal and a π bond from a filled metal d orbital to the vacant antibonding π* orbital of CO.
• This reciprocal donation strengthens both the metal-carbon σ and π bonds.
• This unique bonding pattern contributes significantly to the stability of metal carbonyls.