Chapter 5: Coordination Compounds

Double Salts vs Coordination Compounds

• Double Salts: Dissociate completely into simple ions when dissolved in water. Example: Mohr's salt [FeSO₄·(NH₄)₂SO₄·6H₂O], Potash alum.
• Coordination Compounds: Do not dissociate into simple ions completely in water. They retain their identity in solution. Example: K₄[Fe(CN)₆].

Werner's Theory of Coordination Compounds

Metals possess two types of linkages (valencies):

• Primary Valency: Ionizable, satisfied by negative ions. Corresponds to the oxidation state of the central metal. Represented by dotted lines.
• Secondary Valency: Non-ionizable, satisfied by neutral molecules or negative ions (ligands). Corresponds to the coordination number. Represented by solid lines. Spatial arrangement of secondary valencies gives the complex its specific geometry.

Coordination Entity & Central Atom

Coordination Entity: Constitutes a central metal atom or ion bonded to a fixed number of ions or molecules. Example: [CoCl₃(NH₃)₃].

Central Atom/Ion: The atom/ion to which a fixed number of ligands are bound in a definite geometrical arrangement.

Ligands and Their Types

Ions or molecules bound to the central atom/ion. They act as Lewis bases (electron pair donors).

• Unidentate: Bounds through a single donor atom. Example: Cl⁻, H₂O, NH₃.
• Didentate: Bounds through two donor atoms. Example: ethane-1,2-diamine (en), oxalate (C₂O₄²⁻).
• Polydentate: Bounds through several donor atoms. Example: EDTA⁴⁻ (hexadentate).
• Chelating Ligand: A di- or polydentate ligand uses its two or more donor atoms to bind a single metal ion, forming a ring structure. Imparts extra stability (Chelate effect).
• Ambidentate Ligand: Ligand which can ligate through two different atoms. Example: NO₂⁻ (can bind through N or O), SCN⁻ (can bind through S or N).

General Rules for Naming

• The cation is named first in both positively and negatively charged coordination entities.
• The ligands are named in alphabetical order before the name of the central metal atom/ion.
• Names of anionic ligands end in -o (e.g., chloro, cyanido, oxalato), whereas neutral and cationic ligands are the same except aqua (H₂O), ammine (NH₃), carbonyl (CO), and nitrosyl (NO).
• Prefixes mono, di, tri, etc., are used to indicate the number of individual ligands. If the ligand name includes a numerical prefix, then terms like bis, tris, tetrakis are used.
• Oxidation state of the metal is indicated by Roman numeral in parentheses.
• If the complex ion is an anion, the name of the metal ends with the suffix -ate (e.g., ferrate, cuprate, argentate).

Examples

K₄[Fe(CN)₆] : Potassium hexacyanidoferrate(II) [Co(NH₃)₅Cl]Cl₂ : Pentaamminechloridocobalt(III) chloride

Structural Isomerism

• Ionization Isomerism: Exchange of ions between coordination sphere and ionization sphere. Ex: [Co(NH₃)₅Br]SO₄ and [Co(NH₃)₅SO₄]Br.
• Linkage Isomerism: Arises in complexes containing ambidentate ligands. Ex: [Co(NH₃)₅(NO₂)]²⁺ and [Co(NH₃)₅(ONO)]²⁺.
• Coordination Isomerism: Interchange of ligands between cationic and anionic entities. Ex: [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].
• Hydrate (Solvate) Isomerism: Water molecules act either as a ligand or as water of hydration. Ex: [Cr(H₂O)₆]Cl₃ (violet) and [Cr(H₂O)₅Cl]Cl₂·H₂O (grey-green).

Stereoisomerism

• Geometrical Isomerism: Occurs in heteroleptic complexes. Includes cis (similar ligands adjacent) and trans (similar ligands opposite). Also, fac (facial) and mer (meridional) isomerism in [Ma₃b₃] type octahedral complexes.
• Optical Isomerism: Chiral complexes which are non-superimposable mirror images. They rotate the plane of polarized light (dextro and laevo). Common in octahedral complexes involving didentate ligands like [Co(en)₃]³⁺.

Main Postulates

The metal ion uses its empty (n-1)d, ns, np or ns, np, nd orbitals to undergo hybridization, giving a set of equivalent orbitals of definite geometry (tetrahedral, square planar, octahedral).

Types of Hybridization & Geometries

• sp³: Tetrahedral (C.N. = 4). Example: [NiCl₄]²⁻ (Paramagnetic).
• dsp²: Square Planar (C.N. = 4). Example: [Ni(CN)₄]²⁻ (Diamagnetic).
• sp³d²: Octahedral (Outer orbital complex, high spin). Uses outer nd orbitals. Example: [CoF₆]³⁻.
• d²sp³: Octahedral (Inner orbital complex, low spin). Uses inner (n-1)d orbitals. Example: [Co(NH₃)₆]³⁺.

Magnetic Behavior

Depends on the presence of unpaired electrons (Paramagnetic) or absence (Diamagnetic). Spin-only magnetic moment μ = √(n(n+2)) BM.

Basic Concept

Considers the metal-ligand bond to be purely ionic. Ligands are treated as point charges. When ligands approach the metal ion, the degeneracy of the five d-orbitals is destroyed, splitting them into different energy levels.

Splitting in Octahedral Complexes

The d-orbitals split into two sets:

• t₂g set: Lower energy (dxy, dyz, dxz).
• eg set: Higher energy (dx²-y², dz²).
• The energy difference is Crystal Field Splitting Energy (Δ₀).

Spectrochemical Series

Arrangement of ligands in increasing order of crystal field splitting (Δ₀):

I⁻ < Br⁻ < SCN⁻ < Cl⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO

• Strong Field Ligands: (e.g., CO, CN⁻, en) cause large splitting (Δ₀ > Pairing Energy P). Electrons pair up in t₂g. Form Low Spin complexes.
• Weak Field Ligands: (e.g., Halogens, H₂O) cause small splitting (Δ₀ < P). Electrons jump to eg. Form High Spin complexes.

Colour in Coordination Compounds

Attributed to d-d transition of electrons. An electron from a lower energy d-orbital absorbs light from the visible region and jumps to a higher energy d-orbital. The transmitted light shows the complementary colour.