Coordination Compounds

Coordination Compound: Compound where a central metal atom/ion is bonded to surrounding species (ligands) via coordinate bonds.

Example: $K_4[Fe(CN)_6]$ (potassium ferrocyanide); $[Cu(NH_3)_4]SO_4$.

Terminology:

• Coordination entity: $[Fe(CN)_6]^{4-}$ — the part in square brackets
• Central metal atom/ion: Fe in $[Fe(CN)_6]^{4-}$
• Ligands: Molecules/ions bonded to metal (donating electron pairs)
• Coordination sphere: Inside square brackets
• Counter ion: Outside square brackets (K⁺ in $K_4[Fe(CN)_6]$)
• Coordination number (CN): Number of donor atoms attached to metal

Types of Ligands:

Coordination Numbers: Common: 2, 4, 6. Less common: 3, 5, 7, 8.

IUPAC Nomenclature Rules:

1. Cation first, then anion in name.
2. Within complex ion: ligands alphabetically, then metal.
3. Greek prefixes for multiple identical ligands: di, tri, tetra (for simple); bis, tris, tetrakis (for complex ones like en).
4. Anionic ligands end in -o: chloride → chloro, hydroxide → hydroxo, sulfate → sulfato, cyanide → cyano.
5. Neutral ligands keep their name; exceptions: H₂O = aqua, NH₃ = ammine, CO = carbonyl, NO = nitrosyl.
6. Cationic ligands end in -ium: NO⁺ = nitrosylium.
7. Oxidation state of metal in Roman numerals in parentheses, after metal name.
8. If complex is anionic, metal name ends in -ate: Fe → Ferrate, Cu → Cuprate, Au → Aurate, Ag → Argentate, Pb → Plumbate.

Examples:

Oxidation State (OS) of Metal: Calculated from total charge of complex: $OS_M + \sum OS_{ligands} = $ overall charge of complex.

Werner's Theory (1893): First explanation of coordination compounds.

Postulates:

1. Metals have two types of valences:

• Primary valence (ionizable): Equivalent to oxidation state; satisfied by anions
• Secondary valence (non-ionizable): Equivalent to coordination number; satisfied by ligands; fixed for each metal in a given complex

1. Primary valences are non-directional; secondary are directional → geometry
2. Ligands occupy specific positions around metal at fixed geometry

For $CoCl_3 \cdot 6NH_3$:

• Primary valence: $3$ (from $Co^{3+}$)
• Secondary valence: $6$ (from 6 NH₃)
• Structure: $[Co(NH_3)_6]^{3+}\,3Cl^-$
• All 3 Cl⁻ are ionizable (give AgCl with AgNO₃)

For $CoCl_3 \cdot 5NH_3$:

• Structure: $[Co(NH_3)_5Cl]^{2+}\,2Cl^-$
• Only 2 Cl⁻ ionize (other is in coord. sphere)

For $CoCl_3 \cdot 4NH_3$:

• Structure: $[Co(NH_3)_4Cl_2]^+\,Cl^-$
• Only 1 Cl⁻ ionizes

Valence Bond Theory (VBT): Explains shape and magnetic properties.

Steps:

1. Determine OS of metal and electronic configuration of ion
2. Identify hybridization of metal orbitals to accept ligand electron pairs
3. Predict geometry from hybridization

Strong vs Weak Field Ligands (for VBT):

Strong field ligands (cause pairing of d electrons): $CN^-, CO, en, NH_3, NO_2^-$ Weak field ligands (don't cause pairing): $F^-, Cl^-, Br^-, I^-, OH^-, H_2O$ (intermediate sometimes)

Hybridization-Geometry:

Inner vs Outer Orbital Complex:

• Inner orbital (low spin): Uses inner $d$ orbitals ($(n-1)d$); requires strong field; pairing of e⁻ occurs.
• Outer orbital (high spin): Uses outer $d$ orbitals ($nd$); weak field; no pairing.

Examples:

• $[Fe(CN)_6]^{4-}$: Fe²⁺ ($d^6$). $CN^-$ strong field → pair up → low spin → $d^2sp^3$ → octahedral diamagnetic.
• $[Fe(H_2O)_6]^{2+}$: Fe²⁺ ($d^6$). $H_2O$ weak (for Fe²⁺) → no pairing → high spin → $sp^3d^2$ → octahedral paramagnetic (4 unpaired).
• $[Ni(CN)_4]^{2-}$: Ni²⁺ ($d^8$). $CN^-$ strong → pair to free up $d$ → $dsp^2$ → square planar diamagnetic.
• $[NiCl_4]^{2-}$: Ni²⁺ ($d^8$). $Cl^-$ weak → no pairing → $sp^3$ → tetrahedral paramagnetic.

Isomerism: Compounds with same molecular formula but different structure/properties.

Structural Isomerism:

1. Ionization Isomerism: Same formula, different ions outside coordination sphere.

• $[Co(NH_3)_5SO_4]Br$ vs $[Co(NH_3)_5Br]SO_4$
• First gives Br⁻ (test with AgNO₃ → AgBr), second gives SO₄²⁻ (test with BaCl₂ → BaSO₄).

2. Hydrate (Solvate) Isomerism: Water inside vs outside coordination sphere.

• $[Cr(H_2O)_6]Cl_3$ (violet) vs $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$ (blue-green) vs $[Cr(H_2O)_4Cl_2]Cl \cdot 2H_2O$ (dark green)

3. Linkage Isomerism: Ambidentate ligands attached through different atoms.

• $[Co(NH_3)_5(NO_2)]^{2+}$ (nitro, via N) vs $[Co(NH_3)_5(ONO)]^{2+}$ (nitrito, via O)
• $[Co(NH_3)_5(SCN)]^{2+}$ (S-attached) vs $[Co(NH_3)_5(NCS)]^{2+}$ (N-attached)

4. Coordination Isomerism: Different metals or distributions in cation and anion.

• $[Co(NH_3)_6][Cr(CN)_6]$ vs $[Cr(NH_3)_6][Co(CN)_6]$
• Requires complex with both cationic and anionic parts.

Stereoisomerism:

1. Geometrical (cis-trans) Isomerism:

Square planar $[Ma_2b_2]$:

• cis: same groups adjacent (90° apart)
• trans: same groups opposite (180° apart)
• Example: $[Pt(NH_3)_2Cl_2]$ — both cis and trans forms exist; cis is cisplatin (anti-cancer drug)

Octahedral:

• $[Ma_4b_2]$: cis (b's adjacent) or trans (b's opposite)
• $[Ma_3b_3]$: facial (fac, three a's on triangular face) or meridional (mer, three a's in plane)
• $[M(en)_2Cl_2]$: cis and trans

2. Optical Isomerism:

A complex shows optical isomerism if it's chiral (non-superimposable on its mirror image).

Conditions:

• No improper rotation axis (e.g., plane of symmetry, center of symmetry)
• Common in octahedral with bidentate ligands: $[M(en)_3]^{n+}$ has $\Delta$ (right-handed) and $\Lambda$ (left-handed) forms.

Examples:

• $[Co(en)_3]^{3+}$: only optical isomers (no cis-trans)
• cis-$[Co(en)_2Cl_2]^+$: optical isomers exist (chiral)
• trans-$[Co(en)_2Cl_2]^+$: not chiral (has plane of symmetry)

Tetrahedral with 4 different ligands can be chiral, but few examples.

Crystal Field Theory (CFT): Treats ligands as point charges (negative) producing electrostatic field around metal. The field splits $d$-orbitals.

Octahedral Field:

• 5 $d$-orbitals split into:
• $e_g$ set: $d_{x^2-y^2}, d_{z^2}$ (higher energy; lobes point at ligands)
• $t_{2g}$ set: $d_{xy}, d_{yz}, d_{xz}$ (lower energy; lobes between ligands)

Crystal Field Splitting Energy (CFSE): $\Delta_o$ = energy gap between $e_g$ and $t_{2g}$.

For octahedral:

• $e_g$ at $+0.6\Delta_o$ above mean
• $t_{2g}$ at $-0.4\Delta_o$ below mean

Total CFSE = (electrons in $t_{2g}$)($-0.4\Delta_o$) + (electrons in $e_g$)($+0.6\Delta_o$) + (number of pairs caused by strong field)(P), where P = pairing energy.

Tetrahedral Field:

• 5 $d$-orbitals split into:
• $e$ set ($d_{x^2-y^2}, d_{z^2}$): lower energy
• $t_2$ set ($d_{xy}, d_{yz}, d_{xz}$): higher energy
• $\Delta_t \approx (4/9)\Delta_o$ (smaller splitting)
• No low-spin tetrahedral complexes (splitting too small)

Square Planar: Similar to octahedral but $z$-axis ligands removed.

Spectrochemical Series: Ligands arranged by increasing $\Delta$: $$I^- < Br^- < S^{2-} < SCN^- < Cl^- < NO_3^- < F^- < OH^- < H_2O < NCS^- < NH_3 < en < NO_2^- < CN^- < CO$$

(weak field) ←————————————————→ (strong field)

Spin State for $d^4$ to $d^7$ Octahedral:

Color of Coordination Compounds:

• Complementary to absorbed light (d-d transition)
• $\Delta_o$ depends on metal, ligand, geometry
• Larger $\Delta_o$ → higher energy light absorbed → bluer/redder shift

Examples:

• $[Ti(H_2O)_6]^{3+}$ ($d^1$): absorbs in yellow region → violet color
• $[Cu(H_2O)_4]^{2+}$ ($d^9$): blue
• $[Cu(NH_3)_4]^{2+}$ ($d^9$): deep blue (greater $\Delta_o$, deeper blue)

Crystal Field Stabilization Energy (CFSE) Calculations:

Example: $[Fe(CN)_6]^{4-}$, Fe²⁺ $d^6$ low spin: $t_{2g}^6 e_g^0$. CFSE = $6 \times (-0.4) \Delta_o + 0 \times 0.6 \Delta_o = -2.4 \Delta_o$ (extra stability).

Limitations of CFT:

• Ignores covalent character (ligands treated as pure point charges)
• Cannot explain $\pi$-bonding (e.g., why CO is a strong field ligand)
• More accurate: Ligand Field Theory (LFT) or MOT

Applications of Coordination Compounds: