Unit 6: Transition Metals

Introduction and Characteristics of Transition Metals

Transition metals are elements whose atoms have partially filled d orbitals (d¹ to d⁹) in either the neutral state or in one of their common oxidation states. They occupy the d‑block of the periodic table, spanning Groups 3 to 12.

Key physical and chemical characteristics include:

• High melting and boiling points: Due to strong metallic bonding arising from delocalised d‑electrons.
• Hardness and density: Generally hard, dense metals.
• Good electrical and thermal conductivity: Consequence of mobile electrons.
• Formation of coloured compounds: Result of d‑d electronic transitions.
• Variable oxidation states: Enables diverse chemistry and catalytic activity.
• Catalytic properties: Ability to change oxidation state readily.

Oxidation States of Transition Metals

The ability of transition metals to exhibit multiple oxidation states stems from the relatively small energy difference between the ns and (n‑1)d electrons. Both sets of electrons can be involved in bonding, leading to a variety of stable oxidation states.

General Trends

• The highest oxidation state often equals the group number (e.g., Mn in Group 7 can reach +7).
• Lower oxidation states (+2, +3) are common for the first‑row transition metals.
• Stability of higher oxidation states increases across a period and down a group.

Complex Ions and Metal Complexes

A complex ion consists of a central metal ion surrounded by molecules or anions called ligands that donate electron pairs to the metal. The resulting species can be cationic, anionic, or neutral.

Ligands

Ligands are classified by denticity (number of donor atoms) and charge. Common monodentate ligands include:

• Water: H₂O
• Ammonia: NH₃
• Chloride: Cl⁻
• Cyanide: CN⁻
• Carbon monoxide: CO

Polydentate ligands (chelating agents) such as ethylenediamine (en) and oxalate (C₂O₄²⁻) can occupy multiple coordination sites.

Coordination Number

The coordination number (CN) is the total number of donor atoms attached to the central metal ion. The most frequently observed coordination numbers for transition metals are 4 and 6.

Examples of Complex Ions

1. [Cu(NH3)4]2+ – tetraamminecopper(II) ion, CN = 4.
2. [Fe(CN)6]3− – hexacyanoferrate(III) ion, CN = 6.
3. [Co(NH3)6]3+ – hexaamminecobalt(III) ion, CN = 6.
4. [PtCl4]2− – tetrachloroplatinate(II) ion, CN = 4 (square planar).

Shapes of Complex Ions

The geometry adopted by a complex ion depends primarily on its coordination number and the electronic configuration of the metal centre.

Coordination Number 4

• Tetrahedral: Ligands point toward the corners of a tetrahedron; common with d⁰, d⁵ (high‑spin) and d¹⁰ configurations (e.g., [ZnCl4]2−).
• Square planar: Ligands occupy the corners of a square; frequently observed for d⁸ metal ions such as Ni(II), Pd(II), Pt(II) (e.g., [Pt(NH3)2Cl2]).

Coordination Number 6

• Octahedral: Six ligands positioned at the vertices of an octahedron; the most prevalent geometry for transition‑metal complexes (e.g., [Fe(H2O)6]2+, [Co(NH3)6]3+).

d‑Orbitals in Complex Ions – Crystal Field Theory (CFT)

Crystal Field Theory describes the effect of ligand electric fields on the five degenerate d‑orbitals of a transition metal ion. In an octahedral field, the d‑orbitals split into two sets:

• t_2g set (lower energy): d_xy, d_xz, d_yz
• e_g set (higher energy): d_z², d_x²−y²

The energy gap between these sets is the crystal field splitting energy, denoted Δo (for octahedral) or Δt (tetrahedral).

Whether electrons pair in the lower t_2g orbitals or occupy the higher e_g orbitals depends on the relative magnitude of Δo versus the electron pairing energy (P).

• Weak‑field ligands: Small Δo → high‑spin configurations (maximum unpaired electrons).
• Strong‑field ligands: Large Δo → low‑spin configurations (electron pairing in t_2g).

Example: For an octahedral Fe(II) (d⁶) complex:

• Weak field (e.g., [Fe(H₂O)₆]²⁺): t2g4 eg2 → high‑spin, 4 unpaired electrons.
• Strong field (e.g., [Fe(CN)₆]⁴⁻): t2g6 eg0 → low‑spin, 0 unpaired electrons.

Colour of Transition Metal Compounds

The characteristic colours of many transition‑metal complexes arise from d‑d transitions: promotion of an electron from a lower‑energy t_2g orbital to a higher‑energy e_g orbital (or vice versa) upon absorption of visible light.

The wavelength (and thus colour) absorbed depends on the magnitude of Δo, which is influenced by:

• Oxidation state of the metal: Higher oxidation states increase Δo (greater ligand field strength).
• Nature of the ligand: Ligands are ranked in the spectrochemical series (e.g., I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻ < CO). Stronger field ligands produce larger Δo and shift absorption to higher energy (shorter wavelength).
• Geometry: Tetrahedral complexes have smaller splitting (Δt ≈ 4/9 Δo) and often appear different colours.

Illustrative Examples