Variable Oxidation States in Transition Elements (HL) (DP IB Chemistry): Revision Note

Variable Oxidation States in Transition Elements

Electron Configuration

• The full electronic configuration of the first-row transition metals is shown in the table below

• Following the Aufbau Principle electrons occupy the lowest energy subshells first

  • The 4s overlaps with the 3d subshell so the 4s is filled first

• Remember: You can abbreviate the first five subshells, 1s-3p, to [Ar] representing the configuration of argon (known as the argon core)

Electronic configuration of the first d-series transition elements

• Ti

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²

  • [Ar] 3d² 4s²

• V

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s²

  • [Ar] 3d³ 4s²

• Cr

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹

  • [Ar] 3d⁵ 4s¹

• Mn

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²

  • [Ar] 3d⁵ 4s²

• Fe

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

  • [Ar] 3d⁶ 4s²

• Co

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s²

  • [Ar] 3d⁷ 4s²

• Ni

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s²

  • [Ar] 3d⁸ 4s²

• Cu

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

  • [Ar] 3d¹⁰ 4s¹

Exceptions

• Two first-row d-block elements are exceptions to the Aufbau Principle:

  • Chromium and copper

• In both cases, one 4s electron is promoted to the 3d subshell to create a more stable configuration:

  • Cr: [Ar] 3d⁵ 4s¹ (not [Ar] 3d⁴ 4s²)

  • Cu: [Ar] 3d¹⁰ 4s¹ (not [Ar] 3d⁹ 4s²)

• These configurations are preferred because a half-filled (d⁵) or fully filled (d¹⁰) d-subshell is more stable

• When transition metals form ions, they lose electrons from the 4s subshell first

• This is because, once filled, the 4s orbital is pushed to a higher energy level than the 3d due to electron repulsion

• The 4s becomes the outermost shell, so electrons are removed from it first

• As a result, +2 is a common oxidation state, due to the loss of two 4s electrons

• The ability of transition metals to form variable oxidation states is due to the similar energies of the 4s and 3d orbitals

The common oxidation states of transition elements

• Ti: +3, +4

• V: +3, +5

• Cr: +3, +6

• Mn: +2, +4, +7

• Fe: +2, +3

• Co: +2, +3

• Ni: +2

• Cu: +1, +2

Explaining variable oxidation states using successive ionisation energies

• Using titanium and vanadium as examples, the graph below shows that the first few ionisation energies are relatively small and relatively close together

  • This means that the energy difference associated with removing a small number of electrons enables transition metals to vary their oxidation state with ease

Graph of titanium and vanadium ionisation energies

• Ionisation energies increase with the number of electrons removed

• The +2 and +3 oxidation states are common across all transition elements:

  • +3 is more stable in early transition metals (up to chromium)

  • +2 becomes more stable in later elements

• Transition metal ions with oxidation states of +3 and higher are often polarising

  • They exhibit some covalent character in bonding

    Due to their high charge density, they attract bonding electrons strongly

• The maximum oxidation state for a transition metal corresponds to the total number of electrons in the 4s and 3d orbitals

  • This maximum is reached at manganese (Mn), which can achieve +7

    • Example: the manganate(VII) ion, MnO₄⁻, a strong oxidising agent