Question

Explain why transition metals show variable oxidation states.

Hint:

{Transition metals = ns + (n−1)d electrons} Both participate → Variable oxidation states.

Answer 1

Concept: Transition metals are elements in which the \( (n-1)d \) orbitals are partially filled. Their electronic configuration allows multiple numbers of electrons to participate in bonding, leading to variable oxidation states.
Explanation: Transition metals exhibit variable oxidation states due to the following reasons:

• Comparable energies of orbitals: The energies of \( ns \) and \( (n-1)d \) orbitals are very close, so electrons from both can be lost during bond formation.
• Participation of d-electrons: Unlike s- and p-block elements, transition metals can use d-electrons in addition to outer s-electrons for bonding.
• Incomplete d-subshell: The partially filled d-orbitals allow different numbers of electrons to be removed or shared.
• Stability of multiple configurations: Various oxidation states can be stabilized by ligands or crystal field effects in compounds.

Examples:

• Iron: \( \text{Fe}^{2+} \) and \( \text{Fe}^{3+} \)
• Copper: \( \text{Cu}^{+} \) and \( \text{Cu}^{2+} \)
• Manganese: Shows oxidation states from +2 to +7

Conclusion: Thus, transition metals show variable oxidation states because both \( ns \) and \( (n-1)d \) electrons can participate in bonding due to their similar energies.