In which one of the following transition metal complexes does the metal exhibit the zero oxidation state?
Answer & explanation
Correct answer: option 1
The correct answer is Option (1) → [Ni(CO)4]
The oxidation state x of the metal comes from charge balance: x + (sum of ligand charges) = charge on the complex ion.
\( \ce{[Ni(CO)4]} \): CO is neutral and the complex is neutral, so x + 4(0) = 0 and x = 0.
\( \ce{[Fe(H2O)6]Cl3} \): three \( \ce{Cl^-} \) outside give the ion a charge of +3; water is neutral, so x + 6(0) = +3 and x = +3.
\( \ce{[Co(NH3)6]Cl3} \): three \( \ce{Cl^-} \) outside give +3; ammonia is neutral, so x = +3.
\( \ce{[Fe(H2O)6]SO4} \): one \( \ce{SO4^2-} \) outside gives +2; water is neutral, so x = +2.
Only nickel in \( \ce{[Ni(CO)4]} \) is in the zero oxidation state.
NCERT (2026-27 edition):
4.3.4 Oxidation States – “Low oxidation states are found when a complex compound has ligands capable of π-acceptor character in addition to the σ-bonding. For example, in Ni(CO)4 and Fe(CO)5, the oxidation state of nickel and iron is zero.”