Q.Choose the most correct option.
A compound contains atoms of three elements A, B and C. If the oxidation state of A is +2, B is +5 and that of C is -2, the compound is possibly represented by
a. A2(BC3)2 b. A3(BC4)2 c. A3(B4C)2 d. ABC2

Oxidation number (oxidation state) is the imaginary charge left on an atom once every other atom in its compound is notionally removed, with each atom assigned its usual oxidation state by a fixed set of rules. It gives a systematic way to track electron transfer in any redox reaction, even ones involving neither oxygen nor hydrogen.

Assignment rules:

1. A free (uncombined) element has oxidation state 0 (e.g. each atom in H₂, S₈, Na).
2. A monatomic ion's oxidation state equals its net charge (e.g. Na⁺ is +1).
3. The oxidation states of all atoms in a neutral molecule sum to 0; in a polyatomic ion, they sum to the ion's net charge.
4. Hydrogen is +1 in almost every compound, except −1 in metal hydrides (e.g. NaH).
5. Fluorine is always −1.
6. Oxygen is normally −2, except −1 in peroxides (e.g. H₂O₂), −½ in superoxides (e.g. KO₂), and +2 in compounds with fluorine (e.g. OF₂).
7. Alkali metals are always +1 and alkaline earth metals are always +2 in their compounds.

Method: write an algebraic equation using rule 3, substitute the known oxidation states (from rules 4-7) for every atom except the one you need, and solve for the unknown. Example: for S in H₂SO₄, 2(+1) + x + 4(−2) = 0 ⇒ x = +6.

A reaction in which an element's oxidation number increases is an oxidation; one in which it decreases is a reduction -- this is the generalised, oxidation-number-based version of the older "loss/gain of electron" definition, and it is what lets every redox reaction (Types of Redox Reactions) be classified and every redox equation (Balancing Redox Reactions) be balanced.