In the reaction $\mathrm{H}_2\mathrm{S}(g) + \mathrm{Cl}_2(g) \rightarrow 2\mathrm{HCl}(g) + \mathrm{S}(s)$, which species undergoes oxidation?

1. Track oxidation numbers. In $\mathrm{H}_2\mathrm{S}$, S is $-2$; in product S(s), S is $0$. So S goes from $-2 \rightarrow 0$, an INCREASE in oxidation state. 2. Increase in oxidation state is the modern definition of oxidation. 3. Equivalently (classical view), $\mathrm{H}_2\mathrm{S}$ loses hydrogen to chlorine (a more electronegative element), so it is oxidised. 4. $\mathrm{Cl}_2$ goes from $0 \rightarrow -1$ in HCl, a DECREASE, i.e. reduction (option B reverses this). 5. Since both half-reactions occur simultaneously, the reaction IS redox, ruling out option D. Option C is wrong because only S is oxidised; Cl is reduced. _Source: NCERT Class 11 Chemistry Part 2, Ch 7, Problem 7.1 (i), p. 1–2._