Q.H2S acts only as reductant whereas SO2 acts as both oxidant and reductant. Why?
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🔒 Start your 14-day free trial to unlock the full solution →Concept understanding — Oxidation Reduction
Let’s start with something you already know from everyday life.
The intuition: what does “oxidation” really mean?
Think of a piece of iron left out in the rain. Over time, it turns into reddish-brown rust. Or think of a slice of apple turning brown when you leave it on the table. Or a fire burning wood to ash. In all these cases, something is combining with oxygen — that’s the original meaning of “oxidation.” The iron combines with oxygen from the air to form iron oxide (rust). The apple’s chemicals react with oxygen in the air. The wood burns because carbon in the wood combines with oxygen.
So the first, simplest idea: oxidation = adding oxygen. And the reverse — taking oxygen away — was called reduction. For example, if you heat iron oxide with carbon, the carbon steals the oxygen away, leaving pure iron. That’s reduction: removing oxygen.
But chemists soon realised this was too narrow. Many reactions that look like oxidation-reduction don’t involve oxygen at all. For instance, when sodium metal reacts with chlorine gas to make table salt, no oxygen is involved — yet the sodium clearly “rusts” in a sense, and the chlorine “steals” something from it.
So the definition had to be broadened.
The precise modern definition: electron transfer
Here’s the clean, exam-ready statement:
Oxidation is the loss of electrons by a substance.
Reduction is the gain of electrons by a substance.
They always happen together — you cannot have one without the other. That’s why we call them redox reactions (short for reduction-oxidation).
Let’s see this with the sodium-chlorine example:
- Sodium atom (Na) loses one electron to become Na+. That’s oxidation.
- Chlorine atom (Cl) gains that electron to become Cl− . That’s reduction.
You can write the two halves separately:
Na→Na++e−(oxidation)
Cl+e−→Cl−(reduction)
Add them together:
Na+Cl→Na++Cl−
That’s table salt.
A handy mnemonic: OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
How to spot a redox reaction without seeing electrons
You can’t watch electrons move directly. So chemists use oxidation numbers (also called oxidation states) — a bookkeeping system that tracks electrons.
Rules (simplified for first-time learners):
- An atom in its elemental form has oxidation number 0.
- A monatomic ion has oxidation number equal to its charge (e.g., Na+ is +1, Cl− is -1).
- Oxygen is usually -2 (except in peroxides).
- Hydrogen is usually +1 (except in metal hydrides).
- The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion’s charge.
Then:
- Oxidation = increase in oxidation number.
- Reduction = decrease in oxidation number.
Example: Rusting of iron.
4Fe+3O2→2Fe2O3
- Fe starts at 0 (elemental). In Fe2O3, each Fe is +3. So Fe’s oxidation number goes up from 0 to +3 → oxidation.
- O starts at 0 (in O2). In Fe2O3, each O is -2. So O’s oxidation number goes down from 0 to -2 → reduction.
A common mistake: thinking that “reduction” means something becomes smaller or less. It doesn’t — it’s about gaining electrons (or losing oxygen, in the old sense). The name comes from metallurgy: when you “reduce” iron ore to iron, you’re taking away oxygen, so the mass reduces.
One more way to think about it …
Whether a sulfur compound can act as an oxidant, a reductant, or both depends on where sulfur's oxidation state sits within its full possible range. …
Sulfur's oxidation state is -2 in H2S (its lowest state, so H2S can only reduce others) but +4 in SO2 (an intermediate state, so SO2 can both oxidise and reduce).
Sulfur can exhibit oxidation states ranging from -2 (its lowest, most reduced state) up to +6 (its highest, most oxidised state), with several intermediate states (0, +2, +4) also possible.
In H2S, sulfur has an oxidation state of -2, which is the lowest oxidation state it can have. Since there is no lower state for it to be reduced to, sulfur in H2S can only lose electrons and be oxidised (going to 0, +4, or +6 depending on the reaction), meaning H2S can only act as a reducing agent (reductant) — never as an oxidising agent.
In SO2, sulfur has an oxidation state of +4, which lies in the middle of its possible range. Sulfur here can either:
- lose more electrons and go to a higher oxidation state such as +6 (e.g. SO2 is oxidised to SO3/H2SO4 by an oxidising agent) — in this case SO2 itself acts as a reducing agent, or
- gain electrons and go to a lower oxidation state such as 0 (e.g. SO2 is reduced to S by a strong reducing agent like H2S) — in this case SO2 acts as an oxidising agent. …
Showing the 12 most recent of 39 on this concept.
- CBSE 2026Set ANNUAL1 markMCQQ.Oxidation is defined as :(a) loss of protons(b) gain of electrons(c) loss of electrons(d) gain of protons
›Reveal solutionSolution
Oxidation is the loss of electrons by an atom, ion, or molecule.
In classical terms, oxidation meant addition of oxygen or removal of hydrogen. The modern, more general electronic concept defines oxidation as any process in which a chemical species loses one or more electrons, causing its oxidation number to increase. For example, in Na → Na+ + e-, sodium is oxidised because it loses an electron.
…
- CBSE 2026Set ANNUAL1 markMCQQ.Which of the following is a redox reaction ?(a) NaBr + HCl → NaCl + HBr(b) HBr + AgNO3 → AgBr + HNO3(c) H2 + Br2 → 2HBr(d) Na2O + H2SO4 → Na2SO4 + H2O
›Reveal solutionSolution
H2 + Br2 → 2HBr is the redox reaction (oxidation states change).
A redox reaction requires a change in oxidation numbers.
- (A) NaBr + HCl → NaCl + HBr: no change; double displacement.
- (B) HBr + AgNO3 → AgBr + HNO3: precipitation, no change. …
- CBSE 2026Set ANNUAL1 markMCQQ.Which of the following acts both as oxidising as well as reducing agent ?(a) H2S(b) SO3(c) SO2(d) H2SO4
›Reveal solutionSolution
SO2 acts as both oxidising and reducing agent (S in intermediate +4 state).
An element in an intermediate oxidation state can be both oxidised and reduced.
- In H2S, S = −2 (lowest) → only reducing agent.
- In H2SO4/SO3, S = +6 (highest) → only oxidising agent. …
- CBSE 2026Set ANNUAL1 markMCQQ.The reaction, Fe^3+ + 1e → Fe^2+ is(a) Oxidation(b) Reduction(c) Both oxidation and reduction(d) None of these
›Reveal solutionSolution
Fe^3+ + e → Fe^2+ is reduction (gain of electron).
Reduction is defined as gain of electrons or a decrease in oxidation number. Here Fe^3+ gains one electron to become Fe^2+, and its oxidatio …
- CBSE 2025Set ANNUAL1 markMCQQ.In the reaction CuO(s) + H2(g) -> Cu(s) + H2O(g) oxidising agent is(a) CuO(b) H2(c) Cu(d) H2O
›Reveal solutionSolution
The reactant that gets REDUCED is the oxidising agent (it oxidises the other species while itself being reduced); here CuO loses its oxygen and is reduced to Cu, making CuO the oxidising agent.
Reaction: CuO(s) + H2(g) -> Cu(s) + H2O(g)
Track oxidation states across the reaction:
- Copper: in CuO, Cu is +2 (since O is -2 and the compound is neutral); in Cu(s), Cu is 0. Oxidation state DECREASES from +2 to 0 -> Cu is REDUCED.
- Hydrogen: in H2(g), H is 0; in H2O, H is +1. Oxidation state INCREASES from 0 to +1 -> H is OXIDISED.
Definitions:
- Oxidising agent: the species that gets reduced itself, while causing another species to be oxidised (it 'accepts' electrons) …
- CBSE 2025Set ANNUAL1 markMCQQ.3Mg(s) + N2(g) --(delta)--> Mg3N2(s) is(a) Decomposition reaction(b) Combination reaction(c) Displacement reaction(d) Combustion reaction
›Reveal solutionSolution
Two elements combining to form one compound (A + B -> AB-type) is a combination (synthesis) reaction — here Mg and N2 combine to form Mg3N2.
Reaction: 3Mg(s) + N2(g) --(heat)--> Mg3N2(s)
Classify by the pattern of reactants and products:
- Combination (synthesis) reaction: two or more substances combine to form a single product (A + B -> AB). Here, 3 moles of Mg metal and 1 mole of N2 gas combine directly into one product, Mg3N2 — matching this pattern exactly.
- Decomposition reaction: one compound breaks down into two or more simpler substances (AB -> A + B) — the OPPOSITE of what happens here.
- Displacement reaction: one element displaces another from a compound (A + BC -> AC + B) — not applicable, since there's no compound being displaced from here; both reactants are elements. …
- CBSE 2025Set sz1 markMCQQ.Select the correct one: An oxidation process involves:(a) Increase in oxidation number(b) Decrease in oxidation number(c) Addition of electrons(d) Removal of electronegative element
›Reveal solutionSolution
Oxidation is the process in which the oxidation number of an atom increases, corresponding to a loss of electrons.
In the modern (electronic) view of redox reactions, oxidation is loss of electrons by a species. Since oxidation number is assigned based on the (hypothetical) charge an atom would have if all bonds were fully ionic, losing electrons makes that charge more positive — i.e., the oxidation number increases. The reverse — decrease in oxidation number (gain of electrons) — is reduction, so option (B) and (C) describe reductio …
- CBSE 2025Set ANNUAL1 markMCQQ.The oxidizing agent in the following reaction is Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s).(a) Zn(s)(b) Cu2+(aq)(c) Zn2+(aq)(d) Cu(s)
›Reveal solutionSolution
Cu2+ gains electrons (is reduced) and therefore oxidizes Zn — Cu2+(aq) is the oxidizing agent.
An oxidizing agent is the species that gets reduced (gains electrons) while causing another species to be oxidized. Assigning oxidation states:
- Zn(s)→Zn2+(aq): oxidation state of Zn goes from 0 to +2 → Zn loses 2 electrons → Zn is oxidized → Zn is the reducing agent. …
- CBSE 2025Set ANNUAL1 markQ.The reaction Zn(s) -> Zn2+(s) + 2e- is an example of ______.
›Reveal solutionSolution
Loss of electrons is oxidation; Zn(s) -> Zn2+(s) + 2e- shows zinc being oxidised (its oxidation state rises from 0 to +2).
By the electronic concept of redox reactions, oxidation is defined as the loss of one or more electrons by a species (its oxidation number increases), and reduction is the gain of electrons (oxidation number decreases). Here, elemental zinc (oxidation state 0) loses 2 electrons to form Zn2+ (oxidation state +2), so the oxidation state increases -- thi …
- CBSE 2025Set ANNUAL1 markQ.Decrease in oxidation state of an element is reduction.
›Reveal solutionSolution
The statement is True: reduction is defined as a decrease in the oxidation state of an element (gain of electrons).
By the modern (electronic) concept of redox reactions: oxidation = increase in oxidation number (loss of electrons); reduction = decrease in oxidation number (gain of electrons). So the sta …
- CBSE 2024Set ANNUAL1 markMCQQ.In the reaction, C + H2O → CO + H2, H2O acts as(a) an oxidizing agent(b) a reducing agent(c) both (A) and (B)(d) none of these
›Reveal solutionSolution
Water oxidizes carbon to CO while its own hydrogen is reduced (+1 → 0), so H2O acts as the oxidizing agent here.
In the water-gas reaction C+H2O→CO+H2: carbon goes from oxidation state 0 (in C) to +2 (in CO) — carbon is oxidized. Correspondingly, hydrogen in H2O goes from +1 to 0 (in H2) — hydrogen is reduced, while oxygen stays at −2 throughout (H2O to CO) …
- CBSE 2024Set ANNUAL1 markMCQQ.Which of the following acts both as an oxidizing and reducing agents?(a) H2S(b) SO3(c) H2O2(d) H2SO4
›Reveal solutionSolution
H2O2's oxygen is in the intermediate −1 state, so it can act as both oxidizing and reducing agent depending on the reaction partner.
In H2O2, oxygen has oxidation state −1, midway between −2 (as in H2O) and 0 (as in O2). This means H2O2 can:
- act as an oxidizing agent, being itself reduced to H2O (O: −1 → −2), e.g. oxidizing Fe²⁺ to Fe³⁺ or I⁻ to I2; …
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