While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions

Exercises · 7.8

Q. While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why ?

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Concept understanding — Oxidation State Stability

Oxidation State Stability: The First Meeting

Imagine you have a pile of Lego bricks. Some colours click together easily; others keep falling apart. An atom's oxidation state is like the colour of a Lego brick — it tells you how many electrons the atom has lost (positive state) or gained (negative state) compared to its neutral form. But not all oxidation states are equally happy. Some are rock-solid; others are fragile and want to change.

That's oxidation state stability: why some oxidation states of an element are comfortable and long-lasting, while others are restless and reactive.

The Intuition: Why would an atom "prefer" one state over another?

An atom wants to be stable. In chemistry, stability usually means one of two things:

A full outer shell — like noble gases (8 electrons in the outermost shell, or 2 for helium). Atoms will gain, lose, or share electrons to get there.
A half-filled or fully-filled subshell — this is a special bonus stability. For example, chromium loves having a $3d^5$ configuration (half-filled d-subshell) even if it means sacrificing a 4s electron.

So when an atom takes on an oxidation state, it's essentially deciding how many electrons to give away or accept. The "best" oxidation states are those that leave the atom with a stable electronic configuration.

Note

This is why sodium (Na) almost always shows +1 — losing one electron gives it the same electron configuration as neon (a noble gas). It's not trying to be fancy; it's trying to be comfortable.

The Precise Statement

Oxidation state stability refers to the tendency of an element to exist in a particular oxidation state under given conditions (temperature, pH, presence of other substances). An oxidation state is stable if:

The atom's electronic configuration after gaining/losing electrons is noble-gas-like (octet) or has a half-filled/full-filled d or f subshell.
The energy required to reach that state is low compared to other possible states.
The state does not spontaneously change into another state (e.g., by reacting with air, water, or itself).

Stability order (general trend): For main-group elements: oxidation states that give a noble gas configuration are most stable. For transition metals: +2 and +3 are common, but stability varies with the element and the environment.

Examples to cement the idea

Element	Common stable states	Why?
Sodium (Na)	+1	Losing 1 electron → Ne configuration (2,8)
Chlorine (Cl)	-1	Gaining 1 electron → Ar configuration (2,8,8)
Iron (Fe)	+2, +3	Both leave d-subshell partially filled; +3 is more stable in acidic conditions
Manganese (Mn)	+2, +7	+2 is stable (half-filled d⁵); +7 is stable in permanganate ion (MnO₄⁻) due to strong bonding with oxygen
Carbon (C)	+4, -4	Both give noble gas configuration (He for +4? No — careful: carbon's +4 means losing 4 electrons, leaving 1s², which is He-like. -4 means gaining 4 electrons, giving Ne-like 2,8)

Watch out

Don't confuse "common" with "stable". Carbon's +4 is common but not always stable — CO₂ is stable, but CCl₄ is not very stable in water. Stability depends on the compound, not just the element.

The deeper reason: Why do some states "disproportionate"?

Some oxidation states are so unstable that the atom reacts with itself. This is called disproportionation — one atom in an intermediate state splits into two atoms: one higher, one lower.

Example: Copper(I) in water. Cu⁺ (oxidation state +1) is unstable in aqueous solution. It spontaneously does this:

2\text{Cu}^+ \rightarrow \text{Cu}^{2+} + \text{Cu}

One Cu⁺ gets oxidised to Cu²⁺ (more stable in water), the other gets reduced to Cu metal (even more stable). The +1 state was disproportionation-unstable.

Tip

A quick check: If an element has three consecutive oxidation states (say +1, +2, +3), the middle one (+2) is often prone to disproportionation if the +1 and +3 states are both stable. This is common for elements like copper, mercury, and thallium.

What determines stability in practice? …

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The ability to act as both an oxidant and a reductant depends on whether the element in the compound has an intermediate oxidation state that can be both raised and lowered. SO₂ and H₂O₂ have central atoms in intermediate states (+4 for S in SO₂, –1 for O in H₂O₂), while O₃ and HNO₃ have atoms in their highest possible oxidation states (0 for O in O₃, +5 for N in HNO₃), so they can only be reduced — they act only as oxidants.

The key idea here is oxidation state stability. A substance can act as both an oxidising agent (gets reduced) and a reducing agent (gets oxidised) only if the element responsible for the redox behaviour is in an intermediate oxidation state — one that can be either increased or decreased. If the element is already in its highest possible oxidation state, it can only be reduced (acts only as an oxidant). If it is in its lowest possible oxidation state, it can only be oxidised (acts only as a reductant).

Let’s examine each compound.

Sulphur dioxide ( $\text{SO}_2$ )

Sulphur here has an oxidation state of $+4$ . Sulphur can have oxidation states from $-2$ (e.g., in $\text{H}_2\text{S}$ ) up to $+6$ (e.g., in $\text{H}_2\text{SO}_4$ ).
- $+4$ is intermediate: it can be reduced to lower states (like $0$ in elemental S, or $-2$ in sulphides) — so $\text{SO}_2$ acts as an oxidant.
- It can also be oxidised to $+6$ (in $\text{SO}_3$ or $\text{H}_2\text{SO}_4$ ) — so $\text{SO}_2$ acts as a reductant. Hence $\text{SO}_2$ shows dual behaviour.
Hydrogen peroxide ( $\text{H}_2\text{O}_2$ )

Oxygen here has an oxidation state of $-1$ . Oxygen can range from $-2$ (in water, oxides) to $0$ (in $\text{O}_2$ ).
- $-1$ is intermediate: it can be reduced to $-2$ (forming $\text{H}_2\text{O}$ or $\text{OH}^-$ ) — so $\text{H}_2\text{O}_2$ acts as an oxidant.
- It can also be oxidised to $0$ (evolving $\text{O}_2$ gas) — so $\text{H}_2\text{O}_2$ acts as a reductant. Again, dual behaviour.
Ozone ( $\text{O}_3$ )

Ozone is an allotrope of oxygen. The oxidation state of each oxygen atom in $\text{O}_3$ is $0$ (since it’s an element in its standard state).
- $0$ is effectively the highest oxidation state oxygen reaches in ordinary chemistry — only fluorine can force oxygen positive (OF₂ has O at +2, O₂F₂ at +1), and no fluorine species is involved here.
- Therefore, in practice, oxygen in $\text{O}_3$ cannot be oxidised further — it can only be reduced (to $-2$ in oxides or water, or $-1$ in peroxides). So $\text{O}_3$ acts only as an oxidant.
Nitric acid ( $\text{HNO}_3$ )

Nitrogen here has an oxidation state of $+5$ . Nitrogen can range from $-3$ (in $\text{NH}_3$ ) to $+5$ (in nitrates).
- $+5$ is the highest possible oxidation state for nitrogen. …

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Q.While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in their reactions, ozone and nitric acid act only as oxidants. Why ?

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