Which of the following elements does not show disproportionation tendency? (i) Cl (ii) Br (iii) F (iv) I

NCERT Exemplar · Q11

Q. Which of the following elements does not show disproportionation tendency?

(i) Cl

(ii) Br

(iii) F

(iv) I

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Concept understanding — Disproportionation Reactions

Disproportionation Reactions: The Self-Oxidation-Reduction

Imagine you have a group of friends who are all equally tall. Now imagine that, for no external reason, some of them suddenly grow taller while others shrink shorter — all starting from the same height. That sounds strange, right? But that's exactly what happens in a disproportionation reaction: the same element, in the same starting oxidation state, simultaneously gets oxidised (loses electrons) and reduced (gains electrons).

In other words, one part of the substance acts as the oxidising agent, and another part acts as the reducing agent — on itself.

The Intuition First

Think of a chemical element like chlorine. In its elemental form ( $Cl_2$ ), each chlorine atom has an oxidation state of 0. Now, if you put chlorine gas into water, something odd happens:

Some chlorine atoms gain an electron (reduction) and become $Cl^-$ (oxidation state -1).
Other chlorine atoms lose an electron (oxidation) and become $ClO^-$ (oxidation state +1).

The same starting material ( $Cl_2$ ) produces two different products — one more reduced, one more oxidised. That's disproportionation.

Note

The word "disproportionation" literally means "not in proportion" — the original uniform state splits into two different states.

The Precise Definition

Important

A disproportionation reaction is a redox reaction in which a single substance (element or compound) in an intermediate oxidation state simultaneously undergoes both oxidation and reduction, producing two different products — one with a higher oxidation state and one with a lower oxidation state.

The key conditions are:

One reactant — the same species is both oxidised and reduced.
Intermediate oxidation state — the starting element must be able to go both higher and lower.
Two products — one more oxidised, one more reduced.

The Classic Example: Chlorine in Water

Cl_2 + H_2O \rightarrow HCl + HOCl

Let's track the oxidation states:

Species	Oxidation state of Cl
$Cl_2$	0
$HCl$	-1 (reduced)
$HOCl$	+1 (oxidised)

The chlorine in $Cl_2$ (oxidation state 0) goes to -1 (gain of electron = reduction) and to +1 (loss of electron = oxidation). The same element, same starting state, two different directions.

Watch out

A common mistake is to think that because $Cl_2$ is a single molecule, it must be either oxidised or reduced. But each chlorine atom in $Cl_2$ can behave differently — one gets oxidised, the other gets reduced. The reaction as a whole is a disproportionation.

How to Spot a Disproportionation Reaction

Look for these three clues:

One reactant, two products — especially if the reactant contains an element that can exist in multiple oxidation states.
The same element appears in two different oxidation states in the products — one higher, one lower than the starting state.
No external oxidising or reducing agent — the substance does it to itself.

More Examples

Hydrogen peroxide decomposition:

2H_2O_2 \rightarrow 2H_2O + O_2

Oxygen in $H_2O_2$ has oxidation state -1. In $H_2O$ , oxygen is -2 (reduced). In $O_2$ , oxygen is 0 (oxidised). The -1 state is intermediate between -2 and 0.

Copper(I) chloride in aqueous solution:

2CuCl \rightarrow Cu + CuCl_2

Copper in $CuCl$ is +1. It goes to 0 (reduced) and +2 (oxidised).

Ammonium nitrite decomposition:

NH_4NO_2 \rightarrow N_2 + 2H_2O

Nitrogen in $NH_4^+$ is -3, in $NO_2^-$ is +3. Both go to 0 in $N_2$ — this is actually the reverse (comproportionation), but it shows how the same element can meet in the middle.

The Opposite: Comproportionation …

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Disproportionation reactions require an element to exist in at least three oxidation states. Fluorine, being the most electronegative element, only exhibits a -1 oxidation state in compounds and thus cannot disproportionate from its elemental state. The correct option is (iii).

A disproportionation reaction is a specific type of redox reaction where a single element in a particular oxidation state is simultaneously oxidised to a higher oxidation state and reduced to a lower oxidation state. For such a reaction to occur, the element must be able to exist in at least three different oxidation states: one intermediate state from which it can be both oxidised and reduced, a higher state, and a lower state.

Let's consider the halogens and their typical oxidation states:

Fluorine (F): Fluorine is the most electronegative element in the periodic table. Due to its extremely high electronegativity and the absence of vacant d-orbitals, fluorine can only exhibit a -1 oxidation state in its compounds. Its elemental form is $F_2$ , where its oxidation state is 0. From this 0 state, it can only be reduced to -1 (e.g., in $HF$ ). It cannot be oxidised to any positive oxidation state.
Chlorine (Cl), Bromine (Br), Iodine (I): These halogens are less electronegative than fluorine and possess vacant d-orbitals in their valence shell. This allows them to expand their octet and exhibit a wide range of positive oxidation states in addition to the -1 state. Their common oxidation states include -1, 0, +1, +3, +5, and +7.

For example, in an alkaline medium, chlorine ( $Cl_2$ , oxidation state 0) can disproportionate:

Cl_2 + 2NaOH \to NaCl + NaClO + H_2O

In this reaction:
*   Chlorine in $Cl_2$ (oxidation state 0) is reduced to $Cl^-$ in $NaCl$ (oxidation state -1).
*   Chlorine in $Cl_2$ (oxidation state 0) is oxidised to $ClO^-$ in $NaClO$ (oxidation state +1).

Similar disproportionation reactions are observed for bromine and iodine. For instance, $Br_2$ and $I_2$ also disproportionate in alkaline solutions.

Now, let's apply this understanding to the given options.

Define Disproportionation: A disproportionation reaction involves an element in an intermediate oxidation state simultaneously undergoing both oxidation and reduction. This means the element must be capable of existing in at least three different oxidation states.
Analyse Oxidation States of Halogens:
- Chlorine (Cl): Can exist in -1, 0, +1, +3, +5, +7 oxidation states.
- Bromine (Br): Can exist in -1, 0, +1, +3, +5, +7 oxidation states.
- Fluorine (F): Can only exist in 0 (elemental) and -1 (in compounds) oxidation states. It cannot exhibit positive oxidation states. …

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Q.Which of the following elements does not show disproportionation tendency?

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