Identify the ions and oxidation states of chlorine

• Hypochlorite: $\mathrm{ClO^-}$, chlorine is in oxidation state $+1$.
• Chlorate: $\mathrm{ClO_3^-}$, chlorine is in oxidation state $+5$.
• Perchlorate: $\mathrm{ClO_4^-}$, chlorine is in oxidation state $+7$.

Option A: The hypochlorite ion is the strongest conjugate base.

Compare the parent acids:

• $\mathrm{HClO}$ (hypochlorous acid)
• $\mathrm{HClO_3}$ (chloric acid)
• $\mathrm{HClO_4}$ (perchloric acid)

Acid strength increases strongly with oxidation state / number of oxygens: $\mathrm{HClO_4 > HClO_3 > HClO}$

Therefore, conjugate base strength is the reverse order: $\mathrm{ClO^- > ClO_3^- > ClO_4^-}$

So, hypochlorite ion is the strongest conjugate base.

✅ A is correct.

Option B: The molecular shape of only chlorate ion is influenced by the lone pair of electrons of Cl.

Let us examine shapes:

• Hypochlorite, $\mathrm{ClO^-}$: diatomic ion, so molecular shape discussion in terms of bond angle distortion by lone pair is not meaningful.
• Chlorate, $\mathrm{ClO_3^-}$: central Cl has three bonded oxygens and one lone pair, so electron pair geometry is tetrahedral and molecular shape is trigonal pyramidal. Hence shape is influenced by a lone pair.
• Perchlorate, $\mathrm{ClO_4^-}$: four bonded oxygens, no lone pair on central Cl in VSEPR description; shape is tetrahedral.

Thus, among these three ions, only chlorate ion has its molecular shape influenced by a lone pair on Cl.

✅ B is correct.

Option C: The hypochlorite and chlorate ions disproportionate to give rise to identical set of ions.

Check disproportionation reactions.

• Hypochlorite disproportionation: $3\mathrm{ClO^-} \rightarrow 2\mathrm{Cl^-} + \mathrm{ClO_3^-}$ Products: $\mathrm{Cl^-}$ and $\mathrm{ClO_3^-}$

• Chlorate disproportionation: chlorine in $+5$ state can disproportionate to a lower and higher oxidation state, typically giving $4\mathrm{ClO_3^-} \rightarrow 3\mathrm{ClO_4^-} + \mathrm{Cl^-}$ Products: $\mathrm{Cl^-}$ and $\mathrm{ClO_4^-}$

The product sets are not identical.

❌ C is incorrect.

Option D: The hypochlorite ion oxidises the sulphite ion.

Sulphite ion is $\mathrm{SO_3^{2-}}$, where sulfur is in oxidation state $+4$. It can be oxidised to sulfate, $\mathrm{SO_4^{2-}}$, where sulfur is in oxidation state $+6$.

Hypochlorite ($\mathrm{ClO^-}$, Cl at $+1$) is a known oxidising agent and can be reduced to chloride ($\mathrm{Cl^-}$, Cl at $-1$).

A representative redox process is: $\mathrm{ClO^- + SO_3^{2-} \rightarrow Cl^- + SO_4^{2-}}$

This is balanced in atoms and charge:

• Left charge: $-1 + (-2) = -3$
• Right charge: $-1 + (-2) = -3$

Hence hypochlorite indeed oxidises sulphite to sulfate.

✅ D is correct.

Final selection

Correct statements are: $\boxed{\text{A, B, D}}$

Comparison with stored correct answer

Stored correct answer: A, B, D

My derived answer matches exactly.