Q.Assertion (A): Increasing order of acidity of hydrogen halides is HF < HCl < HBr < HI
Reason (R): While comparing acids formed by the elements belonging to the same group of periodic table, H–A bond strength is a more important factor in determining acidity of an acid than the polar nature of the bond.

Concept: Bond strength and acidity in hydrogen halides

The acidity of a binary acid $\ce{H-A}$ depends on how easily it donates a proton. Two factors compete: bond polarity (which favors ionization) and bond strength (which resists it). For hydrogen halides, as we descend Group 17, the $\ce{H-X}$ bond becomes longer and weaker even though polarity decreases.

Step 1: Down the group $\ce{F} \to \ce{I}$, atomic size increases, so the $\ce{H-X}$ bond length increases and bond dissociation energy decreases sharply ($\ce{HF}: 568\,\text{kJ/mol}$ vs. $\ce{HI}: 298\,\text{kJ/mol}$).

Step 2: Weaker bonds break more easily, releasing $\ce{H^+}$ more readily. Thus $\ce{HI}$ is the strongest acid and $\ce{HF}$ the weakest among the hydrogen halides. …

Why bond strength governs acidity in hydrogen halides

Acidity measures how readily a molecule donates a proton. For hydrogen halides $\text{HX}$, two competing factors determine this:

Polarity of the H–X bond: A more polar bond (larger electronegativity difference) means the hydrogen carries a greater partial positive charge, seemingly making it easier to remove. By this logic, $\text{HF}$ should be the strongest acid because fluorine is the most electronegative.

Strength of the H–X bond: A weaker bond requires less energy to break, making proton release easier. Down the group, as the halogen atom grows larger, the overlap between the hydrogen $1s$ orbital and the halogen's valence orbital decreases. Bond strength falls dramatically: $\text{H–F} \approx 570 \, \text{kJ/mol}$, $\text{H–Cl} \approx 432$, $\text{H–Br} \approx 366$, $\text{H–I} \approx 298$.

The question is: which factor wins?

In aqueous solution, the answer is unambiguous. Bond dissociation energy dominates. The weakest bond ($\text{H–I}$) breaks most easily, making $\text{HI}$ the strongest acid. The strongest bond ($\text{H–F}$) resists breaking, making $\text{HF}$ the weakest. This is why the acidity order runs opposite to electronegativity.

Step-by-step verification

1. Examine the assertion: The order given is $\text{HF} < \text{HCl} < \text{HBr} < \text{HI}$. This matches experimental observations: $\text{HI}$ is a strong acid (completely ionized in water), $\text{HBr}$ and $\text{HCl}$ are also strong, while $\text{HF}$ is a weak acid ($\text{p}K_a \approx 3.2$). The assertion is true.

2. Examine the reason: The statement claims that for acids formed by elements in the same group, H–A bond strength is more important than bond polarity in determining acidity. This is precisely the principle that explains the hydrogen halide trend. As we move down Group 17, bond strength decreases (due to increasing atomic size and poorer orbital overlap), and this decrease in bond strength is the dominant factor that increases acidity. The reason is true. …