Which is the stronger acid, \({\rm{HCl}}{{\rm{O}}_4}\)or \({\rm{HBr}}{{\rm{O}}_4}\)? Why?

An acid, according to Bronsted-Lowry, is a substance that loses or donates a proton \(({{\rm{H}}^ + })\)in solution, whereas a base takes a proton

Chlorine Cl and bromine Br are halogens with valence electron configurations of \({\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{5}}}\)and \({\rm{4}}{{\rm{s}}^{\rm{2}}}{\rm{4}}{{\rm{p}}^{\rm{5}}}\), respectively, that belong to group VII A.

The capacity of an acid to lose a proton determines its strength. The acid dissociation of perchloric \(\left( {{\rm{HCl}}{{\rm{O}}_4}} \right)\)and perbromic \(\left( {{\rm{HBr}}{{\rm{O}}_4}} \right)\)is represented as:

\(\begin{array}{*{20}{l}}{{\rm{HCl}}{{\rm{O}}_4}}&{ \leftrightarrow {{\rm{H}}^ + } + {\rm{ClO}}_4^ - }\\{{\rm{HBr}}{{\rm{O}}_4}}&{ \leftrightarrow {{\rm{H}}^ + } + {\rm{BrO}}_4^ - }\end{array}\)

The electronegativity of elements decreases as they progress through the groups. Because Cl is more electronegative than Br, it will bring the electron density in the \({\rm{H}} - {\rm{O}} - {\rm{Cl}}\) bond closer to it. In comparison to \({\rm{HBr}}{{\rm{O}}_4}\), this would encourage the release of \({\rm{H}} + \)ions more readily in \({\rm{HCl}}{{\rm{O}}_4}\).