Question

A reaction is catalysed by $H^{+}$ ion. In the presence of an acid $HA$, the rate constant is $2\times {10}^{-3}mi{n}^{-1}$ and in the presence of an acid $HB$, the rate constant is $1\times {10}^{-3}mi{n}^{-1}$. Given $HA$ and $HB$ both being strong acids, we may conclude that:

• A. equilibrium constant is 2
• B. $HA$ is stronger acid than $HB$
• C. relative strength of $HA$ and $HB$ is $2$
• D. $HA$ is weaker acid than $HB$ and their relative strength is $0.5$

Answer 1

Answer: (B), (C)

relative strength of $HA$ and $HB$ is $2$
$HA$ is stronger acid than $HB$

The rate constant of the reaction in the presence of $HA$ is $2\times {10}^{-3}mi{n}^{-1}.$

This value is twice the rate constant for the reaction in the presence of $HB$, which is $1\times {10}^{-3}mi{n}^{-1}.$

Greater the value of rate constant of acid, greater is the acidic strength.

Thus, HA is a stronger acid than HB and relative strength of $HA$ and $HB$ is 2.