Question

The rate of disappearance of A at two temperatures is given by $A\rightleftharpoons B$
i. $\frac{-d\left[A\right]}{dt}=2\times {10}^{-2}\left[A\right]-4\times {10}^{-3}\left[B\right]$ at 300 K
ii. $\frac{-d\left[A\right]}{dt}=4\times {10}^{-2}\left[A\right]-16\times {10}^{-4}\left[B\right]$ at 300 K
From the given values of heat of reaction which are incorrect

• A. $3.86kcal$
• B. $6.93kcal$
• C. $1.68kcal$
• D. $1.68\times {10}^{-2}kcal$

Answer 1

Answer: (B), (C), (D)

$1.68\times {10}^{-2}kcal$
$6.93kcal$
$1.68kcal$

$K_1=\frac{K_f}{K_b}=\frac{2\times {10}^{-2}}{4\times {10}^{-3}}=5$ at 300 K
$K_2=\frac{K_f}{K_b}=\frac{4\times {10}^{-2}}{16\times {10}^{-4}}=25$ at 400 K
$\therefore 2.303log\frac{25}{55}=\frac{\Delta H}{2}\times \left[\frac{400-300}{400\times 300}\right]$
or $\Delta H=3.85kcal$

Hence, option A is correct and others are incorrect