Question

Reaction; $A+B⟶C+D$, follows rate law; $rate=k{\left[A\right]}^2{\left[B\right]}^2$. Starting with initial concentration of $1M$ of $A$ and $B$ each, what is the time taken for concentration of $A$ to become $0.25M$?
Given: $k=2.303\times {10}^{-2}{sec}^{-1}$

• A. $300sec$
• B. $600sec$
• C. $900sec$
• D. $1200sec$

Answer 1

Answer: (B)

$600sec$

$K=2.303\times {10}^{-3}{sec}^{-1}$

Since the unit of $K$ is ${sec}^{-1}$, the order of the reaction will be $1$.

$t=\frac{2.303}{K}log\frac{a}{a_t}$

$\Rightarrow t=\frac{2.303}{2.303\times {10}^{-3}}log\frac{1}{0.25}=600sec$

[May be wrong data given here $K=2.303\times {10}^{-3}{sec}^{-1}$ considered]