Question

The experimental data for the reaction; $2A+B_2\to 2AB$, are as follows. The probable rate expression is:

[A]	[$B_2$]	Rate
$mollitr{e}^{-1}$	$mollitr{e}^{-1}$	$mollitr{e}^{-1}se{c}^{-1}$
0.50	0.50	$1.6\times {10}^{-4}$
0.50	1.0	$3.2\times {10}^{-4}$
1.00	1.0	$3.2\times {10}^{-4}$

• A. $dx/dt=K[B_2{]}^1$
• B. $dx/dt=K[B_2{]}^3$
• C. $dx/dt=K[B_2{]}^2$
• D. None of these

Answer 1

The correct option is A $dx/dt=K[B_2{]}^1$

When $\left[A\right]$ is kept constant at $0.50$ M and $\left[B\right]$ is doubled from $0.50$ M to $1.0$ M, the rate of the reaction is doubled from $1.6\times {10}^{-4}mollitr{e}^{-1}se{c}^{-1}$ to $3.2\times {10}^{-4}mollitr{e}^{-1}se{c}^{-1}$.

Hence, the reaction is first-order in $B$.

When $\left[B\right]$ is kept constant at $1.0$ M and $\left[A\right]$ is doubled from $0.50$ M to $1.00$ M, the rate of the reaction remains unchanged at $3.2\times {10}^{-4}mollitr{e}^{-1}se{c}^{-1}$.

Hence, the reaction is of zero-order in A.

The rate expression is $dx/dt=K\left[A{]}^0\right[B_2{]}^1=K[B_2{]}^1$