Question

The experimental data for the reaction
$2A+B_2\to 2AB$ is as follows

[A]	$\left[B_2\right]$	Rate
0.50M	0.50M	$1.6\times {10}^{-4}{s}^{-1}$
0.50M	1.00M	$3.2\times {10}^{-4}{s}^{-1}$
1.00M	1.00M	$3.2\times {10}^{-4}{s}^{-1}$

The most probable equation for the rate of the reaction is :

• A. $k\left[B_2\right][A{]}^2$
• B. $k\left[B_2\right]$
• C. $k[A{]}^2$
• D. $k\left[A\right]\left[B\right]$

Answer 1

The correct option is B $k\left[B_2\right]$

When the concentration of $B_2$ is doubled, the rate is doubled so the order with respect to $B_2$ is one.

When the concentration of A is doubled, the rate remains unaltered so the order with respect to $A$ is 0.

So, $-\frac{dx}{dt}=k\left[A{]}^o\right[B_2{]}^1=k\left[B_2\right]$

Option B is correct.