Rate law for a reaction is $\mathrm{r}=\mathrm{k}[\mathrm{A}]^2[\mathrm{~B}]$. If rate constant is $6.25…

Rate law for a reaction is $\mathrm{r}=\mathrm{k}[\mathrm{A}]^2[\mathrm{~B}]$. If rate constant is $6.25 \mathrm{~mol}^{-2} \mathrm{dm}^6 \mathrm{~s}^{-1}$, what is the rate of reaction when $[\mathrm{A}]=1 \mathrm{~mol} \mathrm{dm}^{-3}$ and $[\mathrm{B}]=0.2 \mathrm{~mol} \mathrm{dm}^{-3}$ ?
  1. $1.250 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{sec}^{-1}$
  2. $2.125 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{sec}^{-1}$
  3. $3.105 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{sec}^{-1}$
  4. $2.0 \times 10^{-2} \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{sec}^{-1}$

Solution

$\begin{aligned} & \mathrm{r}=\mathrm{k}[\mathrm{A}]^2[\mathrm{~B}]=6.25 \times(1)^2 \times 0.2 \\ &=6.25 \times 0.2=1.250 \mathrm{~mol} \mathrm{dm}^{-3s}\end{aligned}$

Asked in: MHT CET 2024 (11 May Shift 2)

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