$$ \text { Match List I with List II : } $$
| List-I (Order of reaction) |
List-II (Unit of rate constant) |
||
|---|---|---|---|
| A. | Zero order | I. | $\mathrm{mol}^{-1} \mathrm{~L} \mathrm{~s}^{-1}$ |
| B. | First order | II. | $\mathrm{mol}^{-2} \mathrm{~L}^2 \mathrm{~s}^{-1}$ |
| C. | Second order | III. | $\mathrm{s}^{-1}$ |
| D. | Third order | IV. | $\mathrm{mol} \mathrm{L}^{-1} \mathrm{~s}^{-1}$ |
For a certain reaction $R \rightarrow$ Product, the plot of concentration $[R]$ vs time has a negative slope as shown. The order of reaction is :
Given below is an expression for the rate constant of a first-order reaction occurring at a certain temperature, $\mathrm{T}(\mathrm{K})$.
$$ \operatorname{lnk}=14.34-\frac{1.25 \times 10^4}{T} $$
The energy of activation in $\mathrm{kcal} \mathrm{mol}^{-1}$ for the reaction is :
(Given: $\mathrm{k}^{-1} \mathrm{~s} \mathrm{~s}^{-1}, \mathrm{R}=1.987 \mathrm{cal} \mathrm{mol}^{-1} \mathrm{~K}^{-1}$ )
If the rate constant of a reaction is $0.03 \mathrm{~s}^{-1}$, how much time does it take for $7.2 \mathrm{~mol} \mathrm{~L}^{-1}$ concentration of the reactant to get reduced to $0.9 \mathrm{~mol} \mathrm{~L}^{-1}$ ? (Given: $\log 2=0.301$ )
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