During Kinetic study of reaction $\mathrm{2 A+B \rightarrow C+D}$, the following results were obtained :
| $\mathrm{A [M]}$ | $\mathrm{B [M]}$ | initial rate of formation of $\mathrm{D}$ | |
|---|---|---|---|
| I | 0.1 | 0.1 | $6.0\times10^{-3}$ |
| II | 0.3 | 0.2 | $7.2\times10^{-2}$ |
| III | 0.3 | 0.4 | $2.88\times10^{-1}$ |
| IV | 0.4 | 0.1 | $2.40\times10^{-2}$ |
Based on above data, overall order of the reaction is _________.
Answer (integer)
3
Solution
<p>$$\begin{aligned}
& \text { Rate }=k[A]^x[B]^y \\
& \frac{6 \times 10^{-}}{2.4 \times 10^{-}}=\left(\frac{0.1}{0.4}\right)^x \Rightarrow x=1 \\
& \frac{7.2 \times 10^{-}}{2.88 \times 10^{-}}=\left(\frac{0.2}{0.4}\right)^y \Rightarrow y=2
\end{aligned}$$</p>
About this question
Subject: Chemistry · Chapter: Chemical Kinetics · Topic: Rate of Reaction
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