The reaction $2 A + B + C \longrightarrow D + E$ is found to be first order in $A$,second order in $B$,and zero order in $C$. What is the effect of increasing the concentration of all reactants twice?
- ARate of reaction increases $8$ times.
- BRate of reaction increases $24$ times.
- CRate of reaction increases $36$ times.
- DRate of reaction remains unaffected.
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Similar Questions
The rate constant for a reaction is $10.8 \times 10^{-5} \ mol \ L^{-1} \ s^{-1}$. The order of the reaction is .......
Easy
View Solution$A_2 + B_2 \to 2AB$; $R.O.R = k[A_2]^a[B_2]^b$
Initial $[A_2]$ Initial $[B_2]$ $R.O.R. (r) \ M s^{-1}$ $0.2$ $0.2$ $0.04$ $0.1$ $0.4$ $0.04$ $0.2$ $0.4$ $0.08$
Order of reaction with respect to $A_2$ and $B_2$ are respectively:
| Initial $[A_2]$ | Initial $[B_2]$ | $R.O.R. (r) \ M s^{-1}$ |
|---|---|---|
| $0.2$ | $0.2$ | $0.04$ |
| $0.1$ | $0.4$ | $0.04$ |
| $0.2$ | $0.4$ | $0.08$ |
Order of reaction with respect to $A_2$ and $B_2$ are respectively:
Medium
View SolutionConsider the following data for the given reaction $2 HI_{(g)} \rightarrow H_{2_{(g)}} + I_{2_{(g)}}$. The order of the reaction is:
Experiment $1$ $2$ $3$ $[HI] \ (mol \ L^{-1})$ $0.005$ $0.01$ $0.02$ Rate $(mol \ L^{-1} \ s^{-1})$ $7.5 \times 10^{-4}$ $3.0 \times 10^{-3}$ $1.2 \times 10^{-2}$
| Experiment | $1$ | $2$ | $3$ |
|---|---|---|---|
| $[HI] \ (mol \ L^{-1})$ | $0.005$ | $0.01$ | $0.02$ |
| Rate $(mol \ L^{-1} \ s^{-1})$ | $7.5 \times 10^{-4}$ | $3.0 \times 10^{-3}$ | $1.2 \times 10^{-2}$ |
DifficultJEE MAIN 2024
View SolutionFor the reaction $2A + B \rightarrow C$,the rate law is given by: $\text{Rate} = k[A][B]$. Which of the following statements is correct regarding this reaction?
Easy
View SolutionWhich one of the following is wrong about molecularity of a reaction?
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