Medium
The rate law for the reaction $A + B \rightarrow \text{product}$ is $\text{rate} = k[A][B]$. When will the rate of reaction increase by a factor of $2$?
- A$[A]$ and $[B]$ both are doubled
- B$[A]$ is doubled and $[B]$ is kept constant
- C$[B]$ is doubled and $[A]$ is halved
- D$[A]$ is kept constant and $[B]$ is halved
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The rate equation for the reaction $2 A + B \longrightarrow$ products is $\text{rate} = k[A][B]^2$. If $k$ at $T \ K$ is $5.0 \times 10^{-6} \ mol^{-2} \ L^2 \ s^{-1}$,the initial rate of the reaction,when $[A] = 0.05 \ mol \ L^{-1}$ and $[B] = 0.1 \ mol \ L^{-1}$ is:
DifficultAP EAMCET 2009
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View SolutionConsider the following two reactions:
$A \to \text{Product}; -\frac{d[A]}{dt} = k_1[A]^0$
$B \to \text{Product}; -\frac{d[B]}{dt} = k_2[B]$
The units of $k_1$ and $k_2$ are expressed in terms of molarity $(M)$ and time $(sec^{-1})$ as:
$A \to \text{Product}; -\frac{d[A]}{dt} = k_1[A]^0$
$B \to \text{Product}; -\frac{d[B]}{dt} = k_2[B]$
The units of $k_1$ and $k_2$ are expressed in terms of molarity $(M)$ and time $(sec^{-1})$ as:
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View SolutionConsider the plots for the types of reaction $nA \to B + C$. These plots respectively correspond to the reaction orders:
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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?
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