(A) Using the Arrhenius equation: $\log \frac{k_2}{k_1} = \frac{E_a}{2.303 R} \left( \frac{T_2 - T_1}{T_1 T_2} \right)$ Given: $k_2 = 2k_1$,$T_1 = 300

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(A) Using the Arrhenius equation: $\log \frac{k_2}{k_1} = \frac{E_a}{2.303 R} \left( \frac{T_2 - T_1}{T_1 T_2} \right)$ Given: $k_2 = 2k_1$,$T_1 = 300 \ K$,$T_2 = 310 \ K$,$R = 8.314 \ J \ K^{-1} \ mol^{-1}$. $\log 2 = \frac{E_a}{2.303 \times 8.314} \left( \frac{310 - 300}{300 \times 310} \right)$ $0.301 = \frac{E_a}{19.147} \left( \frac{10}{93000} \right)$ $E_a = \frac{0.301 \times 19.147 \times 93000}{10} \approx 53598 \ J \ mol^{-1} \approx 53.6 \ kJ \ mol^{-1}$.

Class 12 Chemistry Chemical Kinetics Collision theory, Energy of activation and Arrhenius equation

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The rate of a reaction doubles,when the temperature is changed from $300 \ K$ to $310 \ K$. Activation energy of the reaction is....... $(R=8.314 \ J \ K^{-1} \ mol^{-1}, \log 2=0.301)$

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$53.6 \ kJ \ mol^{-1}$
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$48.6 \ kJ \ mol^{-1}$
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$58.5 \ kJ \ mol^{-1}$
D
$60.5 \ kJ \ mol^{-1}$

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Chemical Kinetics

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Collision theory, Energy of activation and Arrhenius equation

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The Arrhenius equation is represented as $k = A e^{-E_a/RT}$. The activation energy $E_a$ of the reaction can be calculated by plotting:

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In a reaction,for every $10\,^{\circ}C$ rise of temperature,the rate is doubled. If the temperature is increased from $10\,^{\circ}C$ to $100\,^{\circ}C,$ the rate of the reaction will become $.......$ times.

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Which one of the following is true for an exothermic reaction $A \rightleftharpoons B$, if $E_f$ and $E_b$ are the activation energies of forward and backward reactions respectively?

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The rate constant $(k)$ of a reaction is measured at different temperatures $(T),$ and the data are plotted in the given figure. The activation energy of the reaction in $kJ\, mol^{-1}$ is :
($R$ is gas constant)

MediumJEE MAIN 2020

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What is the activation energy $(kJ \, mol^{-1})$ for a reaction if its rate constant doubles when the temperature is raised from $300 \, K$ to $400 \, K$ ? $(R = 8.314 \, J \, mol^{-1} \, K^{-1})$

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The rate of a reaction doubles,when the temperature is changed from $300 \ K$ to $310 \ K$. Activation energy of the reaction is....... $(R=8.314 \ J \ K^{-1} \ mol^{-1}, \log 2=0.301)$

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The Arrhenius equation is represented as $k = A e^{-E_a/RT}$. The activation energy $E_a$ of the reaction can be calculated by plotting:

In a reaction,for every $10\,^{\circ}C$ rise of temperature,the rate is doubled. If the temperature is increased from $10\,^{\circ}C$ to $100\,^{\circ}C,$ the rate of the reaction will become $.......$ times.

Which one of the following is true for an exothermic reaction $A \rightleftharpoons B$, if $E_f$ and $E_b$ are the activation energies of forward and backward reactions respectively?

The rate constant $(k)$ of a reaction is measured at different temperatures $(T),$ and the data are plotted in the given figure. The activation energy of the reaction in $kJ\, mol^{-1}$ is :($R$ is gas constant)

What is the activation energy $(kJ \, mol^{-1})$ for a reaction if its rate constant doubles when the temperature is raised from $300 \, K$ to $400 \, K$ ? $(R = 8.314 \, J \, mol^{-1} \, K^{-1})$

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The rate constant $(k)$ of a reaction is measured at different temperatures $(T),$ and the data are plotted in the given figure. The activation energy of the reaction in $kJ\, mol^{-1}$ is :
($R$ is gas constant)