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(A) According to the Arrhenius equation, the rate constant $k$ is given by $k = A e^{-E_a/RT}$. Since the activation energy $E_a$ is always positive,

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Equilibrium constant may increase or decrease but rate constant always increases with temperature

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(A) According to the Arrhenius equation, the rate constant $k$ is given by $k = A e^{-E_a/RT}$. Since the activation energy $E_a$ is always positive, the rate constant $k$ always increases with an increase in temperature. According to the Van't Hoff equation, $\log \left(\frac{K_2}{K_1}\right) = \frac{\Delta H}{2.303 R} \left[\frac{1}{T_1} - \frac{1}{T_2}\right]$. If the reaction is endothermic $(\Delta H > 0)$, the equilibrium constant $K$ increases with temperature. If the reaction is exothermic $(\Delta H Therefore, the equilibrium constant may increase or decrease depending on the enthalpy change, while the rate constant always increases with temperature.

Which of the following statements is true about the equilibrium constant and rate constant of a single-step chemical reaction?

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The rate constant of a reaction varies with temperature according to the equation: $\log K = \text{constant} - \frac{E_a}{2.303 RT}$. If a plot of $\log K$ versus $1/T$ yields a straight line with a slope of $-5632$,then the activation energy of the reaction is .......... $kJ \ mol^{-1}$.

For the forward reaction $X \rightarrow Y$,the activation energy is $60 \ kJ \ mol^{-1}$ and $\Delta H = -20 \ kJ \ mol^{-1}$. What is the activation energy for the reverse reaction $Y \rightarrow X$ in $kJ \ mol^{-1}$?

$A$ reaction having equal activation energies for forward and reverse reaction has:

The following energy profile diagram is given for the reaction $A + B \to C + D$. The enthalpy change and activation energy for the reverse reaction $C + D \to A + B$ are respectively:

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