The activation energy for the following first-order reaction is 102 kJ/mol.N₂O₅(g) $\rarr$ 2NO₂(g) + ¹/₂O₂(g) The value of the rate constant (k) is 1.35 * 10^-4 s^-1 at 35^$\circ$C. What is the value of k at 0^$\circ$C

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The Arrhenius equation can be used to calculate the rate constant at a different temperature:k = Ae^(-Ea/RT)where:k is the rate constantA is the pre-exponential factorEa is the activation energyR is the gas constant (8.314 J/mol*K)T is the temperature in KelvinWe can use the given values to solve for A:1.35 * 10^-4 s^-1 = A * e^(-102000 J/mol / (8.314 J/mol*K * 308 K))A = 1.05 * 10^11 s^-1Now we can use the Arrhenius equation to calculate the rate constant at 0°C (273 K):k = 1.05 * 10^11 s^-1 * e^(-102000 J/mol / (8.314 J/mol*K * 273 K))k = 8.2 * 10^-7 s^-1Therefore, the correct answer is A.

The Activation Energy for the Following First-Order Reaction Is 102 →\rarr→

The Activation Energy for the Following First-Order Reaction Is 102 $\rarr$