For the first part of the problem, we’re being asked to determine the activation energy (Ea) of the reaction.
We’re given the rate constants at two different temperatures.
This means we need to use the two-point form of the Arrhenius Equation:
k1 = rate constant at T1
k2 = rate constant at T2
Ea = activation energy (in J/mol)
R = gas constant (8.314 J/mol•K)
T1 and T2 = temperature (in K).
Rate constants for the reaction
NO2 (g) + CO (g) → NO (g) + CO2 (g)
are 1.3 M -1 s-1 at 700 K and 23.0 M -1 s-1 at 800 K.
a) What is the value of the activation energy in kJ/mol?
b) What is the rate constant at 730 K?
Express your answer using two significant figures.
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Based on our data, we think this problem is relevant for Professor Donat's class at ODU.