Problem: The reaction between nitrogen dioxide and carbon monoxide is NO2(g) + CO(g) NO(g) + CO2(g). The rate constant at 701 K is measured as 2.57 M-1s-1 and that at 895 K is measured as 567 M-1s-1.Use the value of the activation energy (Ea = 1.50 x 102 kJ/mol) and the given rate constant of the reaction at either of the two temperatures to predict the rate constant at 551 K.        

🤓 Based on our data, we think this question is relevant for Professor Alghoul's class at GSU.

FREE Expert Solution

This means we need to use the two-point form of the Arrhenius Equation:

lnk2k1=-EaR1T2-1T1


where:

k1 = rate constant at T1

k2 = rate constant at T

Ea = activation energy (in J/mol) 

R = gas constant (8.314 J/mol∙K) 

T1 and T2 = temperature (in K).


We first need to convert the activation energy from kJ/mol to J/mol:

1 kJ = 103 J

Ea=1.50×102kJmol×103 J1 kJ=150,000 Jmol


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Problem Details

The reaction between nitrogen dioxide and carbon monoxide is NO2(g) + CO(g) NO(g) + CO2(g). The rate constant at 701 K is measured as 2.57 M-1s-1 and that at 895 K is measured as 567 M-1s-1.


Use the value of the activation energy (E= 1.50 x 102 kJ/mol) and the given rate constant of the reaction at either of the two temperatures to predict the rate constant at 551 K.        

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Based on our data, we think this problem is relevant for Professor Alghoul's class at GSU.