13. 4 & 13. 5 Practice Problems




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13.4 & 13.5 Practice Problems

1. For the reaction

NO + O3  NO2 + O2

The following rate constants have been obtained:

Temperature, °C

k (M-1s-1)

10.0

9.30 x 106

30.0

1.25 x 107

Calculate the activation energy for this reaction in kJ/mol.


2. Calculate the activation energy (in kJ/mol) for a reaction that the rate constant is 4.60 x 10-4/s

at 350°C and 1.87 x 10-4/s at 320°C.


3. The rate constant for a first order reaction is 4.60 x 10-4/s at 250°C. If the activation energy

is 100.0 kJ/mol, calculate the rate constant at 300°C.


4. The gas-phase decomposition of nitrous oxide (N2O) is believed to occur via two elementary steps:

Step 1: N2O  N2 + O

Step 2: N2O + O  N2 + O2


Experimentally the rate law is found to be rate = k[N2O]. (A) Write the equation for the

overall reaction. (B) identify the intermediates. (C) What can you say about the relative

rates of steps 1 and 2 (which is the rate-determining step)?


5. Give the molecularity and the rate law for each of the following elementary reactions:

a. O3 + Cl  O2 + ClO

b. NO2  NO + O

c. Cl + Cl + N2  Cl2 + N2

d. 2NO + Br2  2NOBr

6. The thermal decomposition of nitryl chloride, NO2Cl, is believed to occur by the following

mechanism:

NO2Cl (g)  NO2 (g) + Cl (g)


Cl (g) + NO2Cl (g)  NO2 (g) + Cl2 (g)


  1. What is the overall reaction? (B) What is the molecularity of each of the elementary

steps? (C) What rate law is predicted by this mechanism if the first step is the rate-determining step?


7. The reaction between NO2 and CO to produce NO and CO2 is believed to occur via two steps:

Step 1: NO2 + NO2  NO + NO3


Step 2: NO3 + CO  NO2 + CO2


The experimental rate law is rate = [NO2]2. (A) Write the equation for the

overall reaction. (B) identify the intermediates. (C) What can you say about the relative

rates of steps 1 and 2 (which is the rate-determining step)?

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