How do I solve this problem on a first order reaction?

Question

The reaction 2NO2 (g)+O3 (g)==>N2O5 (s)+O2 (g) is first order with respect to  each reactant

 

A. Write the rate expression for the reaction 

 

B. How does doubling the the concentration of NO2 affect the reaction rate?

 

C. How does tripping the concentration of O3 affect the reaction rate?

Answer 1

You need to know, and understand how to write a rate law.  It isn't that difficult if you know certain facts.  The rate law simply relates the rate of reaction to the concentration of reactants using some rate constant.  It takes the form of 

rate = k [A]ⁿ where n is an experimentally determined value, known as "the order".  So, if n=1, it is first order, if n=2 it is second order etc.

In your example, it is first order in all reactants, so the rate law would be 

rate = k[NO₂][O₃]

(B).  Since it is first order with respect to [NO₂], how do you think changing [NO₂] will affect rate?  

(C.)  Since it is first order with respect to [O₃], again, how do you think tripling it will affect rate?

HINT: Assume [NO2] and [O3] = 1 as the starting point.  Now, to answer part B, leave O3 as 1 and change NO2 to 2.  What does that do to rate in the equation.  (omit k since it is constant in both cases).  Then leave [NO2] as =1 and change [O3] to 3 to answer part C.  What does that do to rate?