Which of the following rate laws is possible for the following two-step reaction?
1) NO(g) + NO(g) N2O2(g); FAST
2) N2O2(g) + O2(g) 2NO2(g); SLOW
Net: 2NO(g) + O2(g) 2NO2(g);
A. k[N2O2][NO]
B. k[NO]2
C. k[N2O2][O2]
D. k[NO][O2]
help!!!!! i don't know how to figure this out
k[NO]^2[O2]
Just took the quiz
Here is one of the best sites I've found that explains how to write the rate law. Be sure and read the first part in order to get the essentials but the second part is the one you want. It explains how to do problems in which the first step is fast and the second step is slow.
https://courses.lumenlearning.com/boundless-chemistry/chapter/reaction-mechanisms/
Post your work if you get stuck but remember that the rate law may NOT contain intermediates which means any equation containing (N2O2) isn't right.
Check your post and the problem you copied. I don't believe the correct answer is listed. Perhaps you made a typo.
These are the only options! Thank you anyway :)
you basically look at the reactants of the slow reaction they should always be in the rate law, their coifficient determines what order they are therefore A is the correct answer
To determine the rate law for a reaction, you need to analyze the reaction mechanism and identify the slowest step (also known as the rate-determining step). The coefficients of the reactants in the rate-determining step will determine the order of the rate law with respect to each reactant.
In this reaction, the slow step is the second step where N2O2 reacts with O2 to form 2NO2. The coefficients of the reactants in this slow step are 1 for N2O2 and 1 for O2. This means that the rate law for this step will have the form:
rate = k[N2O2][O2]
However, we need to consider the overall stoichiometry of the reaction. The balanced equation shows that 2 moles of NO are produced for every mole of N2O2 and O2 reacted. Therefore, the rate of appearance of NO (the product) is twice the rate of disappearance of N2O2 and O2.
So, the rate law for the overall reaction, which includes both steps, will be:
rate = 2k[N2O2][O2]
Comparing this rate law with the given options:
A. k[N2O2][NO] - This rate law does not match the actual rate law because it includes NO instead of O2.
B. k[NO]^2 - This rate law does not match the actual rate law because it includes NO squared instead of N2O2 and O2.
C. k[N2O2][O2] - This rate law matches the actual rate law, as determined earlier.
D. k[NO][O2] - This rate law does not match the actual rate law because it includes NO instead of N2O2.
Therefore, the correct rate law for this two-step reaction is option C, which is k[N2O2][O2].