- #1
kazimmerman
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Homework Statement
This is straight off an old AP Exam, but I can't seem to find it online so I'm hoping you all can help me out.
2 NO(g) + Br2(g) --> 2 NOBr(g)
The following results were obtained in experiments designed to study the rate of the reaction above.
Experiment Initial Concentration (mol/L) Initial Rate of Appearance of NOBr(M/sec)
[NO] [Br2]
1 0.02 0.02 9.6e-2
2 0.04 0.02 3.8e-1
3 0.02 0.04 1.9e-1
(I'm not sure how to correctly display tables, so hopefully you can understand what everything is.
a) Write the rate law for the reaction.
b) Calculate the value of the rate constant, k, for the reaction. Include units.
c) In experiment 2, what was the concentration of NO remaining when half of the original amount of Br2 was consumed?
d) Which of the following reaction mechanisms is consistent with the rate law established in (a)? Explain your choice.
I. NO + NO --> N2O2 (fast)
N2O2 + Br2 --> 2 NOBr (slow)
II. Br2 --> Br + Br (slow)
2 (NO + Br --> NOBr) (fast)
Homework Equations
Rate = k[A]^n^m
The Attempt at a Solution
a) I've figured this out and I know I am correct with the following rate law:
Rate = k[NO]^2[Br2]
b) I also have this one confirmed correct:
k = 1.2e4 L^2 mol^-2 s-1
c) This is where my problems begin. I am sure I should probably be splitting the rate law and using integrated rate laws and/or half-life equations, but I'm not completely sure where to begin. I know for second-order rates, half-life is:
t = (k[A](initial))^-1
and for a first-order reaction:
t = ln(2) k^-1
d) I understand mechanisms a bit, but I guess I don't know how to differentiate between which should occur.
Thanks ahead of time for any help. ;)