The following data was collected at 20 degrees Celsius.
Determination 1, [HBPB^2-] = 7.22x10^-6 mol/L, [OH^-] = 1.00 mol/L, time = 75 s
determination 2, [HBPB^2-] = 7.22x10^-6 mol/L, [OH^-] = 0.25 mol/L, time = 290 s
Determination 3, [HBPB^2-] = 3.63x10^-6 mol/L, [OH^-] = 1.00 mol/L, time =152 s
Assume that the small, constant amount of HBPB^2- consumed in the experiment described above corresponds to an HBPB^2- concentration change of 7.22x10^-7 mol/L. Find the rate constant for this reaction.
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I'm pretty confused by this question. I'm getting conflicting information that isn't helping, so some clarification would be greatly appreciated.
With respect to HBPB, I get a -1 order reaction:
7.22x10^-6/3.63x10^-6 = 1.99^m = 75/152 = .493
Solving for m gives -1. The numbers are slightly off, but a -1 order is the closest I can possibly get.
With respect to OH:
1.00/.25 = 4^m 75/290 = .258
Again, solving for m gives -1.
This gives a rate law of: rate = k[HBPB]^-1[OH]^-1
Assuming I've done everything right up to this point, this is where I am confused by the question. I substituted the concentration change of 7.22x10^-7 M and dived that by the given times for each experiment, which gives me a reaction rate for each experiment. But when I plug the numbers in, using -1 orders for both HPBP and OH, I get different rate constants for each one, and I know this shouldn't be the case.
I realize this question is incredibly involved, but I have no clue what I'm doing wrong. Any help would be much appreciated.