Answer:
Rate = k [HO2]
Rate constant = 0.8456us-1
Explanation:
Time(us) 0 0.6 1.0 1.4 1.8 2.4
[HO2](uM) 8.5 5.1 3.6 2.6 1.1 1.1
The rate law is given as;
Rate = k [HO2]^x
Where x signify the order of reaction.
For an order of reaction, the rate constant is constant for all concentrations. We are going to use this to obtain the order of reaction.
Zero Order:
[A] = [A]o -kt
5.1 = 8.5 - k(0.6)
-k (0.6) = 5.1 - 8.5
k = 5.67
3.6 = 5.1 - k(0.4)
-k (0.4) = 3.6 - 5.1
k = 3.75
The fact that the rate constant was not constant means the reaction is not a zero order reaction.
First Order:
ln[A] = ln[A]o -kt
(5.1) = ln(8.5) - k(0.6)
-k (0.6) = ln(5.1) - ln(8.5)
k = 0.8524
ln(3.6) = ln(5.1) - k(0.4)
-k (0.4) = ln(3.6) - ln(5.1)
k = 0.8708
ln(2.6) = ln(3.6) - k (0.4)
-k (0.4) = ln(2.6) - ln(3.6)
k = 0.8136
From the three calculations we see that the value of the rate constant is fairly constant in the range of 0.8 This means our reaction is a first order reaction.
The rate law is given as;
Rate = k [HO2]
We can represent the rate constant as the average of the three rate constants calculated above;
Rate constant = (0.8136 + 0.8708 + 0.8524 / 3)
Rate constant = 0.8456us-1
