Rate vs concentration data was collected for the following reaction:

 

[ICl]
[H2]
-d[H2]/dt s-1
0.0015
0.0015
3.7 x 10-7
0.0030
0.0015
7.4 x 10-7
0.0030
0.0045
22 x 10-7
0.0047
0.0027
?

Determine the rate law for the above reaction; include the value of the rate constant.

first must determine order of reaction...

0.5 = 0.5y

No need to take logs this time!

y = 1

Hey look the order is different than the stoichiometry!

 

Do the same thing for H2.

so...

2.972 = 3x

another tough one...

x = 1

So, the rate law is

What is the value of the rate constant k? (plug and chug)

What will the initial rate be for the last set of concentration data? (plug and chug)

 

At what rate is HCl is being produced when [H2] = 0.0027 M and [ICl] = 0.0047 M? (you must find d[HCl]/dt)

From the stoichiometry you know that

d[HCl]/dt = 2(&endash;d[H2]/dt)

so...

d[HCl]/dt = 2(2.09x 10&endash;6)

 

or...

d[HCl]/dt = 4.17 x 10&endash;6 => 4.2 x 10&endash;6 M/s

 

What is the rate law if rate = d[HCl]/dt?

The order of the reaction does not change but is k = k'? No, they cannot be equal because the rate of production of HCl is equal to 2 times the rate of consumption of H2. So, k' must be twice as large as k.

So, with k' we have...