Question

The rate constant is 9.50x10^-5 L/mol*s at 195 degrees Celsius and 1.30x10^-3 L/mol*s at 258 degrees Celsius. What is the activation energy of the reaction?


I'm not sure what I'm doing wrong here.
I found the natural log (ln) of: 9.50x10^-5 = -9.26
1.30x10^-3 = -6.65

and took the inverse of the temperatures:
195 = 5.13x10^-3
258 = 3.88x10^-3.

Then:
(-6.65-(-9.26) / ((3.88x10^-3)-(5.13x10^-3))) = -2088K

Ea (activation energy) is equal to slope so I did Ea = -(slope)(R = 8.314) which gave me 17kJ/mol. But I know the answer is SUPPOSED to be 85.8kJ/mol.

So what did I do wrong here? I followed the instructions in my textbook to the T and yet it came out horrible wrong.