Arrhenius Equation Problem?

Question

I am having trouble with the Arrhenius Equation. Here is the problem I must solve. Please explain how do to it.

A reaction rate increases by a factor of 800 in the presence of a catalyst at 34°C.
The activation energy of the original, uncatalysed, pathway is 119 kJ mol-1.
What is the activation energy of the catalyzed pathway, all other factors being equal?

Also, K = Ae^Ea/RT what does the e stand for? This is what is confusing me

Answer 1

e stands for exponential (2.71...)

You have two different activation energies, uncatalyzed is 119, catalyzed is unknown. You know the ratio of rates is 800, therefore

exp(Ea/R*307K) / exp (119/R*307K) = 800.

This uses the fact that 34 C = 273 + 34 K.

R is a constant. Plug in R and solve.

Bozo

Answer 2

k = e ^ [ln A + (Ea/RT)]