Hey everyone, I've been reviewing kinetics and equilibrium material, and I've noticed something that hopefully someone can clear up.
For an equation like this: aA + bB <--> cC + dD
The rate laws for the forward (f) and reverse (r) reactions would be, respectively:
Rate(f) = kf[A]^n^m, and Rate(r) = kr[C]^x[D]^y, where n, m, x, and y are the exponents which are determined experimentally, not from the stoichiometric coefficients.
The relationship between kinetics and equilibrium in this scenario is Rate(f) = Rate(r), therefore kf/kr, which equals Keq. However I was always taught that the equilibrium expression was:
Keq = [C]^c[D]^d / [A]^a^b, where c, d, a, and b are the stoichiometric coefficients.
My question is if kf/kr = Keq, why are the exponents different? I can't seem to find an explanation for this, but the only thing I can think of is that the above equilibrium expression exponents are actually the same as the experimentally determined rate law exponents, but only for single elementary step reactions are they the same as the stoichiometric coefficients.
If this is true, would that mean for a reaction 2A + B <---> C + D, if the forward rate law was determined to be k or k[A], would the correct Keq equations be something like this?
Keq = [C][D]/, or Keq = [C][D]/[A]
Thanks!