Units for dissociation constant and relationship to Gibbs Free Energy How does one "convert" between an equilibrium constant calculated from the Gibbs free energy for a dissociation reaction and experimentally determined dissociation constants?

$$
\Delta G^\circ=-RT\log(K_{eq})
$$

I think I understand the derivation of this equation and why the equilibrium constant should be unitless if it's defined in terms of activities.
But I have a binding reaction that has an estimated free energy change of $-32~\mathrm{kcal/mol}$, which yields a dissociation constant on the order of $10^{-24}$ at $298~\mathrm{K}$.
Is it correct to just multiply the equilibrium constant by the standard concentration (1 molar) to get a more typical dissociation constant with units? A yoctomolar dissociation constant seems kind of strange to me, but I'm not sure where I've made a mistake.
free energy change of -32 kcal/mol, which yields something on the order of 10^-24 at 298K. 


Seems the main mistake is a sign error.  Also "log" must mean "natural logarthim" the way the equation in the question is written to be correct.

If the free energy change is negative, the equillbrium constant is $&gt;1$

Should be $10^{\frac{32\frac{kcal}{mol}}{2.303\times0.00199\frac{kcal}{Kmol}\times{298K}}} = 10^{23}$