[gaudars]

I need assistance regarding whether the reaction NH₃ + H₂S NH₄⁺ + HS^-. My instinct tells me the reaction would favor the reactants, since both reactants only partially ionize. However I calculated the equilibrium constant of the reaction to be 170. This would indicate that the reaction favors the products. I would appreciate it if someone could check over my work. I expressed the given reaction as the sum of two other reactions,
H₂S + H₂O H₃O⁺ + HS^-
NH₃ H₃O⁺ NH₄⁺ + H₂O

The equilibrium constant for the overall reaction is equal to the product of the equilibrium constants of each constituent reaction. The constant for the first reaction is equal to the ionization constant for H₂S, so it is 1.0e-7. I noticed that the equilibrium expression for the second reaction was equal to the reciprocal of the equilibrium expression or the acid-base reaction of ammonium with water. Therefore the equilibrium constant for the second expression is equal to 1/K_a where K_a is the ionization constant for ammonium. Therefore the product of these two reactions is equal to 1.0e-7/5.9e-10= 170. This means the reaction favors the products.

Is my reasoning here correct? Any insight would be much appreciated.
Thank you in advance.

[JustinCh3m]

A Keq = 170 is kind of "in the middle." What I mean is that the dynamic equilibrium doesn't really favor products or reactants per se.  You didn't show the states of the R's and P's -- perhaps there are some entropy issues to be concerned about?  Also, knowing (delta)G would help too.