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December 22, 2014

Homework Help: chem

Posted by Chris on Friday, April 27, 2007 at 5:14pm.

For the following reactions, name the Bronsted-Lowry acids and bases. Then name the conjugate acid and bases.

H3O+(aq) + CN-(aq) <==> HCN(aq) + H2O

I'm really confused on this whole concept even thought it's not really difficult. I said:
acids: H3O+
bases: CN-
conjugate acid: HCN
conjugate base: H2O
However, it seems to me that since this is an equilibrium reaction, there would have to be a conjugate acid/base, etc, on both sides on the equation. Could someone explain this a little better?

I don't know much about chemistry, but I think that your acid and conjugate base are correct. check out Keith's question a little further down on this page, maybe he can help you out

You may be trying to make it too hard OR you may not have the definitions down. Your answers are right.
Just remember the definitions: An acid is a proton donor; a base is a proton acceptor. An acid (the H3O^+) is donating a proton to the CN^- so the base it produces on the other side of the equation(H2O) is its conjugate base. That is, the H3O^+/H2O is the acid/conjugate base pair. Similarly, CN^- is the base because it accepts a proton to become HCN. That makes CN^-/HCN a base/conjugate acid pair. And you are correct, also, that this should exist on both sides of the equation but that is so whether the reaction is at equilibrium or not. We could look at the right side going to the left (the reaction in reverse), in which case, H2O is the base on the right side going to H3O^+, its conjugate acid on the left side. Similarly, HCN is the acid on the right side going on the left to CN^- which is its conjugate base. I think you just need a little practice and confidence and you will do fine. I hope this helps.

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