Equilibrium in Bronsted Acid-Base Systems
In acid-base reactions, equilibrium favours the production of the weaker acid and base. ie; the stronger the reacting acid and base, the more complete the reaction.
eg.                    HClO4(l) + H2O(l) -----> H3O+(aq) + ClO4-(aq)
Think of a Bronsted acid-base reaction as a competition between the 2 bases in the system for protons. The stronger base "wins" and forces the equilibrium in the direction of the weaker acid and base. Above, the H2O and ClO4- ions, both bases, are competing for protons. H2O, the stronger of the two, "wins" and the equilibrium is shifted to the right in the direction of the weaker base, ClO4- ions. HClO4, the stronger acid, has a greater tendency to give it's protons to H2O than H3O+ ions have to give their protons to ClO4- ions. The equilibrium is displaced so far to the right that the reaction is essentially complete.
In contrast to HClO4, HCN is a weakly dissociated acid.
         HCN(g) + H2O(l) <--------=====> H3O+(aq) + CN-(aq)
The reaction above indicates that the equilibrium that is reached has a majority of the HCN molecules unreacted. The H3O+ ions are a much stronger acid than the HCN, and the CN- ions are a much stronger base than H2O. This means that the CN- ions "win" in the competition for protons and force the equilibrium left.
The stronger the acid, the weaker is its conjugate base.
Please note that HClO4, a very strong acid, is the conjugate acid of ClO4- ions. ClO4- is itself a very weak base. The perchlorate ions, are so weak that any base below it in the acid-base table in your databook can take H+ ions away from the ClO4- ions. Thus H2O completely removes the acid hydrogen from the HClO4 molecules and forms hydronium ions, H3O+.
In general, the strong acids in the upper part of the left hand column have the greatest tendency to react with the strong bases in the lower of the table.