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Equilibria of Acids and Bases


We will use the Bronsted-Lowry definitions for acids and bases:

Acids are species that donate a proton (H+).
bases are species that accept a proton.


For the example of water, H2O <--> H+ + OH-, the equilibrium constant is:

      [H+] [OH-]
Keq = ----------
The concentration of water in a water solution is constant and this expression simplifies to:
Kw = (55.56 M)*Keq = [H+] [OH-]
where Kw is called the dissociation constant of water and equals 1.00x10-14 at room temperature. The concentrations of [H+] and [OH-] therefore equal 1.00x10-7 M.


An acid in water will dissociate, that is it donates it proton. We call acids that dissociate completely strong acids and acids that dissociate only partially weak acids.

Strong acid example:
HNO3 (aq) + H2O <--> NO3-(aq) + H3O+(aq)

Keq = a very large number

In this example, HNO3 is an acid and H2O is acting as a base.
NO3- is called the conjugate base of the acid HNO3, and H3O+ is the conjugate acid of the base H2O.


A base in water can accept a proton from water to produce OH-. Strong bases are salts of hydroxide that dissociate completely in water, so this statement is redundant. But weak bases do not have to contain hydroxide themselves, but they produce basic solutions by reacting with water.

Weak base example:
NH3 (aq) + H2O <--> NH4+(aq) + OH-(aq)

K = 1.8x10-5

In this example, NH3 is a base and H2O is acting as an acid. NH4+ is the conjugate acid of the base NH3, and OH- is the conjugate base of the acid H2O.

A compound that can act as either an acid or a base, such as the H2O in the above examples, is called amphiprotic.

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