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.
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.
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.
H2O + H2O H3O+(aq) + OH-(aq)
H2O H+(aq) + OH-(aq)
Kw = [H3O+][OH-] = [H+][OH-] = 1.00x10-14 (at 25oC)
(Using [H3O+] is equivalent to using [H+].)
Kw is called the dissociation constant or ionization constant of water.
In pure water [H+] = [OH-] = 1.00x10-7 M.
pH is a shorthand notation for -log[H+]
pOH is a shorthand notation for -log[OH-].
pH + pOH = 14.
Solutions are called
neutral when pH = 7, [H+] = [OH-] = 1.00x10-7
acidic when pH < 7, [H+] > 1.00x10-7
basic when pH > 7, [H+] < 1.00x10-7
Example: What is the pH of a solution of 0.025 M HNO3? (See example 13.4 in text.)
HNO3 is a strong acid and for all practical purposes dissociates completely.
HNO3(aq) + H2O NO3-(aq) + H3O+(aq)
[H+] = 0.025 M
pH = -log(0.025 M) = 1.6
What is the pOH of this solution? There are 2 ways to calculate pOH:
Kw = 1.00x10-14 = [0.025 M][OH-]
[OH-] = 4.00x10-13
pOH = -log(4.00x10-13) = 12.40
pOH = 14.00 - pH = 14.0 - 1.60 = 12.40
What are pH and pOH for a 0.0025 M solution of HNO3?
pH = -log(0.0025 M) = 2.60
pOH = 14.00 - 2.60 = 11.40
Notice that pH and pOH change by 1 for a factor of 10 change in [H+] and [OH-].