Acids and Bases

Definitions of Acid/Base

	Acid	Base
Bronsted-Lowry	Donates proton	Accepts proton
Arrhenius	Donates proton	Donates OH- (hydroxyl) group
Lewis	Accepts electron pair	Donates electron pair

pH and pOH

pH = -log₁₀[H⁺]

pOH = -log₁₀[OH^-]

K_w

Water dissociates as follows:
H₂O « H⁺ + OH^-

K_w = [H⁺] [OH^-] = 10^-14 (at 25^oC)
\pK_w = pH + pOH = 14


(e.g. A solution with pH=9 has pOH = 14 - 9 = 5)

\In neutral water, [H⁺] = [OH^-] = Ö(10^-14) =10^-7
pH = -log₁₀[H⁺] = 7.0
pOH = -log₁₀[OH^-] = 7.0


K_a
An acid dissociates in water as follows:
HA « H⁺ + A^-

K_a = [H⁺] [A^-] / [HA]
pK_a = -log₁₀K_a

K_b
A base reacts with water as follows:
A^- + H₂O « HA + OH^-

K_b = [HA] [OH^-] / [A^-]


(Note: the concentration of H₂O is usually ignored)
pK_b = -log₁₀K_b

Note, K_a K_b = ([H⁺] [A^-] / [HA]) x ([HA] [OH^-] / [A^-])
= [H⁺] [OH^-]
= K_w
= 10^-14

And since K_a K_b = 10^-14
pK_a + pK_b = 14 = pK_w

Henderson-Hasselbalch Equation
K_a = [H⁺] [A^-] / [HA]
[H⁺] = K_a [HA]/[A^-]
pH = pK_a - log₁₀[HA]/[A^-]
pH = pK_a + log₁₀[A^-]/[HA]

Buffers
Buffers are solutions that resist changes in pH when small amounts of acid or base are added to it.

A buffer usually consists of:


a weak acid and its conjugate base or
a weak base and its conjugate acid.
Such systems have the greatest buffering capability when the two components are in equal concentration.
(e.g. an acetic acid/acetate buffer has greatest buffering capability when [acetic acid] = [acetate], and therefore when
pH = pK_a + log₁₀[acetate]/[acetic acid]
= pK_a + log₁₀(1)
= pK_a + 0
= pK_a = 4.74)