Acids and Bases

The Bronsted-Lowry definition of an acid is a hydrogen ion/proton (H⁺) donator. A B-L bases and a hydrogen ion acceptor.

Strong acids completely dissociate and donate hydrogen ion. Some common strong acids in water are:

HClO₄
HNO₃
HCL
HI
HBr
H₂SO₄

All other acids may be considered weak acids.

The common strong bases are:

NaOH
KOH
RbOH
CsOH
Ca(OH)₂
Sr(OH)₂
Ba(OH)₂

All dissolve completely in water to give OH^- which readily accepts hydrogen ion. All other bases may be considered weak bases.

Some species are capable of accepting or donation hydrogen ions. We call these species amphoteric. For example HCO₃^- is amphoteric.

HCO₃^-(aq) + H₂O(l) <==> H₂CO₃(aq) + OH^-(aq)

HCO₃^-(aq) + H₂O(l) <==> CO₃^2-(aq) + H₃⁺(aq)

Water dissociates/autoionizes to give the hydronium ion and the hydroxide ion.

2 H₂O(l) <==> H₃O⁺(aq) + OH^-(aq)

The is an associated equilibrium constant, and at room temperature it is 1.0E-14.

[H₃O⁺][OH^-] = K_w = 1.0E-14

Ka Kb = Kw

The p function is defined as

pX = -log [X] So
pH = -log[H₃⁺]

pOH = -log[OH^-]

pKw = -log[Kw] = -log[1.0E-14] = 14

pH + pOH = pKw

Acid and Base Strength

We can refer to either the Ka or the pKa to determine acid strength. The larger the Ka the stronger the acid. Or the smaller the pKa the stronger the acid. Conversely we refer to Kb and pKb for bases strenth. See Fig. 6.2 and 6.3.

Acids react with water to produce hydrogen ions. The hydrogen ions piggy-back on another wather molecule and move around and hydronium ions. If the acid is strong it completely reacts with water. If the acid is weak, then with have an equilibrium

Bases react with water and abstract hydrogen ions leaving hydroxide ions.

When weak conjugate acid/base pairs in both in solution they form an equilibrium that resist a change in pH. We call this system a buffer solution. We can adjust the pH of our buffer solution according to

where [HA]_o is the initial weak acid concentration and [A^-]_o is the initial conjugate base concentration.