Acids and Bases

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An Acid is a substance that produces H3O+ (H+) when it is dissolved in water. It is a proton donor and an electron pair acceptor or a species that donates protons. For example: HCl, NH4, AlCl3.
A Base is a substance that produces an OH- when it is dissolved in water (Arrhenius). A proton acceptor (Brønsted), or a electron donor. For example: NaOH, KOH, CH3NH2.
Acids and bases were first identified as specific types of compounds because of their behavior in aqueous solutions.
Acids and bases relate to each other in Conjugate Pairs, somewhat like husbands and wives. For every acid there is a conjugate base; and for every base there is a conjugate acid. Just like every husband has a wife and vice versa. The two members of the conjugate pair are related by the donating and accepting of a single proton.
The equation below, Equation (1), demonstrates a power struggle going on between the two couples and within them. There is a competition for which base, H2O, (keep in mind that H2O can act as an acid or a base because it auto-ionizes itself, meaning it gives protons back and forth within itself, thus acting as both an acid and a base;) See Equation (2). Then A- will get the proton. The winner is the stronger base which has a greater affinity for H+ and everything will go its way. This base will determine whether the equation goes to the right or the left at equilibrium.

HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq)


To determine the strength of an acid or base can be difficult within conjugate pairs. Strong acids have weak conjugate bases so the equilibrium lies far to the right. A weak base has a lower affinity for protons that water. So water wins the H+ ion as in the reaction in Equation (1) above.
Of course, a weak acid has a strong conjugate base so the equilibrium will go to the left, and the acid will not dissociate that much.

Acids


In order to determine how much the H+ will get, or the degree to which a weak monoprotic acid will dissociate, we use:
Ka, the acid dissociation constant: Ka = [H3O+][A-] / [HA]
The weaker the acid, the smaller its Ka and the less it will dissociate. Strong acids completely dissociate into their component ions in aqueous solution.
Note that [H2O is omitted from the Ka expression because the concentration of H2O is so high in an aqueous solution and changes so little, it is basically treated as a constant.
Below are examples of strong and weak acids:


Strong acids:		HCl
			H2SO4
			HNO3
			HClO4

Weak acids:		NH4+
			NCN
			HF
			HNO2

Bases

We can also take a look at this from the base's point of view using the base dissociation constant, Kb. KB is a measure of the degree to which a base will dissociate:

B(aq) + H2O(l) <==> BH+ + OH-(aq)
The weaker the base, the smaller its Kb.
Kb always refers to the reaction of a base with water to from the conjugate acid and the hydroxide ion.
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