Arrhenius Acids:

Substances that when placed in water, will dissociate to produce H⁺ ions.

Arrhenius Bases:

Substances that when placed in water will dissociate to yield OH^- ions.

Bronsted Acids:

Any substance that can transfer a proton (H⁺) to another substance

Bronsted Bases:

Any substance that can accept a proton from another substance

Acid and Base Strength:

In a reaction, look at the bases on both sides of the equation. Think of them as "competing" for protons. More of the protons will get transferred to the stronger base.

Strong Acid:

An acid that is almost 100% dissociated when in solution.

IE: If H—A is a strong acid, the reaction H-A + H₂O à A^- + H₃O⁺ ,

the equilibrium lies almost completely to the right. (There are only a negligible amount of undissociated H—A molecules.

For a weak acid, the solution contains a mixture of undissociated acid and H₃O⁺ ions.

Some typical strong acids: HClO₄, HCl, HBr, HI, H₂SO₄ , HNO₃
Typical weak acids: HNO₂, HF, CH₃COOH

Hydrated Protons:

Dissociation of an Arrhenius acid H—A gives H⁺ ions in aqueous solution. However, the bare proton is much too reactive to exist in aqueous solution where you have all those exposed electrons on the oxygen atoms. So the proton bonds to form H₃O⁺ ions.

Dissociation of Water:

Always remember: the product of the H₃O⁺ and OH^- concentrations in aqueous solution is always 1.0*10^-14 (at 25 deg C).

K_w = [OH^-] [H₃O⁺] = 1*10^-14

If the concentration of one equals the concentration of the other, (ie: both = 1*10^-7), you have a neutral solution. If however, the OH^- is greater, you have a basic solution. If the hydronium ion concentration is greater, you have an acidic solution.
AWOS: If the concentration of one is very high, the other will be very low.

pH = -log (concentration of H₃O⁺)

EG: If you have [H₃O⁺] = 1*10^-3, pH = +3

EG: What is the pH of a 0.025M solution of HNO₃?

Since HNO3 is a strong acid, it dissociates almost completely into H⁺ and NO₃^-. This means that the concentration of hydronium ion is also 0.025M. Therefore, the pH is –log(0.025) = 1.60.

Good EG: See McMurray, p.586 (15.7)

Lewis Acids and Bases:

Lewis proposed an even more general definition of acids and bases. He said that an Acid is anything that can accept electrons. (A base is anything that donates electrons).

For example, when a bare proton forms a bond with a base such as ammonia, it is accepting the free elctrons from the ammonia to form a new covalent bond. That makes the proton an acid.

Lewis acids are more general because they include all molecules that can accept electrons—not just protons. (For example, Al³⁺ would be considered a Lewis acid).

AWOS: Anything with vacant valence orbitals can act as a Lewis Acid.