Acids+and+bases

Acids and bases have a few different definitions. Each definition approaches the concepts of what an acid or base is from a unique perspective.

The most common definition is the **Bronsted-Lowry** definition:
 * A Bronsted acid is defined as a proton donor. The proton is the H+ ion.
 * A Bronsted base is defined as a proton acceptor.

To demonstrate the Bronsted-Lowry definition, look at how hydrochloric acid interacts with water:

HCl (aq) + H 2 O (l) --> H 3 O + (aq) + Cl - (aq)

In this reaction, the acid is HCl and the base is water. Water accepts the H + ion (a proton) from HCl, so water is acting as a base.

Another similar definition of acids and bases comes from **Arrhenius**:
 * an Arrhenius acid is defined as any substance that increases the H + concentration of a solution;
 * an Arrhenius base is defined as any substance that increases the OH - concentration of a solution.

The third acid-base definition comes from **Lewis**:
 * a Lewis acid is defined as an electron-pair acceptor;
 * a Lewis base is defined as an electron-pair donor.

Lewis definitions are most commonly used in organic chemistry.

For example, the Lewis diagram of ammonia (shown below, NH 3 ) has a lone pair of electrons on nitrogen that are capable of forming a bond with boron in BF 3. Because B accepts a pair of electrons, it is a Lewis acid, even though it contains no protons!