• Acids and Bases

The use of acid and base as classifications of chemical substances is an ancient one and predates even alchemy. Acids were originally substances that tasted sour and reacted with bases; bases were alkaline salts such as CaO (lime), NaOH, and KOH, while amphoteric substances were substances which could react with either acids or bases. In modern chemistry, these concepts have taken on considerably more precise meaning. The chemistry of acids and bases is the chemistry of reactions in which protons are transferred.

Arrhenius Theory of Acid-Base Behavior

The approach taken by Arrhenius with acids and bases was in accordance with his general approach to substances in aqueous solution. Materials which dissolved in whole or in part as ions were, and still are, called electrolytes. A strong electrolyte, like sodium chloride, was a substance which dissociated completely into ions while a weak electrolyte remained partially in molecular form. Thus a strong acid is a strong electrolyte which ionized completely to give hydrogen ions in aqueous solution. All of the common strong acids - sulfuric acid, H₂SO₄; nitric acid, HNO₃; hydrochloric acid, HCl; and perchloric acid, HClO₄ - easily fit in this category. Likewise, a strong base is a strong electrolyte which ionized completely to give hydroxyl ions in aqueous solution. The common strong bases, NaOH and KOH, fit in this category. Acid-base reactions are then simply the reaction of the hydrogen ion with the hydroxyl ion as a neutralization reaction, H⁺(aq) + OH^-(aq) --> H₂O. The remaining ions simply continue to exist as a solution of a salt. The reaction including these "bystander ions" could be written as

H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) --> Na⁺(aq) + Cl^-(aq) + H₂O

The Arrhenius concept was not intended to, and does not, deal with acid-base chemistry in solvents other than water. Even in water, however, it has deficiencies. Ammonia, NH3, clearly behaves in a qualitative manner as a base and reacts to neutralize acids. However, ammonia does not contain a hydroxide ion. This is of itself no obstacle because ammonia might well produce hydroxyl ion by reaction with water:

NH₃(aq) + H₂O --> NH₃.HOH --> NH₄OH --> NH₄⁺(aq) + OH^-(aq)

However, no experimental evidence for the existence of NH₄OH has yet been foumd. For these two reasons, the Arrhenius concept is no longer considered an adequate treatment of acids and bases even in aqueous solution.

Bronsted-Lowry Theory of Acid-Base Behavior

Any species which can either accept or give up a proton is said to be amphiprotic. Thus the water molecule is amphiprotic, since it can give up a proton, H₂O --> H⁺ + OH^-, to form the hydroxyl ion OH^-. Alternatively, water can accept a proton to form the hydronium ion H₃O⁺, according to the equation H⁺ + H₂O --> H₃O⁺. The above two equations can be combined to give the dissociation equation for water: 2H₂O --> H₃O⁺ + OH^-.

The Bronsted-Lowry concept is an extension of the Arrhenius concept in that all Arrhenius bases, being sources of hydroxide, can accept protons. Ammonia and amines will also accept protons to form the corresponding ammonium ions, so the existence of NH4OH is no longer necessary to explain the action of ammonia as a base. The Bronsted-Lowry concept also is useful in protonic solvents other than water, such as liquid ammonia or glacial acetic acid, where the Arrhenius concept is not useful. We will, however, generally confine our discussion to aqueous solutions because they are so much more important.

Lewis Explanation of Acid-Base Behavior

Since a base like ammonia, H₃N:, has a lone pair of electrons, it can be considered to "donate" them to a proton in forming the conjugate acid NH₄⁺. The Lewis definitions are used to explain the effect of compounds such as AlCl₃, which acts as an acid in nonaqueous organic solvents, on organic reactions. In protonic solvents, however, the Lewis definitions are far less useful than are the Bronsted definitions.