Introduction to Organic Chemistry Principles:

(ACID-BASE REACTIONS)

Before one can delve into the more intricate reactions of organic chemistry, a general understanding of the nature of electron transfer and bond formation must be established.  One should be able to consider the reactions of organic compounds as acids and bases.

WHAT IS AN ACID?

Two prominent theories that define acids and bases are usually discussed in the course of any general chemistry course.  These theories are Bronsted-Lowry theory and the Lewis theory of acids.  Here we will discuss them both.

The Bronsted-Lowry definition is named for Johannes Bronsted and Thomas Lowry, who independently proposed it in 1923. A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a hydrogen ion (proton). Thus, according to the Bronsted-Lowry definition, acids and bases must come in what is called conjugate pairs. For example, consider acetic acid dissolved in water:

Notice that we have written H2O explicitly in these reactions. The reason is that acid/base dissociation occurs by a proton transfer reaction from an acid species to a specific water molecule. The transfer occurs through a hydrogen bond between the acid molecule and a solvating water molecule.

Here, CHCOOH is a Bronsted-Lowry acid because it can donate a proton, and CHCOO- its conjugate base because it can accept a proton.

Dissociation Constants

An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution.

For a general reaction oA+pB  mC+ nD there exists a equilibrium constant Keq defined as:

In the cases of acids and bases we use the symbol Ka. It is the equilibrium constant for a chemical reaction known as acid-base dissociation.

The equilibrium can be written symbolically as:

 where HA is a generic acid that dissociates by splitting into A−, known as the conjugate base of the acid, and the hydrogen ion or proton, H+, which, in the case of aqueous solutions, exists as a solvated hydronium ion.  The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A−] and [H+]: