Acids, Bases, and pH

REVIEW

Acids and bases are compounds that are important in industry, our everyday lives, and most critically, for the functioning of our bodies. Because biochemistry is aqueous chemistry, the most common definition of an acid is the Bronsted–Lowry definition:

Water can be both a proton donor and a proton acceptor, and a small proportion of water molecules do react with one another:

H2O	+	H2O		OH–	+	H3O+
proton donor		proton acceptor

The ion product constant of this reaction, K_w, is

This expression is useful because it tells us that in water, the product of the OH^– and H₃O⁺ ion concentrations is always a constant value. It also indicates that in pure water, the concentration of both OH^– ions and H₃O⁺ ions is 10^–7 M. Because under physiological conditions, the H₃O⁺ concentrations tend to be small and the numbers cumbersome to deal with, the pH scale was invented to make manipulation of these numbers easier:

Remember that H⁺ is written as H₃O⁺ (the hydronium ion) in the Bronsted–Lowry definition of acids and bases to indicate that free protons (H⁺ ions) always strongly associate with a water molecule.

In the pH scale, which ranges from 0 to 14, pH 7 is neutral, while lower pHs are acidic and higher pHs are basic. What follows from the K_w expression is that

There are only a few strong acids, which dissociate completely (break apart into their constituent ions) completely in water. Most acids are weak, meaning they only partially dissociate in water. The degree of dissociation of a weak acid is given by its K_a, acid dissociation constant. Because K_a values are small for weak acids, the K_a is often converted to a pK_a value.

The larger the K_a value, and the smaller the pK_a value, the stronger the acid.

When a weak acid dissociates in water, it results in the formation of a hydronium ion and a basic molecule that is able to accept a proton. This basic molecule is called the conjugate base of the weak acid, and the weak acid and its conjugate base are said to be a conjugate acid–base pair.

HB	+	H2O		H3O+	+	B–
weak acid		water		hydronium ion		conjugate base

In a conjugate acid–base pair, the conjugate acid has one more proton than its conjugate base.

A solution made with a weak acid and its conjugate base is called a buffer, because it resists changes in pH when OH^– ions or H₃O⁺ ions are added. The Henderson–Hasselbach equation is useful for buffer calculations:

The Henderson–Hasselbalch equation allows the following:

• Determination of the pH of buffered solutions.
• Determination of the ratio of conjugate base to conjugate acid at a given pH .