7.1 Brønsted-Lowry Acids and Bases

Acids may be compounds such as HCl or H₂SO₄, organic acids like acetic acid (CH₃COOH) or ascorbic acid (vitamin C), or H₂O. Anions (such as HSO₄⁻, H₂PO₄⁻, HS⁻, and HCO₃⁻) and cations (such as H₃O⁺, NH₄⁺, and [Al(H₂O)₆]³⁺) may also act as acids. Bases fall into the same three categories. Bases may be neutral molecules (such as H₂O, NH₃, and CH₃NH₂), anions (such as OH⁻, HS⁻, HCO₃⁻, CO₃²⁻, F⁻, and PO₄³⁻), or cations (such as [Al(H₂O)₅OH]²⁺). The most familiar bases are ionic compounds such as NaOH and Ca(OH)₂, which contain the hydroxide ion, OH⁻. The hydroxide ion in these compounds accepts a proton from acids to form water:

H⁺ + OH^– → H₂O

We call the product that remains after an acid donates a proton the conjugate base of the acid. This species is a base because it can accept a proton (to re-form the acid):

acid ⇌ proton + conjugate base
HF ⇌ H⁺ + F^–
H₂SO₄ ⇌ H⁺ + HSO₄^–
H₂O ⇌ H⁺ + OH^–
HSO₄^– ⇌ H⁺ + SO₄^2-
NH₄⁺ ⇌ H⁺ + NH₃

We call the product that results when a base accepts a proton the base’s conjugate acid. This species is an acid because it can give up a proton (and thus re-form the base):

base + proton ⇌ conjugate acid
OH^– + H⁺ ⇌ H₂O
H₂O + H⁺ ⇌ H₃O⁺
NH₃ + H⁺ ⇌ NH₄⁺
S^2- + H⁺ ⇌ HS^–

CO₃^2- + H⁺ ⇌ HCO₃^–
F^– + H⁺ ⇌ HF

In these two sets of equations, the behaviors of acids as proton donors and bases as proton acceptors are represented in isolation. In reality, all acid-base reactions involve the transfer of protons between acids and bases. For example, consider the acid-base reaction that takes place when ammonia is dissolved in water. A water molecule (functioning as an acid) transfers a proton to an ammonia molecule (functioning as a base), yielding the conjugate base of water, OH⁻, and the conjugate acid of ammonia, NH₄⁺:

The reaction between a Brønsted-Lowry acid and water is called acid ionization. For example, when hydrogen fluoride dissolves in water and ionizes, protons are transferred from hydrogen fluoride molecules to water molecules, yielding hydronium ions and fluoride ions:

When we add a base to water, a base ionization reaction occurs in which protons are transferred from water molecules to base molecules. For example, adding ammonia to water yields hydroxide ions and ammonium ions:

Notice that both these ionization reactions are represented as equilibrium processes. The relative extent to which these acid and base ionization reactions proceed is an important topic treated in a later section of this chapter. In the preceding paragraphs we saw that water can function as either an acid or a base, depending on the nature of the solute dissolved in it. In fact, in pure water or in any aqueous solution, water acts both as an acid and a base. A very small fraction of water molecules donate protons to other water molecules to form hydronium ions and hydroxide ions:

This type of reaction, in which a substance ionizes when one molecule of the substance reacts with another molecule of the same substance, is referred to as autoionization.

Pure water undergoes autoionization to a very slight extent. Only about two out of every 10⁹ molecules in a sample of pure water are ionized at 25 °C. The equilibrium constant for the ionization of water is called the ion-product constant for water (K_w):

H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH^–(aq)               K_w = [H₃O⁺][OH^–]

The slight ionization of pure water is reflected in the small value of the equilibrium constant; at 25 °C, K_w has a value of 1.0 × 10⁻¹⁴. The process is endothermic, and so the extent of ionization and the resulting concentrations of hydronium ion and hydroxide ion increase with temperature. For example, at 100 °C, the value for K_w is about 5.1 × 10⁻¹³, roughly 100-times larger than the value at 25 °C.

It is important to realize that the autoionization equilibrium for water is established in all aqueous solutions. Adding an acid or base to water will not change the position of the equilibrium. Example 2 demonstrates the quantitative aspects of this relation between hydronium and hydroxide ion concentrations.