According to the Bronsted-Lowry Theory, an acid and a base react to form a conjugate base and a conjugate acid, as diagrammed below. These reactions are reversible and result in an acid-base equilibrium.

Consider the following example from your textbook (page 724):

In any acid-base equilibrium, there are always two acids and two bases. For example, in the preceding example, CH₃COOH and H₃O⁺ are the acids and H₂O and CH₃COO^- are the bases. You will notice that the acid on the left differs from the base on the right by just one proton. We refer to this pair of substances as a conjugate acid-base pair. Acetic acid is the conjugate acid of the acetate ion, and the acetate ion is the conjugate base of acetic acid. In any Bronsted-Lowry acid-base reaction, however, there are two conjugate acid-base pairs. In this instance, the second conjugate pair involves the hydronium ion and water (as shown in the equation below).

To identify the conjugate acid-base pairs in an acid-base reaction, begin with the reactants. The acid and base are on the reactant side of the equation.

• First, identify which reactant is the acid (proton donor). Look for its conjugate base on the products side of the reaction. The conjugate base is the compound that has one less hydrogen ion than the acidic reactant.
  
  
• Second, identify which reactant is the base (proton acceptor). Look for its conjugate acid on the products side of the reaction. The conjugate acid is the compound that has one more hydrogen ion than the basic reactant.

---

  Read "Conjugate Acids and Bases" on pages 724 to 726 in the textbook to learn more about acid-base pairs.

---

Check Your Understanding

Complete Practice questions 7 and 8 on page 726 of the textbook.