Chemistry Essentials for Dummies

Chapter 11. Acids and Bases

Acid-Base Reactions: Using the Bronsted-Lowry System

The Bronsted-Lowry theory states that acid-base reactions are a competition for a proton. For example, take a look at the reaction of ammonia with water:

Ammonia is a base (it accepts the proton), and water is an acid (it donates the proton) in the forward (left to right) reaction. But in the reverse reaction (right to left), the ammonium ion is an acid and the hydroxide ion is a base.

If water is a stronger acid than the ammonium ion, then there is a relatively large concentration of ammonium and hydroxide ions at equilibrium. If, however, the ammonium ion is a stronger acid, much more ammonia than ammonium ion is present at equilibrium.

Bronsted and Lowry said that an acid reacts with a base to form conjugate acid-base pairs. Conjugate acid-base pairs differ by a single H+. NH3 is a base, for example, and NH4+ is its conjugate acid. H2O is an acid in the reaction between ammonia and water, and OH- is its conjugate base. In this reaction, the hydroxide ion is a strong base and ammonia is a weak base, so the equilibrium is shifted to the left — there’s not much hydroxide at equilibrium.

Acting as either an acid or base: Amphoteric Water

Water can act as either an acid or a base, depending on what it’s combined with. Substances that can act as either an acid or a base are called amphoteric. If you put water with an acid, it acts as a base, and vice versa. For instance, when acetic acid reacts with water, water acts as a base, or a proton acceptor. But in the reaction with ammonia, water acts as an acid, or a proton donor. (See the earlier section “Weak: Ionizing partially” for details on both reactions.)

But can water react with itself? Yes, it can. Two water molecules can react with each other, with one donating a proton and the other accepting it:

This reaction is an equilibrium reaction. A modified equilibrium constant, called the Kw (which stands for water dissociation constant), is associated with this reaction. The Kw has a value of 1.0 x 10-14 and has the following form:

In pure water, the [H3O+] equals the [OH-] from the balanced equation, so [H3O+] = [OH-] = 1.0 x 10-7. The Kw value is a constant.

This value allows you to convert from [H+] to [OH-], and vice versa, in any aqueous solution, not just pure water. In aqueous solutions, the hydronium ion and hydroxide ion concentrations are rarely going to be equal. But if you know one of them, Kw allows you to figure out the other one.

Take a look at the 2.0 M acetic acid solution problem in the section “Acetic acid and other weak acids,” earlier in this chapter. You find that the [H3O+] is 6.0 x 10-3. Now you have a way to calculate the [OH-] in the solution by using the Kw relationship: