15.2

Water: A Bronsted-Lowry Acid and Base

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An amphoteric molecule is capable of acting as both an acid and a base. Water is an example of an amphoteric molecule.

In an aqueous solution of ammonia, water can act as an acid by donating a proton to ammonia, a base. In contrast, water acts as a base in an aqueous solution of hydrochloric acid by accepting a proton from the acid.

Because it is amphoteric, water can undergo autoionization. In this process, one molecule of water acts as the acid by donating a proton to another molecule of water that acts as a base by accepting that proton.

This results in the production of a hydronium ion and a hydroxide ion.

The equilibrium constant for the autoionization of water, K_W, also known as the ion-product constant for water, has a value of 1 × 10⁻¹⁴. Its equilibrium expression is written as the concentration of the hydronium ions times the concentration of the hydroxide ions, with no denominator as both reactants are liquids.

In pure water at 25 °C, the concentrations of hydronium and hydroxide ions are equal—that is 1 × 10⁻⁷ M.

K_W can be used to calculate the concentration of hydronium and hydroxide ions in a solution and to determine whether a solution is acidic or basic.

If a solution is acidic, the concentration of hydronium ions will be higher than hydroxide ions. In contrast, if a solution is basic, the concentration of hydroxide ions will be higher than the hydronium ions.

If the concentration of hydronium ions in solution is 5 × 10⁻⁷ M, the hydroxide ion concentration can be determined using K_W.

As the values of K_W and hydronium ions are known, the concentration of hydroxide ions can be calculated to be 2 × 10⁻⁸ M.

As the hydronium ion concentration, 5 × 10⁻⁷ M, is higher than the hydroxide ion concentration, 2 × 10⁻⁸ M, the solution is acidic.

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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:

Acid-base equilibrium equation for HF and H2O forming H3O+ and F-, chemical reaction diagram.

Base ionization of a species occurs when it accepts protons from water molecules. In the example below, pyridine molecules, C₅NH₅, undergo base ionization when dissolved in water, yielding hydroxide and pyridinium ions:

Acid-base reaction diagram with chemical formulas, showing water and pyridine interaction.

The preceding ionization reactions suggest that water may function as both a base (as in its reaction with hydrogen fluoride) and an acid (as in its reaction with ammonia). Species capable of either donating or accepting protons are called amphiprotic, or more generally, amphoteric, a term that may be used for acids and bases per definitions other than the Brønsted-Lowry one. The equations below show the two possible acid-base reactions for two amphiprotic species, bicarbonate ion, and water:

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The first equation represents the reaction of bicarbonate as an acid with water as a base, whereas the second represents the reaction of bicarbonate as a base with water as an acid. When bicarbonate is added to water, both these equilibria are established simultaneously and the composition of the resulting solution may be determined through appropriate equilibrium calculations. In the liquid state, molecules of an amphiprotic substance can react with one another as illustrated for water in the equations below:

Chemical equilibrium, H2O as acid and base, reversible reaction formula, acid-base equilibrium.

The process in which like molecules react to yield ions is called autoionization. Liquid water undergoes autoionization to a very slight extent; at 25 °C, approximately two out of every billion water molecules are ionized. The extent of the water autoionization process is reflected in the value of its equilibrium constant, the ion-product constant for water, K_W:

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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.

Chemical equilibrium diagram; reaction profile with energy changes, kinetic rates, ΔH‡, ΔH.

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 5.6 × 10⁻¹³, roughly 50 times larger than the value at 25 °C.

The autoionization of water yields the same number of hydronium and hydroxide ions. Therefore, in pure water at 25 °C:

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The concentrations of these ions in a solution are often critical determinants of the solution’s properties and the chemical behaviors of its other solutes, and specific vocabulary has been developed to describe these concentrations in relative terms. A solution is neutral if it contains equal concentrations of hydronium and hydroxide ions; acidic if it contains a greater concentration of hydronium ions than hydroxide ions; and basic if it contains a lesser concentration of hydronium ions than hydroxide ions.

Summary of Relations for Acidic, Basic, and Neutral Solutions

Classification	Relative Ion Concentrations	pH at 25 °C
acidic	[H₃O⁺] > [OH⁻]	pH < 7
neutral	[H₃O⁺] = [OH⁻]	pH = 7
basic	[H₃O⁺] < [OH⁻]	pH > 7

This text is adapted from Openstax, Chemistry 2e, Section 14.1: Brønsted-Lowry Acids and Bases, and Section 14.2: pH and pOH.

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