Any substance whose aqueous solution contains ions is called an electrolyte. Any substance that forms a solution containing no ions is a non-electrolyte. Electrolytes that are present in solution entirely as ions are strong electrolytes, whereas those that are present partly as ions and partly as molecules are weak electrolytes.

Arrhenius recognized that the properties of acidic solutions are due to H+ ions and those of basic solutions are due to OH􀄳 ions. The Brønsted-Lowry concept of acids and bases is more general than the Arrhenius concept and emphasizes the transfer of a proton from an acid to a base.

A Brønsted-Lowry acid is a substance that donates a proton to another substance; a Brønsted-Lowry base is a substance that accepts a proton from another substance. Water is an amphiprotic substance, one that can function as either a Brønsted-Lowry acid or base, depending on the substance with which it reacts. The conjugate base of a Brønsted-Lowry acid is the species that remains when a proton is removed from the acid. The conjugate acid of a Brønsted-Lowry base is the species formed by adding a proton to the base. An acid and its conjugate base (or a base and its conjugate acid) are called a conjugate acid-base pair. The acid-base strengths of conjugate acid-base pairs are related: The stronger an acid, the weaker is its conjugate base; the weaker an acid, the stronger is its conjugate base.

Water ionizes to a slight degree, forming [H⁺] and [OH^–]. The extent of this auto ionization is expressed by the ion product constant for water: K_w = [H⁺] [OH^–]. The K_w expression indicates that the product of [H⁺] and [OH^–] is a constant. Thus, as [H⁺] increases, [OH^–] decreases. Acidic solutions are those that contain more [H⁺] than [OH^–], whereas basic solutions contain more [OH^–] than [H⁺].

The concentration of an acid can be expressed in terms of pH: pH = -log[H⁺]. The pX notation is also used to represent the negative logarithm of other small quantities, as in pOH and pK_w. The pH of a solution can be measured using a pH meter, or it can be estimated using acid–base indicators.

Weak acids are weak electrolytes; only some of the molecules exist in solution in ionized form. The extent of ionization is expressed by the acid-dissociation constant, K_a. The larger the value of K_a, the stronger is the acid. For solutions of the same concentration, a stronger acid also has a larger percent ionization. The concentration of a weak acid and its K_a value can be used to calculate the pH of a solution.

Weak bases include NH₃, amines, and the anions of weak acids. The extent to which a weak base reacts with water to generate the corresponding conjugate acid and [OH^-] is measured by the base-dissociation constant, K_b.

The acid-base properties of salts can be ascribed to the behavior of their respective cations and anions. The reaction of ions with water, with a reacting change in pH, is called hydrolysis.

The Lewis concept of acids and bases emphasizes the shared electron pair rather than the proton. A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. The Lewis concept is more general than the Brønsted-Lowry concept because it can apply to cases in which the acid is some substance other than H⁺.

The common-ion effect tends to suppress the ionization of a weak acid or a weak base. This action can be explained by Le Chatelier’s principle. A buffer solution is a combination of either a weak acid and its conjugate base or a weak base and its conjugate acid. The solution reacts in such a way that the pH of the solution remains nearly constant. Buffer systems play a vital role in maintaining the pH of body fluids.

Acid-base indicators are weak organic acids or bases. They change color near the equivalence point in an acid-base neutralization reaction. Titration is a process in which a solution of known concentration is used to determine the concentration of an unknown solution. The pH at the equivalence point of an acid-base titration depends on the extent of hydrolysis of the salt formed in the neutralization reaction. For strong acid strong base titrations, the pH at the equivalence point is 7. For weak acid strong base titrations, the pH at equivalence point is greater than 7. For weak base–strong acid titrations, the pH at the equivalence point is less than 7.