Acids and Bases: Brønsted-Lowry Theory

FIELDS OF STUDY: Physical Chemistry; Inorganic Chemistry; Chemical Engineering

ABSTRACT

The basic properties of acids and bases according to the Brønsted-Lowry theory are elaborated. Acidity and basicity (or alkalinity) are measured using the pH scale, which charts the concentration of H⁺ ions relative to pure water.

The Nature of the Brønsted-Lowry Acids and Bases

In 1923, chemists Johannes Nicolaus Brønsted (1879–1947) and Thomas Martin Lowry (1874–1936) defined acids and bases in a straightforward way. Simply stated, an acid, a proton donor, is any compound that can release a proton from its molecular structure, as

AH → A⁻ + H⁺

A base is a proton acceptor, any compound that can accept a proton from another compound, as

B + H⁺ → BH⁺

Acids that act in this way are termed "protic" acids, since the H⁺ cation is identically a proton. For most practical purposes, this definition is sufficient, as most water-soluble acids are either protic or amphiprotic compounds. (Amphoterism is the ability of a compound to act as either an acid or a base, depending on its environment and the other materials present.)

The Brønsted-Lowry definition refines those formulated before the development of the modern atomic theory. In 1838, Justus von Liebig (1803–73) described acids as materials containing hydrogen that could be replaced by a metal, which indeed they are. The acidic hydrogen atom on compounds such as the inorganic sulfuric acid and the organic acetic acid is easily replaced by metal ions such as magnesium to form magnesium sulfate (MgSO₄) or sodium to form sodium acetate (C₂H₃NaO₂). The corresponding equations are

Mg²⁺ + H₂SO₄ → MgSO₄ + 2H⁺

and

Na⁺ + CH₃COOH → CH₃COONa + H⁺

Svante Arrhenius (1859–1927) first elucidated the definition of acids and bases as specifically forming hydrogen ions and hydroxide ions in aqueous solutions. This statement seems to limit the identity of acids and bases to protic acids and hydroxide salts. The Brønsted-Lowry definition, however, recognizes that acidic and alkalinic behaviors are complementary actions of the same fundamental molecular structure. That is, a compound that releases a proton in acting as an acid produces a corresponding anion as its conjugate base. An acid and its conjugate base are a conjugate pair and do not exist in isolation from each other in an equilibrium system. Accordingly, both the acid and base equations above should be properly written as,

AH ⇌ A⁻ + H⁺

and

B + H⁺ ⇌ BH⁺

These equations more clearly demonstrate the conjugate relationship of the various components, and they include such compounds as ammonia, which do not dissociate into ions in solution. When dissolved in water, ammonia accepts a proton from the water molecules, according to the equation

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

and so acts as a base without itself dissociating to release a hydroxide ion, as would be required by the Arrhenius definition.

Measuring Acids and Bases: The pH Scale

Equal molar quantities of different acids do not produce equally acidic solutions. The conjugate nature of the components of an acid dissociation equilibrium is an important factor in determining the acidity of the resulting solution. Hydrogen chloride and nitric acid, for example, are completely dissociated into ions in aqueous solutions, but acetic acid and other organic acids generally do not dissociate completely in solution. The resulting solutions all have different concentrations of H⁺ ions and, thus, different acidities. As the most common material on the planet, water is the standard by which H⁺ ion concentrations are measured. The structure of the water molecule allows it to dissociate into its component ions in solution without the participation of a second material (a process known as "autolysis"), according to the equation

H₂O ⇌ H⁺ + OH⁻

In pure water, this equilibrium produces equal quantities of both H⁺ and OH⁻. The quantity of each has been determined experimentally to be 10⁻⁷ moles per liter (M). Solutions in which the concentration of H⁺, [H⁺], is greater than 10⁻⁷M are acidic, while those in which [H⁺] is less than 10⁻⁷M are alkaline. The pH scale was developed as a simple means of communicating the [H⁺] of an aqueous solution. The pH of an aqueous solution is defined as

pH = −log[H⁺]

Because the [H⁺] of pure water is 10⁻⁷M, the pH of pure water is defined as +7. A solution of HCl in water is hydrochloric acid, characterized by the complete dissociation of HCl into H⁺ and Cl⁻ ions. Because hydrochloric acid has an [HCl] of 0.01M, it also has a [H⁺] of 0.01M, or 10⁻²M. Therefore, the corresponding pH is +2. Similarly, a solution in which the [H⁺] is 10⁻⁹M, for example, has a pH of +9.

A complementary scale, the pOH scale, can be used for alkaline solutions, based on the identical relationship

pOH = −log[OH⁻]

Since the concentration of H₂O in pure water is always a strictly constant value, the relationship [H⁺][OH⁻] must also be as strictly constant. In pure water, this p-scale value is the equilibrium constant of the autolysis reaction, with the value (10⁻⁷)(10⁻⁷), or 10⁻¹⁴. Therefore, in any aqueous solution the product [H⁺][OH⁻] must also always be 10⁻¹⁴. If the [H⁺] is greater than 10⁻⁷M, the equilibrium of the autolysis reaction requires that the [OH⁻] be reduced by the corresponding amount. (When two numbers are multiplied together, their logarithmic values combine.) Thus, the value of the autolysis equilibrium is 7 + 7 = 14 (the pH and the pOH of a solution must always add up to 14). Normally, aqueous acid solutions have pH between 1 and 7, corresponding to [H⁺] in the range of 10⁻¹ to 10⁻⁷. A pH with a value of 0 or less is also possible. A 10M solution of HCl, for example, has [H⁺] of 10¹M and a corresponding pH of −1.

Acid-Base Neutralization Reactions

The goal of a neutralization reaction is to bring the pH of a particular solution to the neutral pH of 7. The product of any neutralization reaction is a salt. This is easily seen in the neutralization reaction of hydrochloric acid, HCl, with sodium hydroxide, NaOH:

HCl + NaOH → H₂O + NaCl

By adding hydrochloric acid to a solution of sodium hydroxide to the point at which the pH is 7, one ends up with a solution of plain "salt water," NaCl in H₂O. This process is called "titration" and is a common laboratory method for determining precisely the [H⁺] of a solution. The technique can also be used to analyze many other materials by determining the end point or equivalence point of a specific reaction. The technique depends on a means of monitoring the change in pH of the solution as acid or base is added. This has traditionally been done by using litmus paper or another indicator material that changes color according to the pH, but this method has been superseded by electronic devices that can measure the pH directly and display the measured values in graphic form.

PRINCIPAL TERMS

• amphoterism: the ability of a compound to act as either an acid or a base, depending on its environment and the other materials present.
• conjugate pair: an acid and the conjugate base that is formed when it donates a proton, or a base and the conjugate acid that is formed when it accepts a proton.
• hydronium ion: a polyatomic ion with the formula H₃O⁺, formed by the addition of the hydrogen cation (H⁺) to a molecule of water; also called oxonium (IUPAC preference) or hydroxonium.
• proton acceptor: a compound or part of a chemical compound that has the ability to accept a proton (H⁺) from a suitably acidic material in a chemical reaction.
• proton donor: a compound or part of a chemical compound that has the ability to relinquish a proton (H⁺) to a suitably basic material in a chemical reaction.

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