The Brønsted-Lowry theory of acids and bases explains how acids and bases behave by focusing on the transfer of hydrogen ions between substances. According to this theory, an acid is a substance that donates a proton, while a base is a substance that accepts a proton. This definition provides a broader understanding of acid-base reactions than the earlier model based on hydrogen ions and hydroxide ions alone.
Introduced independently in 1923 by Danish chemist Johannes Nicolaus Brønsted and British chemist Thomas Martin Lowry, the theory helps explain familiar chemical processes, including the reaction between hydrochloric acid and water, the behavior of ammonia in water, and the chemical systems that help maintain the pH of blood. Its central idea is simple: an acid-base reaction involves the transfer of a proton from one chemical species to another.
How the Brønsted-Lowry theory defines acids and bases
A Brønsted-Lowry acid is a substance that donates a proton to another substance. A Brønsted-Lowry base is a substance that accepts that proton. A proton is the positively charged particle found in an ordinary hydrogen atom’s nucleus. When a hydrogen atom loses its only electron, what remains is a hydrogen ion, H⁺, which is essentially a proton.
Consider hydrochloric acid, HCl, reacting with water. Hydrochloric acid donates a proton to a water molecule, producing a hydronium ion, H₃O⁺, and a chloride ion, Cl⁻.
HCl + H₂O → H₃O⁺ + Cl⁻
In this reaction, HCl acts as the acid because it donates a proton. Water acts as the base because it accepts that proton. The resulting hydronium ion is a water molecule with an additional proton, while the chloride ion is what remains of hydrochloric acid after the proton is transferred.
This example illustrates an important feature of the theory: whether a substance acts as an acid or a base depends on the reaction in which it participates. Acidity and basicity are defined by what happens during a chemical interaction, rather than simply by the name or composition of a substance.
Why the theory goes beyond earlier definitions
An earlier approach, known as the Arrhenius theory, defines an acid as a substance that increases the concentration of hydrogen ions in water and a base as a substance that increases the concentration of hydroxide ions, OH⁻, in water. This model explains many common reactions, including the behavior of hydrochloric acid and sodium hydroxide in aqueous solutions.
However, the Arrhenius definition is limited because it focuses on substances dissolved in water and on the production of particular ions. Some important acid-base reactions do not require hydroxide ions or even water.
The Brønsted-Lowry theory addresses this limitation by concentrating on proton transfer. For example, hydrogen chloride gas and ammonia gas can react directly to form solid ammonium chloride.
HCl + NH₃ → NH₄Cl
Hydrogen chloride donates a proton, making it the acid. Ammonia accepts the proton, making it the base. The products are ammonium ions, NH₄⁺, and chloride ions, Cl⁻, arranged in an ionic solid.
Because this reaction can occur without water, it demonstrates why a more general definition of acids and bases is useful. The Brønsted-Lowry theory applies to proton-transfer reactions in gases, liquids, aqueous solutions, and other chemical environments where such transfers are possible.
How water can act as both an acid and a base
One of the most useful features of the Brønsted-Lowry theory is its recognition that some substances can act as either acids or bases. Water is a familiar example of this behavior.
When water reacts with hydrochloric acid, it accepts a proton and acts as a base. When water reacts with a substance that donates a proton less readily than water does, water can instead donate a proton and act as an acid.
For example, ammonia accepts a proton from water in an aqueous solution. Water donates the proton, producing ammonium and hydroxide ions.
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
Here, ammonia is the base, and water is the acid. The double arrow indicates that the reaction can proceed in both directions, with the relative amounts of reactants and products depending on the chemical equilibrium.
A substance that can act as either a Brønsted-Lowry acid or a base is called amphiprotic. Water is amphiprotic, as are substances such as bicarbonate, HCO₃⁻, which can either donate or accept a proton.
This flexibility is important in chemistry because it shows that a substance does not always have one fixed acid-base role. Its behavior depends on the other substances present and on the conditions of the reaction.
What conjugate acid-base pairs are
When an acid donates a proton, it becomes a species that can accept a proton in the reverse reaction. When a base accepts a proton, it becomes a species that can donate a proton in the reverse reaction. These related substances are called a conjugate acid-base pair.
In the reaction between hydrochloric acid and water, HCl donates a proton and becomes Cl⁻. Hydrochloric acid and chloride are therefore a conjugate acid-base pair. Water accepts the proton and becomes H₃O⁺, making water and hydronium another conjugate acid-base pair.
The complete reaction can be represented as:
HCl + H₂O → H₃O⁺ + Cl⁻
The conjugate acid of a base has one more proton than the original base, while the conjugate base of an acid has one fewer proton than the original acid. In these pairs, the two members differ by a single proton.
Conjugate pairs help chemists track proton movement and understand how reactions behave in both directions. They also explain why a species that acts as an acid in one reaction may act as a base in another.
For example, the bicarbonate ion can accept a proton to form carbonic acid, H₂CO₃, or donate a proton to form carbonate, CO₃²⁻. Its conjugate acid is carbonic acid, and its conjugate base is carbonate.
Recognizing conjugate pairs is especially useful when analyzing chemical equilibrium, comparing acid and base strengths, and understanding how solutions resist changes in acidity.
How strong acids and bases behave
The Brønsted-Lowry theory also provides a framework for understanding acid and base strength. The strength of an acid depends on how readily it donates a proton to a suitable base under specified conditions. The strength of a base depends on how readily it accepts a proton from an acid.
A strong acid transfers its proton to water very extensively when dissolved in water. Hydrochloric acid is a common example. In dilute aqueous solution, it ionizes essentially completely, producing hydronium and chloride ions.
A weak acid, by contrast, transfers its proton to water only partially. Acetic acid, the acid responsible for much of vinegar’s characteristic acidity, establishes an equilibrium between the undissociated acid and the ions it forms.
CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
In this reaction, acetic acid donates a proton, and water accepts it. The acetate ion, CH₃COO⁻, is the conjugate base of acetic acid.
Weak acids are not necessarily harmless, nor are strong acids necessarily more concentrated. Acid strength describes the extent to which an acid donates protons under specified conditions. Concentration describes how much of a substance is present in a given volume of solution. These are different properties.
The same distinction applies to bases. A strong base reacts very extensively with an appropriate proton donor, whereas a weak base accepts protons to a more limited extent under comparable conditions. Ammonia is a common weak base in water because only a fraction of its molecules accept protons from water at equilibrium.
The strengths of conjugate acid-base pairs are related. A strong acid generally has a very weak conjugate base, while a strong base generally has a very weak conjugate acid. This relationship helps explain which direction an acid-base reaction tends to favor.
How proton transfer determines the direction of a reaction
Acid-base reactions are governed not only by whether proton transfer is possible but also by how favorable that transfer is. Chemical equilibrium describes the balance between forward and reverse reactions when the overall composition no longer changes with time.
In a Brønsted-Lowry reaction, proton transfer generally favors the side containing the weaker acid and the weaker base. This principle helps predict the direction in which a reaction will proceed under given conditions.
For example, when hydrochloric acid reacts with water, the proton transfer strongly favors the formation of hydronium and chloride ions. Hydrochloric acid is a much stronger acid than hydronium in water’s usual acid-base context, and chloride is a very weak base. Consequently, the reverse reaction is not favored to an appreciable extent in dilute aqueous solution.
For a weak acid such as acetic acid, proton transfer to water is less extensive. Both the acid and its conjugate base remain present in appreciable amounts at equilibrium.
The tendency of an acid to donate protons is commonly expressed through its acid dissociation constant, Kₐ. A larger Kₐ generally indicates a stronger acid in a given solvent under comparable conditions. The corresponding pKₐ is defined as −log₁₀(Kₐ), so a lower pKₐ generally indicates a stronger acid.
These quantities allow chemists to compare acid strengths and estimate whether a proton-transfer reaction is likely to favor products or reactants. However, acid-base strength depends on the solvent and chemical environment, so comparisons must be made in an appropriate context.
How the theory explains pH
The Brønsted-Lowry theory provides the chemical basis for understanding many processes involving pH. In dilute aqueous solutions, pH is commonly defined as the negative base-10 logarithm of the hydronium ion activity. For many introductory calculations, hydronium concentration is used as an approximation.
pH = −log₁₀
The brackets indicate the molar concentration of hydronium ions when concentration is used in this approximation.
At about 25 °C, a neutral aqueous solution has a pH close to 7. Solutions with lower pH values are acidic, while those with higher pH values are basic. The exact neutral pH depends on temperature because the equilibrium between water molecules and their ions changes with temperature.
When an acid donates protons to water, it generally increases the hydronium concentration and lowers the pH. When a base accepts protons from water, it can increase the hydroxide concentration, which generally reduces the hydronium concentration and raises the pH.
However, pH is not simply a measure of how much acid or base has been added. It depends on chemical equilibria, concentration, temperature, and the other substances present. A concentrated weak acid and a dilute strong acid can have different pH values than their strengths alone might suggest.
The theory also explains why an aqueous base does not need to contain hydroxide ions in its original chemical formula. Ammonia, for instance, accepts protons from water, producing ammonium and hydroxide ions. Its basic behavior follows from proton transfer rather than from the direct release of hydroxide ions by ammonia itself.
Why buffer solutions are important
A buffer is a solution that resists large changes in pH when modest amounts of acid or base are added. Many buffers contain a weak acid and its conjugate base, allowing the solution to respond to proton additions or removals.
When a small amount of acid is added to an acetate buffer, acetate ions can accept the added protons and form acetic acid. When a small amount of base is added, acetic acid can donate protons, helping neutralize the added base. These reactions reduce the change in hydronium concentration compared with what would occur in an unbuffered solution.
The Brønsted-Lowry framework explains the process directly: the conjugate acid-base pair acts as a proton donor and proton acceptor, helping keep the solution’s acid-base balance relatively stable.
Buffers are important in biological systems, chemical manufacturing, laboratory experiments, and many environmental processes. Blood, for example, relies on several interacting buffer systems to help maintain a pH compatible with normal physiological function. The bicarbonate system is particularly important because it links acid-base chemistry with dissolved carbon dioxide and the body’s respiratory and metabolic processes.
Buffers have limits. Once a substantial fraction of the available acid or base component has reacted, the solution becomes less effective at resisting pH changes. Their performance depends on the amounts of the conjugate pair, their acid-base properties, and the surrounding conditions.
What the Brønsted-Lowry theory cannot explain by itself
Although the Brønsted-Lowry theory is broadly useful, it does not encompass every reaction that chemists classify as acid-base chemistry. Its central requirement is the transfer of a proton, so it cannot directly describe acid-base reactions in which no proton is transferred.
For example, boron trifluoride, BF₃, can accept an electron pair from ammonia to form a chemical bond. Boron trifluoride has no proton to donate, yet it behaves as an acid in the broader Lewis theory of acids and bases. In the Lewis model, an acid accepts an electron pair, and a base donates one.
The Lewis theory therefore includes many reactions that fall outside the Brønsted-Lowry definition. The two theories are not contradictory; they describe acid-base behavior from different perspectives. The Brønsted-Lowry theory is especially useful when tracking proton transfer, while the Lewis theory provides a broader framework for reactions involving electron pairs.
The Brønsted-Lowry model also does not, by itself, predict every quantitative detail of a reaction. To determine equilibrium composition or pH accurately, chemists may need acid and base dissociation constants, concentrations, temperature, solvent properties, and other relevant information.
Why the Brønsted-Lowry theory remains useful
The enduring value of the Brønsted-Lowry theory lies in its simple, flexible explanation of acid-base reactions. By defining acids as proton donors and bases as proton acceptors, it unifies many chemical processes that might otherwise appear unrelated.
The theory explains why water can act as either an acid or a base, how conjugate acid-base pairs form, why acids and bases differ in strength, and how proton transfer shapes chemical equilibrium. It also provides a foundation for understanding pH, buffer solutions, and acid-base reactions in biological and industrial settings.
Most importantly, it shifts attention from memorizing lists of acids and bases to understanding what substances do when they interact. Whenever a proton moves from one chemical species to another, the Brønsted-Lowry framework offers a clear way to identify the acid, the base, and their conjugate partners—and to begin explaining why the reaction proceeds as it does.
