Acid–base chemistry is a cornerstone of chemical science, underpinning everything from laboratory experiments to the processes that occur inside the human body.
One of the most useful frameworks for understanding acids and bases is the Brønsted-Lowry theory, which was introduced in the early 20th century.
This approach broadened the traditional definition of acids and bases, making it more versatile across different chemical systems. At its core, the Brønsted-Lowry model defines acids as proton donors and bases as proton acceptors.
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Key Takeaways
- A Brønsted-Lowry base is a proton (H⁺) acceptor, typically containing a lone pair of electrons
- This definition is broader than the older Arrhenius model and applies to a wide variety of systems
- Brønsted-Lowry bases are often also Lewis bases, but the two theories are not identical
- Common household and medical examples include ammonia, soaps, milk of magnesia, and bleach
What Defines a Brønsted-Lowry Base?
The Brønsted-Lowry base is defined by its ability to accept a proton (H⁺ ion). This is different from older definitions, such as the Arrhenius model, which restricted bases to substances that produce hydroxide ions (OH⁻) in water.
By focusing instead on proton transfer, Brønsted and Lowry created a more universal way of classifying bases.
A base’s ability to accept a proton usually comes from the presence of a lone pair of electrons on atoms such as nitrogen or oxygen. This lone pair can form a bond with the incoming proton, creating what is called a conjugate acid.
For instance, when ammonia (NH₃) accepts a proton, it becomes ammonium (NH₄⁺). This demonstrates not only how bases function but also introduces the concept of conjugate acid–base pairs.
The power of this definition is its flexibility. It allows us to describe reactions that occur in non-aqueous systems, such as in organic solvents or in biological molecules.
It also lets chemists analyse complex proton transfer reactions in ways that the Arrhenius model simply cannot.
Understanding the Brønsted-Lowry Theory
The Brønsted-Lowry theory was proposed independently by Johannes Brønsted and Thomas Lowry in 1923.
Their aim was to address the limitations of earlier acid–base definitions and to provide a framework that worked across a wider variety of chemical environments. By shifting the focus to proton transfer, the theory became a more powerful tool for both chemists and educators.
In a Brønsted-Lowry reaction, there are always two pairs at play: an acid with its conjugate base, and a base with its conjugate acid.
Consider the reaction of hydrochloric acid (HCl) with ammonia (NH₃):
HCl + NH₃ → NH₄⁺ + Cl⁻
Here, HCl donates a proton and becomes Cl⁻, while NH₃ accepts the proton and becomes NH₄⁺. Each species has a direct partner—HCl and Cl⁻ form one pair, NH₃ and NH₄⁺ form the other. This dual-pair concept is central to understanding the theory.
Because the Brønsted-Lowry approach is not limited to water-based reactions, it also applies to systems like liquid ammonia or organic solvents.
For example, acids and bases in petroleum chemistry, biological enzymes, or industrial catalysts can all be described under this framework, making it a vital part of modern chemistry.
Brønsted-Lowry vs. Lewis Theory
While the Brønsted-Lowry theory focuses on proton transfer, the Lewis theory, proposed by Gilbert N. Lewis in 1923, takes a different perspective.
A Lewis base is defined as any substance that can donate a pair of electrons to form a bond, while a Lewis acid is one that accepts an electron pair.
This difference may seem subtle, but it has significant consequences. Many Brønsted-Lowry bases, such as ammonia, are also Lewis bases because they donate an electron pair when accepting a proton.
However, the Lewis theory is broader, covering situations that do not involve protons at all. For example, reactions involving metal ions and ligands are best explained using Lewis definitions.
For students and chemists alike, it’s important to remember that neither theory is “wrong.” Instead, each offers a useful lens through which we can view acid–base behaviour.
Examples of Brønsted-Lowry Bases in Everyday Life
The Brønsted-Lowry theory may seem abstract, but its relevance in everyday life is clear once we look at the substances around us that make use of its principles.
Many of the chemicals we use daily, whether in the kitchen, bathroom, or medicine cabinet, act as Brønsted-Lowry bases.
Ammonia
One of the simplest and most familiar Brønsted-Lowry bases is ammonia (NH₃).
Found in both household cleaning chemical agents and fertilisers, ammonia is widely used due to its ability to accept a proton and form ammonium (NH₄⁺).
In water, ammonia reacts to form ammonium and hydroxide ions, making the solution basic:
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
This property explains why ammonia-based cleaners are effective at cutting through grease and grime, as the hydroxide ions increase the solution’s alkalinity.
From an industrial perspective, ammonia is one of the most produced chemicals worldwide, with global production exceeding 180 million tonnes per year. It plays a crucial role in agriculture through fertiliser contract manufacturing and in chemical synthesis more broadly.
Soaps
Soaps provide an excellent everyday example of Brønsted-Lowry bases. They are typically made by reacting fats or oils with strong bases such as sodium hydroxide in a process known as saponification. The result is a mixture of long-chain carboxylate salts, which act as bases in water.
These soap molecules contain negatively charged carboxylate groups that can accept protons. When dissolved in water, they slightly increase the concentration of hydroxide ions, giving soap its characteristic slippery and alkaline feel.
Beyond chemistry, this property has a practical application: soaps are effective emulsifiers, meaning they can break down oils and fats into small droplets that can be rinsed away.
Without their base-like behaviour, the cleaning power of soap would be much less effective.
Milk of Magnesia
A well-known medicinal example of a Brønsted-Lowry base is milk of magnesia, which is a suspension of magnesium hydroxide [Mg(OH)₂] in water. It is widely used as an antacid to relieve symptoms of indigestion and heartburn.
Magnesium hydroxide accepts protons from excess stomach acid (HCl), forming water and magnesium chloride in the process:
Mg(OH)₂ + 2 HCl → MgCl₂ + 2 H₂O
By neutralising stomach acid, it alleviates discomfort and restores balance in the digestive system.
Pharmacological studies note that milk of magnesia has a pH around 10.5, making it strongly basic.
Its effectiveness and safety have made it a common over-the-counter remedy for decades, illustrating how Brønsted-Lowry bases contribute directly to human health.
Household Bleach
Household bleach, usually containing sodium hypochlorite (NaOCl), is another familiar Brønsted-Lowry base.
Used for disinfecting and whitening, bleach solutions are highly alkaline, with a pH close to 12.6.
In solution, sodium hypochlorite can accept protons, producing hydroxide ions that contribute to the strong basicity of bleach:
OCl⁻ + H₂O ⇌ HOCl + OH⁻
This high alkalinity helps break down organic stains and effectively kill bacteria and viruses, making bleach a powerful cleaning agent.
While bleach is a valuable base in domestic and industrial settings, its caustic nature also makes it hazardous. Safe handling guidelines stress the importance of protective gloves and proper ventilation when using bleach solutions.
Conclusion
A Brønsted-Lowry base is a proton acceptor, typically possessing a lone electron pair. It extends beyond traditional Arrhenius definitions and encompasses many common substances. From ammonia to milk of magnesia, these bases play vital roles in cleaning, neutralisation, health, and industrial chemistry.
