What is Bronsted-Lowry acid?

Short Answer

A Bronsted–Lowry acid is a substance that can donate a proton (H⁺ ion) to another substance. According to this theory, acids are not defined only in water; they can act as acids in many different solvents. Any molecule or ion that gives away a proton is considered a Bronsted–Lowry acid.

For example, hydrochloric acid (HCl), sulphuric acid (H₂SO₄), and even water (in some reactions) can act as Bronsted–Lowry acids because they donate protons. This theory gives a broader and more flexible explanation of acids compared to the Arrhenius theory.

Detailed Explanation :

Bronsted–Lowry Acid

The Bronsted–Lowry concept was introduced by two scientists, Johannes Bronsted and Thomas Lowry, in 1923. This theory expanded the idea of acids beyond the Arrhenius theory by explaining acid–base behaviour in terms of proton (H⁺) transfer. Because of this, it provides a much wider understanding of how acids behave in different chemical environments, not just in water.

According to the Bronsted–Lowry theory:

A Bronsted–Lowry acid is a substance that donates a proton (H⁺ ion).
A Bronsted–Lowry base is a substance that accepts a proton.

This simple idea helps explain many reactions that the Arrhenius theory could not describe.

Meaning of a Bronsted–Lowry Acid

A Bronsted–Lowry acid is defined as a proton donor. In any reaction where a substance gives away an H⁺ ion, it behaves as an acid. This definition does not depend on water. The acid can donate a proton in a liquid other than water or even in the gas phase.

For example:

  • HCl → H⁺ + Cl⁻
  • H₂SO₄ → H⁺ + HSO₄⁻
  • NH₄⁺ → H⁺ + NH₃

In all these reactions, the substance releases a proton, showing acidic behaviour.

This wider definition allows many more substances to be considered acids.

How Bronsted–Lowry Acids Work

In the Bronsted–Lowry concept, acids and bases always appear in pairs, called conjugate acid–base pairs.
When a Bronsted–Lowry acid donates a proton, it becomes a conjugate base.

Example:
HCl (acid) → donates H⁺ → forms Cl⁻ (conjugate base)

Another example:
H₂O (acid) → donates H⁺ → forms OH⁻ (conjugate base)

This helps explain why water can act both as an acid and a base, depending on the reaction.

Examples of Bronsted–Lowry Acids

  1. Mineral Acids – HCl, H₂SO₄, HNO₃
    These acids donate hydrogen ions easily.
  2. Weak Acids – CH₃COOH (acetic acid), H₂CO₃ (carbonic acid)
    These donate protons but less readily than strong acids.
  3. Ammonium Ion (NH₄⁺)
    This behaves as an acid because it can release H⁺.
  4. Water (H₂O)
    Water behaves as an acid when it donates a proton to another substance.

The Bronsted–Lowry definition shows that acids are not limited to substances that release hydrogen ions in water. Any proton donor qualifies as an acid.

Why Bronsted–Lowry Theory Is Important

The Bronsted–Lowry theory is important for several reasons:

  1. Works in Any Solvent

The Arrhenius theory works only in water, but Bronsted–Lowry theory explains acid–base behaviour in:

  • Water
  • Alcohol
  • Ammonia
  • Organic solvents
  • Even gases

This makes it more general and useful.

  1. Explains Many More Reactions

The theory explains reactions where no hydroxide ions (OH⁻) are present. For example, ammonia acts as a base even though it does not contain OH⁻. Bronsted–Lowry theory correctly describes it because ammonia accepts a proton.

  1. Helps Understand Conjugate Pairs

The theory introduces the idea of conjugate acid–base pairs, which helps explain reversible reactions and equilibrium in chemistry.

  1. Explains Amphoteric Substances

Some substances, like water, can behave as both acids and bases.

  • When water donates H⁺ → it behaves as an acid.
  • When water accepts H⁺ → it behaves as a base.

This dual behaviour is easily explained using the Bronsted–Lowry theory.

Reaction Examples Explained Through Bronsted–Lowry Theory

  1. HCl + NH₃ → NH₄⁺ + Cl⁻
  • HCl donates a proton → Bronsted–Lowry acid
  • NH₃ accepts a proton → Bronsted–Lowry base
  1. H₂O + HCl → H₃O⁺ + Cl⁻
  • HCl donates proton → acid
  • Water accepts proton → base

The hydronium ion (H₃O⁺) forms because water accepts the proton from HCl.

These examples show how the theory describes reactions that the Arrhenius idea could not fully explain.

Conclusion

A Bronsted–Lowry acid is any substance that donates a proton (H⁺) to another substance. This theory gives a broader and more flexible explanation of acids compared to the Arrhenius theory because it applies to many solvents and chemical environments. It also introduces the concept of conjugate acid–base pairs and explains amphoteric behaviour. Understanding Bronsted–Lowry acids helps in studying acid–base reactions more accurately and in understanding how substances behave in different chemical systems.