The Brönsted-Lowry Acids and Bases

The Brønsted–Lowry Theory of Acids and Bases

The older idea of acids and bases — that acids are substances that taste sour, turn blue litmus red, and liberate hydrogen gas with certain metals, while bases taste bitter, feel soapy, and turn red litmus blue — is purely descriptive. It tells you what happens, but not why. A deeper understanding requires a theory that explains the actual chemical behaviour.

The Brønsted–Lowry theory, proposed independently by Johannes Brønsted and Thomas Lowry in 1923, provides exactly that. It shifts the focus from the substance itself to the role it plays in a chemical reaction.

The core idea is simple: an acid is a proton ($\text{H}^+$) donor, and a base is a proton acceptor. A proton, in this context, is simply a hydrogen ion — a hydrogen atom that has lost its electron, leaving just a bare nucleus.

This definition is much more powerful than the old one. It explains why a substance behaves as an acid in one situation and a base in another — it depends on what it is reacting with.

Conjugate Acid-Base Pairs

The key to the Brønsted–Lowry theory is that proton transfer is a two-way process. When an acid donates a proton, what remains is a species that can potentially accept a proton back — it is a base. Similarly, when a base accepts a proton, it becomes a species that can donate that proton — it is an acid.

Consider the reaction of hydrochloric acid with water:

Here, $\text{HCl}$ donates a proton to $\text{H}_2\text{O}$. So, $\text{HCl}$ is the acid and $\text{H}_2\text{O}$ is the base. After the transfer, the $\text{Cl}^-$ ion is what remains of the acid. It can accept a proton to reform $\text{HCl}$, so it is a base. The $\text{H}_3\text{O}^+$ ion is what the water molecule becomes after accepting a proton. It can donate a proton to reform $\text{H}_2\text{O}$, so it is an acid.

This gives us two pairs of substances that are related by the gain or loss of a single proton. These are called conjugate acid-base pairs.

• Pair 1: $\text{HCl}$ (acid) and $\text{Cl}^-$ (its conjugate base).
• Pair 2: $\text{H}_2\text{O}$ (base) and $\text{H}_3\text{O}^+$ (its conjugate acid).

The two species in a conjugate pair differ only by a single proton ($\text{H}^+$). The conjugate base of a strong acid is a weak base, and the conjugate acid of a strong base is a weak acid.

Amphoteric Substances

A substance that can act as both an acid and a base is called amphoteric. Water is the most common example. In the reaction with $\text{HCl}$, water acted as a base (accepted a proton). But water can also act as an acid, as in its reaction with ammonia:

Here, water donates a proton to ammonia. So, water is the acid and ammonia is the base. The conjugate pairs are $\text{H}_2\text{O}$ (acid) / $\text{OH}^-$ (conjugate base) and $\text{NH}_3$ (base) / $\text{NH}_4^+$ (conjugate acid).

Relative Strengths of Acids and Bases

The Brønsted–Lowry theory also provides a way to understand the relative strengths of acids and bases. The strength of an acid is measured by its tendency to donate a proton. A strong acid, like $\text{HCl}$, donates its proton almost completely. A weak acid, like acetic acid ($\text{CH}_3\text{COOH}$), donates its proton only partially.

This tendency is directly related to the stability of the conjugate base. A strong acid has a weak conjugate base (one that holds onto its electrons tightly and is stable as an anion). A weak acid has a strong conjugate base (one that has a high tendency to accept a proton).

The Role of the Solvent

The solvent plays a crucial role in determining the apparent strength of an acid or base. Water, with its high dielectric constant, is very effective at stabilizing the ions formed during proton transfer. This is why strong acids like $\text{HCl}$ and $\text{HNO}_3$ are completely ionized in water. …