Welcome to A-Level acid–base chemistry! At GCSE, you learned that acids release $H^{+}$ ions in water and alkalis release ${OH}^{-}$ ions. While this is true, it is a slightly limited view of how chemical reactions actually work. In this topic, we will look at a much more powerful way to understand acids and bases.

What you'll learn:

• The Brønsted–Lowry definitions of acids and bases.
• Why an $H^{+}$ ion is simply referred to as a "proton".
• How acid–base reactions in solution form dynamic equilibria.
• How to identify conjugate acid–base pairs in a chemical equation.

The Brønsted–Lowry definitions

In 1923, two chemists—Johannes Nicolaus Brønsted and Thomas Martin Lowry—independently proposed a new way to define acids and bases. Instead of just looking at what substances release into water, they looked at how molecules interact with each other during a reaction.

They realised that acid–base chemistry is entirely about the transfer of $H^{+}$ ions from one species to another.

Why do we call the $H^{+}$ ion a "proton"? Think about a neutral hydrogen atom. It consists of exactly one proton in the nucleus and one electron orbiting it. If that hydrogen atom loses its only electron to become an $H^{+}$ ion, there is nothing left but the bare proton. Therefore, in chemistry, donating an $H^{+}$ ion is exactly the same as donating a proton.

Proton transfer in aqueous solution

Protons are highly reactive—they have a very high charge density and do not just float around on their own in aqueous solutions. When an acid donates a proton, there must be a base ready to accept it immediately.

Let's look at what happens when hydrogen chloride ($HCl$) gas dissolves in water:

$$HCl(g)+H_{2}O(l)\to H_3{O}^{+}(aq)+{Cl}^{-}(aq)$$

In this reaction:

• $HCl$ donates a proton to water. Therefore, $HCl$ is the acid.
• $H_{2}O$ accepts the proton using one of the lone pairs on its oxygen atom. Therefore, water is acting as the base.

The resulting $H_3{O}^{+}$ ion is called the oxonium ion (sometimes called the hydronium ion). Whenever we write $H^{+}(aq)$ in a chemical equation, it is actually a shorthand for $H_3{O}^{+}(aq)$.

Acid–base equilibria and conjugate pairs

Many acid–base reactions do not go to completion. Weak acids and weak bases only partially dissociate in water, establishing a dynamic equilibrium.

Because these reactions are reversible, proton transfer happens in both directions. Let's examine the equilibrium established when ethanoic acid (${CH}_{3}COOH$) is dissolved in water:

$${CH}_{3}COOH(aq)+H_{2}O(l)\rightleftharpoons {CH}_3{COO}^{-}(aq)+H_3{O}^{+}(aq)$$

Let's break down the two directions:

• The forward reaction: ${CH}_{3}COOH$ donates a proton to $H_{2}O$. So, ${CH}_{3}COOH$ is an acid, and $H_{2}O$ is a base.
• The reverse reaction: $H_3{O}^{+}$ donates a proton back to the ethanoate ion (${CH}_3{COO}^{-}$). Therefore, $H_3{O}^{+}$ is acting as an acid, and ${CH}_3{COO}^{-}$ is acting as a base.

Notice how the ethanoic acid (${CH}_{3}COOH$) turns into the ethanoate ion (${CH}_3{COO}^{-}$) by losing a proton. These two species are intimately linked and are known as a conjugate acid–base pair.

In any acid–base equilibrium, there will always be two conjugate pairs.

The amphoteric nature of water

You might have noticed something interesting in the examples above.

• When reacting with $HCl$ or ethanoic acid, water accepted a proton (acted as a base).
• When reacting with ammonia, water donated a proton (acted as an acid).

Species that can behave as either an acid or a base depending on what they are reacting with are described as amphoteric (specifically, amphiprotic, since they can either donate or accept a proton).