What Are Acids and Bases?

Acids and bases are two of the most important classes of chemical compounds in chemistry. You encounter them every day — the sour taste of lemon juice, the sting of vinegar, the slippery feel of soap, and the burn of stomach acid are all related to acids and bases.

But what exactly makes something an acid or a base? Scientists have developed several theories over time to explain their behaviour, and each theory is more powerful and broad than the one before it.

In everyday terms:

Acids taste sour, react with metals to produce hydrogen gas, and turn blue litmus paper red
Bases feel slippery or soapy, neutralise acids, and turn red litmus paper blue

These are useful observations — but chemistry needs more precise definitions. That's where the theories come in.

Analogy

Think of acids as "givers" and bases as "receivers" — much like passing a ball in a game. The acid passes a hydrogen ion (the ball) to the base, which catches it. This giving-and-receiving idea is central to the most widely used acid-base theory.

At MYP 5 level, you will explore two main theories of acids and bases:

The Arrhenius theory — the earliest and simplest model
The Brønsted-Lowry theory — a broader, more useful model

Each theory gives a formal definition of what an acid and a base actually are. You'll meet those definitions in detail in the sections ahead.

Understanding both theories will help you explain why substances behave the way they do in chemical reactions.

The Arrhenius Theory

In 1884, Swedish chemist Svante Arrhenius proposed the first modern scientific definition of acids and bases. His theory was groundbreaking at the time because it connected the behaviour of acids and bases to the ions they produce in water.

Arrhenius Acid

A substance that dissociates (breaks apart) in water to produce hydrogen ions, H⁺.

Arrhenius Base

A substance that dissociates in water to produce hydroxide ions, OH⁻.

Key examples of Arrhenius acids:

Hydrochloric acid: $HCl \to H^{+} + Cl^{-}$
Nitric acid: $HNO_{3} \to H^{+} + NO_{3}^{-}$
Sulfuric acid (simplified, first dissociation): $H_{2} SO_{4} \to H^{+} + HSO_{4}^{-}$

Note

Sulfuric acid actually dissociates in two steps. The first step (shown above) is essentially complete. The second step — $HSO_{4}^{-} ⇌ H^{+} + SO_{4}^{2 -}$ — is only partial. For simplicity, you will sometimes see the overall equation $H_{2} SO_{4} \to 2 H^{+} + SO_{4}^{2 -}$ used as an approximation, but be aware this is a simplification.

Key examples of Arrhenius bases:

Sodium hydroxide: $NaOH \to Na^{+} + OH^{-}$
Potassium hydroxide: $KOH \to K^{+} + OH^{-}$
Calcium hydroxide: $Ca(OH)_{2} \to Ca^{2 +} + 2 OH^{-}$

Note

In water, the H⁺ ion doesn't actually float around freely — it immediately bonds to a water molecule to form the hydronium ion, H₃O⁺. You may see H⁺ and H₃O⁺ used interchangeably in different textbooks. Both refer to the same thing in aqueous chemistry. You'll see H₃O⁺ appear frequently in the Brønsted-Lowry section ahead.

According to Arrhenius, when an acid and a base react together, their H⁺ and OH⁻ ions combine to form water:
$H^{+} + OH^{-} \to H_{2} O$
This reaction is called neutralisation. In practice, the full equation also produces a salt — for example:
$HCl + NaOH \to NaCl + H_{2} O$
Here, sodium chloride (NaCl) is the salt produced alongside water.

Note

Notice that the dissociation equations for strong acids and bases above use a one-way arrow (→), indicating essentially complete dissociation. Later, when you see Brønsted-Lowry equations for weak acids, you'll notice an equilibrium arrow (⇌) instead — this signals that the reaction does not go to completion and both forward and reverse reactions occur simultaneously.

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