Acids and bases · Lesson 1

What makes an acid an acid

An acid is a proton donor. That one idea explains conjugate pairs, amphiprotic ions, and why two acids at the same concentration can behave so differently.

In 1884 a young Swedish chemist, Svante Arrhenius, submitted a doctoral thesis claiming that acids, bases and salts break up into charged particles when they dissolve in water. His examiners were unconvinced and gave it a low pass. Nineteen years later the same idea won him the Nobel Prize in Chemistry.

Arrhenius described an acid as a substance that produces hydrogen ions in water, and a base as one that produces hydroxide ions. That works for hydrochloric acid and sodium hydroxide, but it cannot explain ammonia, a base with no hydroxide in its formula. Nor can it explain the white smoke that appears when hydrogen chloride gas meets ammonia gas, with no water anywhere.

In 1923 Johannes Brønsted in Copenhagen and Thomas Lowry in Cambridge, working independently, proposed a simpler and more powerful idea: acid–base reactions are about the transfer of a single proton. It is the definition you will use for the rest of this course.

Acids donate protons; bases accept them

A hydrogen atom that has lost its only electron is just a proton, written H⁺. According to the Brønsted–Lowry theory:

an acid is a proton donor;
a base is a proton acceptor.

A bare proton never floats free in water: it is always held by water molecules. We write this hydrated proton as the hydrogen ion, H⁺(aq). So when hydrogen chloride dissolves, it gives its proton to the water: HCl(g) → H⁺(aq) + Cl⁻(aq). HCl is the acid (the proton donor). On this site, and in your answers, an acid in water is shown releasing H⁺(aq).

Ammonia is a base because it accepts a proton, from water: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq). This time water is the acid. And the white smoke? HCl(g) + NH₃(g) → NH₄Cl(s): a proton moves from HCl to NH₃, with no water involved at all.

Conjugate acid–base pairs

When an acid gives away a proton, what remains can accept a proton back. That species is its conjugate base. A conjugate pair always differs by exactly one H⁺:

to find the conjugate base, remove one H and lower the charge by one: H₂CO₃ → HCO₃⁻;
to find the conjugate acid, add one H and raise the charge by one: NH₃ → NH₄⁺.

Every Brønsted–Lowry reaction contains two conjugate pairs. In CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq), the pairs are CH₃COOH/CH₃COO⁻ and H⁺/H₂O. In the reverse reaction, H⁺ acts as the acid and CH₃COO⁻ as the base.

Amphiprotic species

Some species can either donate or accept a proton, depending on what they react with. They are amphiprotic. Water is the most important example: it acts as a base towards HCl and as an acid towards NH₃. The hydrogencarbonate ion is another:

as an acid: HCO₃⁻(aq) ⇌ CO₃²⁻(aq) + H⁺(aq)
as a base: HCO₃⁻(aq) + H₂O(l) ⇌ H₂CO₃(aq) + OH⁻(aq)

Others include HSO₄⁻, H₂PO₄⁻ and HPO₄²⁻. An amphiprotic species must contain a hydrogen it can lose and be able to accept one more.

Strong and weak are not the same as concentrated and dilute

The strength of an acid describes how completely it releases protons (H⁺) in water:

a strong acid ionises completely, written with a single arrow. Examples: HCl, HNO₃, and H₂SO₄ (its first proton);
a weak acid ionises only partly and sets up an equilibrium, written with ⇌. Examples: ethanoic acid, carbonic acid, phosphoric acid, HF.

Bases are classified the same way. Ionic hydroxides such as NaOH and KOH are strong bases: they dissociate completely, giving OH⁻ ions directly. Ammonia and the carbonate ion are weak bases: they accept protons from water only partly.

The concentration of an acid is a different idea entirely: how many moles of acid were dissolved per litre. A solution can be any combination:

Solution	Strength	Concentration	What is in it
2.0 mol L⁻¹ HCl	strong	concentrated	lots of ions, almost no HCl molecules
0.001 mol L⁻¹ HCl	strong	dilute	few ions, still no HCl molecules
17 mol L⁻¹ ethanoic acid (glacial)	weak	concentrated	mostly CH₃COOH molecules, a small fraction ionised
0.01 mol L⁻¹ ethanoic acid	weak	dilute	few molecules, even fewer ions

Loading the 3D model…

Strong and weak acids

Compare the two acids at the same concentration, then compare dilute and concentrated. The bulb shows how well each solution conducts electricity.

Because a strong acid gives far more ions, it conducts electricity much better than a weak acid of the same concentration, and its pH is lower. 0.10 mol L⁻¹ HCl has a pH of 1.0; 0.10 mol L⁻¹ ethanoic acid, only about 1.3% ionised, has a pH of about 2.9.

Measuring strength: the acidity constant

For a weak acid HA, the ionisation is an equilibrium, so it has an equilibrium constant, the acidity constant:

\[\mathrm{HA(aq) + H_2O(l) \rightleftharpoons A^-(aq) + H_3O^+(aq)} \qquad K_\mathrm{a} = \frac{[\mathrm{H^{+}}][\mathrm{A^{-}}]}{[\mathrm{HA}]}\]

Water is left out because it is the liquid solvent. The larger K_a, the further the equilibrium lies to the right, and the stronger the acid.

Acid	Conjugate base	K_a at 25 °C (approximate)
HSO₄⁻	SO₄²⁻	1.2 × 10⁻²
H₃PO₄	H₂PO₄⁻	7.5 × 10⁻³
HF	F⁻	6.8 × 10⁻⁴
HCOOH (methanoic acid)	HCOO⁻	1.8 × 10⁻⁴
CH₃COOH (ethanoic acid)	CH₃COO⁻	1.8 × 10⁻⁵
H₂CO₃	HCO₃⁻	4.3 × 10⁻⁷
NH₄⁺	NH₃	5.6 × 10⁻¹⁰
HCN	CN⁻	6.2 × 10⁻¹⁰

Two patterns are worth remembering. First, the stronger an acid, the weaker its conjugate base: Cl⁻ has essentially no tendency to take a proton back, while CN⁻ (the conjugate base of the very weak acid HCN) is a reasonably good base. Second, acid–base equilibria lie on the side of the weaker acid and weaker base.

Polyprotic acids

A polyprotic acid can donate more than one proton, one step at a time. Phosphoric acid is triprotic:

H₃PO₄ ⇌ H₂PO₄⁻ + H⁺, K_a = 7.5 × 10⁻³
H₂PO₄⁻ ⇌ HPO₄²⁻ + H⁺, K_a = 6.2 × 10⁻⁸
HPO₄²⁻ ⇌ PO₄³⁻ + H⁺, K_a = 4.8 × 10⁻¹³

Each step is much weaker than the one before, because it is harder to pull a positive proton away from an ion that is already negative, and more negative each time. Sulfuric acid is diprotic: its first ionisation is complete (strong), its second is not (HSO₄⁻ is a weak acid).

Dilution and a weak acid

Diluting a weak acid lowers every concentration. By Le Chatelier’s principle the equilibrium shifts towards the side with more dissolved particles (the ions), so the percentage ionised increases, even though [H⁺] itself falls. You can see this in the 3D model: the dilute weak acid has a higher percentage ionised.

Reactions of acids, as ionic equations

Spectator ions are left out of ionic equations. Strong acids are written as H⁺(aq) (or H⁺); weak acids are written as molecules, because that is mostly what is present.

Reaction	Ionic equation	Observation
acid + reactive metal	Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g)	bubbles; the metal dissolves
acid + carbonate	CaCO₃(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂O(l) + CO₂(g)	bubbles; the solid dissolves
strong acid + hydroxide	H⁺(aq) + OH⁻(aq) → H₂O(l)	no visible change; the mixture warms
acid + metal oxide	CuO(s) + 2H⁺(aq) → Cu²⁺(aq) + H₂O(l)	black solid dissolves; blue solution
weak acid + hydroxide	CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l)	no visible change; the mixture warms

A lovely consequence of weak-acid equilibria: react excess magnesium with equal volumes of 0.10 mol L⁻¹ HCl and 0.10 mol L⁻¹ ethanoic acid. The ethanoic acid fizzes far more slowly, because its [H⁺] is much lower. But it produces exactly the same volume of hydrogen in the end. As H⁺ is used up, the ethanoic acid equilibrium keeps shifting to the right, releasing more H⁺, until every molecule has reacted.

Excess magnesium with 50 mL of each acid at 0.10 mol L⁻¹: different rates, same total volume.

Conjugate pairs without mistakes

Find the species that lost a proton (the acid) and the species that gained it (the base) in the forward reaction.
Pair each with its partner on the other side: the two members of a pair differ by exactly one H⁺, and their charges differ by exactly one.
Check: one pair has the acid on the left, the other has the acid on the right.

Common traps

H⁺/OH⁻ is not a conjugate pair: they differ by two protons.
“Strong” never means “concentrated”. A strong acid can be very dilute.
Weak acids are written as molecules in ionic equations, not as H⁺.

Now put it to workFive worked examples, twenty practice questions and ten harder problems on this lesson.

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