Brønsted–Lowry Acids and Bases

Acids and bases are among the most familiar substances in chemistry, but what actually makes something an acid? The Brønsted–Lowry theory answers in one line: acids give away protons and bases accept them. This note covers the common acids and alkalis you must know by name and formula, the Brønsted–Lowry definitions and conjugate pairs, the difference between strong and weak acids and bases (and how to tell them apart experimentally), the pH scale, and neutralisation and salt formation. It underpins the titration curves in the next note and all of the acid–base chemistry in A Level.

Common acids and alkalis

An alkali is a base that dissolves in water to give hydroxide ions, $\ce{OH-(aq)}$. All alkalis are bases, but not all bases are alkalis: copper(II) oxide is a base (it neutralises acids) but is insoluble in water, so it is not an alkali.

The Brønsted–Lowry theory

A "proton" here means a hydrogen ion, $\ce{H+}$: a hydrogen atom that has lost its only electron is just a proton.

When hydrogen chloride dissolves in water, it donates a proton to a water molecule:

$$\ce{HCl(g) + H2O(l) -> H3O+(aq) + Cl-(aq)}$$

HCl is the acid (proton donor) and water is the base (proton acceptor). The oxonium ion, $\ce{H3O+}$, is what is really present in acidic solutions; $\ce{H+(aq)}$ is shorthand for it. The water's oxygen uses a lone pair to form a coordinate bond to the proton.

When ammonia dissolves in water, the roles are reversed:

$$\ce{NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)}$$

Ammonia is the base (it accepts a proton, using the lone pair on nitrogen) and water is the acid (it donates a proton). The hydroxide ions formed make the solution alkaline.

So water can act as either an acid or a base, depending on what it reacts with. A substance that can do both is described as amphoteric (or, for proton transfer specifically, amphiprotic).

Conjugate acid–base pairs

Every Brønsted–Lowry reaction is a proton transfer, and it is reversible in principle. When an acid donates a proton, what remains can accept one back: it is a base. These two species form a conjugate acid–base pair, differing by one $\ce{H+}$.

$$\ce{CH3COOH + H2O <=> CH3COO- + H3O+}$$

$\ce{CH3COOH}$	$\ce{H2O}$	$\ce{CH3COO-}$	$\ce{H3O+}$
acid 1	base 2	base 1	acid 2

• $\ce{CH3COOH}$ and $\ce{CH3COO-}$ are a conjugate pair (acid 1 and its conjugate base).
• $\ce{H2O}$ and $\ce{H3O+}$ are a conjugate pair (base 2 and its conjugate acid).

Strong and weak acids and bases

Hydrochloric acid is strong. Essentially every HCl molecule donates its proton:

$$\ce{HCl(aq) -> H+(aq) + Cl-(aq)}$$

Ethanoic acid is weak. In a $0.1\ \text{mol dm}^{-3}$ solution only about one molecule in a hundred is dissociated at any moment; the rest remain as $\ce{CH3COOH}$ molecules. An equilibrium is set up that lies well to the left:

$$\ce{CH3COOH(aq) <=> CH3COO-(aq) + H+(aq)}$$

Similarly, sodium hydroxide is a strong base (fully dissociated into $\ce{Na+}$ and $\ce{OH-}$), while ammonia is a weak base (only a small proportion reacts with water to form $\ce{NH4+}$ and $\ce{OH-}$).

Sulfuric acid is strong in its first dissociation, $\ce{H2SO4 -> H+ + HSO4-}$. At AS, treat it as a strong acid that can release two protons, giving sulfate salts with excess alkali.

The pH scale

The pH of a solution is a measure of its acidity: the higher the concentration of $\ce{H+(aq)}$, the lower the pH.

pH can be measured with a pH meter (most precise) or estimated with universal indicator, which gives a range of colours from red (pH 1) through green (pH 7) to purple (pH 14).

Telling strong and weak acids apart

Compare $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and $0.100\ \text{mol dm}^{-3}$ ethanoic acid. They have the same concentration of acid, but the strong acid has a much higher concentration of $\ce{H+}$ ions because it is fully dissociated.

test	hydrochloric acid (strong)	ethanoic acid (weak)	reason
pH (pH meter)	about 1	about 3	higher $[\ce{H+}]$ in the strong acid
universal indicator	red	orange	higher $[\ce{H+}]$ in the strong acid
electrical conductivity	high	low	more ions in solution to carry the current
reaction with magnesium ribbon	vigorous fizzing, magnesium dissolves quickly	slow fizzing	higher $[\ce{H+}]$, so more frequent collisions with Mg
reaction with calcium carbonate	rapid effervescence	slower effervescence	as above
total volume of $\ce{H2}$ from excess Mg	same	same	the same total moles of acid; as $\ce{H+}$ is used up, more ethanoic acid dissociates (equilibrium shifts right)
volume of NaOH to neutralise	same	same	the same moles of acid available to react

The last two rows are the trap. A weak acid reacts more slowly but not less: given time, every molecule of ethanoic acid can donate its proton, because removing $\ce{H+}$ shifts the dissociation equilibrium to the right.

Neutralisation and salts

In every reaction between an acid and an alkali in aqueous solution, the essential change is the same:

A salt is the compound formed when the hydrogen of an acid is replaced by a metal ion or an ammonium ion. The acid determines the type of salt:

acid	salt	example
hydrochloric acid	chloride	$\ce{HCl + NaOH -> NaCl + H2O}$
sulfuric acid	sulfate	$\ce{H2SO4 + 2KOH -> K2SO4 + 2H2O}$
nitric acid	nitrate	$\ce{HNO3 + NH3 -> NH4NO3}$
ethanoic acid	ethanoate	$\ce{CH3COOH + NaOH -> CH3COONa + H2O}$

Other reactions of acids also form salts:

• acid + metal oxide or hydroxide: salt + water, for example $\ce{CuO + H2SO4 -> CuSO4 + H2O}$;
• acid + carbonate: salt + water + carbon dioxide, for example $\ce{CaCO3 + 2HCl -> CaCl2 + H2O + CO2}$;
• acid + reactive metal: salt + hydrogen, for example $\ce{Mg + 2HCl -> MgCl2 + H2}$ (this is a redox reaction, not neutralisation);
• acid + ammonia: an ammonium salt, with no water formed, for example $\ce{NH3 + HCl -> NH4Cl}$.

Worked examples

Practice

Answers

1. Acid: proton donor. Base: proton acceptor.
2. (a) $\ce{HNO3(aq) + H2O(l) -> H3O+(aq) + NO3-(aq)}$. (b) $\ce{NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)}$.
3. $\ce{HCO3-}$ (acid) / $\ce{CO3^2-}$ (base); $\ce{H2O}$ (acid) / $\ce{OH-}$ (base).
4. As an acid it donates a proton: $\ce{HCO3- + OH- -> CO3^2- + H2O}$. As a base it accepts a proton: $\ce{HCO3- + H+ -> H2O + CO2}$ (via $\ce{H2CO3}$). It is amphoteric.
5. A strong acid is fully dissociated in aqueous solution; a concentrated acid has a large amount of acid per unit volume. They describe different things. Example: $0.001\ \text{mol dm}^{-3}$ hydrochloric acid is dilute but strong.
6. Sodium hydroxide is a strong base, fully dissociated, so $[\ce{OH-}] = 0.1\ \text{mol dm}^{-3}$. Ammonia is a weak base: only a small fraction reacts with water to form $\ce{OH-}$, so $[\ce{OH-}]$ is much lower and the solution is less alkaline (lower pH).
7. (a) $\ce{H2SO4 + 2KOH -> K2SO4 + 2H2O}$. (b) $\ce{2NH3 + H2SO4 -> (NH4)2SO4}$. (c) $\ce{2CH3COOH + Na2CO3 -> 2CH3COONa + H2O + CO2}$.
8. (a) HCl gives a faster initial rate: it is fully dissociated, so $[\ce{H+}]$ is higher and collisions between $\ce{H+}$ and Mg are more frequent. (b) The same total volume: each contains $0.0500\ \text{mol}$ of acid; as $\ce{H+}$ is used up, more ethanoic acid dissociates until all has reacted. $\ce{Mg + 2H+ -> Mg^2+ + H2}$: $n(\ce{H2}) = 0.0250\ \text{mol}$; volume $= 0.0250 \times 24.0 = 0.600\ \text{dm}^3$ ($600\ \text{cm}^3$) in each case.
9. An acid can only donate a proton if there is a base to accept it. Methylbenzene has no lone pairs able to accept a proton, so HCl stays as molecules. Water is a base: a lone pair on oxygen accepts the proton, forming $\ce{H3O+}$ and $\ce{Cl-}$, so the solution is acidic.
10. In water, $\ce{2H2O <=> H3O+ + OH-}$: $\ce{H3O+}$ is the acidic species and $\ce{OH-}$ the basic species. By analogy, in liquid ammonia $\ce{NH4+}$ is the acid (proton donor) and $\ce{NH2-}$ the base (proton acceptor). Neutralisation: $\ce{NH4+ + NH2- -> 2NH3}$ (the solvent, like water in aqueous neutralisation). Overall: $\ce{NH4Cl + NaNH2 -> NaCl + 2NH3}$; the salt is sodium chloride.