Introduction to Acid-Base Reactions: Study Guide (Chemistry 4.8)

Acids and bases as proton donors and acceptors

In Topic 4.7 you met acid-base reactions as reactions that move a proton from one species to another. The Brønsted-Lowry model gives the two roles names:

• A Brønsted-Lowry acid is a proton donor. It gives away an H⁺.
• A Brønsted-Lowry base is a proton acceptor. It takes an H⁺, using a lone pair of electrons to form a new bond to it.

A hydrogen atom that has lost its electron is just a proton, so “H⁺” and “proton” mean the same thing here. An acid must therefore contain a hydrogen atom that it can give away. A base needs a lone pair to hold the incoming proton. Notice that a base does not have to contain OH⁻. Ammonia, NH₃, and the carbonate ion, CO₃²⁻, are both bases because they can accept a proton.

The model describes a reaction. A species is an acid or a base because of what it does in that reaction. You will see below that the same species can be an acid in one reaction and a base in another.

The course focuses on acid-base reactions in aqueous solution. Another model, the Lewis model, describes acids and bases in terms of electron pairs; it is not assessed in this course.

Conjugate acid-base pairs

When an acid gives away its proton, what is left can take a proton back. So the leftover species is a base, called the acid’s conjugate base. In the same way, when a base accepts a proton it becomes the conjugate acid of that base.

A conjugate acid-base pair is two species that differ by exactly one H⁺:

• Conjugate base = acid minus one H⁺. The charge goes down by 1.
• Conjugate acid = base plus one H⁺. The charge goes up by 1.

For example, the conjugate base of NH₄⁺ is NH₃, and the conjugate acid of OH⁻ is H₂O.

Every proton-transfer reaction has two conjugate pairs. A good habit is to label all four species, as in Figure 1. The reverse reaction is also a proton transfer: H₃O⁺ can give a proton back to NO₂⁻. That is why H₃O⁺ is called an acid and NO₂⁻ a base.

Water: an acid or a base

Water plays a special role in aqueous solutions because its structure lets it do both jobs:

• The O atom in H₂O has two lone pairs. One can bond to an incoming H⁺, so water can accept a proton and become H₃O⁺. It acts as a base.
• The O–H bonds are polar, and an H can leave as H⁺ when a base is present to take it. Water then donates a proton and becomes OH⁻. It acts as an acid.

Which role water plays depends on its partner. With an acid such as HNO₂, water accepts the proton (Figure 1). With a base such as methylamine, water donates one:

CH₃NH₂(aq) + H₂O(l) ⇌ CH₃NH₃⁺(aq) + OH⁻(aq)

Water molecules can even transfer a proton to each other: H₂O + H₂O ⇌ H₃O⁺ + OH⁻. This happens only to a tiny extent, but it means pure water always contains a few H₃O⁺ and OH⁻ ions. Unit 8 makes this quantitative.

A species that can act as either a proton donor or a proton acceptor is called amphiprotic. Water is the most important one. Ions such as HSO₃⁻ and HCO₃⁻ are also amphiprotic: they have an H to give away and a lone pair to accept one.

Worked example 1: label the acid, base and conjugate pairs

Question. For each reaction, identify the Brønsted-Lowry acid, the base, the conjugate acid and the conjugate base. State the role water plays.

(a) HNO₂(aq) + H₂O(l) ⇌ H₃O⁺(aq) + NO₂⁻(aq)

(b) CH₃NH₂(aq) + H₂O(l) ⇌ CH₃NH₃⁺(aq) + OH⁻(aq)

(a) Find the species that loses an H⁺: HNO₂ becomes NO₂⁻. So HNO₂ is the acid and NO₂⁻ is its conjugate base. H₂O gains the H⁺ and becomes H₃O⁺, so H₂O is the base and H₃O⁺ is its conjugate acid. Water acts as a base.

(b) H₂O loses an H⁺ and becomes OH⁻, so H₂O is the acid and OH⁻ is its conjugate base. CH₃NH₂ gains the H⁺ and becomes CH₃NH₃⁺, so CH₃NH₂ is the base and CH₃NH₃⁺ is its conjugate acid. Water acts as an acid.

Check. In each pair the formulas differ by one H and the charges differ by 1: HNO₂ (0) and NO₂⁻ (−1); CH₃NH₃⁺ (+1) and CH₃NH₂ (0). Water appears as a base in (a) and as an acid in (b).

Worked example 2: an amphiprotic ion

Question. The hydrogen sulfite ion, HSO₃⁻, is amphiprotic.

(a) Write its conjugate base and its conjugate acid. (b) Write an equation in which HSO₃⁻ acts as an acid with hydroxide ions, and one in which it acts as a base with hydronium ions.

(a) Conjugate base: remove one H⁺ and lower the charge by 1, giving SO₃²⁻. Conjugate acid: add one H⁺ and raise the charge by 1, giving H₂SO₃.

(b) As an acid (donating H⁺ to OH⁻):

HSO₃⁻(aq) + OH⁻(aq) → SO₃²⁻(aq) + H₂O(l)

As a base (accepting H⁺ from H₃O⁺):

HSO₃⁻(aq) + H₃O⁺(aq) → H₂SO₃(aq) + H₂O(l)

Check. Count charge on each side. First equation: −1 + (−1) = −2 on the left; −2 + 0 = −2 on the right. Second: −1 + 1 = 0; 0 + 0 = 0. Both balance. A common slip is to write the conjugate base as “SO₃⁻”, removing the H but forgetting to change the charge.

Strong and weak acids, and their conjugate bases

When an acid dissolves in water, it donates protons to water molecules. Acids differ in how completely this happens:

• A strong acid ionises essentially completely: almost every acid molecule gives its proton to water. Examples are HCl, HBr, HI, HNO₃, HClO₄ and H₂SO₄ (for its first proton).
• A weak acid ionises only partly: at any moment most of the acid molecules are still intact. Examples are HNO₂, HF and ethanoic acid.

Now think about the conjugate bases. The extent of ionisation is a competition for the proton. If an acid gives up its proton completely, its conjugate base must have almost no pull on a proton. So:

The stronger the acid, the weaker its conjugate base.

Cl⁻, the conjugate base of the strong acid HCl, has a negligible tendency to accept a proton in water. NO₂⁻, the conjugate base of the weak acid HNO₂, accepts protons noticeably. That is why so much HNO₂ stays as molecules: the NO₂⁻ ions keep taking protons back.

The same idea works the other way round for bases: the stronger a base, the weaker its conjugate acid.

Worked example 3: comparing acids from a particle model

Question. Figure 2 shows particle models of two solutions made with equal amounts of two acids, HX and HY. Water molecules are not shown.

(a) What percentage of each acid has ionised? (b) Which is the stronger acid? (c) Which is the stronger base, X⁻ or Y⁻? (d) How does the H₃O⁺ concentration of the HX solution compare with that of the HY solution?

(a) Work out how many acid molecules there were at the start. Each ionised molecule gives one H₃O⁺ and one anion, so count either.

• HX: 5 X⁻ ions and 0 HX molecules, so 5 of 5 ionised: 100%.
• HY: 1 Y⁻ ion and 4 HY molecules, so 1 of 5 ionised: 20%.

(b) HX ionises completely, so HX is the stronger acid (it behaves as a strong acid). HY is a weak acid.

(c) In the HY solution, most protons stay attached to Y. So Y⁻ holds on to protons more strongly than X⁻ does. Y⁻ is the stronger base. This matches the rule: the weaker acid has the stronger conjugate base.

(d) The HX box has 5 H₃O⁺ ions and the HY box has 1, in the same volume. The H₃O⁺ concentration in the HX solution is 5 times that in the HY solution, even though the two acids were added in equal amounts.

Interpretation. Equal amounts of acid do not give equal H₃O⁺ concentrations. Strength describes how fully an acid ionises, not how much acid you added.

Common misconceptions

• “A base must contain OH⁻.” In the Brønsted-Lowry model a base is any proton acceptor: NH₃, CH₃NH₂, CO₃²⁻ and H₂O can all be bases.
• “Strong means concentrated.” Strength is the fraction that ionises. A dilute solution of a strong acid is still fully ionised; a concentrated weak acid is still mostly molecules.
• “The conjugate base of a strong acid is a strong base.” The opposite is true. Cl⁻ is an extremely weak base because HCl gives up its proton so readily.
• Changing the formula but not the charge. The conjugate base of HSO₃⁻ is SO₃²⁻, not SO₃⁻.
• Pairing species that differ by two protons. H₂SO₃ and SO₃²⁻ are not a conjugate pair; each pair differs by exactly one H⁺.
• “Water is just the solvent.” In many aqueous acid-base reactions water is a reactant, acting as the acid or the base.
• “H⁺ floats around on its own in water.” Protons are always attached to water molecules. H₃O⁺(aq) and H⁺(aq) are two ways of writing the same thing.

Where this leads

Next, Topic 4.9 turns from proton transfer to electron transfer, with balancing redox reactions using half-reactions. Unit 8 develops the ideas here into pH, acid ionisation constants and buffers, and uses them to explain acid-base titration curves, building on the titrations of Topic 4.6. Try the practice questions now, then use the revision notes and the checklist to consolidate.