← Unit 3: Equilibrium, acids and redox reactions

QLDChemistrySyllabus dot point

Topic 1: Chemical equilibrium systems

Describe acids and bases using the Bronsted-Lowry model, including the identification of conjugate acid-base pairs, amphiprotic species, and the distinction between strong and weak acids and bases

A focused answer to the QCE Chemistry Unit 3 dot point on Bronsted-Lowry acids and bases. Defines proton donors and acceptors, walks through conjugate acid-base pairs, identifies amphiprotic species (water, hydrogencarbonate, dihydrogenphosphate), and contrasts strong with weak acids and bases using Ka and Kb in equilibrium terms.

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What this dot point is asking

QCAA wants you to define acids and bases under the Bronsted-Lowry model (proton donors and acceptors), identify conjugate acid-base pairs in equations, recognise amphiprotic species, and explain the distinction between strong and weak acids/bases as a position-of-equilibrium claim. The dot point sits beneath every acid-base IA1 stimulus and underpins the pH calculations expected later in the topic.

The answer

Under the Bronsted-Lowry model, an acid is a proton (H+) donor and a base is a proton acceptor. Every acid-base reaction is a transfer of a proton from one species to another, producing a conjugate base and a conjugate acid.

The Bronsted-Lowry framework

For a generic reaction:

HA+Bβ‡ŒAβˆ’+HB+HA + B \rightleftharpoons A^- + HB^+

  • HA is the acid (donates a proton).
  • B is the base (accepts the proton).
  • A- is the conjugate base of HA.
  • HB+ is the conjugate acid of B.

The acid and its conjugate base differ by exactly one proton. The base and its conjugate acid likewise differ by one proton. There are always two conjugate acid-base pairs in a Bronsted-Lowry equation.

Conjugate acid-base pairs

Worked example. For HCl dissociating in water:

HCl(aq)+H2O(l)β†’Cl(aq)βˆ’+H3O(aq)+HCl_{(aq)} + H_2O_{(l)} \rightarrow Cl^-_{(aq)} + H_3O^+_{(aq)}

Pair 1 Pair 2
HCl (acid) H2O (base)
Cl- (conjugate base) H3O+ (conjugate acid)

For ammonia reacting with water:

NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH_4^+_{(aq)} + OH^-_{(aq)}

Pair 1 Pair 2
H2O (acid) NH3 (base)
OH- (conjugate base) NH4+ (conjugate acid)

Notice water is the acid in one equation and the base in the other. Water is the most common amphiprotic species.

Amphiprotic species

An amphiprotic species can either donate or accept a proton, depending on the partner.

Common amphiprotic species:

  • H2O. Acid: donates H+ to NH3. Base: accepts H+ from HCl.
  • HCO3- (hydrogencarbonate). Acid: donates H+ to give CO3^2-. Base: accepts H+ to give H2CO3.
  • H2PO4- (dihydrogenphosphate). Acid: gives HPO4^2-. Base: gives H3PO4.
  • HPO4^2- (hydrogenphosphate). Acid: gives PO4^3-. Base: gives H2PO4-.
  • HSO4- (hydrogensulfate). Acid: gives SO4^2-. Base: gives H2SO4.
  • NH3 (ammonia). Acid: gives NH2- (rare in aqueous solution). Base: gives NH4+.
  • Amino acids in their zwitterion form (Unit 4 context).

QCAA past papers regularly ask for two balanced equations (one as acid, one as base) and the term amphiprotic. Note: amphiprotic specifically describes proton transfer; amphoteric is the broader Lewis-style term for both acidic and basic behaviour.

Strong vs weak acids and bases

The Bronsted-Lowry model frames acid strength as the position of the dissociation equilibrium.

Strong acids and bases dissociate essentially completely. The equilibrium lies overwhelmingly to the right; in QCE Chemistry we typically write a single arrow.

  • Common strong acids: HCl, HBr, HI, HNO3, H2SO4 (first dissociation), HClO4.
  • Common strong bases: NaOH, KOH (group 1 hydroxides), Ca(OH)2 (slightly less soluble, but fully ionised).

Weak acids and bases dissociate only partially. The equilibrium lies to the left; both molecular and ionic forms are present at significant concentration. Written with a double arrow.

  • Common weak acids: CH3COOH (ethanoic), HF, H2CO3, HCN, NH4+, H3PO4 (and stepwise).
  • Common weak bases: NH3, organic amines, conjugate bases of weak acids (CH3COO-, F-).

The strength of a weak acid is quantified by the acid dissociation constant Ka:

HA+H2Oβ‡ŒAβˆ’+H3O+HA + H_2O \rightleftharpoons A^- + H_3O^+

Ka=[Aβˆ’][H3O+][HA]K_a = \frac{[A^-][H_3O^+]}{[HA]}

Smaller Ka means weaker acid. Ka < 10^-2 is considered weak; Ka > 1 effectively strong.

Similarly Kb for weak bases:

B+H2Oβ‡ŒBH++OHβˆ’B + H_2O \rightleftharpoons BH^+ + OH^-

Kb=[BH+][OHβˆ’][B]K_b = \frac{[BH^+][OH^-]}{[B]}

Strong vs concentrated (an exam trap)

Strength (fraction dissociated) and concentration (mol/L) are independent.

Acid Strength Concentration Resulting pH
0.001 mol/L HCl strong dilute about 3
1.0 mol/L CH3COOH weak concentrated about 2.4
1.0 mol/L HCl strong concentrated about 0
0.001 mol/L CH3COOH weak dilute about 3.9

QCAA stimulus regularly pairs these to test whether students conflate the terms.

Conjugate strength relationship

The conjugate of a strong acid is a very weak base (Cl- has essentially no proton-accepting tendency in water). The conjugate of a weak acid is a measurable base (CH3COO- accepts H+ to a small but real extent in water; CH3COO- solutions are basic).

This explains why solutions of salts of weak acids (e.g. sodium ethanoate, sodium hydrogencarbonate) are mildly basic, and why salts of weak bases (e.g. ammonium chloride) are mildly acidic.

Linking back to Le Chatelier and Kc

The weak acid equilibrium responds to disturbances per Le Chatelier:

  • Add water (dilute): equilibrium shifts right; fraction ionised rises. (Ostwald dilution.)
  • Add A- (e.g. as the sodium salt): equilibrium shifts left; fraction ionised falls. This is the basis of buffers, addressed in a separate dot point.
  • Increase temperature: depends on Delta H of dissociation; for most weak acids slightly endothermic, so Ka rises modestly.

This continuity is one reason QCAA cross-links the acid-base topic to the earlier equilibrium dot points in EA Paper 2.

Worked example: identifying conjugate pairs from a stimulus equation

Equation given: HF(aq) + NH3(aq) leftrightarrow F-(aq) + NH4+(aq).

Identify each conjugate pair and predict the position of equilibrium.

Pair 1 HF (acid) F- (conjugate base)
Pair 2 NH3 (base) NH4+ (conjugate acid)

Ka(HF) = 6.6 x 10^-4 (moderately weak). Ka(NH4+) = 5.6 x 10^-10 (very weak). Since HF is a stronger acid than NH4+, equilibrium lies to the right (HF donates its proton to NH3 readily).

Predict: position of equilibrium is to the right; Kc > 1.

This kind of comparative-strength reasoning appears in QCAA EA short response.

Common traps

Calling water "neutral" when it acts as an acid. Water is amphiprotic. Whether it acts as acid or base depends on the partner.

Forgetting to write the conjugate acid of the base. Every Bronsted-Lowry equation has two pairs; if you label only one, you have only half the answer.

Mixing up strength and concentration. A 0.001 mol/L solution of HCl is dilute but the acid is still strong (fully ionised). Likewise concentrated ethanoic acid is still weak.

Using "acidic" for the conjugate of a strong base. Cl- is a spectator; its presence does not lower pH. A weak conjugate has measurable pH effect; a vanishingly weak one does not.

Treating the H3O+ vs H+ notation as a substantive distinction. Both refer to the proton in water; QCAA accepts either but H3O+ is preferred in Bronsted-Lowry contexts because it reinforces the proton-transfer mechanism.

In one sentence

Under the Bronsted-Lowry model, an acid is a proton donor and a base is a proton acceptor; every acid-base equation contains two conjugate pairs that differ by a single proton, water and HCO3- are common amphiprotic examples, and the distinction between strong and weak acids and bases is the position of the dissociation equilibrium (essentially complete for strong, far to the left for weak), quantified by Ka or Kb.

Past exam questions, worked

Real questions from past QCAA papers on this dot point, with our answer explainer.

2023 QCAA-style4 marksHydrogencarbonate ion (HCO3-) can act as both an acid and a base in aqueous solution. (a) Write balanced equations showing HCO3- acting as a Bronsted-Lowry acid and as a Bronsted-Lowry base, in each case identifying the conjugate species. (b) Name the property HCO3- exhibits.
Show worked answer β†’

A 4-mark answer needs both equations with conjugates labelled and the term named.

(a) HCO3- as an acid.

HCO_3^-_{(aq)} + H_2O_{(l)} \rightleftharpoons CO_3^{2-}_{(aq)} + H_3O^+_{(aq)}

Conjugate base of HCO3- is CO3^2-. Conjugate acid of H2O is H3O+.

HCO3- as a base.

HCO_3^-_{(aq)} + H_2O_{(l)} \rightleftharpoons H_2CO_{3(aq)} + OH^-_{(aq)}

Conjugate acid of HCO3- is H2CO3. Conjugate base of H2O is OH-.

(b) Property. HCO3- is amphiprotic (it can donate or accept a proton).

Markers reward both balanced equations with correctly identified conjugates, and the term amphiprotic (or amphoteric, with amphiprotic preferred for proton-transfer behaviour specifically).

2022 QCAA-style4 marksEthanoic acid (CH3COOH) has Ka = 1.8 x 10^-5 at 25 degrees C. Hydrochloric acid is fully ionised in water. Both are 0.10 mol/L solutions. (a) Write the dissociation equation for each acid in water. (b) Explain in terms of position of equilibrium why ethanoic acid is classified as a weak acid and hydrochloric acid as a strong acid. (c) Predict which solution will have the lower pH.
Show worked answer β†’

A 4-mark answer needs both equations, the equilibrium-based distinction, and the pH comparison with justification.

(a) Dissociation equations.

CH3COOH(aq)+H2O(l)β‡ŒCH3COO(aq)βˆ’+H3O(aq)+CH_3COOH_{(aq)} + H_2O_{(l)} \rightleftharpoons CH_3COO^-_{(aq)} + H_3O^+_{(aq)}

HCl(aq)+H2O(l)β†’Cl(aq)βˆ’+H3O(aq)+HCl_{(aq)} + H_2O_{(l)} \rightarrow Cl^-_{(aq)} + H_3O^+_{(aq)}

(Single arrow for HCl, double for the weak acid.)

(b) Equilibrium position. Ethanoic acid's Ka = 1.8 x 10^-5 is small (Kc << 1), so the equilibrium lies far to the left; most molecules remain undissociated and only a small fraction donate protons. Hydrochloric acid has effectively no equilibrium (single arrow); essentially every molecule donates its proton to water. The fraction ionised determines acid strength under the Bronsted-Lowry model.

(c) pH comparison. 0.10 mol/L HCl will have the lower pH. HCl gives [H3O+] approximately 0.10 mol/L, so pH approximately 1.0. Ethanoic acid gives [H3O+] much less than 0.10 mol/L (around 1.3 x 10^-3, pH around 2.9) because only a fraction dissociates.

Markers reward correct equations (arrows in particular), explicit reference to position of equilibrium and Ka, and a numerical pH comparison.

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