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

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 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:

For ammonia reacting with water:

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:

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

Similarly Kb for weak bases:

Strong vs concentrated (an exam trap)

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

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.

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.

Examples in context

Example 1. Ammonia neutralisation at Townsville sugar mill stack scrubbers. Wilmar's stack-gas scrubbers contact flue gas with dilute sulfuric acid to remove ammonia: $\text{NH}_3(g) + \text{H}_2 \text{SO}_4(aq) \rightarrow \text{NH}_4^+(aq) + \text{HSO}_4^-(aq)$. The reaction is classic Bronsted-Lowry: $\text{NH}_3$ accepts a proton (base), $\text{H}_2 \text{SO}_4$ donates one (acid). The conjugate pairs are $\text{NH}_3 / \text{NH}_4^+$ and $\text{H}_2 \text{SO}_4 / \text{HSO}_4^-$. Scrubbed gas dropping below the EPA limit of $20 \, \text{ppm}$ ammonia keeps Townsville downwind communities safe; the resulting $\text{(NH}_4)_2 \text{SO}_4$ solution is sold as fertiliser, closing the loop on nitrogen.

Example 2. Bicarbonate dosing in Brisbane River pH control. Brisbane City Council adds sodium bicarbonate to pump stations to neutralise acidic infiltration: $\text{HCO}_3^- + \text{H}^+ \rightarrow \text{H}_2 \text{CO}_3 \rightarrow \text{H}_2 \text{O} + \text{CO}_2$. Bicarbonate is amphiprotic: it can also accept a proton or donate one. In the second role $\text{HCO}_3^- + \text{OH}^- \rightarrow \text{CO}_3^{2-} + \text{H}_2 \text{O}$, where bicarbonate acts as a Bronsted acid donating $\text{H}^+$ to hydroxide. The two conjugate pairs ($\text{HCO}_3^- / \text{H}_2 \text{CO}_3$ and $\text{HCO}_3^- / \text{CO}_3^{2-}$) explain why bicarbonate is the central buffer of natural waters.

Try this

Q1. Define a Bronsted-Lowry acid and base. Give a balanced equation in which $\text{HSO}_4^-$ acts as a base. [3 marks]

• Cue. Acid = proton donor; base = proton acceptor. $\text{HSO}_4^- + \text{H}_3 \text{O}^+ \rightarrow \text{H}_2 \text{SO}_4 + \text{H}_2 \text{O}$.

Q2. Write the equation for the reaction of acetic acid with water and identify both conjugate acid-base pairs. [3 marks]

• Cue. $\text{CH}_3 \text{COOH} + \text{H}_2 \text{O} \rightleftharpoons \text{CH}_3 \text{COO}^- + \text{H}_3 \text{O}^+$. Pairs: $\text{CH}_3 \text{COOH} / \text{CH}_3 \text{COO}^-$ and $\text{H}_2 \text{O} / \text{H}_3 \text{O}^+$.

Q3. Consider phosphoric acid in cane-juice neutralisation. (a) Write three successive dissociations. (b) Identify all conjugate pairs. (c) Predict which dissociation has the lowest $K_a$ and justify. [3+2+2 marks]

• Cue. (a) $\text{H}_3 \text{PO}_4 \rightleftharpoons \text{H}_2 \text{PO}_4^- \rightleftharpoons \text{HPO}_4^{2-} \rightleftharpoons \text{PO}_4^{3-}$. (c) Third lowest: removing $\text{H}^+$ from $\text{HPO}_4^{2-}$ separates against increasing negative charge.