Chemistry • Year 12 • Module 7 • Lesson 18
Organic Acids & Bases: pKa, Strength & Reactions
Apply pKa reasoning to real quantitative data, Australian wine chemistry, and pharmaceutical drug absorption — interpreting graphs and predicting reactions from first principles.
1. Interpret the pKa comparison — carboxylic acid substituent effects
The bar chart below shows the pKa values of five organic acids relevant to Australian industry and pharmaceuticals. Use the data to answer the questions. 9 marks
1.1 Using data from the chart, identify which acid is the strongest and explain, using the concept of the inductive effect, why chloroacetic acid has a lower pKa than ethanoic acid. 3 marks
1.2 The TGA regulates the absorption of non-steroidal anti-inflammatory drugs (NSAIDs). Ibuprofen (pKa 4.91) is absorbed in the stomach (pH ≈ 2) but not in the small intestine (pH ≈ 7.4). Using pKa reasoning, explain why ibuprofen is predominantly in its neutral, protonated form in the stomach. 3 marks
1.3 Aspirin (pKa 3.5) and ibuprofen (pKa 4.91) are both carboxylic acids. Calculate the ratio of Ka values for aspirin to ibuprofen, showing your working. What does this ratio tell you about the relative acid strengths? 3 marks
2. AWRI data — organic acids in Australian wine
The Australian Wine Research Institute (AWRI) monitors the concentration of organic acids in wine because they directly control wine pH and flavour stability. The table below shows three key wine acids with their pKa values. 7 marks
| Acid | pKa₁ | pKa₂ | Typical concentration in wine | Role |
|---|---|---|---|---|
| Tartaric acid | 2.98 | 4.34 | 3–9 g/L | Primary natural acid; stabilises wine pH |
| Malic acid | 3.46 | 5.10 | 1–4 g/L | Contributes crisp, green-apple character |
| Citric acid | 3.13 | 4.76 | 0.1–0.5 g/L | Minor acid; adds freshness; antimicrobial |
Source: Adapted from AWRI technical review data. Wine pH typically 3.1–3.8.
2.1 Rank the three acids from strongest to weakest (using pKa₁). Justify your ranking using pKa values. 2 marks
2.2 Tartaric acid and malic acid are both diprotic carboxylic acids. Identify the Brønsted–Lowry acid–base equilibrium for the first ionisation of tartaric acid, using HA for the acid and A⁻ for the conjugate base. Explain why pKa₁ < pKa₂ for both acids. 3 marks
2.3 A winemaker adds sodium hydrogen carbonate (NaHCO₃) to a wine to lower its acidity. Using pKa reasoning, predict whether this addition would reduce the concentration of tartaric acid in the wine. Write the ionic equation. 2 marks
3. Cause-and-effect chain — electron-withdrawing groups and pKa
Complete the cause-and-effect chain by filling in the empty effect boxes. Each cause leads to the next effect. 5 marks (1 per effect)
3.4 Using the completed chain, explain why adding a second Cl to the carbon (e.g. CHCl₂COOH, pKa ≈ 1.48) further lowers the pKa compared to ClCH₂COOH (pKa 2.86). 2 marks
4. Applied scenario — amino acid nutrition labelling
Food Standards Australia New Zealand (FSANZ) requires nutritional labelling of amino acids in protein supplements. Glycine is the simplest amino acid; it has an amino group (–NH₂, pKb ≈ 4.4) and a carboxyl group (–COOH, pKa ≈ 2.35). At physiological pH (7.4) glycine exists as a zwitterion: both groups are fully ionised. 5 marks
4.1 Identify which group acts as a Brønsted–Lowry acid and which acts as a Brønsted–Lowry base when glycine dissolves in water at pH 7.4. Write the ionisation equation for each group. 3 marks
4.2 FSANZ nutritional analysis is carried out at pH 7.4. Explain whether the carboxyl group of glycine would be protonated (–COOH) or deprotonated (–COO⁻) at this pH, using pKa reasoning. 2 marks
Q1.1 — Strongest acid and inductive effect (3 marks)
Chloroacetic acid (pKa 2.86) is the strongest acid — it has the lowest pKa value [1]. The chlorine atom is highly electronegative and withdraws electron density inductively through the sigma bonds toward the –COOH group [1]. This stabilises the conjugate base (ClCH₂COO⁻) — the negative charge is better dispersed, making the conjugate base a weaker base, shifting the equilibrium further right (lower pKa = stronger acid) [1].
Q1.2 — Ibuprofen absorption (3 marks)
Ibuprofen has pKa = 4.91. At stomach pH ≈ 2, the pH is much lower than the pKa [1]. When pH < pKa, the protonated (neutral, HA) form of the acid predominates — the equilibrium HA ⇌ A⁻ + H⁺ lies to the left because the high H⁺ concentration suppresses ionisation [1]. The neutral (HA) form is lipid-soluble and can cross the stomach lining, enabling absorption; the ionised (A⁻) form is water-soluble and cannot as easily cross the lipid membrane [1].
Q1.3 — Ka ratio calculation (3 marks)
Ka(aspirin) = 10⁻³·⁵ = 3.16 × 10⁻⁴; Ka(ibuprofen) = 10⁻⁴·⁹¹ = 1.23 × 10⁻⁵ [1].
Ratio = Ka(aspirin) / Ka(ibuprofen) = (3.16 × 10⁻⁴) / (1.23 × 10⁻⁵) ≈ 25.7 [1].
Aspirin is approximately 26 times more acidic than ibuprofen in terms of the equilibrium constant. At equal concentrations, aspirin produces significantly more H₃O⁺ in solution than ibuprofen [1].
Q2.1 — Ranking wine acids (2 marks)
Ranking by pKa₁ (strongest to weakest): Tartaric acid (pKa₁ 2.98) > Citric acid (pKa₁ 3.13) > Malic acid (pKa₁ 3.46) [1 for correct order, 1 for citing pKa values as justification]. Note: lower pKa = stronger acid.
Q2.2 — First ionisation and diprotic pKa reasoning (3 marks)
First ionisation of tartaric acid: HA + H₂O ⇌ A⁻ + H₃O⁺ (where HA = C₂H₂(OH)₂(COOH)COO⁻ simplified) [1].
pKa₁ < pKa₂ because: after the first proton is donated, the conjugate base A⁻ carries a negative charge that stabilises the remaining –COOH group through electrostatic inductive effects, making the second proton harder to remove [1]. The negatively charged A⁻ opposes the formation of the doubly-charged A²⁻; more energy is required (the equilibrium lies further left for the second ionisation) [1].
Q2.3 — NaHCO₃ and tartaric acid (2 marks)
Yes — tartaric acid (pKa₁ 2.98) is stronger than H₂CO₃ (pKa 6.4). The reaction proceeds left to right [1].
Ionic equation: C₄H₅O₆⁻ (tartrate) + HCO₃⁻ + H⁺ → C₄H₅O₆²⁻ + H₂O + CO₂(g) [simplified net ionic: RCOOH + HCO₃⁻ → RCOO⁻ + H₂O + CO₂] [1].
Q3 — Cause-and-effect chain (5 marks)
Effect 1: Chlorine withdraws electron density inductively through the C–C sigma bonds toward the carboxyl group, reducing electron density on the oxygen atoms of –COO⁻. [1]
Effect 2: The resulting conjugate base (ClCH₂COO⁻) has its negative charge better stabilised/dispersed (less concentrated on the oxygens) compared to CH₃COO⁻. [1]
Effect 3: A more stable conjugate base is a weaker base — the equilibrium ClCH₂COOH + H₂O ⇌ ClCH₂COO⁻ + H₃O⁺ lies further to the right → Ka is larger. [1]
3.4: A second Cl atom provides an additional inductive electron-withdrawal effect, stabilising the conjugate base even more than one Cl [1]. The cumulative effect of two Cl atoms withdrawing electron density further stabilises CHCl₂COO⁻, making it an even weaker base, further increasing Ka and lowering pKa [1].
Q4.1 — Glycine zwitterion ionisations (3 marks)
The carboxyl group acts as a Brønsted–Lowry acid: –COOH + H₂O ⇌ –COO⁻ + H₃O⁺ [1].
The amino group acts as a Brønsted–Lowry base: –NH₂ + H₂O ⇌ –NH₃⁺ + OH⁻ [1].
At pH 7.4, both groups are ionised — the carboxyl has donated its proton (–COO⁻) and the amino has accepted a proton (–NH₃⁺), forming the zwitterion ⁺NH₃–CH₂–COO⁻ [1].
Q4.2 — Protonation state at pH 7.4 (2 marks)
The pKa of the carboxyl group of glycine is 2.35. At pH 7.4, the pH is much greater than the pKa (7.4 >> 2.35) [1]. When pH > pKa, the equilibrium HA ⇌ A⁻ + H⁺ lies heavily to the right — the deprotonated form (–COO⁻) overwhelmingly predominates. The carboxyl group is fully deprotonated at pH 7.4 [1].