Chemistry • Year 12 • Module 6 • Lesson 13
Buffers: Mechanism, Calculations & Natural Systems
Synthesise buffer theory, clinical data, and environmental chemistry to construct and evaluate extended evidence-based responses at Band 5–6.
1. Multi-step calculation — phosphate buffer preparation at Westmead Hospital
Westmead Hospital’s pathology laboratory prepares intracellular-pH simulation buffers for enzyme assay work. A technician needs a phosphate buffer at pH 7.30 using NaH₂PO₄ (providing H₂PO₄⁻) and Na₂HPO₄ (providing HPO₄²⁻). The second ionisation of phosphoric acid gives Ka2(H₂PO₄⁻ → H⁺ + HPO₄²⁻) = 6.2 × 10−8. 8 marks
(a) Identify the weak acid and conjugate base in this buffer, and write the ionisation equation that defines the Ka2 given. 2 marks
(b) Calculate the pKa2 and use Henderson-Hasselbalch to find the [HPO₄²⁻]/[H₂PO₄⁻] ratio required to give pH 7.30. Show all working. 2 marks
(c) The technician decides to prepare the buffer by partial neutralisation: she starts with 200 mL of 0.300 mol/L NaH₂PO₄ and adds solid Na₂HPO₄ separately to achieve the required ratio. Instead of this method, she could add 0.200 mol/L NaOH to the NaH₂PO₄ solution. Calculate what volume of 0.200 mol/L NaOH must be added to 200 mL of 0.300 mol/L NaH₂PO₄ to give pH 7.30. Assume volume changes are negligible. 3 marks
(d) Interpret the biological significance of this buffer pH. Why is pH 7.30 a meaningful intracellular simulation value, and what would happen to enzyme-catalysed reactions if cytoplasmic pH fell to 6.80? 1 mark
2. Data + scenario evaluation — Great Barrier Reef carbonate buffer system (AIMS data)
The Australian Institute of Marine Science (AIMS) has monitored ocean chemistry at the Great Barrier Reef (GBR) since the 1990s. The table below shows measured pH and dissolved CO₂ data at a representative reef site over four decades. The carbonate system involves: CO₂(aq) + H₂O ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺ ⇌ CO₃²⁻ + 2H⁺. At the observed pH range, the dominant carbonate equilibrium for buffering is H₂CO₃/HCO₃⁻ (pKa1 = 6.35). 8 marks
| Decade | Average ocean pH (GBR surface) | Dissolved CO₂ (ppm atmospheric) | [HCO₃⁻]/[H₂CO₃] ratio (calculated) |
|---|---|---|---|
| 1980s | 8.18 | 340 | 676 |
| 1990s | 8.12 | 360 | 589 |
| 2000s | 8.08 | 385 | 540 |
| 2010s | 8.03 | 405 | 479 |
Adapted from AIMS Long-Term Monitoring Program GBR ocean chemistry records; atmospheric CO₂ from NOAA (Mauna Loa). [HCO₃⁻]/[H₂CO₃] ratio calculated via Henderson-Hasselbalch: ratio = 10(pH − pKa1); pKa1 = 6.35.
Using the data, your knowledge of buffer chemistry, and the Henderson-Hasselbalch equation, write an extended response that:
- Verifies the 1980s [HCO₃⁻]/[H₂CO₃] ratio of 676 using Henderson-Hasselbalch (pH = 8.18; pKa1 = 6.35);
- Describes the trend in ocean pH and [HCO₃⁻]/[H₂CO₃] ratio from the 1980s to the 2010s, and explains this trend using Le Chatelier’s principle and the Henderson-Hasselbalch equation;
- Evaluates the significance of the pH change from 8.18 to 8.03 for coral organisms that build CaCO₃ shells, using the equilibrium CaCO₃(s) + H⁺ ⇌ Ca²⁺(aq) + HCO₃⁻(aq);
- Reaches an evidence-based conclusion about whether the GBR carbonate buffer system can prevent continued ocean acidification if atmospheric CO₂ keeps rising, referencing the concept of buffer capacity.
Write your extended response here (approximately 200–250 words). Refer to the data table in your response.
3. Source critique — a student’s explanation of blood buffering
Read the following student response submitted in a Year 12 Chemistry exam. It contains a subtle scientific flaw. Your task is to identify the flaw, explain the correct chemistry, and describe how the error could be detected through a calculation or experimental observation. 7 marks
“Blood maintains a pH of 7.4 using the H₂CO₃/HCO₃⁻ buffer system. The pKa of carbonic acid is 6.10, and blood pH is 7.40, so the [HCO₃⁻]/[H₂CO₃] ratio is 20:1 by the Henderson-Hasselbalch equation. This means the buffer operates at maximum capacity because the ratio of 20:1 is very large, giving the buffer the most resistance to pH change. When lactic acid (from exercise) enters the bloodstream, HCO₃⁻ neutralises the H⁺ produced, preventing any change in blood pH. The buffer keeps blood pH exactly at 7.40 at all times.”
3.1 Identify all scientific flaws or inaccuracies in the student’s response. List each flaw clearly. 3 marks (1 per flaw)
3.2 Explain the correct chemistry for each flaw identified. Your answer must include: (i) the correct definition of maximum buffer capacity and when it occurs; (ii) the correct description of the buffer mechanism when H⁺ is added — including the equation and what happens to pH; (iii) the correct range of blood pH maintenance, not a single fixed value. 3 marks
3.3 Describe one calculation or experimental approach that would reveal that the student’s claim about “no change in blood pH” is incorrect. 1 mark
Q1(a) — Weak acid, conjugate base, ionisation equation
Weak acid: H₂PO₄⁻ (dihydrogen phosphate ion). Conjugate base: HPO₄²⁻ (hydrogen phosphate ion). Ionisation equation: H₂PO₄⁻(aq) ⇌ H⁺(aq) + HPO₄²⁻(aq), Ka2 = 6.2 × 10−8.
Q1(b) — pKa2 and [HPO₄²⁻]/[H₂PO₄⁻] ratio
pKa2 = −log(6.2 × 10−8) = 8 − log(6.2) = 8 − 0.792 = 7.21. pH = pKa + log([HPO₄²⁻]/[H₂PO₄⁻]) ⇒ 7.30 = 7.21 + log(ratio) ⇒ log(ratio) = 0.09 ⇒ ratio = 100.09 = 1.23. So [HPO₄²⁻]/[H₂PO₄⁻] = 1.23 (approximately 5:4).
Q1(c) — Volume of NaOH required
n(H₂PO₄⁻) initial = 0.300 × 0.200 = 0.0600 mol. Let n(NaOH) = x mol. Reaction: H₂PO₄⁻ + OH⁻ → HPO₄²⁻ + H₂O. n(HPO₄²⁻) formed = x; n(H₂PO₄⁻) remaining = 0.0600 − x. Required ratio: x/(0.0600 − x) = 1.23. x = 1.23(0.0600 − x) = 0.0738 − 1.23x. 2.23x = 0.0738. x = 0.0331 mol. Volume = 0.0331/0.200 = 165 mL of 0.200 mol/L NaOH. Check: n(NaOH) = 0.0331 < n(H₂PO₄⁻) = 0.0600 ✓ buffer forms.
Q1(d) — Biological significance
pH 7.30 is within the physiological intracellular range (pH 6.8–7.4 in most mammalian cells). The phosphate buffer (pKa = 7.21) is near-optimal for this range. If cytoplasmic pH fell to 6.80, ionisable side chains on enzyme active sites (histidine imidazole, pKa ~6.0; lysine amine, pKa ~10.5) would change protonation state, altering binding geometry and catalytic activity. Many metabolic enzymes (e.g. phosphoglycerate kinase, hexokinase) would show dramatically reduced rates, impairing glycolysis and cellular energy production.
Q2 — Extended response marking guide (8 marks)
(1) Verify 1980s ratio (1 mark): ratio = 10(pH − pKa) = 10(8.18 − 6.35) = 101.83 = 676 ✓. Matches the table value.
(2) Trend description + Le Chatelier + HH (3 marks): Ocean pH decreases from 8.18 (1980s) to 8.03 (2010s), a fall of 0.15 units corresponding to a 41% increase in [H⁺] (100.15 − 1 = 0.41). The [HCO₃⁻]/[H₂CO₃] ratio falls from 676 to 479 as atmospheric CO₂ rises from 340 to 405 ppm. Mechanism: increased atmospheric CO₂ dissolves in seawater; equilibrium CO₂(aq) + H₂O ⇌ H₂CO₃ shifts right (Le Chatelier); [H₂CO₃] increases, raising the denominator in pH = pKa + log([HCO₃⁻]/[H₂CO₃]); the log term decreases and pH falls.
(3) Significance for CaCO₃ organisms (2 marks): Increased [H⁺] shifts CaCO₃(s) + H⁺ ⇌ Ca²⁺(aq) + HCO₃⁻(aq) to the right. Shell dissolution rate increases; organisms expend more energy precipitating CaCO₃ against higher [H⁺]; at pH 8.03, the saturation state of aragonite (coral’s CaCO₃ polymorph) is significantly reduced compared to pre-industrial pH, threatening GBR coral calcification and structural integrity.
(4) Buffer capacity conclusion (2 marks): The data show that despite the carbonate buffer, ocean pH has already fallen by 0.15 units in 40 years — demonstrating that the buffer’s capacity is being exceeded by the rate of CO₂ addition. Buffer capacity is finite; as atmospheric CO₂ continues to rise, the ratio [HCO₃⁻]/[H₂CO₃] will continue to decrease. The ocean carbonate buffer can slow but cannot prevent ocean acidification because the buffer’s capacity (total moles of CO₃²⁻ + HCO₃⁻) is insufficient to absorb all the additional CO₂ generated by fossil fuel combustion.
Q3.1 — Flaws in the student response (3 marks)
Flaw 1 (1 mark): The student claims a ratio of 20:1 represents “maximum capacity.” Maximum buffer capacity occurs when [A⁻] = [HA] — a ratio of exactly 1:1 (or the half-equivalence point) — not when the ratio is 20:1. A ratio of 20:1 means only a small fraction of HA remains; the buffer can absorb almost no further acid before failing.
Flaw 2 (1 mark): The student says HCO₃⁻ “neutralises” H⁺ and prevents “any change in blood pH.” A buffer resists but does not prevent pH change; the reaction H⁺ + HCO₃⁻ → H₂CO₃ produces a small but real pH decrease (e.g. pH 7.40 → 7.30 during intense exercise). “Prevents any change” is incorrect; “neutralises” overstates the reaction.
Flaw 3 (1 mark): The student says blood pH is maintained “exactly at 7.40 at all times.” Normal blood pH is a range (7.35–7.45) and it does fluctuate within this range (e.g. falling slightly during exercise, returning to ~7.40 at rest). Blood pH is not a fixed constant value.
Q3.2 — Correct chemistry (3 marks)
(i) Maximum buffer capacity (1 mark): Buffer capacity is maximised when [A⁻] = [HA] (ratio = 1:1), i.e. at the half-equivalence point during titration. At this point, equal moles of weak acid and conjugate base are present, so the buffer can absorb the maximum amount of either additional acid (consuming A⁻) or additional base (consuming HA) before one component is exhausted. A ratio of 20:1 represents low residual capacity to absorb further acid (only 1 part HA remains per 20 parts A⁻).
(ii) Correct mechanism (1 mark): When H⁺ is added to blood, it reacts with HCO₃⁻: H⁺ + HCO₃⁻ → H₂CO₃ → CO₂ + H₂O. This reduces [HCO₃⁻] and increases [H₂CO₃], lowering the [HCO₃⁻]/[H₂CO₃] ratio. By Henderson-Hasselbalch, pH falls slightly. Blood pH does not remain at exactly 7.40 — it decreases (e.g. to ~7.30 in metabolic acidosis).
(iii) pH range (1 mark): Blood pH is maintained within a range of 7.35–7.45 (not fixed at exactly 7.40). Values below 7.35 = acidosis; above 7.45 = alkalosis. The buffer system, combined with respiratory and renal compensation, maintains pH within this range rather than at a single fixed value.
Q3.3 — Calculation or experiment to detect the error
Apply Henderson-Hasselbalch after exercise: if lactic acid adds 0.003 mol H⁺/L to blood ([HCO₃⁻] falls from 24 to ~21 mmol/L; [H₂CO₃] rises from ~1.2 to ~1.5 mmol/L), then new pH = 6.10 + log(21/1.5) = 6.10 + log(14) = 6.10 + 1.15 = 7.25 — a measurable fall from 7.40. A direct pH meter measurement of arterial blood before and after intense exercise shows a decrease of ~0.05–0.15 pH units, demonstrating the buffer resists but does not prevent pH change.