Chemistry • Year 12 • Module 6 • Lesson 7

Conjugate Pairs, Amphiprotic Substances & Water's Role

Develop HSC Band 5–6 technique: synthesise data with theory, evaluate claims, and justify multi-step chemical reasoning.

Master · Extended Response — Band 5–6

1. Extended response — phosphate buffer in biological research (Band 5–6)

8 marks Band 5–6

Scenario. Biological researchers at a Sydney university use a phosphate-buffered saline (PBS) solution at pH 7.4 and 25°C. The buffer is prepared from a mixture of Na₂HPO₄ and KH₂PO₄. Both salts are listed explicitly in the NSW Chemistry Stage 6 Syllabus as examples of amphiprotic salts. The buffer must maintain pH 7.4 when small amounts of either acid or base are added to cell cultures.

Table 1.1. Relevant K_a values for the phosphoric acid system at 25°C.
Equilibrium	K_a	pK_a
H₃PO₄ ⇌ H⁺ + H₂PO₄⁻	7.5 × 10⁻³	2.12
H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻	6.2 × 10⁻⁸	7.21
HPO₄²⁻ ⇌ H⁺ + PO₄³⁻	4.8 × 10⁻¹³	12.32

Q1. Using the data and your understanding of amphiprotic species and K_w, respond to the following. In your response you must:

Identify the amphiprotic ion from each of the two NESA-named salts and explain why each ion is amphiprotic (with equations).
Explain which of the two amphiprotic ions predominates in its acid mode and which predominates in its base mode at pH 7.4, using K_a values as evidence.
Calculate [OH⁻] in the buffer at pH 7.4 and 25°C, and state whether [H₃O⁺] or [OH⁻] is greater.
Use the phosphate chain diagram concept (from Card 3 of the lesson) to justify why neither H₃PO₄ nor PO₄³⁻ are amphiprotic, but both H₂PO₄⁻ and HPO₄²⁻ are.

Plan first: start with the definition of amphiprotic, then write equations for each ion, then compare K_a3 to K_b (= K_w/K_a2) to determine which mode dominates at pH 7.4. Then do the K_w calculation.

2. Source critique — evaluate a media claim about ocean acidification (Band 5–6)

7 marks Band 5–6

Source. The following passage is adapted from a popular science article about Great Barrier Reef chemistry published in a Sydney newspaper:

“As the ocean absorbs more CO₂, the water becomes acidic. This happens because CO₂ is itself an acid that directly lowers the pH of seawater. Since HCO₃⁻ (bicarbonate) is also an acid — as anyone who has seen baking soda dissolve in vinegar knows — the increasing HCO₃⁻ levels caused by CO₂ absorption will make reef waters even more acidic over time. The neutral pH of seawater is always 7, so any measurement below 7 is acidic.”

Q2. This passage contains three distinct scientific errors. For each error:

Identify the specific claim that is incorrect.
Explain the correct chemistry using precise Module 6 terminology.
For at least one error, describe how a chemist would detect or confirm the correct science experimentally.

Look for: (1) CO₂ is NOT itself the acid that lowers pH — it forms H₂CO₃ which dissociates; (2) HCO₃⁻ is amphiprotic (not just an acid); increasing HCO₃⁻ doesn’t automatically acidify; (3) neutral pH = 7 is temperature-dependent, not universal.

Answers — Do not peek before attempting

Q1 — Marking criteria (8 marks)

Amphiprotic ions identified with equations [2 marks]:

Na₂HPO₄ → amphiprotic ion = HPO₄²⁻. Acting as acid: HPO₄²⁻(aq) ⇌ H⁺(aq) + PO₄³⁻(aq) (K_a3 = 4.8 × 10⁻¹³). Acting as base: HPO₄²⁻(aq) + H⁺(aq) ⇌ H₂PO₄⁻(aq) (reverse of K_a2).

KH₂PO₄ → amphiprotic ion = H₂PO₄⁻. Acting as acid: H₂PO₄⁻(aq) ⇌ H⁺(aq) + HPO₄²⁻(aq) (K_a2 = 6.2 × 10⁻⁸). Acting as base: H₂PO₄⁻(aq) + H⁺(aq) ⇌ H₃PO₄(aq) (reverse of K_a1).

Award 1 mark for each ion with both acid and base equations correct (max 2).

Dominant mode at pH 7.4 [2 marks]:

For HPO₄²⁻: K_a3 = 4.8 × 10⁻¹³ is extremely small; K_b(HPO₄²⁻ as base) = K_w/K_a2 = 1.0 × 10⁻¹⁴ / 6.2 × 10⁻⁸ = 1.6 × 10⁻⁷ ≫ K_a3. At pH 7.4, HPO₄²⁻ predominantly acts as a base. [1]

For H₂PO₄⁻: K_a2 = 6.2 × 10⁻⁸; K_b(H₂PO₄⁻ as base) = K_w/K_a1 = 1.0 × 10⁻¹⁴ / 7.5 × 10⁻³ = 1.3 × 10⁻¹² < K_a2. At pH 7.4, H₂PO₄⁻ predominantly acts as an acid. [1]

K_w calculation at pH 7.4 [2 marks]:

[H₃O⁺] = 10^−7.4 = 3.98 × 10⁻⁸ mol L⁻¹. [OH⁻] = K_w/[H₃O⁺] = 1.0 × 10⁻¹⁴ / 3.98 × 10⁻⁸ = 2.51 × 10⁻⁷ mol L⁻¹. [OH⁻] > [H₃O⁺] → buffer is basic at 25°C. [1 calculation; 1 comparison and classification]

Chain-navigation justification [2 marks]:

H₃PO₄ can only act as an acid (it is the start of the chain; it cannot accept a proton to move further left) → not amphiprotic [1]. PO₄³⁻ has no ionisable protons; it can only accept protons (base only; moves left to HPO₄²⁻) → not amphiprotic [1]. H₂PO₄⁻ and HPO₄²⁻ are intermediate species: each retains at least one ionisable proton (can move right = acid) and can accept a proton (can move left = base) → both amphiprotic.

Q2 — Marking criteria (7 marks)

Error 1 — “CO₂ is itself an acid” [2 marks]: CO₂ is not itself a Brønsted-Lowry acid. CO₂ dissolves in water to form carbonic acid (H₂CO₃), which then dissociates to give H⁺ + HCO₃⁻. The acidification arises from H₂CO₃ donating a proton, not from CO₂ directly. A chemist could confirm this by measuring [H₃O⁺] in pure CO₂(g) versus CO₂ dissolved in water: only the aqueous solution shows pH change because H₂CO₃ is the actual proton donor. [1 for identification; 1 for correct chemistry + experimental detection]

Error 2 — “HCO₃⁻ is also an acid” (incomplete claim) [2 marks]: HCO₃⁻ is amphiprotic — it can act as both an acid (donating H⁺ to produce CO₃²⁻) and as a base (accepting H⁺ to produce H₂CO₃). The article’s claim that increasing HCO₃⁻ will “make waters even more acidic” is wrong: HCO₃⁻ actually acts as a base to buffer against further acidification. The net effect of CO₂ absorption is that HCO₃⁻ increases while CO₃²⁻ decreases — pH falls, but the change is moderated by the buffering action of HCO₃⁻ itself. [1 identification; 1 correct chemistry]

Error 3 — “Neutral pH is always 7” [3 marks]: Neutral pH equals 7.00 only at 25°C. Neutrality is defined by [H₃O⁺] = [OH⁻], which gives pH = −log(√K_w). Because K_w is temperature-dependent and increases at higher temperatures (endothermic equilibrium), neutral pH is below 7 at elevated temperatures (e.g. 6.81 at 37°C, 6.51 at 60°C). Ocean surface temperature also affects K_w and therefore the neutral pH reference point. [1 identification; 1 correct definition of neutrality using [H₃O⁺] = [OH⁻]; 1 linking K_w to temperature dependence with supporting values]