Brønsted-Lowry acid-base theory explained: HSC Chemistry Module 5

A 4 mark answer needs the acid/base assignment, the conjugate pairs, and a clear demonstration that $HCO_3^-$ is amphiprotic.

In the forward direction:

• Water donates a proton to $HCO_3^-$, so $H_2O$ is the Brønsted-Lowry acid.
• The $HCO_3^-$ ion accepts a proton, so $HCO_3^-$ is the Brønsted-Lowry base.

In the reverse direction:

• The $H_2CO_3$ molecule donates a proton, so it is the conjugate acid of $HCO_3^-$.
• The $OH^-$ ion accepts a proton, so it is the conjugate base of $H_2O$.

Conjugate pairs: $HCO_3^- / H_2CO_3$ and $H_2O / OH^-$.

Amphiprotic means a species can act either as a Brønsted-Lowry acid (donating $H^+$) or as a Brønsted-Lowry base (accepting $H^+$). In the equation above, $HCO_3^-$ acts as a base. But $HCO_3^-$ can also donate a proton, for example $HCO_3^- + H_2O \rightleftharpoons CO_3^{2-} + H_3O^+$, where it acts as an acid. Because it can do both, $HCO_3^-$ is amphiprotic.

Markers reward (1) correct assignment in the forward and reverse directions, (2) explicit naming of the two conjugate pairs, (3) the definition of amphiprotic with two equations showing $HCO_3^-$ in both roles.

Arrhenius defined an acid as a substance that produces $H^+$ in aqueous solution and a base as a substance that produces $OH^-$. This definition is limited to aqueous solutions and cannot explain basic behaviour without hydroxide ions.

Brønsted-Lowry defines an acid as a proton donor and a base as a proton acceptor. This extends the theory in two ways:

1. It works in non-aqueous solvents (for example, $NH_3$ acting as a base toward $HCl$ in liquid ammonia).
2. It explains the basic behaviour of species like $NH_3$, $CO_3^{2-}$ and $HCO_3^-$ that contain no $OH^-$ but still accept protons in water.

Markers reward (1) clearly stating both definitions, (2) at least one specific extension Brønsted-Lowry accounts for that Arrhenius cannot.

A 3-mark identify needs the correct acid/base labels plus both conjugate pairs.

Forward direction. $H_2S$ donates a proton to become $HS^-$, so $H_2S$ is the Brønsted-Lowry acid. Water accepts that proton to become $H_3O^+$, so $H_2O$ is the Brønsted-Lowry base.

Conjugate pairs. $H_2S / HS^-$ (acid / conjugate base) and $H_3O^+ / H_2O$ (conjugate acid / base).

Marking criteria: 1 mark for correctly identifying the acid, 1 mark for correctly identifying the base, 1 mark for both conjugate pairs named correctly (differing by one $H^+$ each).

Acting as an acid (donates $H^+$ to water):

Acting as a base (accepts $H^+$ from water):

Because $H_2PO_4^-$ can both donate and accept a proton depending on its partner, it is amphiprotic, in the same way as $HCO_3^-$ and water itself.

Marking criteria: 1 mark for the correct acid-role equation, 1 mark for the correct base-role equation (both with state symbols), 1 mark for explicitly naming and explaining "amphiprotic".

Step 1: write the dissociation equation. A monoprotic acid donates one proton per molecule:

One mole of HA gives one mole of $H_3O^+$ when fully dissociated (a strong acid assumption).

Step 2: concentration of HA.

Step 3: $[H_3O^+]$. Full dissociation with a 1:1 mole ratio means:

Step 4: pH.

Step 5: round to 2 decimal places (matching the precision requested).

Marking criteria: 1 mark for the correct 1:1 dissociation relationship, 1 mark for the correct concentration of HA, 1 mark for correctly equating $[H_3O^+]$ to that concentration (full dissociation), 1 mark for the correct pH to 2 decimal places. Note this method assumes HA behaves as a strong (fully dissociating) Brønsted-Lowry acid, as stated in the question.

(a) Strong or weak. The curve shows a low, flat initial pH (around pH 1) before the equivalence point, and a very steep, near-vertical jump in pH exactly at the equivalence point (20.0 mL). A weak acid would start at a higher initial pH and show a much more gradual rise (a buffering region) before a smaller, less abrupt jump. The low starting pH plus the large steep jump indicate this is a strong acid.

(b) Proton-transfer reaction at equivalence. At the equivalence point, moles of $OH^-$ added exactly equal the initial moles of $H_3O^+$ present. In Brønsted-Lowry terms, hydroxide ion (behaving as a Brønsted-Lowry base, accepting a proton) reacts with the hydronium ion (behaving as a Brønsted-Lowry acid, donating a proton):

Because the conjugate base of a strong acid is negligibly basic, the resulting solution at equivalence is neutral, pH ≈ 7, matching the curve.

Marking criteria: (a) 1 mark for correctly identifying "strong", 1 mark for justification using BOTH the low starting pH and the steep/large jump. (b) 1 mark for identifying $H_3O^+$ as the Brønsted-Lowry acid and $OH^-$ as the base, 1 mark for the correct neutralisation equation, 1 mark for linking this to the neutral equivalence pH shown on the curve.

This is a 6-mark ASSESS: markers reward a judgement supported by specific evidence, not a simple description of both theories.

Band 6 PLAN.

• Thesis: Brønsted-Lowry did not make Arrhenius theory obsolete so much as generalise it; Arrhenius remains correct and useful within its aqueous domain, but Brønsted-Lowry is required to fully explain species like $NH_3$ and $HCO_3^-$.
• Evidence 1 (ammonia): Arrhenius cannot classify $NH_3$ as a base directly, since $NH_3$ contains no $OH^-$ to "produce"; only after reacting with water ($NH_{3(aq)} + H_2O_{(l)} \rightleftharpoons NH_4^+{}_{(aq)} + OH^-_{(aq)}$) does Arrhenius indirectly explain the resulting basicity. Brønsted-Lowry explains it directly: $NH_3$ accepts a proton from water, so it is a base regardless of solvent.
• Evidence 2 (bicarbonate, amphiprotic behaviour): $HCO_3^-$ acts as an acid toward $OH^-$ and as a base toward $H_3O^+$. Arrhenius theory has no mechanism at all for a single species switching roles; Brønsted-Lowry explains both roles with one consistent proton-transfer definition.
• Evidence 3 (non-aqueous systems): Brønsted-Lowry acid-base chemistry occurs in solvents such as liquid ammonia, where Arrhenius theory (defined only for aqueous solution) cannot apply at all.
• Judgement: for everyday aqueous strong-acid/strong-base chemistry, Arrhenius's simpler definition is still an accurate special case; therefore Brønsted-Lowry theory extends and subsumes Arrhenius rather than making it "obsolete" in the sense of incorrect.

Model paragraph (excerpt). The Brønsted-Lowry theory is best understood as a generalisation of Arrhenius theory rather than a replacement that renders it wrong. Arrhenius correctly predicts that $HCl$ is acidic and $NaOH$ is basic in water, and Brønsted-Lowry agrees, classifying $HCl$ as a proton donor and $NaOH$'s hydroxide as a proton acceptor. The real gap Arrhenius leaves is species like ammonia and bicarbonate: $NH_3$ has no hydroxide to "produce" yet is unambiguously basic, and $HCO_3^-$ behaves as an acid in one reaction and a base in another, a role-switching behaviour ("amphiprotic") that a fixed-ion-production model cannot describe at all. Because the proton-transfer definition covers every case the ion-production definition covers, plus these extra cases, and even extends beyond water to non-aqueous solvents like liquid ammonia, Brønsted-Lowry theory is more accurately described as generalising Arrhenius theory than as making it obsolete.

Marker's note: top-band answers (1) state an explicit judgement (not "both theories are useful" as a non-answer), (2) use at least two specific chemical examples with correct equations, (3) explicitly address the non-aqueous extension, and (4) conclude by characterising the RELATIONSHIP between the theories (generalisation, not obsolescence) rather than simply listing differences.