Year 12 ChemistryModule 6, Acid/Base Reactions⏱ ~45 minLesson 1 of 19IQ1
Acid-Base Models, Arrhenius to Brønsted-Lowry
In 1923, Johannes Brønsted in Copenhagen and Thomas Lowry in London published the same proton-transfer theory of acids and bases simultaneously, in different journals, in different languages, 1,200 km apart, neither aware of the other. The single molecule that forced both of them to abandon Arrhenius's model was ammonia: a base with no OH⁻ ion at all.
🧪
Today's hook, In 1923, two chemists on opposite sides of Europe independently published the same new acid-base theory on the same day, without ever contacting each other. The molecule that forced both of them to rewrite chemistry? Ammonia: a base with no OH⁻ ion at all.
0/5QUESTS
1
You’re here
Get oriented
Worksheets
Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.
Ammonia (NH₃) is one of the most important industrial chemicals on Earth, used to make the fertilisers that feed roughly half the global population. But here's the problem: dissolve ammonia in water and the solution is unmistakably basic. Litmus turns blue. pH climbs above 7. Every test confirms it is a base. According to one of the most famous acid-base theories in the history of chemistry, that should be impossible because ammonia contains no hydroxide ions whatsoever.
Question 1: Before reading on, write down your best guess, what do you think a substance needs to contain or do in order to make a solution basic?
Question 2: The same problem applies to this gas-phase reaction: HCl(g) + NH₃(g) → NH₄Cl(s). No water, no ions in solution, yet chemists call this an acid-base reaction. How do you think that's possible?
Hold your answers. You will come back to test them, and almost certainly revise them, at the end of this lesson.
→
📐 Key Relationships
Arrhenius acid → produces H⁺ in aqueous solution
e.g. HCl → H⁺ + Cl⁻ (in water)Limitation: aqueous solutions only; cannot explain NH₃ as a base
Arrhenius base → produces OH⁻ in aqueous solution
e.g. NaOH → Na⁺ + OH⁻ (in water)Limitation: cannot classify NH₃, it has no OH⁻
Brønsted-Lowry acid → proton (H⁺) donor in any reaction
Works in aqueous AND non-aqueous contextse.g. HCl + H₂O → H₃O⁺ + Cl⁻ (HCl donates H⁺ to H₂O)
Brønsted-Lowry base → proton (H⁺) acceptor in any reaction
e.g. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (NH₃ accepts H⁺ from H₂O)
Conjugate base = acid − H⁺ | Conjugate acid = base + H⁺
Pairs always on opposite sides of the equation, differ by exactly one H⁺
H⁺(aq) + H₂O(l) → H₃O⁺(aq) (hydronium ion)
Bare H⁺ cannot exist in water, instantly bonds to H₂OH⁺(aq) and H₃O⁺(aq) represent the same species
No calculation formulas this lesson, models and definitions are conceptual.
Learning Intentions
Know
Key facts
The four acid-base models: Lavoisier, Davy, Arrhenius, Brønsted-Lowry
Each model's definition of an acid and base
The specific limitation that caused each model to be replaced
What H₃O⁺ (hydronium ion) is and why we write it
Understand
Concepts
Why ammonia is a base in Brønsted-Lowry but cannot be classified by Arrhenius
Why the gas-phase reaction HCl + NH₃ → NH₄Cl is an acid-base reaction under Brønsted-Lowry
The inverse strength relationship between conjugate pairs
When to write H⁺ vs H₃O⁺ and why it matters in extended responses
Can do
Skills
Identify the Brønsted-Lowry acid and base in any given equation
Identify both conjugate acid-base pairs from a full ionic equation
Evaluate a student claim about acid-base models using specific examples
Write balanced ionic equations showing proton transfer with correct arrow notation
Scan these before reading
Arrhenius acid
A substance that produces H⁺ (hydrogen) ions when dissolved in water.
Arrhenius base
A substance that produces OH⁻ (hydroxide) ions when dissolved in water.
Brønsted-Lowry acid
A proton (H⁺) donor in an acid-base reaction.
Brønsted-Lowry base
A proton (H⁺) acceptor in an acid-base reaction.
Conjugate acid-base pair
Two species that differ by exactly one proton (H⁺), interconverting in an acid-base equilibrium.
Amphiprotic substance
A substance that can act as both a proton donor and a proton acceptor (e.g., water, HCO₃⁻).
2
Why the models kept changing
Core Content
Cross-lesson links: The Brønsted-Lowry framework introduced here underpins every calculation in this module, Ka and Kb (L05, L09), Kw and pH (L07–L08), buffer action (L12–L13), and titration curve interpretation (L14–L17). Return here whenever a calculation asks you to identify the proton donor or acceptor.
1
A History of Wrong Models, Why Acid-Base Theory Kept Changing
Each model replaced by the one anomaly it could not explain
Every scientific model is replaced not when scientists grow bored of it, but when a single observation appears that it cannot explain, the history of acid-base theory shows this process operating three times in just over a century.
Late 1700sLavoisierAcids contain oxygen
→
1810DavyAcids contain hydrogen
→
1884ArrheniusH⁺(aq) / OH⁻(aq), aqueous only
→
1923Brønsted-LowryProton donor / acceptor, any medium
Each step above was forced by one observation the previous model could not explain. Read the chain as a sequence of anomalies, not a list of names.
Lavoisier → Davy
Lavoisier studied sulfuric (H₂SO₄), nitric (HNO₃) and phosphoric (H₃PO₄) acid, so "all acids contain oxygen" was a reasonable induction. The anomaly: HCl is strongly acidic and contains no oxygen at all. Davy's revision, in 1810, was that the essential ingredient is hydrogen.
Davy → Arrhenius
Davy's model still had no definition for bases, and could not say which hydrogen compounds are acidic. The anomaly: CH₄ contains hydrogen but is not acidic. In 1884 Arrhenius gave the first quantitative model: acids produce H⁺(aq), bases produce OH⁻(aq).
Arrhenius → Brønsted-Lowry
Arrhenius worked beautifully for laboratory acids and alkali metal hydroxides, and underpinned the first quantitative neutralisation calculations. The limit was structural: it was locked to aqueous solution and to bases that contain OH⁻.
What broke it
Three cases at once: ammonia is basic with no OH⁻; gaseous acid-base reactions occur with no solution at all; and carbonate and bicarbonate ions are basic without containing OH⁻.
In 1923, Johannes Brønsted and Thomas Lowry independently proposed the model that resolves every one of those cases: an acid is a proton donor and a base is a proton acceptor, in any reaction and in any medium.
Remember!
A model is replaced by the anomaly it cannot explain, not by being "wrong". Naming the anomaly is what earns the mark.
HSC Tip
In any extended response comparing or evaluating acid-base models, you must state each model's definition AND its specific limitation. Writing only the definition earns partial marks. The limitation is worth as many marks as the definition in most marking guidelines.
Common Error
Students say Arrhenius was "wrong" or "disproved." He was not, his model correctly describes every strong acid and strong base in aqueous solution and is still used for basic pH calculations. The correct language is: his model is "limited to aqueous systems" or "cannot account for bases that do not contain OH⁻." A model that is incomplete is not the same as a model that is incorrect.
Insight
The gaseous reaction HCl(g) + NH₃(g) → NH₄Cl(s) is a complete acid-base reaction producing a salt, yet it occurs with no water, no solvent, and no ions in solution at any stage. Arrhenius cannot classify either reactant in this reaction. Brønsted-Lowry handles it without modification: HCl donates a proton to NH₃, producing NH₄⁺ and Cl⁻. The model works because it describes what actually happens, a proton transfer, rather than what ions appear in solution.
Acid-Base Model Timeline, each model replaced by the one anomaly it could not explain
Acid-base models evolved through four stages, Lavoisier (acids contain O), Davy (acids contain H), Arrhenius (acid → H⁺(aq); base → OH⁻(aq), aqueous only), each replaced by one anomaly it couldn't explain; the Brønsted-Lowry model (acid = H⁺ donor; base = H⁺ acceptor) works in any medium and correctly classifies NH₃ as a base.
Pause, copy the highlighted definition into your book before moving on.
Which acid-base model was overturned because it could not explain why ammonia (NH₃) is a base?
3
Proton transfer, Brønsted-Lowry
2
Brønsted-Lowry Acids and Bases, Proton Transfer at the Molecular Level
Model before formula, picture the proton moving before writing any equation
We just saw that the Brønsted-Lowry model defines acids and bases by proton donation/acceptance in any medium. That raises a question: What does a proton transfer actually look like at the molecular level? This card answers it → by tracing exactly which H⁺ moves, from which atom, to which partner, in real equations.
Before writing any equation, picture what is physically happening: a proton, a bare hydrogen nucleus, detaches from one species and bonds to another, and that single event is the entire definition of a Brønsted-Lowry acid-base reaction.
In the Brønsted-Lowry model, every acid-base reaction is a proton transfer event. The species that releases the proton is the acid; the species that receives it is the base.
These roles are defined by what happens in that specific reaction, not by the identity of the substance in isolation. The same molecule can be an acid in one reaction and a base in another.
Consider HCl dissolving in water: HCl donates H⁺ to H₂O, making HCl the acid and H₂O the base. The product H₃O⁺ forms because water has accepted the proton. Since HCl ionises completely, we write a single forward arrow (→):
HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
H⁺ vs H₃O⁺
Use this rule consistently. In a molecular or mechanistic explanation (showing proton transfer and conjugate pairs), write H₃O⁺ and include water as the proton acceptor, e.g. HCl + H₂O → H₃O⁺ + Cl⁻, because a bare H⁺ does not exist alone in water. In a pH or equilibrium calculation, the shorthand H⁺ is acceptable (pH = −log[H⁺]); [H⁺] and [H₃O⁺] refer to the same species. Markers expect H₃O⁺ wherever you are explaining the Brønsted-Lowry mechanism.
Now consider ammonia dissolving in water. Here the roles swap: water donates the proton, so water is the acid and NH₃ is the base.
Where the OH⁻ comes from
The OH⁻ is not released by NH₃. It is what remains of the water molecule after water donated its proton. Ammonia contains no OH⁻ at all.
Why the arrow is ⇌
NH₃ only partially ionises, so the reaction sits at equilibrium. Compare HCl above, which ionises completely and takes a single forward arrow.
This is exactly why ammonia defeated the Arrhenius model: it is a base because it accepts a proton, not because it contains OH⁻.
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Remember!
Across the two reactions above, water acted as a base with HCl and as an acid with NH₃. That dual behaviour is what amphiprotic means, and it underpins pH, Kw and buffers in later lessons.
HCl + H₂O → H₃O⁺ + Cl⁻
What the acid does:HCl loses H⁺ → Cl⁻
Acid:HCl
What the base does:H₂O gains H⁺ → H₃O⁺
Base:H₂O
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
What the acid does:H₂O loses H⁺ → OH⁻
Acid:H₂O
What the base does:NH₃ gains H⁺ → NH₄⁺
Base:NH₃
CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
What the acid does:CH₃COOH loses H⁺
Acid:CH₃COOH
What the base does:H₂O gains H⁺ → H₃O⁺
Base:H₂O
HSC Tip
When asked to identify a Brønsted-Lowry acid and base, always explicitly state which proton is transferred, which species it moves FROM, and which species it moves TO. "HCl is the acid" is incomplete. "HCl is the Brønsted-Lowry acid because it donates a proton (H⁺) to water, forming H₃O⁺" is complete and earns full marks.
Common Error
"Ammonia is a base because it produces OH⁻ in solution." This is the Arrhenius explanation and it is incomplete, worse, it implies NH₃ contains OH⁻, which it does not. The Brønsted-Lowry explanation is: NH₃ is a base because it accepts a proton from water. The OH⁻ produced is the remains of the water molecule after it donated its proton to NH₃.
Brønsted-Lowry proton transfer: HCl donates H⁺ and water accepts it, forming H₃O⁺ and Cl⁻.
A Brønsted-Lowry acid donates H⁺ and a base accepts H⁺ in any reaction and any medium: HCl + H₂O → H₃O⁺ + Cl⁻ (strong, single arrow); NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (weak, equilibrium arrow), water is amphiprotic, acting as base with HCl and as acid with NH₃.
Add the highlighted point to your notes before the check below.
In the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, which species is the Brønsted-Lowry acid?
4
Conjugate acid-base pairs
3
Conjugate Acid-Base Pairs, Identifying and Using Them
Differ by exactly one H⁺ · Always on opposite sides of the equation
We just saw that every Brønsted-Lowry reaction involves a proton transfer from an acid to a base. That raises a question: What happens to those species after the proton moves, do they just disappear? This card answers it → the donor becomes its conjugate base and the acceptor becomes its conjugate acid, forming pairs linked by exactly one H⁺.
Every Brønsted-Lowry acid-base reaction produces a new acid and a new base on the product side, and these conjugate species are always related to their parent by exactly one proton, nothing more.
Acid → conjugate base
An acid donates a proton. What remains has one fewer H and one more negative charge. HA → A⁻ + H⁺
Base → conjugate acid
A base accepts a proton. What forms has one more H and one less negative charge. B + H⁺ → BH⁺
A conjugate acid-base pair consists of two species on opposite sides of the equation that differ by exactly one proton.
Consider the full equation: CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
Pair 1: CH₃COOH (acid, left) and CH₃COO⁻ (conjugate base, right), differ by one H⁺ ✓
Pair 2: H₂O (base, left) and H₃O⁺ (conjugate acid, right), differ by one H⁺ ✓
A crucial strength relationship flows from this pairing, and it runs inversely.
Strong acid → very weak conjugate base
If an acid donates its proton very readily, what remains has almost no tendency to accept it back.
Weak acid → stronger conjugate base
If the acid holds its proton, the conjugate base retains a real tendency to take it back, which is why the equilibrium does not run to completion.
Remember!
This inverse relationship governs which direction an acid-base equilibrium favours, and it underpins buffer theory in L13.
Reaction
Acid (left)
Conjugate base (right)
Base (left)
Conjugate acid (right)
HCl + H₂O → H₃O⁺ + Cl⁻
HCl
Cl⁻
H₂O
H₃O⁺
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
H₂O
OH⁻
NH₃
NH₄⁺
CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
CH₃COOH
CH₃COO⁻
H₂O
H₃O⁺
HCO₃⁻ + H₂O ⇌ H₂CO₃ + OH⁻
H₂O
OH⁻
HCO₃⁻
H₂CO₃
HPO₄²⁻ + H₂O ⇌ H₂PO₄⁻ + OH⁻
H₂O
OH⁻
HPO₄²⁻
H₂PO₄⁻
Method
To find a conjugate base: remove one H⁺ and add one negative charge. To find a conjugate acid: add one H⁺ and remove one negative charge. Check your answer by verifying the formula differs by exactly H⁺ and that the species appear on opposite sides of the equation.
Common Error
Students list two species on the SAME side of the equation as a conjugate pair, for example, calling CH₃COOH and H₂O a conjugate pair because both appear on the left. Conjugate pairs are always one species from the left AND one from the right. Another common error: identifying SO₄²⁻ as the conjugate base of H₂SO₄. This is wrong, the conjugate base of H₂SO₄ is HSO₄⁻ (remove only one proton at a time).
Insight
The strength relationship between conjugate pairs has a direct consequence for equilibrium direction. An acid-base reaction always proceeds to favour the weaker acid and weaker base, the equilibrium lies on the side of the weaker pair. HCl + H₂O → H₃O⁺ + Cl⁻ proceeds to completion (→) because HCl is a far stronger acid than H₃O⁺, and H₂O is a far stronger base than Cl⁻.
A conjugate base is formed by removing one H⁺ from the acid (one fewer H, one more negative charge); a conjugate acid by adding one H⁺ to the base, conjugate pairs always appear on opposite sides of the equation, differ by exactly one H⁺, and have inversely related strengths (strong acid → very weak conjugate base).
Pause, write the highlighted definition into your book.
In the reaction HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O, what is the conjugate base of HCO₃⁻?
5
H₃O⁺ and the limits of the model
4
H⁺ in Water, Why We Write H₃O⁺ and When It Matters
The hydronium ion · high-mark molecular-level explanations require H₃O⁺
We just saw that conjugate pairs differ by exactly one H⁺, which means protons move between species. That raises a question: If H⁺ is being transferred, can a bare proton actually exist in solution? This card answers it → a free H⁺ has a lifetime of less than 10⁻¹³ s in water; it instantly bonds to H₂O to form H₃O⁺, which is why notation matters in extended responses.
Writing H⁺(aq) is a chemist's shorthand that hides a physical reality, a bare proton cannot exist in water for even a fraction of a nanosecond before bonding to a water molecule, and understanding this is essential for any molecular-level explanation.
A hydrogen ion, H⁺, is simply a proton: a hydrogen atom stripped of its single electron, leaving only the nucleus. That gives it an extraordinarily high charge density, making it one of the most reactive species in chemistry.
So in water it cannot stay free. A bare H⁺ immediately forms a coordinate covalent bond with a lone pair on a water molecule, producing the hydronium ion:
H⁺(aq) + H₂O(l) → H₃O⁺(aq)
Remember!
A free H⁺ in water lasts under 10⁻¹³ seconds, so it effectively does not exist on its own. That is why Brønsted-Lowry equations show H₃O⁺ as the product, not H⁺ alone.
In HSC contexts, H⁺(aq) is an accepted shorthand used extensively in calculations, and it represents the same species as H₃O⁺(aq).
However, when a question asks what happens at the molecular or ionic level as an acid dissolves in water, writing H₃O⁺ and explaining the proton transfer to water is required for high-mark responses.
Acceptable notation
H⁺(aq) or H₃O⁺(aq)
H₃O⁺ preferred in Brønsted-Lowry equations
H₃O⁺ required in molecular-level extended responses
H⁺ acceptable in pH calculation contexts
Why
Both represent the same species; H⁺ is shorthand
Shows water acting as base (proton acceptor)
Demonstrates understanding of proton transfer to water
Arrhenius model predates H₃O⁺ notation
HSC Tip
In extended response questions asking you to describe what happens "at the molecular level" or "using the Brønsted-Lowry model" when an acid dissolves in water, you must explicitly state that H⁺ bonds to H₂O to form H₃O⁺ and identify water as the proton acceptor (Brønsted-Lowry base). Writing only "HCl → H⁺ + Cl⁻" omits the role of water and will not earn full marks.
Common Error
Students write HCl → H⁺ + Cl⁻ in a Brønsted-Lowry context. This equation shows dissociation in isolation, it does not show the proton transfer to water that the Brønsted-Lowry model requires. The correct Brønsted-Lowry equation is HCl + H₂O → H₃O⁺ + Cl⁻, which shows HCl donating H⁺ to H₂O (the base), producing H₃O⁺ (the conjugate acid of water) and Cl⁻ (the conjugate base of HCl).
Arrhenius describes acids and bases by the ions they produce in water. Brønsted-Lowry explains acid-base behaviour by proton transfer, so it can handle NH₃ and non-aqueous reactions.
H⁺(aq) and H₃O⁺(aq) represent the same species, a bare H⁺ bonds instantly to H₂O (lifetime <10⁻¹³ s) to form the hydronium ion H₃O⁺; either notation is acceptable in pH calculations, but H₃O⁺ is required in Brønsted-Lowry equations and any molecular-level extended response.
Pause, write the highlighted principle into your book.
H⁺(aq) and H₃O⁺(aq) represent different chemical species with different properties.
Ammonia (NH₃) is the chemical that broke Arrhenius's model, and understanding why it is basic is not just a theoretical curiosity. The Haber process synthesises NH₃ from N₂ and H₂ at high temperature and pressure. This ammonia is then reacted with sulfuric acid in a Brønsted-Lowry acid-base reaction: NH₃ + H₂SO₄ → (NH₄)₂SO₄ (ammonium sulfate, a major nitrogen fertiliser). Without Brønsted-Lowry's insight that NH₃ accepts protons (rather than producing OH⁻), chemists would have struggled to explain and optimise the reactions that produce the fertilisers feeding approximately half the global population today.
The gas-phase reaction HCl(g) + NH₃(g) → NH₄Cl(s), white solid ammonium chloride forming as a smoke when the gases mix, was the definitive experimental proof that acid-base reactions do not require water. It was observations like this that made Arrhenius's aqueous-only model untenable and drove Brønsted and Lowry to develop their proton-transfer framework.
Beyond the syllabus. The Arrhenius and Brønsted–Lowry models, their limitations and how the model changed over time are core. Lewis acid–base theory and BF3 chemistry go beyond the syllabus — treat them as a preview of university chemistry, not exam material.
!
Misconceptions to Fix
✗ "Arrhenius was proved wrong."
✓ Arrhenius was not wrong, he was incomplete. His model correctly predicts the behaviour of strong acids and bases in aqueous solution and is still used in pH calculations. The correct language is "limited to aqueous systems" or "cannot account for bases without OH⁻."
✗ "NH₃ is a base because it contains OH⁻."
✓ NH₃ contains no OH⁻ whatsoever. It is a base because it accepts a proton (H⁺) from water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. The OH⁻ in solution comes from the water molecule that donated its proton to NH₃, not from NH₃ itself.
✓ In the Brønsted-Lowry model, you must show the proton transfer to water: HCl + H₂O → H₃O⁺ + Cl⁻. Writing HCl → H⁺ + Cl⁻ in a Brønsted-Lowry context omits water's role as the proton acceptor (base) and will not earn full marks.
✗ "SO₄²⁻ is the conjugate base of H₂SO₄."
✓ Remove only one proton at a time. The conjugate base of H₂SO₄ is HSO₄⁻ (H₂SO₄ − H⁺). SO₄²⁻ is the conjugate base of HSO₄⁻. A conjugate pair always differs by exactly one H⁺.
✗ "Conjugate pairs are two species on the same side of the equation."
✓ Conjugate pairs are always one species from the left AND one from the right, related by the loss or gain of exactly one H⁺. Species on the same side are not conjugate pairs.
6
Worked examples and activities
Worked Examples
Worked Example 1, Identifying conjugate pairs and roles in a full equation
For the reaction HPO₄²⁻ + H₂O ⇌ H₂PO₄⁻ + OH⁻, identify (a) the Brønsted-Lowry acid and base on the left side; (b) the conjugate acid and conjugate base on the right side; (c) both conjugate pairs.
1
(a) Identify which species donates a proton. Compare left and right: H₂O (left) → OH⁻ (right). H₂O has lost one H⁺ to become OH⁻. Therefore H₂O is the Brønsted-Lowry acid. HPO₄²⁻ (left) → H₂PO₄⁻ (right). HPO₄²⁻ has gained one H⁺ to become H₂PO₄⁻. Therefore HPO₄²⁻ is the Brønsted-Lowry base.
2
(b) Identify conjugate species on the right. Conjugate acid = the base after it gains a proton = H₂PO₄⁻ (HPO₄²⁻ + H⁺). Conjugate base = the acid after it loses a proton = OH⁻ (H₂O − H⁺).
3
(c) State both conjugate pairs explicitly. Conjugate pair 1, H₂O (acid, left) and OH⁻ (conjugate base, right): differ by one H⁺, on opposite sides ✓. Conjugate pair 2, HPO₄²⁻ (base, left) and H₂PO₄⁻ (conjugate acid, right): differ by one H⁺, on opposite sides ✓.
Worked Example 2, Evaluating model limitations with a real reaction
(a) Explain why the Arrhenius model cannot describe HCl(g) + NH₃(g) → NH₄Cl(s). (b) Using Brønsted-Lowry, identify the acid, base, conjugate acid, and conjugate base, and write the proton-transfer equation. (c) Explain at the molecular level why NH₃ acts as a base in this reaction.
1
(a) Why Arrhenius fails: The Arrhenius model defines an acid as a substance that produces H⁺ in aqueous solution, and a base as a substance that produces OH⁻ in aqueous solution. This reaction occurs entirely in the gas phase, there is no solvent, no aqueous solution, and no ions produced in water. Furthermore, NH₃ contains no OH⁻, it cannot satisfy the Arrhenius definition of a base under any circumstances. The Arrhenius model is therefore completely inapplicable to this reaction.
2
(b) Brønsted-Lowry analysis: Identify the proton transfer: HCl → Cl⁻ (HCl loses H⁺) → HCl is the acid; Cl⁻ is the conjugate base. NH₃ → NH₄⁺ (NH₃ gains H⁺) → NH₃ is the base; NH₄⁺ is the conjugate acid.
HCl(g) + NH₃(g) → NH₄⁺(aq) + Cl⁻(aq) → NH₄Cl(s)
3
(c) Molecular-level explanation: The nitrogen atom in NH₃ has a lone pair of electrons. When HCl and NH₃ molecules collide with sufficient energy in the gas phase, the lone pair on N forms a coordinate covalent bond with the H of HCl, the proton transfers from Cl to N. This leaves Cl with the bonding electrons as Cl⁻, and produces NH₄⁺. NH₃ acts as a Brønsted-Lowry base because it accepts the proton via its lone pair, no OH⁻ is involved at any stage.
Answer: (a) Arrhenius requires aqueous solution and OH⁻ for bases, neither applies here. (b) Acid = HCl; conjugate base = Cl⁻; base = NH₃; conjugate acid = NH₄⁺. Equation: HCl(g) + NH₃(g) → NH₄Cl(s). (c) NH₃ accepts H⁺ via its nitrogen lone pair, Brønsted-Lowry base by proton acceptance, no OH⁻ involved.
Worked Example 3, High-Mark Extended Response: Model Comparison
(6 marks) A chemistry student states: "The Arrhenius model and the Brønsted-Lowry model are essentially the same, they both just describe acids producing H⁺ and bases producing OH⁻ in water." Evaluate this statement using at least two specific chemical examples.
1
Identify what the student got partially right: For strong acids and bases in aqueous solution, both models give identical predictions. HCl in water, Arrhenius: produces H⁺; Brønsted-Lowry: donates H⁺ to water producing H₃O⁺. Both correctly identify HCl as an acid. NaOH in water, Arrhenius: produces OH⁻; Brønsted-Lowry: OH⁻ is the base (it accepts H⁺). The student is not entirely wrong for this class of substances.
2
Where the models diverge, Example 1 (NH₃): Arrhenius cannot classify NH₃ as a base, it contains no OH⁻ and produces none directly. The Arrhenius model would predict NH₃ is not a base, contradicting experimental evidence. Brønsted-Lowry correctly identifies NH₃ as a base: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, where NH₃ accepts H⁺ from water. The OH⁻ comes from water, not from NH₃. The two models make different predictions, Arrhenius fails here.
3
Where the models diverge, Example 2 (gas phase): HCl(g) + NH₃(g) → NH₄Cl(s) is a complete acid-base reaction producing a salt. Arrhenius cannot describe it at all, no aqueous solution exists. Brønsted-Lowry describes it straightforwardly: HCl donates H⁺ to NH₃ in the gas phase. The models are NOT equivalent in scope, Brønsted-Lowry is a broader framework.
4
Evaluate the key conceptual difference: Arrhenius defines acids and bases by what ions they produce in water, a macroscopic, product-based definition. Brønsted-Lowry defines acids and bases by what they do, proton transfer, a mechanistic, process-based definition that applies in any medium. Brønsted-Lowry also explains WHY bases make solutions basic: by accepting H⁺ from water, they leave OH⁻ behind. Both models also have their own limitations, Brønsted-Lowry cannot describe Lewis acid-base reactions (e.g. BF₃ + F⁻ → BF₄⁻) that involve no proton transfer.
Answer: The statement is an oversimplification. For aqueous strong acid-base reactions both models agree, but they diverge for NH₃ (Arrhenius fails; Brønsted-Lowry correctly identifies NH₃ as a proton acceptor) and for non-aqueous reactions (Arrhenius cannot apply; Brønsted-Lowry applies universally). The models differ in scope, in their definition of bases, and in mechanistic depth. Both retain limitations beyond which a further model (Lewis) is required. 6 marks: 1 mark for agreement in aqueous strong acids/bases; 2 marks for NH₃ divergence with equation; 2 marks for gas-phase reaction divergence; 1 mark for evaluating mechanistic difference.
Using → for a weak acid is a critical error, check every equation
🔬Predict, Then Reveal+8 XP
Ammonia (NH₃) dissolves in water and the solution turns litmus blue. According to the Arrhenius model, this should be impossible. Before revealing, explain WHY the Arrhenius model fails here.
Your predictionExpert answerCompare
The Arrhenius model defines a base as a substance that produces OH⁻ ions in aqueous solution. NH₃ contains no OH⁻ and releases none directly, it only produces OH⁻ indirectly when it accepts H⁺ from water. The Arrhenius model cannot classify any base that doesn't contain OH⁻, making it structurally incapable of explaining NH₃. The Brønsted-Lowry model resolves this by defining a base as any proton acceptor, NH₃ fits that definition perfectly.
Good comparison. Keep the reasoning that matched the expert answer.
Re-read the expert answer and identify the missing reasoning step.
Interactive · Acid/Base Theory Matcher
Use the matcher. Ammonia acts as a base by accepting a proton, and the molecule itself contains no OH− to release. Which definitions classify it as a base?
Complete the Learn phase to unlock Practice.
✓
Practice questions
✏️ Activities
Activity 1, Classifying Reactions Under Each Model
For each reaction below: (A) State whether Arrhenius can classify it as an acid-base reaction. (B) State whether Brønsted-Lowry can classify it. (C) Identify the Brønsted-Lowry acid and base where applicable.
Reaction
Arrhenius can classify?
BL can classify?
BL acid / BL base
HNO₃(aq) + H₂O(l) → H₃O⁺(aq) + NO₃⁻(aq)
Your answer
Your answer
Your answer
NH₃(g) + HCl(g) → NH₄Cl(s)
Your answer
Your answer
Your answer
NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)
Your answer
Your answer
Your answer
CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq)
Your answer
Your answer
Your answer
BF₃ + F⁻ → BF₄⁻
Your answer
Your answer
Your answer
Check Your Understanding
01
Multiple Choice
+5 XP
A fresh set drawn from this lesson's question bank, with feedback shown immediately. +5 XP per correct · +25 XP all correct
Short Answer Questions
Q6. (5 marks) (a) State the Arrhenius definitions of an acid and a base. (b) Explain why the Arrhenius model cannot classify ammonia (NH₃) as a base. (c) Write the Brønsted-Lowry equation that correctly explains why NH₃ produces a basic solution in water. Identify the acid and base in your equation.
Q7. (4 marks) For the reaction: H₂PO₄⁻(aq) + CO₃²⁻(aq) ⇌ HPO₄²⁻(aq) + HCO₃⁻(aq)
(a) Identify both conjugate acid-base pairs. (b) Identify the Brønsted-Lowry acid and base on the left side. (c) Verify that your identified conjugate pairs each differ by exactly one H⁺ and appear on opposite sides of the equation.
Q8. (5 marks) Real-World Application: The industrial synthesis of ammonium sulfate fertiliser [(NH₄)₂SO₄] involves reacting ammonia gas with sulfuric acid: 2NH₃(g) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)
(a) Using the Brønsted-Lowry model, explain why this is classified as an acid-base reaction. Identify the proton donor and proton acceptor. (2 marks) (b) A student says this reaction cannot be described by Arrhenius. Is this correct? Explain your reasoning with reference to the specific definitions involved. (3 marks)
Row 5 (BF₃ + F⁻): Arrhenius: No, no H⁺ or OH⁻ involved. BL: No, no proton transfer; F⁻ donates an electron pair to BF₃ (Lewis acid-base, not Brønsted-Lowry). This reaction requires the Lewis model.
Q6 Model Answer (5 marks)
(a) Arrhenius acid: a substance that produces H⁺ ions in aqueous solution [1]. Arrhenius base: a substance that produces OH⁻ ions in aqueous solution [1]. (b) Arrhenius cannot classify NH₃ as a base because NH₃ contains no OH⁻ ions and does not produce OH⁻ directly, it has the formula NH₃ with no oxygen present [1]. (c) NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) [1]. Acid = H₂O (donates H⁺ to NH₃); Base = NH₃ (accepts H⁺ from H₂O). The OH⁻ in solution comes from the water molecule that donated its proton, not from NH₃ [1].
Q7 Model Answer (4 marks)
(a) Pair 1: H₂PO₄⁻ (acid, left) and HPO₄²⁻ (conjugate base, right) [1]. Pair 2: CO₃²⁻ (base, left) and HCO₃⁻ (conjugate acid, right) [1]. (b) BL acid on left = H₂PO₄⁻ (loses H⁺ → HPO₄²⁻); BL base on left = CO₃²⁻ (gains H⁺ → HCO₃⁻) [1]. (c) Pair 1: H₂PO₄⁻ and HPO₄²⁻ differ by one H and one negative charge = one H⁺ ✓. On opposite sides ✓. Pair 2: CO₃²⁻ / HCO₃⁻, differ by one H⁺; on opposite sides ✓ [1].
Q8 Model Answer (5 marks)
(a) This is a Brønsted-Lowry acid-base reaction because a proton (H⁺) is transferred from H₂SO₄ to NH₃ [1]. Proton donor (acid) = H₂SO₄; proton acceptor (base) = NH₃ [1]. (b) Partially correct [1]. Arrhenius defines an acid as producing H⁺ in aqueous solution, H₂SO₄ satisfies this [1]. However, Arrhenius defines a base as producing OH⁻ in aqueous solution. NH₃ contains no OH⁻ and cannot be classified as an Arrhenius base. So the Arrhenius model CAN classify H₂SO₄ as an acid but CANNOT classify NH₃ as a base, making the Arrhenius framework incomplete (not totally inapplicable) for this reaction [1].
✓
Retrieve and reflect
Check what actually stuck
Review
What is the Brønsted-Lowry definition of an acid?
A proton (H⁺) donor in any reaction, in any medium.
Why can't the Arrhenius model classify NH₃ as a base?
NH₃ contains no OH⁻ ions. The Arrhenius model requires a base to produce OH⁻ in aqueous solution, NH₃ does not do this directly.
What are conjugate acid-base pairs and how do they differ?
Two species on opposite sides of an equation that differ by exactly one proton (H⁺). Acid loses H⁺ → conjugate base. Base gains H⁺ → conjugate acid.
When must you write H₃O⁺ rather than H⁺ in a chemistry response?
In any molecular-level explanation using the Brønsted-Lowry model, you must write H₃O⁺ to show that H⁺ bonds to H₂O. In Brønsted-Lowry equations, H₃O⁺ is preferred. In pH calculations, either is acceptable.
What is the fatal limitation of the Arrhenius model?
It is restricted to aqueous solutions and requires bases to contain OH⁻. It cannot classify NH₃ as a base, cannot describe gas-phase reactions like HCl(g) + NH₃(g) → NH₄Cl(s), and cannot explain the basicity of carbonate ions.
Revisit Think First
Go back to your Think First response at the top of this lesson. In 1923, Brønsted and Lowry each solved the ammonia problem independently, NH₃ accepts a proton (H⁺) from an acid, acting as a base without ever producing OH⁻. Now that you've studied the Brønsted-Lowry model:
Q1: You were asked what a substance needs to contain or do to make a solution basic. Was your answer Arrhenius-style (contains OH⁻) or Brønsted-Lowry style (accepts a proton)? What is the more complete answer now?
Q2: You were asked how HCl(g) + NH₃(g) → NH₄Cl(s) can be an acid-base reaction. Can you now write the full Brønsted-Lowry explanation, identifying the proton donor and acceptor?
What was the single most surprising thing you learned in this lesson?
