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Theories of Acids and Bases
Choose the definition that explains the reaction: aqueous ions, proton transfer or electron-pair donation.
A-Level 9476 (2026-2027)
The Arrhenius definition is about aqueous solution
Acids increase hydrogen-ion concentration; bases increase hydroxide-ion concentration.
An Arrhenius acid increases the concentration of hydrogen ions in water. An Arrhenius base increases the concentration of hydroxide ions. The H+(aq) notation is shorthand for a proton solvated by water; H3O+ is a useful explicit representation.
| Substance | What happens in water | Classification |
|---|---|---|
| HCl | HCl + H2O → H3O+ + Cl-: the water accepts a proton. | Acid |
| CH3COOH | CH3COOH + H2O ⇌ H3O+ + CH3COO-: only a fraction ionises. | Weak acid |
| NaOH | NaOH(s) → Na+(aq) + OH-(aq): dissolution releases hydroxide ions. | Base and soluble alkali |
| NH3 | NH3 + H2O ⇌ NH4+ + OH-: reaction with water generates hydroxide. | Aqueous basic behaviour, despite no OH group in its formula |
Neutralisation between a strong aqueous acid and a strong aqueous alkali has the net ionic equation H+(aq) + OH-(aq) → H2O(l). Counter-ions such as Na+ and Cl- remain in solution and do not belong in the net ionic equation.
This definition is convenient for aqueous behaviour but is not a general explanation of every acid-base reaction. A reaction between gaseous ammonia and hydrogen chloride can produce ammonium chloride without water being present. Following proton transfer explains it more broadly.
Check your understanding
Ammonia has no OH group. Why is it still an Arrhenius base in water?
Show the answer
It reacts with water to form OH- ions, raising [OH-]
NH3 + H2O <=> NH4+ + OH-. An Arrhenius base increases the hydroxide-ion concentration in water, however the OH- is produced.
The acid donates H+; the base accepts it
A conjugate pair differs by exactly one proton and one unit of charge.
A Bronsted-Lowry acid is a proton donor; a Bronsted-Lowry base is a proton acceptor. An acid becomes its conjugate base after losing H+. A base becomes its conjugate acid after gaining H+. The role belongs to the species in the particular reaction, not simply to a familiar name.
- Find the proton transfer
A hydrogen nucleus moves from water to ammonia.
- Identify the donor and acceptor
H2O is the acid; NH3 is the base.
- Pair each species with its proton-change partner
H2O/OH- is one pair; NH4+/NH3 is the other.
- Check the reverse reaction
NH4+ can donate a proton to OH-, regenerating NH3 and H2O.
Worked exampleConjugate charge follows proton transfer
Give the conjugate acid and conjugate base of HCO3-.
Try it first. Then show the worked stepsHide the worked steps
- Add H+ to obtain the conjugate acid: HCO3- + H+ → H2CO3.
- Remove H+ to obtain the conjugate base: HCO3- → CO32- + H+.
- Check charges: -1 + 1 = 0 in the first equation; -2 + 1 = -1 in the second.
Conjugate acid: H2CO3. Conjugate base: CO32-. Hydrogencarbonate can act as either a proton donor or acceptor.
A species able to donate or accept a proton is amphiprotic. Water accepts a proton from HCl but donates one to NH3. Hydrogencarbonate accepts H+ from an acid and donates H+ to a sufficiently strong base. Amphiprotic concerns proton transfer; amphoteric is the wider description of reacting as an acid or a base.
For a general equilibrium HA + B ⇌ A- + BH+, conjugate pairs are HA/A- and BH+/B. A stronger acid has a weaker conjugate base. The equilibrium tends towards the side containing the weaker acid and weaker base; numerical extent can be evaluated with the equilibrium constants introduced later.
Check your understanding
Are H2SO4 and SO42- a conjugate acid-base pair?
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No. They differ by two protons; H2SO4/HSO4- and HSO4-/SO42- are the pairs.
A conjugate pair differs by exactly one proton. H2SO4 loses one H+ to form HSO4-, which loses another to form SO42-.
The Lewis base supplies an electron pair
A Lewis acid accepts a pair; a Lewis base donates one.
A Lewis acid accepts an electron pair to form a covalent bond. A Lewis base donates a pair. This definition includes proton-transfer reactions because H+ can accept a lone pair, but it also explains reactions with no proton transfer.
Ammonia donates to boron trifluoride
A lone pair on the nitrogen of NH3 is directed towards the electron-deficient boron of BF3. The arrow starts at the electron pair and ends at boron. No hydrogen ion is transferred, and no water is required.
In BF3, boron has only three shared pairs in its simple bonding description and can accept another pair. Nitrogen in NH3 has a lone pair. Donation forms an N-B bond in the neutral adduct H3N→BF3. Both centres then have four bonding regions. The arrow records the origin of the bonding pair; the final bond is covalent.
| Reaction feature | Acid | Base |
|---|---|---|
| H+ accepts a lone pair from water. | H+ is a Lewis acid. | H2O is a Lewis base. |
| NH3 transfers a proton from water to itself. | H2O is the Bronsted-Lowry donor. | NH3 is the proton acceptor and uses its lone pair. |
| BF3 accepts the nitrogen pair from NH3. | BF3 is a Lewis acid. | NH3 is a Lewis base. |
| A metal ion binds a water ligand through oxygen. | The metal ion accepts the pair. | The water ligand donates a lone pair. |
Watch out for this
A student says: "NH3 donates a lone pair to BF3, so NH3 is a reducing agent and BF3 is reduced." What is wrong with this?
Show the answer
An electron pair becomes shared in a new bond. Do not automatically call its donor a reducing agent. Acid-base classification follows proton or electron-pair donation; oxidation and reduction require a change in oxidation state.
Check your understanding
Why does BF3 + NH3 -> F3B-NH3 need the Lewis definition of acids and bases?
Show the answer
No proton moves; BF3 accepts a lone pair from NH3 to form a B-N bond.
Bronsted-Lowry needs a proton transfer, and there is none here. The Lewis definition covers electron-pair acceptors (BF3) and donors (NH3).
Quick revision
Revisit the essentials, then return to an explanation when you need it.
| Term | Definition to write |
|---|---|
| Arrhenius acid and base | An Arrhenius acid produces H+ ions in aqueous solution; an Arrhenius base produces OH- ions in aqueous solution. |
| Brønsted-Lowry acid and base | An acid is a proton (H+) donor; a base is a proton acceptor. |
| Lewis acid and base | A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor. |
| Conjugate acid-base pair | Two species that differ by one proton, such as CH3COOH and CH3COO-. |
| Theory | Ask |
|---|---|
| Arrhenius | Does it increase H+(aq) or OH-(aq)? |
| Bronsted-Lowry | Which species donates a proton, and which accepts it? |
| Lewis | Where is the donated electron pair, and which species accepts it? |
A conjugate pair differs by one H+. Water can be acid or base depending on its partner. BF3 accepting a lone pair from NH3 shows why acid-base chemistry is broader than aqueous neutralisation.
Scope
Learning outcomes
4. Theories of Acids and Bases. Use the outcome map to find the explanation for a particular syllabus requirement.
See the learning outcome map
4(a) Understand and apply the Arrhenius theory.
- Acids and hydrogen ions in water
- Bases and hydroxide ions in water
- Aqueous scope and neutralisation
4(b) Understand and apply proton-transfer theory.
- Bronsted-Lowry donors and acceptors
- Conjugate acids and bases
- One-proton/one-charge distinction and amphiprotic examples
4(c) Understand and apply the Lewis theory.
- Electron-pair donors and acceptors
- Non-aqueous systems
- Named BF3 and NH3 reaction