Last updated: 17 August 2026
At O-Level, it is often enough to think of acids as substances that produce hydrogen ions and alkalis as substances that produce hydroxide ions in water. That picture remains useful, but at A-Level we quickly encounter reactions for which it is too narrow.
Ammonia, NH₃, is an obvious example. It contains no OH group, yet it behaves as a base. BF₃ presents an even bigger challenge: it acts as an acid towards NH₃ even though no H⁺ is present anywhere in the reaction.
The solution is not to discard the earlier definition. Instead, Chemistry uses three related theories. Arrhenius theory focuses on aqueous solutions, Brønsted–Lowry theory follows the movement of a proton, and Lewis theory follows the movement of an electron pair. Once you know what each theory is tracking, the three definitions fit together naturally.
Try these five questions before reading the chapter. They test the ideas about ions, proton transfer and lone pairs that will be developed below.
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Why are there three theories of acids and bases?
The three theories are best viewed as different levels of description rather than competing definitions. Each asks a different question about the reaction.
| Theory | What it focuses on | Acid | Base |
|---|---|---|---|
| Arrhenius | Ions produced in aqueous solution | Produces H⁺(aq) | Produces OH⁻(aq) |
| Brønsted–Lowry | Transfer of H⁺ | Proton donor | Proton acceptor |
| Lewis | Electron-pair donation and acceptance | Electron-pair acceptor | Electron-pair donor |
Arrhenius theory is the narrowest because it is concerned specifically with aqueous solution. Brønsted–Lowry theory removes that restriction and asks whether H⁺ is transferred. Lewis theory is broader still because an acid–base reaction can occur simply through electron-pair donation, even when no proton is present.
Think of them as three ways of looking at acid–base behaviour. Arrhenius follows ions in water, Brønsted–Lowry follows H⁺, and Lewis follows an electron pair. Choosing the correct theory becomes much easier when you first identify what is actually moving in the reaction.
Arrhenius theory: acids and bases in water
Arrhenius theory provides the familiar starting point for acid–base chemistry. In the simple Arrhenius treatment, acids and bases are identified by the ions they produce when dissolved in water.
Arrhenius acids
An Arrhenius acid produces H⁺ ions in aqueous solution.
For example, hydrogen chloride behaves as an acid when dissolved in water:
However, a completely free H⁺ ion does not simply float around independently in water. It is accepted by a water molecule to form the hydroxonium ion, H₃O⁺.
H⁺(aq) is therefore commonly used as convenient shorthand for the acidic proton present in hydrated form in aqueous solution.
Arrhenius bases
In the simple Arrhenius model, an Arrhenius base produces OH⁻ ions when it dissolves in water.
Sodium hydroxide provides a straightforward example:
Calcium hydroxide produces two hydroxide ions per formula unit:
Neutralisation in the Arrhenius picture
When an aqueous acid reacts with an aqueous alkali, H⁺ and OH⁻ combine to form water. The spectator ions can be removed to give the ionic equation:
If the acidic species is written explicitly as H₃O⁺, the same process becomes:
Why Arrhenius theory is limited
Arrhenius theory works very well for familiar acids and alkalis in water, but its dependence on aqueous H⁺ and OH⁻ makes it too narrow for many other acid–base reactions.
Ammonia illustrates the limitation. NH₃ contains no OH⁻ ions that can simply dissociate from it, yet aqueous ammonia is basic:
The OH⁻ is produced because NH₃ reacts with water. Brønsted–Lowry theory explains the process much more naturally: NH₃ accepts a proton from H₂O.
Arrhenius theory also cannot provide a useful description of reactions such as:
There is no aqueous solution and no H⁺ or OH⁻ transfer. Lewis theory is needed to describe this reaction.
Too narrow: “A base must contain an OH group.”
Better reasoning: Ionic hydroxides such as NaOH release OH⁻ directly in water, but substances such as NH₃ can also behave as bases even though their formulas contain no OH group.
Brønsted–Lowry theory: follow the proton
Brønsted–Lowry theory shifts attention away from particular ions in water and towards the transfer of a proton from one species to another.
Why is H⁺ called a proton?
The most common isotope of hydrogen contains one proton and one electron. When a hydrogen atom loses its electron to form H⁺, what remains is simply its nucleus — a proton.
This is why the expressions H⁺ transfer and proton transfer mean the same thing in this topic.
The acid and base must act together
If one species donates a proton, another species must receive it. Acids and bases therefore work as partners in a Brønsted–Lowry reaction.
A general reaction can be represented as:
- HA donates H⁺ and therefore acts as the acid.
- B accepts H⁺ and therefore acts as the base.
- HA becomes A⁻ after losing H⁺.
- B becomes BH⁺ after gaining H⁺.
Do not classify a species simply because its formula “looks acidic” or “looks basic”. Compare the reactants with the products. The species that loses H⁺ is the acid, while the species that gains H⁺ is the base.
Example: hydrogen chloride and water
Compare each reactant with the corresponding product:
- HCl loses H⁺ to form Cl⁻.
- H₂O gains H⁺ to form H₃O⁺.
Therefore, HCl is the Brønsted–Lowry acid and H₂O is the Brønsted–Lowry base.
Example: ammonia and water
This time:
- NH₃ gains H⁺ to form NH₄⁺, so NH₃ is the base.
- H₂O loses H⁺ to form OH⁻, so H₂O is the acid.
Notice that NH₃ does not need to contain OH⁻ in order to be a base. Its basic behaviour comes from its ability to accept H⁺.
Conjugate acid–base pairs
Proton transfer creates a relationship between the species before and after the transfer.
When an acid donates H⁺, the species left behind is its conjugate base.
When a base accepts H⁺, the species formed is its conjugate acid.
This gives you a reliable way to identify the pairs. Compare formulas rather than trying to guess which substances “look related”.
Finding the conjugate pairs in NH₃ + H₂O
| Species | What happens? | Partner | Relationship |
|---|---|---|---|
| NH₃ | Accepts H⁺ | NH₄⁺ | Base / conjugate acid |
| H₂O | Donates H⁺ | OH⁻ | Acid / conjugate base |
The two conjugate acid–base pairs are therefore:
NH₄⁺ / NH₃
H₂O / OH⁻
A reliable method for finding conjugate pairs
- Compare each reactant with the products.
- Find the species that loses one H⁺. It is the acid.
- Its product is the conjugate base.
- Find the species that gains one H⁺. It is the base.
- Its product is the conjugate acid.
Worked example: HSO₄⁻ and NH₃
HSO₄⁻ loses H⁺ to become SO₄²⁻.
HSO₄⁻ = acid
SO₄²⁻ = conjugate base
NH₃ gains H⁺ to become NH₄⁺.
NH₃ = base
NH₄⁺ = conjugate acid
The conjugate pairs are therefore HSO₄⁻/SO₄²⁻ and NH₄⁺/NH₃.
Incorrect idea: HSO₄⁻ and NH₃ form a conjugate pair because one is the acid and the other is the base.
Correct reasoning: They are reaction partners. A conjugate pair consists of two species that differ from each other by exactly one H⁺.
Amphiprotic species can donate or accept H⁺
Some species can behave as either a Brønsted–Lowry acid or a Brønsted–Lowry base, depending on the species with which they react. Such species are described as amphiprotic.
Water acting as a base
Consider water reacting with HCl:
H₂O accepts H⁺ and becomes H₃O⁺. It therefore acts as a base.
Water acting as an acid
Now consider water reacting with NH₃:
H₂O donates H⁺ and becomes OH⁻. It therefore acts as an acid.
Water is consequently amphiprotic.
Other important amphiprotic species include HCO₃⁻, HSO₄⁻ and H₂PO₄⁻.
It means that the species has the ability to donate H⁺ in one reaction and accept H⁺ in another. Its actual role must be determined from the particular equation.
Lewis theory: follow the electron pair
Brønsted–Lowry theory is powerful, but it still requires proton transfer. Lewis theory removes even this requirement and focuses instead on the electron pair used to form a bond.
A Lewis base therefore needs an available pair of electrons that it can donate. This is commonly a lone pair.
A Lewis acid needs an electron-deficient centre or an available orbital that can accept the electron pair.
BF₃ is electron-deficient
In BF₃, boron forms three covalent bonds. There are therefore only six electrons around the boron atom.
Boron does not have a complete octet and has an available orbital that can accept an electron pair. BF₃ can therefore act as a Lewis acid.
NH₃ has an available lone pair
Nitrogen in NH₃ forms three N–H bonds and has one lone pair of electrons.
This lone pair can be donated to an electron-pair acceptor. NH₃ can therefore act as a Lewis base.
Reaction between BF₃ and NH₃
The nitrogen atom in NH₃ donates its lone pair to the electron-deficient boron atom in BF₃.
- NH₃ donates the electron pair, so NH₃ is the Lewis base.
- BF₃ accepts the electron pair, so BF₃ is the Lewis acid.
The new N→B bond is a co-ordinate, or dative covalent, bond because both electrons in the shared pair originally came from nitrogen.
For NH₃ + BF₃, the electron pair originates on nitrogen. The arrow therefore points from N towards B, not from B towards N.
Another Lewis acid–base reaction: AlCl₃ and Cl⁻
Cl⁻ possesses lone pairs and donates an electron pair to aluminium. It is therefore the Lewis base.
AlCl₃ accepts the electron pair at its electron-deficient aluminium centre and is therefore the Lewis acid.
Metal ions can also act as Lewis acids
Positive metal ions can accept electron pairs from ligands when complex ions form. For example:
NH₃ donates lone pairs and acts as a Lewis base. Ag⁺ accepts the electron pairs and acts as a Lewis acid.
How Brønsted–Lowry and Lewis theories are connected
A proton-transfer reaction can also be examined at the electron-pair level. Consider:
From the Brønsted–Lowry viewpoint, NH₃ accepts H⁺ and is therefore a base.
Now look more closely at how the N–H bond forms. The lone pair on nitrogen is donated to H⁺.
- NH₃ donates an electron pair and is therefore a Lewis base.
- H⁺ accepts the electron pair and is therefore a Lewis acid.
Brønsted–Lowry proton transfer can therefore also be understood as an electron-pair interaction.
Brønsted–Lowry view: follow the H⁺ being transferred.
Lewis view: follow the electron pair used by the base to form a new bond to that H⁺.
Comparing the scope of the three theories
The essential difference between the theories can now be seen clearly.
| Feature | Arrhenius | Brønsted–Lowry | Lewis |
|---|---|---|---|
| Main idea | H⁺ and OH⁻ in aqueous solution | Transfer of H⁺ | Electron-pair donation and acceptance |
| Acid | Produces H⁺ in water | Donates H⁺ | Accepts an electron pair |
| Base | Produces OH⁻ in water | Accepts H⁺ | Donates an electron pair |
| Requires aqueous solution? | Yes | No | No |
| Requires proton transfer? | Acid behaviour centres on H⁺ in water | Yes | No |
| Can explain BF₃ + NH₃? | No | No | Yes |
It is useful to think about the increasing scope at the reaction level. Reactions described using the Arrhenius picture can generally also be examined as proton-transfer reactions. Proton-transfer reactions can in turn be examined at the electron-pair level.
However, the reverse is not necessarily possible. The Lewis reaction between BF₃ and NH₃ involves electron-pair donation but no H⁺ transfer, so it cannot be described using Brønsted–Lowry theory.
Connection to organic chemistry
Lewis theory becomes especially important when you later study organic reaction mechanisms.
A nucleophile donates an electron pair to an electron-deficient centre. It therefore behaves as a Lewis base.
An electrophile accepts an electron pair. It therefore behaves as a Lewis acid.
| Lewis description | Electron-pair behaviour | Organic Chemistry term |
|---|---|---|
| Lewis acid | Electron-pair acceptor | Electrophile |
| Lewis base | Electron-pair donor | Nucleophile |
This is why understanding Lewis acids and bases now makes curved-arrow mechanisms much easier later: the arrow simply tracks the movement of an electron pair from a donor towards an acceptor.
Worked examples using the three theories
Example 1: identifying conjugate pairs
Consider:
CH₃COOH loses H⁺ to form CH₃COO⁻.
CH₃COOH = acid
CH₃COO⁻ = conjugate base
H₂O gains H⁺ to form H₃O⁺.
H₂O = base
H₃O⁺ = conjugate acid
Example 2: recognising amphiprotic behaviour
Consider:
CO₃²⁻ gains H⁺ to form HCO₃⁻, so CO₃²⁻ is the base.
H₂O loses H⁺ to form OH⁻, so H₂O is the acid.
This is one example of water displaying its amphiprotic character.
Example 3: identifying Lewis acid and base
Each NH₃ molecule provides a lone pair to form a co-ordinate bond with Ag⁺.
NH₃ = electron-pair donor = Lewis base
Ag⁺ = electron-pair acceptor = Lewis acid
Example 4: choosing an appropriate theory
Suppose HCl(aq) reacts with NaOH(aq). Arrhenius theory gives a perfectly useful description because H⁺ and OH⁻ are present in aqueous solution.
Now consider BF₃ reacting with NH₃ in the gas phase. No proton is transferred and no water is required.
Lewis theory is the appropriate description because NH₃ donates an electron pair and BF₃ accepts it.
Common mistakes with acid–base theories
| Common mistake | Why it is wrong | Correct approach |
|---|---|---|
| Saying that every acid must contain hydrogen | Lewis acids such as BF₃ can accept electron pairs without containing hydrogen | State which acid–base theory you are using |
| Saying that every base must contain OH | NH₃ behaves as a base even though it contains no oxygen | For Brønsted–Lowry theory, ask whether the species accepts H⁺ |
| Defining a Lewis acid as an electron-pair donor | The donor is the Lewis base | Lewis acid accepts; Lewis base donates |
| Calling any acid and base in the same equation a conjugate pair | Conjugate partners must differ by exactly one H⁺ | Compare the formulas before and after proton transfer |
| Assuming H₂O is always a base | H₂O can either accept or donate H⁺ | Determine its role from the particular reaction |
| Calling BF₃ a Brønsted–Lowry acid in BF₃ + NH₃ | BF₃ does not donate H⁺ | BF₃ is the Lewis acid because it accepts an electron pair |
| Drawing the co-ordinate-bond arrow from BF₃ towards NH₃ | The electron pair originates from NH₃ | Draw the arrow from N towards B |
| Starting a curly arrow at a positive charge | A curly arrow shows movement of an electron pair and must begin where the electrons are | Start at a lone pair or bond and point towards the electron-pair acceptor |
| Thinking a co-ordinate bond remains a fundamentally different weak bond | “Co-ordinate” describes how the shared pair originated | After formation, treat it as a covalent bond |
Acid–base theory checklist
Choose the theory
- Use Arrhenius for simple H⁺ and OH⁻ behaviour in water.
- Use Brønsted–Lowry when H⁺ is transferred.
- Use Lewis when an electron pair is donated and accepted.
- Check whether the reaction actually requires water or H⁺.
Track the proton
- The acid donates H⁺.
- The base accepts H⁺.
- A conjugate pair differs by exactly one H⁺.
- Check the equation before assigning the role of an amphiprotic species.
Track the electron pair
- The Lewis base supplies the electron pair.
- The Lewis acid accepts the electron pair.
- Start an electron-flow arrow at a lone pair or bond.
- Point towards the electron-deficient centre.
Write the explanation
- Name the species and state its role.
- Explain whether H⁺ is gained or lost.
- Identify both conjugate pairs when required.
- For Lewis reactions, state who donates and who accepts the electron pair.
These ten questions move beyond simple definitions and test proton transfer, conjugate pairs, amphiprotic behaviour and Lewis electron-pair reasoning.
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Frequently asked questions
Questions students often ask about acid–base theories
These questions connect the three theories and clear up distinctions that become increasingly important in acid–base equilibria, inorganic chemistry and organic reaction mechanisms.
Does an acid always have to contain hydrogen?
Not under Lewis theory. A Brønsted–Lowry acid must be able to donate H⁺, but a Lewis acid only needs to accept an electron pair. BF₃ and AlCl₃ are therefore Lewis acids even though they contain no hydrogen.
Does a base always need to contain OH⁻?
No. Ionic hydroxides such as NaOH release OH⁻ directly in water, but a Brønsted–Lowry base is defined by its ability to accept H⁺. NH₃ is a base because its nitrogen lone pair can accept a proton to form NH₄⁺.
Why do some equations use H⁺ while others use H₃O⁺?
H⁺ is convenient shorthand for the proton involved in acid–base reactions. In aqueous solution, a bare proton is strongly solvated and can be represented as H₃O⁺ when bonded to a water molecule.
HCl + H₂O → H₃O⁺ + Cl⁻ therefore gives a more explicit particle-level description than simply writing HCl → H⁺ + Cl⁻.
What is the difference between amphiprotic and amphoteric?
Amphiprotic specifically describes a species that can both donate and accept a proton according to Brønsted–Lowry theory.
Amphoteric is the broader term for a substance capable of behaving as either an acid or a base. Amphiprotic behaviour is therefore a particular type of amphoteric behaviour.
Why is BF₃ a Lewis acid when every fluorine atom has an octet?
The electron arrangement around fluorine does not determine whether the boron centre is electron-deficient. Boron forms only three bonds in BF₃, giving it six electrons around itself.
Boron therefore has an available orbital capable of accepting an electron pair, allowing BF₃ to behave as a Lewis acid.
Is a co-ordinate bond weaker than an ordinary covalent bond?
Not simply because it was formed by lone-pair donation. The term “co-ordinate” describes the origin of the bonding pair: both electrons were supplied by the same atom.
Once the bond has formed, it is a covalent bond. Its strength depends on the atoms and their bonding environment rather than on which atom originally supplied the electrons.
Are Brønsted–Lowry acid–base reactions also Lewis acid–base reactions?
They can be interpreted that way at the reaction level. A Brønsted–Lowry base accepts H⁺ by donating an electron pair to that proton. The base is therefore also acting as a Lewis base, while H⁺ accepts the electron pair and acts as a Lewis acid.
The reverse is not generally true. BF₃ + NH₃ is a Lewis acid–base reaction but not a Brønsted–Lowry acid–base reaction because no proton is transferred.
Why is Brønsted–Lowry theory so important for acid–base equilibria?
Acid–base equilibria involve reversible proton transfer. Brønsted–Lowry theory therefore identifies the acid, base, conjugate acid and conjugate base directly.
This becomes especially useful when studying weak acids, weak bases, Kₐ, Kᵦ, buffers and other aqueous equilibria.
Can NH₃ be both a Brønsted–Lowry base and a Lewis base?
Yes. When NH₃ reacts with H⁺, it accepts the proton and is therefore a Brønsted–Lowry base.
At the same time, nitrogen donates its lone pair to H⁺ to form the new N–H bond. NH₃ is therefore also behaving as a Lewis base.
Is being an acid or base a permanent property of a species?
Not always. The role of a species depends on the reaction and on the theory being used.
Water, for example, is a Brønsted–Lowry base when it accepts H⁺ from HCl, but a Brønsted–Lowry acid when it donates H⁺ to NH₃. This is why analysing the actual reaction is more reliable than relying on labels alone.
Arrhenius, Brønsted–Lowry and Lewis theories describe acid–base behaviour at increasingly broad levels. Arrhenius theory gives the familiar aqueous picture, Brønsted–Lowry theory reveals proton transfer and conjugate pairs, and Lewis theory reveals the electron-pair interaction underlying an even wider range of reactions. Once you can confidently follow either H⁺ or an electron pair from one species to another, later topics such as acid–base equilibria, complex-ion formation and organic mechanisms become much easier to understand.