Three Definitions, Increasing in Scope
Arrhenius, Brønsted-Lowry, and Lewis theories of acids and bases
Chemistry uses more than one formal definition of "acid" and "base" because each successive theory was developed to explain reactions the previous definition couldn't account for. The oldest and narrowest is the Arrhenius definition: an acid is a substance that increases the concentration of H⁺ ions when dissolved in water, and a base is a substance that increases the concentration of OH⁻ ions when dissolved in water. This definition works well for classic examples like HCl and NaOH, but it's limited specifically to aqueous solutions and doesn't explain why substances like ammonia (NH₃), which contains no OH⁻ group at all, still behave as a base.
The Brønsted-Lowry definition, developed later, broadens the scope considerably: an acid is any substance that donates a proton (H⁺), and a base is any substance that accepts a proton. This definition no longer requires water as the solvent, and it correctly explains ammonia's basic behavior — NH₃ accepts a proton from water to form NH₄⁺, making it a Brønsted-Lowry base even though it contains no hydroxide group. Brønsted-Lowry theory also introduces the concept of conjugate acid-base pairs: when an acid donates a proton, what's left behind is its conjugate base; when a base accepts a proton, what results is its conjugate acid.
The Lewis definition is the broadest of the three: a Lewis acid is any substance that accepts a pair of electrons, and a Lewis base is any substance that donates a pair of electrons. This definition doesn't require a proton to be involved at all, which means it correctly classifies reactions with no H⁺ transfer whatsoever as acid-base reactions — most famously, the reaction between boron trifluoride (BF₃) and ammonia (NH₃), where BF₃ (missing a full octet, and therefore able to accept an electron pair) acts as a Lewis acid, and NH₃ (with a lone pair on nitrogen) acts as a Lewis base, even though no proton changes hands in the reaction at all.
💡 Every Brønsted-Lowry Acid Is Also a Lewis Acid — But Not the Reverse
Because Lewis theory is the broadest of the three definitions, every reaction that qualifies as a Brønsted-Lowry acid-base reaction also qualifies as a Lewis acid-base reaction — but the reverse is not true. There are Lewis acid-base reactions that involve no proton transfer at all, and therefore don't fit the Brønsted-Lowry definition, meaning they can only be correctly described using Lewis theory.
Think of the three theories as concentric circles rather than competing, mutually exclusive definitions: the Arrhenius definition (aqueous H⁺/OH⁻ behavior) sits inside the Brønsted-Lowry definition (proton donation/acceptance in any context), which itself sits inside the Lewis definition (electron pair donation/acceptance, the broadest category of all). A substance's classification as an 'acid' or 'base' can depend on which theoretical lens you're using — H⁺ itself is an acid under all three definitions, but BF₃ is only recognizably an acid under Lewis theory, since it has no proton to donate and doesn't behave as an Arrhenius acid in water in the classic sense.
Arr
Arrhenius theory
Arrhenius theory, the earliest formal acid-base definition, states that an acid increases [H⁺] when dissolved in water, and a base increases [OH⁻] when dissolved in water. This definition is intuitive and works well for many classic acids (HCl, H₂SO₄) and bases (NaOH, KOH), but its scope is limited: it applies only to reactions in aqueous solution, and it cannot explain why substances lacking an OH⁻ group (like ammonia) still act as bases, since strictly speaking Arrhenius theory requires a base to directly release hydroxide ions upon dissolving.
NaOH is a textbook Arrhenius base — it dissociates directly in water to release Na⁺ and OH⁻, immediately raising [OH⁻] in solution.
BL
Brønsted-Lowry theory and conjugate pairs
Brønsted-Lowry theory defines an acid as a proton (H⁺) donor and a base as a proton acceptor, applicable in any solvent (or even without a solvent), not just water. When an acid (HA) donates a proton, it becomes its conjugate base (A⁻); when a base (B) accepts a proton, it becomes its conjugate acid (BH⁺). Every Brønsted-Lowry acid-base reaction actually involves two conjugate acid-base pairs reacting simultaneously — the original acid/its conjugate base, and the original base/its conjugate acid.
Water itself is amphoteric (or amphiprotic) under this theory — capable of acting as either an acid or a base depending on what it's reacting with, since it can both donate and accept a proton. This dual behavior is central to water's role as the near-universal solvent for acid-base chemistry.
In the reaction HCl + H₂O → H₃O⁺ + Cl⁻, HCl is the acid (donates a proton) and its conjugate base is Cl⁻; water is the base (accepts a proton) and its conjugate acid is H₃O⁺ (hydronium).
Lewis
Lewis theory — electron pairs, not protons
Lewis theory shifts the focus entirely away from protons and onto electron pairs: a Lewis acid accepts an electron pair, and a Lewis base donates an electron pair. This reframing captures reactions that have nothing to do with H⁺ transfer at all. A classic example is a metal cation like Fe³⁺ acting as a Lewis acid when it accepts electron pairs from surrounding water molecules or other ligands to form a coordination complex — this kind of reaction is central to a huge amount of inorganic and coordination chemistry, none of which fits neatly into the Brønsted-Lowry proton-transfer framework.
Substances that are electron-deficient (missing a full octet, like BF₃, or having an empty orbital, like many metal cations) tend to act as Lewis acids, since they have an available "slot" to accept an incoming electron pair. Substances with an available lone pair of electrons (like NH₃, or the oxygen in water) tend to act as Lewis bases, since they have electrons available to donate.
BF₃ + NH₃ → F₃B–NH₃ is a textbook Lewis acid-base reaction with no proton transfer whatsoever — BF₃'s boron atom (only 6 electrons, missing a full octet) accepts the lone pair from nitrogen in NH₃, forming a new covalent bond.
🔬 Applied Scenario — Choosing the Right Theory for the Reaction
Recognizing which acid-base theory actually applies to a given reaction — and understanding why more than one might apply simultaneously — is a practical skill this topic is really testing.
A
A reaction in water involving a hydroxide-releasing compound. If a substance dissolves in water and directly releases OH⁻ ions (like NaOH or KOH), all three theories agree it's a base — it satisfies Arrhenius (raises [OH⁻] directly), Brønsted-Lowry (the released OH⁻ can go on to accept a proton), and Lewis (oxygen's lone pairs can accept an electron pair) simultaneously.
B
A nitrogen-containing base with no hydroxide group, like ammonia. Arrhenius theory alone cannot classify NH₃ as a base, since it contains no OH⁻ to release. Brønsted-Lowry theory correctly identifies it as a base because it accepts a proton from water. This is the classic case demonstrating why Brønsted-Lowry theory was needed as an advance beyond Arrhenius theory.
C
A metal ion forming a coordination complex. When a transition metal cation like Fe³⁺ or Cu²⁺ bonds with surrounding water molecules or other ligands by accepting electron pairs, this reaction has no proton transfer at all and cannot be described using Brønsted-Lowry theory. Only Lewis theory, focused on electron pair donation and acceptance rather than protons, correctly frames this as an acid-base reaction, with the metal cation as the Lewis acid.
D
Choosing the theory that matches the reaction's mechanism. In practice, chemists select whichever theory offers the clearest explanatory lens for a given reaction — proton-transfer reactions in solution are usually described with Brønsted-Lowry language (acid, base, conjugate pair), while reactions involving coordination complexes, electron-deficient species, or organic mechanisms are more naturally described with Lewis acid-base language (electrophile-like behavior, nucleophile-like behavior).
⚠️ Most Common Acid-Base Theory Mistakes
Every Brønsted-Lowry acid is a Lewis acid, but not every Lewis acid is a Brønsted-Lowry acid — students frequently reverse this relationship. Lewis theory is the broader, more inclusive definition. A substance like BF₃ is a Lewis acid but has no proton to donate, so it doesn't fit the Brønsted-Lowry definition at all — it's not that BF₃ fails to be an acid, it's that only the broader Lewis definition captures it.
A conjugate base is not simply "the opposite" of an acid — it's specifically what remains after the acid donates a proton. Students sometimes confuse a conjugate base with a completely unrelated base. The conjugate base of HCl is specifically Cl⁻ (what's left after HCl loses its proton), not some arbitrary basic substance.
Ammonia (NH₃) is a base despite containing no hydroxide group — a frequently missed point when only Arrhenius theory is considered. Students taught only the Arrhenius definition sometimes struggle to classify NH₃ as a base, since it releases no OH⁻ directly. Brønsted-Lowry theory resolves this: NH₃ accepts a proton from water, making it a base under that broader definition.
✓ Quick Self-Test
1. What are the three major acid-base theories, and in what order of increasing scope do they apply?
2. What is the key limitation of Arrhenius theory that Brønsted-Lowry theory resolves?
3. What is a conjugate acid-base pair, and how is a conjugate base formed?
4. How does Lewis theory define acids and bases, and what kind of reaction can it explain that Brønsted-Lowry theory cannot?
5. Why is water considered amphoteric under Brønsted-Lowry theory?
Answers:
1. The three theories, in order of increasing scope, are: Arrhenius (acid increases [H⁺], base increases [OH⁻], limited to aqueous solutions), Brønsted-Lowry (acid = proton donor, base = proton acceptor, works in any solvent), and Lewis (acid = electron pair acceptor, base = electron pair donor, the broadest of the three).
2. Arrhenius theory can only classify a base if it directly releases OH⁻ ions when dissolved in water, which means it cannot explain why substances like ammonia (NH₃), which contain no hydroxide group, behave as bases. Brønsted-Lowry theory resolves this by defining a base as any proton acceptor, which correctly classifies NH₃ as a base because it accepts a proton from water to form NH₄⁺.
3. A conjugate acid-base pair consists of two species that differ from each other by exactly one proton (H⁺). A conjugate base is formed when an acid donates a proton — for example, when HCl donates a proton, its conjugate base is Cl⁻.
4. Lewis theory defines an acid as any substance that accepts a pair of electrons, and a base as any substance that donates a pair of electrons. This allows it to explain reactions with no proton transfer at all, such as the reaction between BF₃ (a Lewis acid, missing a full octet) and NH₃ (a Lewis base, donating its lone pair), which Brønsted-Lowry theory cannot describe since there is no proton being transferred.
5. Water is considered amphoteric (or amphiprotic) because it can act as either a Brønsted-Lowry acid (donating a proton) or a Brønsted-Lowry base (accepting a proton), depending on what it's reacting with — this dual capability is why water is such a versatile solvent for acid-base chemistry.