Arrhenius vs Brønsted-Lowry Acids: The Difference

Ammonia turns red litmus blue, neutralises acids and behaves like a base in every way you can test. It also contains no OH group whatsoever. That awkward fact is the reason chemistry has more than one definition of a base.

The short answer: the Arrhenius definition says an acid produces H⁺ ions in water and a base produces OH⁻ ions in water. The Brønsted–Lowry definition is broader: an acid is any proton donor and a base any proton acceptor, which works outside water and explains substances like ammonia that have no hydroxide to give.

Quick comparison at a glance

Feature	Arrhenius	Brønsted–Lowry
Proposed	1884	1923
Acid is	A substance producing H⁺ in water	A proton (H⁺) donor
Base is	A substance producing OH⁻ in water	A proton (H⁺) acceptor
Requires water?	Yes	No
Covers ammonia as a base?	Not under the original definition	Yes
Covers gas-phase reactions?	No	Yes
Introduces conjugate pairs?	No	Yes
Scope	Narrower	Broader — includes all Arrhenius acids and bases

What the Arrhenius definition says

Svante Arrhenius proposed the first modern definition in 1884, and it's still the one most students meet first because it's concrete and it works for the common cases.

• An Arrhenius acid dissociates in water to release H⁺: HCl → H⁺ + Cl⁻
• An Arrhenius base dissociates in water to release OH⁻: NaOH → Na⁺ + OH⁻

Clean, testable, and enough to handle HCl, HNO₃, H₂SO₄, NaOH and KOH — most of a first-year syllabus.

Its limits are the price of that simplicity. It's tied to water: a reaction with no water isn't covered at all. And it defines a base by containing and releasing hydroxide, which leaves ammonia stranded. NH₃ has no OH group, yet an ammonia solution is unmistakably basic.

What the Brønsted–Lowry definition says

In 1923 Johannes Brønsted and Thomas Lowry (working separately, and now permanently hyphenated) shifted the focus from what a substance contains to what it does.

• A Brønsted–Lowry acid donates a proton.
• A Brønsted–Lowry base accepts a proton.

"Proton" here just means H⁺ — a hydrogen atom stripped of its only electron is a bare proton.

Ammonia now fits without strain:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

Ammonia accepts a proton from water, so ammonia is the base and water is the acid. No hydroxide had to be in the ammonia to begin with; the OH⁻ is simply what water leaves behind.

The definition also frees acids and bases from water entirely. Hydrogen chloride gas meeting ammonia gas produces a white smoke of ammonium chloride — a genuine acid–base reaction with no solvent in sight, invisible to Arrhenius.

Conjugate pairs — the useful part

Because Brønsted–Lowry is about transfer, every reaction has two acids and two bases. When an acid donates its proton, what remains is a base — its conjugate base. When a base accepts one, what forms is its conjugate acid.

CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺

Reading left to right, that's acid + base ⇌ conjugate base + conjugate acid. The pairs differ by exactly one H⁺: CH₃COOH/CH₃COO⁻ and H₂O/H₃O⁺.

This is the machinery behind buffers, which are nothing more than a weak acid sitting alongside its conjugate base. It also explains amphoteric substances like water, which donates a proton to ammonia and accepts one from acetic acid depending on the company it keeps.

How to tell which definition a question wants

If the reaction is in water and involves H⁺ or OH⁻, either definition works. Use Arrhenius; it's simpler.

If ammonia, a carbonate, or any base without an OH group appears, you need Brønsted–Lowry.

If there's no water, it has to be Brønsted–Lowry.

If the question mentions conjugate pairs or proton transfer, it's Brønsted–Lowry by definition.

The two aren't rivals. Every Arrhenius acid is also a Brønsted–Lowry acid — HCl releases H⁺ in water precisely by donating a proton to it. Brønsted–Lowry is a wider circle drawn around the same idea. (A third, wider still, is the Lewis definition — acid as electron-pair acceptor — which brings in species with no hydrogen at all.)

Worked examples

Identify the acid and base before reading on.

• HCl + H₂O → H₃O⁺ + Cl⁻ — HCl donates, water accepts. Acid: HCl. Base: H₂O. Both definitions agree.
• NH₃ + HCl → NH₄Cl (as gases) — ammonia accepts, HCl donates. Brønsted–Lowry only; no water.
• HCO₃⁻ + H₂O ⇌ H₂CO₃ + OH⁻ — bicarbonate accepts a proton, so here it's the base.
• HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O — same ion, now donating. Bicarbonate is amphoteric.
• What is the conjugate base of H₂SO₄? Remove one H⁺: HSO₄⁻.
• What is the conjugate acid of NH₃? Add one H⁺: NH₄⁺.

Common mistakes to avoid

• Thinking the definitions contradict each other. They don't. Brønsted–Lowry extends Arrhenius; anything the older definition calls an acid, the newer one does too.
• Adding or removing more than one H⁺ for a conjugate. Conjugate pairs differ by exactly one proton. H₂SO₄'s conjugate base is HSO₄⁻, not SO₄²⁻.
• Assuming a base must contain OH. Ammonia, carbonates and bicarbonates are all bases with no hydroxide group at all.

FAQ

What is the difference between an Arrhenius and a Brønsted–Lowry acid?
An Arrhenius acid produces H⁺ ions in water; a Brønsted–Lowry acid donates a proton to anything. The Brønsted–Lowry definition is broader and doesn't require water.

Why is ammonia a base if it has no OH⁻?
Because it accepts a proton. Under the Brønsted–Lowry definition that's all a base has to do — the OH⁻ in ammonia solution comes from the water molecule that donated the proton.

What is a conjugate acid–base pair?
Two species that differ by exactly one proton, such as CH₃COOH and CH₃COO⁻. The acid becomes its conjugate base when it donates a proton.

Which definition should I use?
Arrhenius is fine for typical reactions in water. Use Brønsted–Lowry when there's no water, when the base has no hydroxide group, or when the question asks about proton transfer or conjugate pairs.

The takeaway

Arrhenius asks what a substance releases in water; Brønsted–Lowry asks what it does with a proton. The second is the more useful lens, because it explains ammonia, works without water, and gives you conjugate pairs — the idea that everything from buffers to titration curves rests on.

Background → Acids vs Bases: What's the Difference? and What Is an Ion? Cations, Anions, and Charges. Related → [Strong vs Weak Acids] (sibling) is about how far a proton transfer goes, and [What Is a Buffer?] (sibling) is conjugate pairs put to work.

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