TL;DR:
Arrhenius: acids make H⁺ in water, bases make OH⁻. Brønsted-Lowry: acids donate protons, bases accept them. Lewis: acids accept an electron pair, bases donate one.
Each definition is wider than the one before it.
In
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, the conjugate pairs are NH₄⁺/NH₃ and H₂O/OH⁻.
Vinegar is an acid and so is stomach fluid. Ammonia in a cleaner is a base and so is baking soda. But what actually makes something an acid? Chemists have given three answers over the years, each one wider than the last, and you need all three because different questions call for different ones.
Definition 1: Arrhenius
The earliest definition focuses on water. An Arrhenius acid is a substance that produces hydrogen ions, H⁺, when dissolved in water. An Arrhenius base produces hydroxide ions, OH⁻.
Hydrochloric acid fits: HCl → H⁺ + Cl⁻. Sodium hydroxide fits: NaOH → Na⁺ + OH⁻. The limit is that it only covers water solutions, and it cannot explain why ammonia, which has no OH in its formula, behaves as a base.
Definition 2: Brønsted-Lowry
A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. A proton here means H⁺, a hydrogen atom without its electron. This definition works in any solvent, and it focuses on what is being passed from one species to another.
After an acid donates a proton, what remains is its conjugate base. After a base accepts a proton, the result is its conjugate acid. Each acid-base pair differs by exactly one H⁺.
Definition 3: Lewis
A Lewis acid accepts a pair of electrons. A Lewis base donates a pair of electrons. This is the widest definition, because it does not need hydrogen at all. It describes the formation of a new bond where one partner supplies both electrons.
Worked example: ammonia in water, seen three ways
Take NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.
Arrhenius view: the reaction produces OH⁻ in water, so ammonia counts as a base in water.
Brønsted-Lowry view: one H⁺ moves from H₂O to NH₃. Water donates it, so water is the acid. NH₃ accepts it, so ammonia is the base.
Conjugate pairs:
H₂O(acid) loses H⁺ to becomeOH⁻(its conjugate base).NH₃(base) gains H⁺ to becomeNH₄⁺(its conjugate acid). The pairs are H₂O/OH⁻ and NH₄⁺/NH₃.Lewis view: the nitrogen in NH₃ has a lone pair. It donates that pair to the incoming H⁺ to form the N-H bond, so NH₃ is a Lewis base.
Check the bookkeeping. Atoms: left has 1 N, 5 H, 1 O. Right: NH₄⁺ has 1 N and 4 H, OH⁻ has 1 O and 1 H, so 1 N, 5 H, 1 O. Charge: left 0, right (+1) + (−1) = 0. And each conjugate pair differs by exactly one H⁺: NH₄⁺ − H⁺ = NH₃, H₂O − H⁺ = OH⁻.
A reaction only the Lewis definition covers
Boron trifluoride and ammonia react without any proton moving: BF₃ + NH₃ → F₃B-NH₃. Boron has an empty orbital and nitrogen has a lone pair. Nitrogen donates the pair, making a new bond. BF₃ is a Lewis acid and NH₃ is a Lewis base, yet neither the Arrhenius nor Brønsted-Lowry definition applies.
Water itself can play either role in Brønsted-Lowry terms. With ammonia it acts as an acid, and with HCl it acts as a base: HCl + H₂O → H₃O⁺ + Cl⁻. A substance that can do both is called amphoteric.
Acids and bases at home: Never mix household cleaners. Bleach mixed with ammonia or with an acid can release toxic gases. Concentrated acids and bases burn skin and eyes. In a lab, wear goggles and gloves, and never taste or sniff a chemical to identify it.
Strength is a separate idea from the definition. A strong acid such as HCl donates its proton almost completely in water. A weak acid such as acetic acid donates only a small fraction of the time, so the reaction runs both ways. Conjugates show the link: the conjugate base of a strong acid is a very weak base, and the conjugate base of a weak acid is a stronger base than that.
Use a three-step routine on any problem. First, ask which species gives up H⁺ and which takes it. Second, write each conjugate by moving one H⁺ and adjusting the charge by one. Third, confirm that atoms and charge balance. If no proton moves, check for an electron pair being donated and use the Lewis idea.
How to check your answer
One proton. Conjugate pairs differ by exactly one H⁺.
Charge. The charge on the conjugate must be one lower for the acid-to-base direction.
Balance. Atoms and charge balance on both sides.
Pick the definition. Does the question mention water only, protons, or electron pairs?
Mistakes that cost marks
Writing the pair backwards: the conjugate base has one fewer H⁺, not one more.
Saying a base must contain OH. Ammonia does not.
Forgetting that water can be an acid or a base.
Misspelling the name: write Brønsted-Lowry with a hyphen.
Assuming every acid is strong. Strength is how fully it donates protons, not how dangerous it is.
Where you will use it next
These definitions underpin pH, buffers and titrations. Read how strength shows up numerically in pH and the logarithmic scale, and stay safe around them with lab safety rules and the reasons behind them.
Frequently asked questions
Is every Brønsted-Lowry acid also a Lewis acid?
Not exactly. A Brønsted-Lowry acid donates H⁺, and the H⁺ itself is the Lewis acid that accepts an electron pair. So the definitions overlap but describe different things.
Which definition should I use?
Use the narrowest one that answers the question. For water solutions, Brønsted-Lowry usually suffices, and Lewis is for reactions with no protons.
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