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“An acid is something that tastes sour.” That was the working definition for centuries, and for obvious reasons it isn’t used in labs any more. Over the past 140 years chemists have proposed three increasingly general definitions of acids and bases. Each one is still in use, because each is the most convenient for a different kind of problem.
1. Arrhenius (1884): acids and bases in water
The Swedish chemist Svante Arrhenius proposed that:
- an acid is a substance that produces hydrogen ions (H⁺) in water, and
- a base is a substance that produces hydroxide ions (OH⁻) in water.
HCl → H⁺ + Cl⁻ (acid) NaOH → Na⁺ + OH⁻ (base)
Neutralization is then just H⁺ + OH⁻ → H₂O.
This was a big step — it explained why all acids behave alike (they all make H⁺) — and it earned Arrhenius the 1903 Nobel Prize in Chemistry. But it has two limits:
- It only works in water.
- It can’t explain bases that contain no hydroxide. Ammonia, NH₃, clearly makes solutions basic, yet it has no OH⁻ to release.
2. Brønsted-Lowry (1923): proton transfer
In the same year, Johannes Brønsted in Denmark and Thomas Lowry in England independently proposed a broader idea. A hydrogen ion is just a bare proton, so:
- an acid is a proton donor, and
- a base is a proton acceptor.
Now ammonia makes sense. It takes a proton from water:
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
Ammonia accepts a proton, so it’s a base. Water donates one, so here water is acting as an acid.
This definition also works outside water. Hydrogen chloride gas reacting with ammonia gas to form a white smoke of ammonium chloride is an acid-base reaction, with no water anywhere:
HCl + NH₃ → NH₄Cl
Conjugate pairs
Every Brønsted-Lowry reaction involves two pairs. When an acid gives away a proton, what remains can take one back, so it’s a base — the acid’s conjugate base. Likewise, a base that has accepted a proton becomes its conjugate acid.
| Acid | ⇌ | Conjugate base |
|---|---|---|
| HCl | Cl⁻ | |
| CH₃COOH | CH₃COO⁻ | |
| NH₄⁺ | NH₃ | |
| H₂O | OH⁻ | |
| H₃O⁺ | H₂O |
Conjugate pairs differ by exactly one H⁺. The stronger an acid, the weaker its conjugate base — see strong vs. weak acids.
Amphoteric substances
Water appears as both an acid (giving OH⁻) and a base (giving H₃O⁺) in the table. Substances that can do both are called amphoteric (or amphiprotic). Water, the hydrogen carbonate ion HCO₃⁻ and amino acids are important examples. Some metal oxides and hydroxides, like aluminium hydroxide, are also amphoteric: they dissolve in both strong acids and strong bases.
3. Lewis (1923): electron pairs
Gilbert N. Lewis — the same Lewis of Lewis dot structures — looked at the electrons instead of the proton. When a base accepts a proton, it does so by using a lone pair of electrons to form a new bond. So Lewis defined:
- a Lewis base as an electron-pair donor, and
- a Lewis acid as an electron-pair acceptor.
Every Brønsted-Lowry base is also a Lewis base. But the Lewis definition includes reactions with no protons at all. Boron trifluoride has an empty orbital and accepts a lone pair from ammonia:
BF₃ + NH₃ → F₃B–NH₃
BF₃ is a Lewis acid, even though it has no hydrogen to donate. Metal ions in solution are Lewis acids too: when Cu²⁺ attracts water or ammonia molecules around it to form a complex ion, each of those molecules is donating a lone pair.
Which definition should I use?
| Definition | Acid | Base | Best for |
|---|---|---|---|
| Arrhenius | makes H⁺ in water | makes OH⁻ in water | simple aqueous neutralization, introductory problems |
| Brønsted-Lowry | proton donor | proton acceptor | pH, equilibria, conjugate pairs, weak acids and bases |
| Lewis | electron-pair acceptor | electron-pair donor | complex ions, organic mechanisms, catalysis |
Each definition contains the one before it. All Arrhenius acids are Brønsted-Lowry acids, and all Brønsted-Lowry bases are Lewis bases. You pick the simplest one that covers the reaction in front of you.
Quick answers
Is H⁺ really free in water? No. A bare proton immediately attaches to a water molecule to make hydronium, H₃O⁺ (and larger clusters). Writing H⁺ is a convenient shorthand.
Is water an acid or a base? Both. It’s amphoteric, and in pure water it does both to itself very slightly, which is where pH 7 comes from. See the pH scale explained.
Why is ammonia a base if it has no OH⁻? Because it accepts a proton from water, producing OH⁻. It’s a base by the Brønsted-Lowry definition, not the Arrhenius one.
Try it
The pH calculator works with both strong and weak acids and bases, including weak bases like ammonia defined by their Kb.
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