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Squeeze a lemon onto your tongue and your face does something involuntary. That sharp, puckering sourness is your taste buds detecting acid. Long before anyone knew what an atom was, people grouped substances by exactly this kind of behaviour: acids taste sour, they dissolve some metals, they fizz with chalk and limestone, and they turn certain plant dyes red.
Those observations are still a good starting point. But modern chemistry can tell you why acids behave the way they do, and it turns out almost all of it comes down to one tiny particle: the hydrogen ion.
The short answer
An acid is a substance that releases hydrogen ions (H⁺) when it dissolves in water.
Take hydrogen chloride gas and bubble it into water. The molecules split apart:
HCl → H⁺ + Cl⁻
The solution you get, hydrochloric acid, is full of free hydrogen ions. Those ions are responsible for the sour taste, the reaction with metals, the colour change in indicators, and the low pH reading.
A hydrogen atom is one proton and one electron. Remove the electron and all that’s left is a bare proton. That’s why chemists often describe acids as proton donors. “Hydrogen ion” and “proton” mean the same thing here.
What really happens in water
A bare proton is far too reactive to float around on its own. In water it immediately attaches to a water molecule and forms the hydronium ion, H₃O⁺:
HCl + H₂O → H₃O⁺ + Cl⁻
Textbooks switch freely between writing H⁺(aq) and H₃O⁺(aq). Both describe the same thing: the acidic particle in the solution. H⁺ is just shorthand.
This detail explains something that puzzles many students. Pure hydrogen chloride, as a dry gas or dissolved in a non-polar solvent like methylbenzene, doesn’t turn blue litmus red and doesn’t react with magnesium. The HCl molecules are all there, but without water they don’t release protons. An acid only shows acidic behaviour once it’s in water, or some other solvent that can accept its protons.
The properties of acids
Every classic property of acids is really a property of H⁺ ions.
They taste sour. Sourness is literally how your tongue detects hydrogen ions. Citric acid in lemons, acetic acid in vinegar, lactic acid in yoghurt and phosphoric acid in cola all taste sour for the same reason. (Never taste lab chemicals to test this.)
They change the colour of indicators. Acids turn blue litmus red, turn universal indicator red, orange or yellow, and leave phenolphthalein colourless. Indicators are dyes whose structure and colour change when they gain or lose a proton.
They react with reactive metals to give hydrogen gas. Drop magnesium ribbon into dilute hydrochloric acid and it fizzes:
Mg + 2HCl → MgCl₂ + H₂
The metal gives electrons to the H⁺ ions, which pair up into hydrogen gas. Metals below hydrogen in the reactivity series, such as copper, silver and gold, don’t do this with ordinary dilute acids.
They react with carbonates to give carbon dioxide. This is why vinegar fizzes on baking soda and why acid rain slowly dissolves marble statues:
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
They neutralise bases. Acid plus base gives a salt and water. The H⁺ from the acid and the OH⁻ from the base combine: H⁺ + OH⁻ → H₂O. This is how antacid tablets calm indigestion.
Their solutions conduct electricity. The dissolved ions carry charge, so acidic solutions are electrolytes.
They have a pH below 7. The pH scale measures hydrogen ion concentration. More H⁺ means lower pH. For a full explanation, see our guide to the pH scale.
Three definitions of an acid
As chemistry progressed, the definition of an acid was widened twice. Each new version includes everything the older one covered and adds more.
| Definition | An acid is… | Proposed |
|---|---|---|
| Arrhenius | a substance that produces H⁺ ions in water | 1884, Svante Arrhenius |
| Brønsted–Lowry | a proton (H⁺) donor | 1923, Johannes Brønsted and Thomas Lowry |
| Lewis | an electron-pair acceptor | 1923, Gilbert N. Lewis |
Arrhenius is the version most people learn first, and it’s the one used above. Its limitation is that it only works in water.
Brønsted–Lowry focuses on the transfer of a proton from one particle to another, in any solvent. It also explains why ammonia is a base even though it contains no hydroxide: ammonia accepts a proton from water.
Lewis looks at electrons rather than protons. Any particle that can accept a pair of electrons counts as an acid, even with no hydrogen at all. Boron trifluoride (BF₃) and metal ions such as Fe³⁺ are Lewis acids.
For most school chemistry you’ll work with the first two. Our article on acid–base definitions compares all three in detail.
Strong and weak acids
Not every acid gives up its protons equally readily.
A strong acid ionises almost completely in water. In dilute hydrochloric acid, essentially every HCl molecule has handed its proton to water.
A weak acid only partly ionises. In vinegar, the great majority of ethanoic (acetic) acid molecules stay intact at any moment. Only a small fraction have released a proton, and the reaction runs both ways:
CH₃COOH ⇌ CH₃COO⁻ + H⁺
“Strong” and “concentrated” mean different things. A strong acid can be very dilute, and a weak acid can be very concentrated. Strength describes how completely the acid ionises; concentration describes how much acid is dissolved in a given volume. We explain the difference fully in strong vs weak acids.
How acids are named
A few patterns help you recognise acids from their formulas.
- Binary acids contain hydrogen and one other non-metal: HCl (hydrochloric acid), HBr (hydrobromic acid), HF (hydrofluoric acid). The name uses hydro- plus the non-metal root plus -ic.
- Oxyacids contain hydrogen, oxygen and another element: H₂SO₄ (sulfuric acid), HNO₃ (nitric acid), H₃PO₄ (phosphoric acid). Their names follow the polyatomic ion: sulfate → sulfuric acid, nitrite → nitrous acid.
- Carboxylic acids are organic acids containing the –COOH group: methanoic (formic) acid, ethanoic (acetic) acid, citric acid.
Not every hydrogen in a formula is acidic. Ethanoic acid, CH₃COOH, has four hydrogen atoms but only one of them, the one attached to oxygen, is released as H⁺. The three hydrogens bonded to carbon stay put.
Acids you meet every day
| Acid | Where you find it | Approximate pH |
|---|---|---|
| Hydrochloric acid | Stomach juice | 1.5–3.5 |
| Citric acid | Lemons, limes, sweets | Lemon juice ≈ 2–2.6 |
| Ethanoic (acetic) acid | Vinegar | ≈ 2.4–3 |
| Phosphoric acid | Cola drinks | ≈ 2.5 |
| Carbonic acid | Fizzy water | ≈ 4 |
| Lactic acid | Yoghurt, sour milk | ≈ 4–4.6 |
| Sulfuric acid | Car batteries | below 1 |
| Ascorbic acid | Vitamin C | depends on concentration |
Your own body depends on acids too. The hydrochloric acid in your stomach helps unfold proteins so enzymes can break them down, and it kills many of the microbes you swallow with food. Your stomach lining protects itself with a layer of mucus and bicarbonate.
Staying safe with acids
Concentrated strong acids are corrosive. They damage skin and eyes, and they can react violently when mixed with water carelessly. Two rules matter in any lab:
- Wear eye protection whenever acids are in use.
- When diluting, add acid to water, never water to acid. Dilution releases a lot of heat, and adding a little water to a lot of concentrated acid can make the mixture boil and spit.
Even weak acids deserve respect. Hydrofluoric acid is a weak acid, but it’s extremely dangerous because the fluoride ion passes through skin and attacks tissue underneath.
Common misconceptions
- “All acids are dangerous.” Plenty of acids are safe to eat. Citric, malic, lactic and ascorbic acids are in everyday food.
- “Acids burn because they’re hot.” Corrosive damage is chemical, not thermal. It happens at room temperature.
- “A low pH means a strong acid.” pH depends on concentration as well as strength. A concentrated weak acid can have a lower pH than a very dilute strong acid.
- “Every compound with H in it is an acid.” Methane (CH₄) and water are not acids in the everyday sense. What matters is whether the hydrogen can leave as H⁺.
Key takeaways
- An acid releases H⁺ ions (protons) in water. In solution these exist as H₃O⁺.
- The familiar properties of acids come from H⁺: sour taste, red litmus, fizzing with metals and carbonates, neutralising bases, and a pH below 7.
- Chemists use three definitions: Arrhenius (H⁺ producer), Brønsted–Lowry (proton donor) and Lewis (electron-pair acceptor).
- Strong acids ionise completely; weak acids only partly. Strength and concentration are different ideas.
- To work out the pH of an acidic solution, try our pH calculator.
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