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To most people, “salt” means the white crystals on the dinner table. To a chemist, table salt (sodium chloride) is just one member of an enormous family. Chalk is a salt. So is the copper sulfate that makes a blue solution in the lab, the potassium nitrate in fertiliser, the calcium phosphate in your bones and the baking soda in your cupboard.
This article explains what chemists mean by a salt, how salts form, and why they behave the way they do.
The chemical definition
A salt is an ionic compound made of positive ions (cations) and negative ions (anions), other than hydrogen ions (H⁺) and hydroxide ions (OH⁻) as the only ions present.
In practical terms, a salt is the compound formed when the hydrogen of an acid is replaced by a metal ion or an ammonium ion.
| Acid | Hydrogen replaced by | Salt |
|---|---|---|
| HCl (hydrochloric acid) | Na⁺ | NaCl (sodium chloride) |
| H₂SO₄ (sulfuric acid) | Cu²⁺ | CuSO₄ (copper(II) sulfate) |
| HNO₃ (nitric acid) | NH₄⁺ | NH₄NO₃ (ammonium nitrate) |
| H₃PO₄ (phosphoric acid) | Ca²⁺ | Ca₃(PO₄)₂ (calcium phosphate) |
| CH₃COOH (ethanoic acid) | Na⁺ | CH₃COONa (sodium ethanoate) |
The cation usually comes from a base (or a metal), and the anion comes from an acid.
How salts form
There are several reactions that make salts. They’re the backbone of practical inorganic chemistry.
1. Acid + base (neutralisation) HCl + NaOH → NaCl + H₂O
2. Acid + metal oxide H₂SO₄ + CuO → CuSO₄ + H₂O
3. Acid + metal carbonate 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂
4. Acid + reactive metal Mg + 2HCl → MgCl₂ + H₂
5. Acid + ammonia NH₃ + HCl → NH₄Cl
6. Direct combination of elements 2Na + Cl₂ → 2NaCl
7. Precipitation (two soluble salts swapping partners) AgNO₃ + NaCl → AgCl(s) + NaNO₃
For a full guide to the neutralisation reactions, see acids vs bases.
Structure: the ionic lattice
In the solid state, a salt isn’t made of separate molecules. Its ions are arranged in a huge, repeating three-dimensional lattice, each ion surrounded by ions of opposite charge. In sodium chloride, every Na⁺ is surrounded by six Cl⁻ ions, and every Cl⁻ by six Na⁺.
This is why the formula NaCl is a formula unit, not a molecule: it tells you the ratio of ions (1 : 1), not a discrete particle. For more on bonding in salts, see ionic vs covalent bonds.
Properties of salts
The strong electrostatic attraction between oppositely charged ions explains most salt properties:
| Property | Typical behaviour | Why |
|---|---|---|
| Melting point | high (NaCl: 801 °C) | lots of energy needed to break strong ionic bonds throughout the lattice |
| State at room temperature | crystalline solid | ions locked in a regular lattice |
| Hardness | hard but brittle | shifting a layer brings like charges together, which repel and split the crystal |
| Conductivity as a solid | none | ions can’t move |
| Conductivity when molten or dissolved | good | ions are free to move and carry charge |
| Solubility | many dissolve in water, many don’t | depends on the balance between lattice energy and hydration energy |
Salts dissolved in water are electrolytes. That’s why sports drinks contain salts, why seawater conducts electricity, and why molten salts can be electrolysed to extract reactive metals such as sodium and aluminium.
Soluble and insoluble salts
Salts vary hugely in how well they dissolve. A few simple rules cover most common cases:
- All sodium, potassium and ammonium salts are soluble.
- All nitrates are soluble.
- Most chlorides are soluble, except silver chloride and lead(II) chloride.
- Most sulfates are soluble, except barium sulfate, lead(II) sulfate and (only slightly soluble) calcium sulfate.
- Most carbonates are insoluble, except those of sodium, potassium and ammonium.
The full rules are in solubility rules.
Solubility decides how a salt is made in the lab:
- Soluble salts are made by reacting an acid with an excess of an insoluble base, carbonate or metal, filtering off the excess, then evaporating and crystallising.
- Insoluble salts are made by precipitation: mixing two solutions whose ions combine to form the insoluble product, then filtering, washing and drying it.
Hydrated salts
Many salts crystallise with water molecules built into their lattice, called water of crystallisation. Their formulas show this with a dot:
- CuSO₄·5H₂O: blue copper(II) sulfate pentahydrate
- Na₂CO₃·10H₂O: washing soda
- MgSO₄·7H₂O: Epsom salts
- CaSO₄·2H₂O: gypsum
Heating a hydrated salt drives off the water. Blue copper sulfate crystals turn to a white powder of anhydrous CuSO₄; adding water turns it blue again, which is used as a test for water. See hydrates and water of crystallisation.
Acidic, basic and neutral salts
Not all salts give neutral solutions. It depends on the acid and base they came from:
- Strong acid + strong base (NaCl, KNO₃): neutral.
- Weak acid + strong base (sodium ethanoate, sodium carbonate): basic.
- Strong acid + weak base (ammonium chloride): acidic.
The ions of weak acids and bases react with water. This is explained fully in salt hydrolysis.
Chemists also use “acid salt” for salts that still contain a replaceable hydrogen, such as sodium hydrogensulfate (NaHSO₄) and sodium hydrogencarbonate (NaHCO₃). Confusingly, an “acid salt” isn’t always acidic in solution: sodium hydrogencarbonate is mildly alkaline.
Salts you meet every day
| Salt | Formula | Where you find it |
|---|---|---|
| Sodium chloride | NaCl | table salt, seawater, road de-icer |
| Sodium hydrogencarbonate | NaHCO₃ | baking soda, antacids |
| Calcium carbonate | CaCO₃ | chalk, limestone, marble, eggshells |
| Calcium sulfate | CaSO₄·2H₂O | plaster, plasterboard |
| Magnesium sulfate | MgSO₄·7H₂O | Epsom bath salts |
| Potassium chloride | KCl | low-sodium salt substitute |
| Ammonium nitrate | NH₄NO₃ | fertiliser |
| Calcium phosphate | Ca₅(PO₄)₃OH | bones and tooth enamel (as hydroxyapatite) |
| Sodium fluoride | NaF | toothpaste |
| Copper(II) sulfate | CuSO₄·5H₂O | fungicide, school labs |
| Silver bromide | AgBr | traditional photographic film |
| Monosodium glutamate | C₅H₈NO₄Na | flavour enhancer |
Why is table salt called “salt”?
The word “salt” came first, from Latin sal, and chemists borrowed it. In the 18th century, as chemists began to understand that many substances formed from acids and bases resembled common salt (crystalline, soluble, often with a salty or bitter taste), they extended the word to the whole family. Guillaume-François Rouelle, a French chemist, is often credited with defining salts in 1744 as the products of acids combining with bases.
Common misconceptions
- “All salts taste salty.” Many taste bitter or sweetish, and many are toxic. Never taste lab chemicals.
- “All salts dissolve in water.” Silver chloride, barium sulfate and calcium carbonate barely dissolve at all.
- “Salt solutions are always neutral.” Sodium carbonate solution has a pH of about 11.
- “Salts are molecules.” They’re lattices of ions; the formula gives the ratio.
Key takeaways
- A salt is an ionic compound formed when the hydrogen of an acid is replaced by a metal or ammonium ion.
- Salts form from acids reacting with bases, oxides, carbonates, metals or ammonia, and by precipitation.
- Their ionic lattices give high melting points, brittleness and conductivity when molten or dissolved.
- Solubility varies; it decides whether a salt is made by crystallisation or precipitation.
- Learn to name them in naming salts.
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