What a Salt Is
A salt is the ionic compound formed when the replaceable hydrogen ion(s) of an acid are swapped out for a metal or ammonium ion — typically through neutralisation. Salts are electrically neutral overall: the numbers of cations and anions in the formula balance out.
Types of Salts
| Type | Description | Example |
|---|---|---|
| Normal salt | All the acid’s replaceable hydrogens are replaced. Neutral to litmus (pH 7) — though a few, like sodium carbonate, aluminium chloride and sodium sulphide, hydrolyse in water to give an acidic or alkaline solution anyway. | NaCl(aq) + H₂O(l), from HCl(aq) + NaOH(aq) →; also H₂SO₄(aq) + ZnO(s) → ZnSO₄(aq) + H₂O(l) |
| Acid salt | Only some of a di- or tri-basic acid’s hydrogens are replaced — formed when there’s not enough metal/base available. Turns blue litmus red. | H₂SO₄(aq) + KOH(aq) → KHSO₄(aq) + H₂O(l); H₃PO₄(aq) + NaOH(aq) → NaH₂PO₄(aq) + H₂O(l) |
| Basic salt | Contains leftover hydroxide ions — formed when there’s not enough acid to fully neutralise a base. Turns red litmus blue. | Ca(OH)₂(aq) + HCl(aq) → Ca(OH)Cl(aq) + H₂O(l) |
| Double salt | Formed from two salts combined; ionises in water to give three (or more) distinct kinds of ion. | Ammonium iron(II) sulphate hexahydrate, (NH₄)₂Fe(SO₄)₂·6H₂O; aluminium potassium sulphate dodecahydrate, KAl(SO₄)₂·12H₂O |
| Complex salt | Contains a complex ion — a central (usually transition) metal ion bonded to surrounding ligands. | Na₂Zn(OH)₄ (sodium tetrahydroxozincate(II)) → 2Na⁺(aq) + [Zn(OH)₄]²⁻(aq); K₄Fe(CN)₆ (potassium hexacyanoferrate(II)) → 4K⁺(aq) + [Fe(CN)₆]⁴⁻(aq) |
Solubility Rules
| Salt family | General rule | Exceptions (insoluble) |
|---|---|---|
| Nitrates | All soluble | None |
| Sulphates | All soluble | Lead(II) sulphate, barium sulphate, calcium sulphate |
| Chlorides | All soluble | Lead(II) chloride, silver chloride, mercury(II) chloride |
| Carbonates | All insoluble | Potassium, sodium and ammonium carbonate |
| Ammonium salts | All soluble | None |
Preparing Soluble Salts
- Acid + reactive metal (more reactive than hydrogen — calcium, magnesium, zinc, iron; sodium and potassium are too reactive to do safely this way, copper/mercury/silver/gold too unreactive to work at all): Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g); Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)
- Acid + alkali (neutralisation, by titration): HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
- Acid + insoluble base: warm the acid, add the base gradually until no more dissolves, then filter off the excess. 2HCl(aq) + CuO(s) → CuCl₂(aq) + H₂O(l)
- Acid + carbonate: add the carbonate gradually until it stops dissolving, then filter off the excess. ZnCO₃(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂O(l) + CO₂(g)
Recovering the dissolved salt: heating to dryness (only for salts that survive dry heat — most chlorides except ZnCl₂ and FeCl₃) leaves the salt behind as a residue in an evaporating dish; crystallisation (for salts that would decompose if fully dried — most nitrates, carbonates, sulphates) boils off some water, then cools the concentrated solution slowly (a scratch on the container, or a seed crystal, can help crystals start forming) so pure crystals can be filtered off, washed and dried between filter paper.
Preparing Insoluble Salts
Double decomposition (precipitation): mixing two soluble salts — one supplying the wanted cation, the other the wanted anion — precipitates the insoluble salt directly, recovered by filtration.
NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s); K₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2KCl(aq)
Solubility and Solubility Curves
Solubility is the mass (in grams) of a solute that will saturate 100g of solvent at a given temperature. A saturated solution holds as much dissolved solute as it can at that temperature. Plotting solubility against temperature gives a solubility curve.
Most ionic solids (e.g. potassium nitrate) get more soluble as temperature rises — sometimes sharply (KNO₃), sometimes only gradually (NaCl, whose solubility barely changes with temperature). A few show more complex behaviour: sodium sulphate’s solubility rises up to about 36°C, then drops sharply, because below 36°C the stable form is a hydrate, while above it, the anhydrous salt is stable instead — and the two forms simply have different solubilities.
Using solubility curves to separate mixtures: if two dissolved salts have very different solubility-temperature curves, cooling a hot, saturated mixed solution lets the less soluble one crystallise out first (removable by filtration) while the more soluble one stays in solution — a technique called fractional crystallisation. This is how sodium chlorate(V) (used in weedkillers and matches, made by reacting NaOH with chlorine gas, which also produces NaCl as a by-product) is separated from the sodium chloride formed alongside it: sodium chloride’s solubility barely changes with temperature, while sodium chlorate(V)’s rises much more steeply, so cooling a hot saturated mixture crystallises out the chloride first, leaving the chlorate(V) to be recovered afterward.
Solubility curves are also useful more generally for choosing the right solvent for a substance at a given temperature — including, in the pharmaceutical industry, for working out suitable solvents and dosing for particular drugs.
Action of Heat on Salts
Carbonates: potassium, sodium and ammonium carbonate are the only common exceptions — everything else decomposes to the metal oxide and CO₂ on heating (mercury and silver carbonate go further, all the way to the free metal).
| Carbonate | Decomposition |
|---|---|
| K₂CO₃, Na₂CO₃ | Does not decompose |
| CaCO₃, MgCO₃, ZnCO₃, PbCO₃, CuCO₃ | → metal oxide + CO₂ (e.g. CaCO₃ → CaO + CO₂) |
| Al₂(CO₃)₃, Fe₂(CO₃)₃ | → 2 metal oxide + 3CO₂ (e.g. Al₂(CO₃)₃ → Al₂O₃ + 3CO₂) |
| 2HgCO₃ | → 2Hg + 2CO₂ + O₂ |
| 2Ag₂CO₃ | → 4Ag + 2CO₂ + O₂ |
| (NH₄)₂CO₃ | → 2NH₃ + CO₂ + H₂O |
Nitrates: all decompose on heating, but into different products depending on the metal’s reactivity (the same three-tier pattern already established under nitrogen chemistry): the very reactive Group I metals (except lithium) give the nitrite plus oxygen; moderately reactive metals give the oxide plus NO₂ and oxygen; the least reactive metals (mercury, silver) give the free metal plus NO₂ and oxygen; ammonium nitrate is the odd one out, giving dinitrogen oxide and water.
| Nitrate | Decomposition |
|---|---|
| 2KNO₃, 2NaNO₃ | → 2 nitrite + O₂ |
| 2Ca(NO₃)₂, 2Mg(NO₃)₂, 2Zn(NO₃)₂, 2Cu(NO₃)₂, 2Pb(NO₃)₂ | → 2 oxide + 4NO₂ + O₂ |
| 4Al(NO₃)₃, 4Fe(NO₃)₃ | → 2 oxide (M₂O₃) + 12NO₂ + 3O₂ |
| Hg(NO₃)₂ | → Hg + 2NO₂ + O₂ |
| 2AgNO₃ | → 2Ag + 2NO₂ + O₂ |
| NH₄NO₃ | → N₂O + 2H₂O |
Sulphates: most (sodium, potassium, calcium) are thermally stable and don’t decompose at all. Among the ones that do:
| Sulphate | Decomposition |
|---|---|
| ZnSO₄, CuSO₄ | → metal oxide + SO₃ |
| Fe₂(SO₄)₃ | → Fe₂O₃ + 3SO₃ |
| 2FeSO₄·7H₂O | → Fe₂O₃ + SO₂ + SO₃ + 14H₂O (here the iron is also being oxidised from Fe²⁺ to Fe³⁺, which is why — unlike the anhydrous examples above — a reducing SO₂ by-product appears alongside the SO₃) |
| (NH₄)₂SO₄ | → 2NH₃ + H₂SO₄ |
Chlorides: all are thermally stable except ammonium chloride, which decomposes reversibly: NH₄Cl(s) → NH₃(g) + HCl(g).
Efflorescence, Deliquescence and Hygroscopy
| Term | Meaning | Examples |
|---|---|---|
| Efflorescent | Loses (some or all of) its water of crystallisation to the air at room temperature, forming a lower hydrate or the anhydrous salt | Washing soda, Na₂CO₃·10H₂O, loses 9 of its 10 water molecules: Na₂CO₃·10H₂O(s) → Na₂CO₃·H₂O(s) + 9H₂O(g). Na₂SO₄·10H₂O loses all of its water. |
| Deliquescent | Absorbs so much moisture from the air that it dissolves into a solution | NaOH, KOH, CaCl₂, MgCl₂, FeCl₃, P₄O₁₀ |
| Hygroscopic | Absorbs moisture from the air without fully dissolving (solids become damp/sticky; concentrated sulphuric acid, a hygroscopic liquid, can dilute itself to roughly three times its original volume this way) | CaO, NaNO₃, CuO; concentrated H₂SO₄ |
Hydrated salts contain water of crystallisation loosely bound in their structure (e.g. CuSO₄·5H₂O, FeSO₄·7H₂O), lost on heating; a salt without any is anhydrous. CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(g)
Drying Agents (for Gases)
A drying agent must not react with the gas it’s meant to dry.
| Drying agent | Suitable for |
|---|---|
| Concentrated sulphuric acid | All gases except ammonia and hydrogen sulphide (both react with it) |
| Fused calcium chloride | All gases except ammonia (forms a complex with it) |
| Calcium oxide (quicklime) | Ammonia specifically |
| Phosphorus(V) oxide | All gases except ammonia |
| Silica gel | All gases (a purely physical desiccant, unreactive) |
