What is a Salt?
In chemistry, a salt is an ionic compound made of:
- Positive ions (cations) – often from a metal or the ammonium ion (NH₄⁺)
- Negative ions (anions) – usually from an acid
These ions are held together by electrostatic attraction (ionic bonding) in a regular giant lattice.
Key idea: A salt is formed when the H⁺ ions of an acid are replaced by metal ions or ammonium ions.
- Sodium chloride, NaCl – table salt (from HCl and NaOH or Na metal).
- Potassium nitrate, KNO₃ – fertiliser (from HNO₃ and KOH).
- Calcium sulfate, CaSO₄ – plaster, gypsum (from H₂SO₄ and CaCO₃ or CaO).
- Ammonium chloride, NH₄Cl – in some fertilisers and dry cells (from NH₃ and HCl).
Salts in Solution – Neutral, Acidic or Basic?
- When a salt dissolves in water, its ions interact with water.
- Depending on the strength of the acid and base it came from, the solution can be:
- Neutral – salt of strong acid + strong base (e.g. NaCl, KNO₃).
- Basic – salt of weak acid + strong base (e.g. Na₂CO₃, CH₃COONa).
- Acidic – salt of strong acid + weak base (e.g. NH₄Cl).
- For MYP level, remember:
- Salts from strong acid + strong base → usually neutral solutions.
- Salts from weak acid + strong base → basic solutions.
- Salts from strong acid + weak base → acidic solutions.
How are Salts Formed?
Acid-Base Neutralization
- A neutralization reaction between an acid and a soluble base (alkali) produces salt and water.
- General word equation: $$\text{Acid} + \text{Base} \to \text{Salt} + \text{Water}$$
Sodium sulfate from sulfuric acid and sodium hydroxide:
$$\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq})+2 \mathrm{NaOH}(\mathrm{aq}) \rightarrow \mathrm{Na}_2 \mathrm{SO}_4(\mathrm{aq})+2 \mathrm{H}_2 \mathrm{O}(\mathrm{l})$$
- This method is used especially when both reactants and the salt are soluble.
- To get the solid salt, you evaporate water and allow crystals to form (crystallisation).
Metal and Acid Reaction
- Acids can react with reactive metals to form salt and hydrogen gas.
- General word equation: $$\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen gas}$$
Magnesium and hydrochloric acid:
$$\text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g)$$
This method is suitable for:
- Metals above hydrogen in the reactivity series (e.g. Mg, Zn, Fe).
- Not suitable for very reactive metals (too dangerous) or unreactive metals (no reaction).
To prepare a salt:
- Use excess metal to ensure all acid is used up.
- Filter off unreacted metal.
- Evaporate the solution to obtain salt crystals.
Acid + Insoluble Base (Metal Oxide or Metal Carbonate)
- Some bases are insoluble in water (e.g. CuO, CaCO₃), but they still neutralize acids.
- General word equation: $$\text{Metal oxide} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$$
Copper(II) oxide and sulfuric acid:
$$\text{CuO}(s) + 2\text{HCl}(aq) \rightarrow \text{CuCl}_2(aq) + \text{H}_2\text{O}(l)$$
Calcium carbonate and hydrochloric acid:
$$\mathrm{CaCO}_3(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{CaCl}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})$$
Method (for insoluble base/carbonate + acid):
- Warm the acid gently.
- Add insoluble base or carbonate in excess until no more dissolves / reaction stops.
- Filter to remove the excess solid.
- Evaporate the filtrate to form crystals.
This is especially useful for preparing salts of metals that don’t react directly with acids, like copper.
Ammonia and Acid Reactions
- Ammonia solution (NH₃(aq)) is a weak base.
- It reacts with acids to form ammonium salts.
Ammonium chloride:
$$\mathrm{NH}_3(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \rightarrow \mathrm{NH}_4 \mathrm{Cl}(\mathrm{aq})$$
Ammonium salts (e.g. NH₄Cl, NH₄NO₃) are ionic solids that form crystals at room temperature but may decompose on strong heating, so they are usually crystallised by gentle evaporation (no strong heating).
Precipitation Reaction: Insoluble Salts
If the salt you want is insoluble in water, you can make it by mixing two soluble salts so that an insoluble product (precipitate) forms.
Soluble salt A + soluble salt B → insoluble salt C (precipitate) + soluble salt D
To prepare barium sulfate (BaSO₄):
- Mix solutions of barium chloride (BaCl₂) and sodium sulfate (Na₂SO₄): $$\text{BaCl}_2(aq) + \text{Na}_2\text{SO}_4(aq) \rightarrow \text{BaSO}_4(s) + 2\text{NaCl}(aq)$$
- The insoluble BaSO₄ precipitates out.
- Filter, wash and dry to obtain pure solid BaSO₄.
Choosing the Best Method to Prepare a Salt
To decide which method to use, ask:
- Is the salt soluble or insoluble in water?
- Soluble salt → use neutralisation (acid + soluble base) or metal/insoluble base + acid, then crystallisation.
- Insoluble salt → use precipitation (two soluble salts).
- Is the base soluble?
- Soluble base (alkali) (e.g. NaOH, KOH) → use titration for accurate neutralisation.
- Insoluble base (e.g. CuO, CaCO₃) → add in excess, then filter.
- Is the metal reactive with acid?
- If yes (e.g. Mg, Zn, Fe) → can use metal + acid.
- If not (e.g. Cu) → use oxide or carbonate + acid instead.
- Soluble salt + both reactants soluble → titration (acid + alkali).
- Soluble salt + insoluble base/carbonate → warm acid, add excess solid, filter, crystallise.
- Insoluble salt → precipitation from two soluble solutions.
Properties of Salts and Their Uses
Salts are not just abstract particles – their physical properties determine where they are used in real life.
Crystal Structure and Shape
- Most salts are crystalline solids with a regular lattice structure.
- NaCl forms cubic crystals.
- CuSO₄·5H₂O forms bright blue crystals.
- Some salts are hydrated (contain water molecules in their crystal structure), others are anhydrous.
- Uses linked to crystals:
- Copper(II) sulfate crystals – used in education to show crystallisation and as a fungicide.
- Gypsum (CaSO₄·2H₂O) – used in plaster and construction because it sets in a controlled way.
Solubility
- Many salts (e.g. NaCl, KNO₃, Na₂SO₄) are highly soluble in water.
- Some are sparingly soluble or insoluble, like BaSO₄ or AgCl.
Uses:
- Soluble salts as fertilisers:
- KNO₃, NH₄NO₃, (NH₄)₂SO₄ dissolve easily in soil water and supply plants with nitrate and potassium.
- Insoluble salts in medicine:
- Barium sulfate (BaSO₄) is used as a contrast agent in X-ray imaging of the digestive system.
- It is opaque to X-rays and insoluble, so it passes through the body without dissolving.
- De-icing salts:
- Salts like NaCl or CaCl₂ are spread on roads in winter because they are soluble and lower the freezing point of water.
Electrical Conductivity
- In the solid state, salts do not conduct electricity because the ions are fixed in the lattice.
- When molten or dissolved in water, the ions are free to move, so they conduct electricity.
Uses:
- Electrolytes in the human body (e.g. Na⁺, K⁺, Cl⁻) allow nerve impulses and muscle function.
- Salts in electrolysis, e.g. molten NaCl used to produce sodium metal and chlorine gas.
Thermal Stability
- Some salts decompose on heating:
- Many carbonates (except Group 1) decompose to oxides and CO₂.
- Ammonium salts (e.g. NH₄Cl) decompose on strong heating to give gases (e.g. NH₃ and HCl).
- This affects:
- How they are prepared (e.g. ammonium salts crystallised by gentle evaporation, not strong heating).
- Their storage and industrial use.
- What is a salt in chemical terms? Explain where the cation and anion usually come from.
- List three methods to prepare a salt and give an example reaction for each.
- Which method would you use to prepare an insoluble salt like BaSO₄, and why?
- Why are salts overall neutral in charge even though they contain ions?
- Explain why Na₂CO₃ solution is basic and NH₄Cl solution is acidic.
- Give two examples of how solubility affects the use of salts in everyday life.
- Give one example where the crystalline nature of salts is important in practice.
