Preparation of SaltsCambridge IGCSE Chemistry: Revision notes
Section 1
How are soluble salts prepared from acids and alkalis?
Soluble salts can be prepared by reacting an acid with an alkali. This is a neutralisation reaction where hydrogen ions from the acid react with hydroxide ions from the alkali to form water and a salt.
The general equation is: Acid + Alkali → Salt + Water
For example: HCl + NaOH → NaCl + H₂O
Titration method:
- Use a burette to dispense a precise volume of acid (or alkali) into the flask
- Add an indicator (such as phenolphthalein) to show when neutralisation is complete
- The endpoint is when the indicator changes colour, showing acid and alkali have reacted in exact proportions
- Record the volume of acid used
- Repeat the process without indicator to produce the salt solution
- Evaporate the solution to crystallise the salt
Why use titration?
- Ensures exact stoichiometric proportions (no excess acid or alkali remains)
- Produces pure salt crystals with no impurities
- Only suitable for preparing soluble salts
Examiners want to see that you understand titration produces no excess acid or alkali, making it ideal for pure salt preparation. State that the indicator change marks the endpoint, not that the solution is 'neutral'.
To prepare copper chloride: CuCl₂ + 2HCl would be wrong. Instead: CuCl₂ + 2HCl → CuCl₂ + H₂O is incorrect. The correct reaction is 2HCl + Cu(OH)₂ → CuCl₂ + 2H₂O (acid + base). Using titration ensures no excess acid remains in the final salt.
Section 2
How are soluble salts prepared using excess metal or base?
When an excess of a reactant is used, the acid is completely consumed. Excess reactant must then be removed by filtration and washing before crystallisation.
Preparation using excess metal:
- Add excess metal powder (such as zinc or iron) to dilute acid
- The metal reacts until all acid is consumed
- Filter to remove excess unreacted metal
- Wash the residue with water to remove dissolved salt
- Evaporate the filtrate to crystallise the salt
Example: Zn + H₂SO₄ → ZnSO₄ + H₂
Preparation using excess insoluble base:
- Add excess insoluble base (such as copper oxide or zinc oxide) to dilute acid
- Stir frequently; the base reacts until all acid is consumed
- Filter to remove excess unreacted base
- Wash the filter paper and residue with distilled water
- Evaporate the filtrate to crystallise the salt
Example: CuO + 2HCl → CuCl₂ + H₂O
Preparation using excess insoluble carbonate:
- Add excess carbonate powder (such as calcium carbonate) to dilute acid
- Stir; carbon dioxide gas is produced and the carbonate reacts until all acid is consumed
- Filter to remove excess unreacted carbonate
- Wash the filter paper and residue with distilled water
- Evaporate the filtrate to crystallise the salt
Example: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Key difference from titration: Using excess reactant is simpler as no careful volume measurement is needed, but produces impure salt if not washed thoroughly.
Students often forget to wash the residue after filtration. If you don't wash, dissolved salt from the filtrate remains on the filter paper and is lost, reducing your yield and purity.
When describing excess metal or base methods, examiners expect you to mention filtration to remove unreacted solid, then washing with distilled water, then evaporation to crystallise. These three steps are essential for a complete answer.
Using excess reactant is like adding more washing powder than needed to clean clothes—you know all dirt will be removed, but you need to rinse thoroughly to wash away the extra powder (excess reactant).
Section 3
What are the solubility rules for ionic compounds?
Solubility rules determine which salts are soluble and which are insoluble in water. These rules are essential for predicting whether a salt can be prepared by dissolving or will precipitate.
| Compound Type | Solubility | Exceptions |
|---|---|---|
| Sodium salts | Soluble | None |
| Potassium salts | Soluble | None |
| Ammonium salts | Soluble | None |
| Nitrates | Soluble | None |
| Chlorides | Soluble | Lead chloride (PbCl₂), Silver chloride (AgCl) |
| Sulfates | Soluble | Barium sulfate (BaSO₄), Calcium sulfate (CaSO₄), Lead sulfate (PbSO₄) |
| Carbonates | Insoluble | Sodium carbonate (Na₂CO₃), Potassium carbonate (K₂CO₃), Ammonium carbonate ((NH₄)₂CO₃) |
| Hydroxides | Insoluble | Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Ammonium hydroxide (NH₃·H₂O), Calcium hydroxide (Ca(OH)₂ — slightly soluble) |
How to use these rules:
- To prepare a soluble salt (like sodium chloride or potassium nitrate), use methods that keep all products soluble
- To prepare an insoluble salt (like barium sulfate or silver chloride), use precipitation
- When an acid reacts with a carbonate, check if the salt formed is soluble before deciding which preparation method to use
The three groups that are always soluble (sodium, potassium, ammonium, nitrates, chlorides) should be memorised as a set. For all other anions (sulfates, carbonates, hydroxides), learn the exceptions.
Is barium carbonate soluble? Carbonates are insoluble except Na₂CO₃, K₂CO₃, and (NH₄)₂CO₃. Barium is not in the exceptions, so BaCO₃ is insoluble and will precipitate.
Section 4
How are insoluble salts prepared by precipitation?
Precipitation is used to prepare insoluble salts that cannot be obtained by crystallising from solution. Two soluble salts are mixed, and an insoluble salt forms immediately as a solid.
General method:
- Prepare two solutions of soluble salts that will react to form the insoluble salt
- Mix the two solutions slowly, stirring
- The insoluble salt precipitates immediately as a solid
- Filter the mixture to collect the precipitate (insoluble salt)
- Wash the residue thoroughly with distilled water to remove impurities
- Dry the salt by pressing between filter paper or heating gently
Example reactions:
-
Silver nitrate + Sodium chloride → Silver chloride (white precipitate) + Sodium nitrate AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
-
Barium chloride + Sodium sulfate → Barium sulfate (white precipitate) + Sodium chloride BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
-
Lead nitrate + Potassium iodide → Lead iodide (yellow precipitate) + Potassium nitrate Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
Why use precipitation?
- Insoluble salts cannot be obtained by crystallising because they don't dissolve in water
- Both reactants must be soluble so they can be mixed as solutions
- The reaction is rapid and complete
- The insoluble salt is easily separated by filtration
Examiners want to see that you choose two soluble salts whose anions and cations recombine to form an insoluble product. Always check solubility rules to confirm both starting salts are soluble and the product is insoluble.
Students sometimes try to prepare insoluble salts by evaporating solutions. This will not work because insoluble salts don't dissolve in water—they will precipitate immediately when solutions are mixed, not crystallise on heating.
To prepare barium sulfate: Mix BaCl₂ solution + Na₂SO₄ solution. Both are soluble (Ba²⁺ and SO₄²⁻ are from the soluble salt rules). BaSO₄ is insoluble (sulfates insoluble except Ba, Ca, Pb are exceptions that are insoluble). A white precipitate forms immediately.
Section 5
What is water of crystallisation and how does it differ from anhydrous compounds?
Hydrated salts contain water of crystallisation—water molecules chemically bonded within the crystal structure. Anhydrous salts contain no water.
Key definitions:
- Hydrated substance: A compound that contains water molecules chemically combined within its structure. The water is written in the formula as a dot followed by the number of water molecules.
- Anhydrous substance: A compound that contains no water molecules.
- Water of crystallisation: Specific water molecules present in hydrated crystals that are part of the crystal structure.
Common examples:
| Hydrated Formula | Name | Number of Water Molecules |
|---|---|---|
| CuSO₄·5H₂O | Copper sulfate pentahydrate | 5 |
| CoCl₂·6H₂O | Cobalt chloride hexahydrate | 6 |
| MgSO₄·7H₂O | Magnesium sulfate heptahydrate | 7 |
| Na₂CO₃·10H₂O | Sodium carbonate decahydrate | 10 |
Properties:
- Hydrated salts often have coloured crystals (e.g. CuSO₄·5H₂O is blue)
- Anhydrous salts are often white (e.g. CuSO₄ is white)
- The water is not just absorbed on the surface but is part of the crystal lattice
- Heating hydrated salts drives off the water, converting them to anhydrous form
- The water is written in the formula to show the exact molar ratio (e.g. one mole of CuSO₄·5H₂O contains 5 moles of H₂O)
Why this matters in salt preparation:
- When you crystallise a salt solution, water of crystallisation may be incorporated into the crystals
- The mass of the final product includes the mass of water
- This must be considered when calculating the theoretical yield
When describing the difference between hydrated and anhydrous, examiners expect you to state that hydrated compounds have water chemically combined in the crystal structure, not just water absorbed on the surface. Include an example like CuSO₄·5H₂O (blue) vs anhydrous CuSO₄ (white).
Copper sulfate pentahydrate (CuSO₄·5H₂O) contains 5 water molecules per formula unit. If you heat it gently, the water is driven off and you obtain anhydrous copper sulfate (CuSO₄), a white powder. The blue colour is due to the hydrated form.
Must Know
- Titration produces pure soluble salts with no excess acid or alkali because reactants are added in exact stoichiometric proportions. Use an indicator to mark the endpoint.
- Excess metal, base, or carbonate methods require filtration to remove unreacted solid, washing with distilled water to remove impurities, then evaporation to crystallise the salt.
- Solubility rules: Sodium, potassium, ammonium salts, nitrates, and chlorides are soluble; exceptions are PbCl₂ and AgCl. Sulfates are soluble except BaSO₄, CaSO₄, and PbSO₄. Carbonates and hydroxides are insoluble except Na₂CO₃, K₂CO₃, (NH₄)₂CO₃, NaOH, KOH, NH₃·H₂O, and Ca(OH)₂ (slightly soluble).
- Precipitation prepares insoluble salts by mixing two soluble ionic solutions; both starting salts must be soluble and the product must be insoluble.
- Hydrated salts contain water of crystallisation chemically bonded in the crystal structure (e.g. CuSO₄·5H₂O); anhydrous salts contain no water. Hydrated salts are often coloured; anhydrous forms are often white.
- Use solubility rules to identify which salts are soluble (prepare by crystallisation from solution) and which are insoluble (prepare by precipitation).
That's the notes covered.
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