Salts
Salts: definitions, formulae, solubility rules, precipitate prediction and selecting excess-solid, titration or precipitation methods.
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This lesson connects three decisions: identify the ions in a salt, balance their charges to write its formula, then use solubility rules to predict whether it dissolves or forms a precipitate. Those decisions also guide the preparation methods in the following lessons.
1. Definition
A. Salt
A salt contains positive and negative ions. In the acid–base reactions studied here, replaceable hydrogen in an acid is replaced by a metal ion or an ammonium ion, NH₄ +. The resulting formula has no overall charge.
2. Key Ideas
- Salt names are made from:
- the cation (metal or ammonium), and
- the anion from the acid (chloride, sulfate, nitrate, ethanoate, carbonate, etc.).
The acid supplies the salt anion in these reactions:
| Acid | Formula | Salt name ends with | Anion |
|---|---|---|---|
| hydrochloric acid | HCl | chloride | Cl⁻ |
| nitric acid | HNO₃ | nitrate | NO₃⁻ |
| sulfuric acid | H₂SO₄ | sulfate | SO₄²⁻ |
| ethanoic acid | CH₃COOH | ethanoate | CH₃COO⁻ |
- Solubility rules let you predict whether a salt remains dissolved or forms a precipitate; practise applying the exceptions rather than guessing.
3. Detailed Explanations
- A salt contains positive and negative ions, with no overall charge.
- Its name combines the cation and anion names.
- Common nitrates, ammonium salts and Group 1 salts are generally soluble.
- Apply the named chloride, sulfate and carbonate exceptions in the table below.
A. Where the “two parts” come from
In many salt-making reactions:
- the cation comes from the metal/base, and
- the anion comes from the acid.
Example: KOH(aq) + HCl(aq) → KCl(aq) + H₂O(l)
Here, K⁺ comes from KOH and Cl⁻ comes from HCl, so the salt is potassium chloride, KCl.
B. Common ways salts are formed (O-Level)
| Reaction type | What you write | Example |
|---|---|---|
| acid + metal | salt + hydrogen | Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) |
| acid + metal oxide/hydroxide | salt + water | CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l) |
| acid + alkali | salt + water | NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l) |
| acid + carbonate | salt + water + carbon dioxide | CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) |
For the metal reaction, use a suitable metal above hydrogen with dilute hydrochloric or sulfuric acid. Nitric acid has different oxidising behaviour. The carbonate and hydroxide examples show named reagents, rather than a promise that every possible acid–solid pair is a suitable preparation route.
Do not write “acid + base → salt + water” as an absolute rule.
Ammonia forms ammonium salts without water:
NH₃(aq) + HCl(aq) → NH₄Cl(aq)
C. Solubility rules for common salts
Solubility rules describe how readily a substance dissolves in water. The table gives general school-level rules for common salts; apply any more specific solubility information supplied in a question.
“Insoluble” means very little dissolves, rather than absolutely none. Lead(II) chloride and calcium sulfate have low solubility; temperature and concentration can affect whether a visible solid forms.
A salt being soluble does not mean it can only be written as aqueous. Sodium chloride crystals are NaCl(s); when dissolved, they supply Na + (aq) and Cl⁻(aq). Choose state symbols for the sample actually described.
| Salt type | Usually soluble? | Common exceptions |
|---|---|---|
| nitrates (NO₃⁻) | yes | none among the common salts studied |
| ammonium salts (NH₄ +) | yes | none among the common salts studied |
| Group 1 salts (Li⁺, Na⁺, K⁺) | yes | none among the common salts studied |
| chlorides (Cl⁻) | yes | AgCl, PbCl₂ |
| sulfates (SO₄²⁻) | yes | BaSO₄, PbSO₄, CaSO₄ (sparingly soluble) |
| carbonates (CO₃²⁻) | no | Group 1 carbonates + ammonium carbonate are soluble |
| hydroxides (OH⁻) | no | Group 1 hydroxides are soluble; Ca(OH)₂ is sparingly soluble |
Solubility rules appear constantly in qualitative analysis (precipitate formation):
Qualitative Analysis (QA)
.
D. Choosing a preparation method
The best method depends on whether the salt is soluble or insoluble.
| Salt type | Preparation idea | Why |
|---|---|---|
| soluble salt | use excess insoluble solid + acid, then filter + crystallise | easy to remove excess solid |
| soluble salt (acid + alkali) | titration then crystallise | you cannot “filter out excess alkali” |
| insoluble salt | mix two solutions to form a precipitate, then filter + wash + dry | salt forms as a solid |
Use Preparing soluble salts for excess-solid and titration methods, or Preparing insoluble salts for precipitation and washing.
E. Keep the three representations separate
For the precipitation of silver chloride:
- Macroscopic: a white precipitate is seen.
- Particle level: aqueous Ag⁺ and Cl⁻ ions form an insoluble ionic lattice; spectator ions remain in solution.
- Symbolic: Ag + (aq) + Cl⁻(aq) → AgCl(s).
An observation describes what is seen. It should not be replaced by a particle explanation or an equation.
4. Common Mistakes
- Thinking “salt” means only NaCl.
- Using the wrong solubility exception (learn the chloride/sulfate/carbonate exceptions).
- Saying “do titration” for every salt (wrong): titration is specifically for soluble salts made from acid + alkali.
5. Exam Tips
State the solubility rule, then identify the precipitate, if any. For preparation, explain why each separation works: for example, an insoluble excess reagent can be removed by filtration while the required salt stays dissolved.
- If asked to predict a precipitate when two solutions are mixed: write the insoluble salt and label it (s).
- Always include state symbols if the question expects them.
6. Worked Examples
Modelled example 1
Identify the salt from the reactants
Problem
Study the worked solution
Identify the cation
Method
Take sodium ions from sodium hydroxide.Reason
The base supplies the positive ion in the salt.Working
Cation: Na⁺.Identify the acid anion
Method
Take sulfate ions from sulfuric acid.Reason
The acid determines the salt-family ending.Working
Anion: SO₄²⁻; salt family: sulfate.Balance charges and name
Method
Combine two sodium ions with one sulfate ion.Reason
Total positive and negative charges must cancel.Working
Sodium sulfate, Na₂SO₄.
Guided practice 2
Predict if a precipitate forms (chloride exception)
Problem
Apply the exception, not only the general rule
Hints
Hint 1: chloride rule
Hint 2: combine the ions
View solution step by step
Identify the possible products
Method
Pair Ag⁺ with Cl⁻.Reason
Mixing the solutions brings those ions together.Working
Candidate product: AgCl.Apply the exception
Method
Classify silver chloride as insoluble.Reason
It is an exception to the usual solubility of chlorides.Working
Silver chloride, AgCl(s), precipitates.
Common misconception 3
Carbonate rule
Learner claim
Check the allowed exceptions
View solution step by step
State the carbonate rule
Method
Begin with carbonates being insoluble.Reason
Only Group 1 and ammonium carbonates are the stated soluble exceptions.Working
Carbonates: insoluble except Group 1 and NH₄ + salts.Test calcium against the exceptions
Method
Place calcium in Group 2.Reason
It matches neither soluble exception.Working
CaCO₃ is insoluble.
Examiner practice 4
Sulfate exception
Examination question
Name and formulate the insoluble product
View solution step by step
Name the precipitate
1 markMethod
Combine barium and sulfate names.Reason
Those are the ions that form the insoluble product.Working
Barium sulfate.Give formula and state
1 markMethod
Write a 1:1 neutral formula with (s).Reason
Ba²⁺ and SO₄²⁻ balance, and a precipitate is solid.Working
BaSO₄(s).
Self-mark with the mark scheme
Compare your response with each mark point. Select a point only when your response contains that evidence.
Self-mark name and state-labelled formula.
Challenge 5
Choose the preparation method (acid + alkali)
Method-selection transfer
Pure sodium nitrate crystals are required from nitric acid and sodium hydroxide. Select the preparation method and justify it.
Use the separation constraint
Hints
Hint 1: can excess be filtered
Dissolved nitric acid and sodium hydroxide both pass through filter paper.
Hint 2: find exact volumes
Titration identifies the neutralising volumes before crystallisation.
View solution step by step
Identify the separation problem
Method
Recognise that both reactants are aqueous.Reason
An excess of either would remain dissolved and contaminate the salt solution.
Working
Excess acid/alkali cannot be removed by filtration.
Select and complete the method
Method
Use titration, repeat without indicator, then concentrate and crystallise.
Reason
Exact neutralising volumes leave no excess reactant in the soluble salt solution.
Working
Titration → repeat the reacting volumes without indicator → concentrate and crystallise NaNO₃.
7. Mind Stretchers
Mind stretcher 1: Fix the student’s claimExtension
Question: A student writes: “All chlorides are insoluble.” Correct this with one exception example.
Show answer
Chlorides are usually soluble. Exceptions include silver chloride, AgCl(s), and lead(II) chloride, PbCl₂(s).
Mind stretcher 2: Extraction logicExtension
Question: Why is “acid + excess insoluble base” a good method for preparing some soluble salts?
Show answer
Because the excess base is a solid and can be filtered off, leaving the salt solution to be crystallised, provided suitable clean reactants were used. This avoids excess acid contaminating the product.
Try independently: A sample labelled ammonium chloride is a soluble salt, and its measured pH is below 7. Is either property a contradiction? Write its formula and explain what the two observations tell you.
Show answer and reasoning
The formula is NH₄Cl: one NH₄ + balances one Cl⁻. Common ammonium salts are soluble, so dissolving is expected. “Salt” does not mean “neutral solution”: the measured pH shows this sample is acidic. Solubility and pH describe different properties.
8. Quiz
Practise naming, solubility and method selection in the Acid–Base Chemistry topic check.
Syllabus and review details
- SEC G3 Pure Chemistry 2027 · 2027
Content structure and subject content, PDF pages 9–24
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