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2.7Inorganic Chemistry

Acids, Bases and Salt Preparations — Edexcel Linear - OL (Chemistry) Revision Notes & Topical Questions

10 min read 14 parts Lesson 14 of 23 Unit quiz

📚 Edexcel Linear - OL (Chemistry) Hub / Topical Revision
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Curriculum Equation Key معادلة تفاعل تكوين الإستر (الأسترة) - Esterification Reaction EquationFast Exam Solution
R-COOH + R2-OH ⇌ (conc. H₂SO₄) R-COO-R2 + H₂O
Word Equation & Conditions: حمض كربوكسيلي + كحول (في وجود حمض الكبريتيك المركز كعامل حفاز ونازع للماء) ⇌ إستر + ماء
Location & Specifications: الكيمياء العضوية للثانوية العامة والتحصيلي و Edexcel Chemistry.
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Lesson summary

Acids, Bases and Making Soluble Salts — EDEXCEL LINEAR – OL (CHEMISTRY) notes: Write equations for the reactions of acids with metals, bases and carbonates and describe how to prepare a pure dry sample of a soluble salt.A salt forms when the hydrogen of an acid is replaced by a m…

  • Worked example 1
  • Worked example 2
  • Worked example 3
  • Worked example 4
  • Crystallisation
  • Effervescence

Section 1 of 14Reactions of acids

A salt forms when the hydrogen of an acid is replaced by a metal or ammonium ion. The salt name has two parts: the metal comes from the base or metal and the second part from the acid (hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate).

1
Acid + metal → salt + hydrogen: Zn + H2SO4 → ZnSO4 + H2 (only for metals above hydrogen, and not too reactive).
2
Acid + metal oxide (base) → salt + water: CuO + H2SO4 → CuSO4 + H2O.
3
Acid + metal hydroxide / alkali → salt + water: NaOH + HCl → NaCl + H2O.
4
Acid + metal carbonate → salt + water + carbon dioxide: CaCO3 + 2HCl → CaCl2 + H2O + CO2.
5
Acid + ammonia → ammonium salt: NH3 + HNO3 → NH4NO3, used as a fertiliser.

A base is a substance that neutralises an acid; an alkali is a base that dissolves in water. Carbonates fizz because CO2 is released, and the gas turns limewater milky.

Section 2 of 14Solubility rules

Deciding the preparation method depends on solubility: all sodium, potassium and ammonium salts are soluble; all nitrates are soluble; chlorides are soluble except silver and lead(II); sulfates are soluble except lead(II), barium and calcium; carbonates and hydroxides are insoluble except those of sodium, potassium and ammonium (calcium hydroxide is sparingly soluble).

Section 3 of 14Preparing a soluble salt from an insoluble base or metal

Examplecopper(II) sulfate crystals from copper(II) oxide and sulfuric acid.

Warm about 50 cm3 of dilute sulfuric acid gently in a beaker (do not boil).
Excess
Add copper(II) oxide a little at a time, stirring, until it is in excess (black solid stops dissolving) so that all the acid is used up.
Filter off the excess oxide into an evaporating basin; the filtrate is blue copper(II) sulfate solution.
Heat to evaporate about half the water until the crystallisation point (a drop on a glass rod forms crystals when cooled),
  • then leave to cool slowly so crystals form.
Filter off crystals, wash with a little cold distilled water and dry between filter papers.

Using excess of a solid is key because the excess is easily filtered. The same method works for zinc oxide, magnesium carbonate or zinc metal with acids; effervescence stops when all the acid is used.

Section 4 of 14Preparing a soluble salt from an alkali: titration

For sodium, potassium or ammonium salts, both reactants are soluble, so no excess can be filtered off.

1
The exact volumes are found by titration with an indicator.
2
Then the experiment is repeated with the same volumes but without the indicator; the solution is evaporated to crystallise the salt.
3
Example: NaOH + HCl → NaCl + H2O gives sodium chloride.
4
Final crystals are dried in a warm oven or between filter papers.

Section 5 of 14Choosing the correct preparation method

Starting materialsMethodWhy
Insoluble oxide/carbonate + acidAdd excess solid, filter, crystalliseExcess removable by filtering
Reactive-but-safe metal (Zn, Mg) + acidAdd excess metal, filter, crystalliseHydrogen given off, visible end point
Soluble alkali + acidTitration, then evaporateNeither can be filtered off

Do not use sodium or potassium metal with acid because the reaction is dangerously violent, and do not use copper because it does not react. Always write the word equation, then the balanced symbol equation, and check the solubility of the salt before deciding on the method.

Section 6 of 14Worked examples

  1. Write a balanced equation for the reaction of zinc oxide with sulfuric acid. [2]2 marks

  2. Why is excess copper(II) oxide added when preparing copper sulfate? [2]2 marks

  3. Describe how to prepare crystals of sodium nitrate from sodium hydroxide solution and nitric acid. [4]4 marks

  4. Name the gas formed when magnesium carbonate reacts with acid and give the test. [2]2 marks

Section 7 of 14The precipitation method

An insoluble salt is made by mixing solutions of two soluble salts that between them contain the ions of the desired salt. The insoluble salt forms as a solid precipitate:

soluble salt 1 + soluble salt 2 → insoluble salt + soluble salt 3.

The solid salt obtained must be insoluble in water for the method to work, so you must be able to use the solubility rules.

The precipitation method
Example: lead(II) sulfate.
Lead(II) nitrate solution and potassium sulfate solution are mixed: Pb(NO3)2(aq) + K2SO4(aq) → PbSO4(s) + 2KNO3(aq).
The white precipitate is the insoluble product.

Section 8 of 14Ionic equations and spectator ions

In an ionic equation only the ions that form the precipitate are shown: Pb2+(aq) + SO42-(aq) → PbSO4(s).

Spectator ions
The potassium and nitrate ions remain dissolved and unchanged; they are spectator ions.
Other examples: silver chloride, Ag+(aq) + Cl-(aq) → AgCl(s); barium sulfate, Ba2+ + SO42- → BaSO4; calcium carbonate, Ca2+ + CO32- → CaCO3.
State symbols are required: (aq), (s), (l), (g).
To choose reactants for a given insoluble salt, pick one soluble salt of the metal (nitrates are always soluble) and one soluble salt of the negative ion (sodium or potassium salts are always soluble).

Section 9 of 14Practical: lead(II) sulfate

1
Measure 25 cm3 of 0.5 mol/dm3 lead(II) nitrate solution into a small beaker.
2
Add 25 cm3 of 0.5 mol/dm3 potassium sulfate solution and mix with a stirring rod; a white precipitate forms immediately.
3
Filter the mixture using a funnel and filter paper.
4
Wash the residue with distilled water to remove soluble impurities (potassium nitrate).
5
Dry the solid between filter papers or in a warm oven.

The mixed amounts are equal moles, in a 1:1 ratio, so both reactants are fully used. The residue left on the filter paper is the product; the filtrate contains the soluble by-product. Lead compounds are toxic, so gloves are worn and waste is disposed of correctly.

Section 10 of 14Comparing methods and the titration approach

Salt typeMethodKey idea
Soluble, from insoluble base, carbonate or metalAdd excess solid, filter, crystalliseExcess is easily removed
Soluble, from alkali (Na, K, NH4)Titration with indicator, repeat without, crystalliseExact volumes needed
InsolublePrecipitation: mix two soluble salts, filter, wash, dryProduct is the residue

For copper(II) sulfate by titration or insoluble base methods, evaporate until the crystallisation point, leave to cool, and dry the crystals; a pure sample is hydrated, CuSO4.5H2O, which loses its water of crystallisation if heated strongly.

Section 11 of 14Predicting precipitates and identifying ions

Precipitation reactions also help in the lab to identify ions.

1
Sodium hydroxide gives coloured hydroxide precipitates with metal ions: copper(II) blue, iron(II) green, iron(III) red-brown.
2
Silver nitrate with halide ions gives precipitates: chloride white, bromide cream, iodide yellow.
3
Barium chloride with sulfate ions gives white barium sulfate.
4
To predict whether a precipitate forms, list the ions in each mixed solution, pair each positive ion with the other negative ion,
  • and check the solubility rules: if either new compound is insoluble, it forms a precipitate.

Examplemixing magnesium sulfate solution and sodium hydroxide solution pairs Mg2+ with OH-, and magnesium hydroxide is insoluble, so a white precipitate forms: Mg2+ + 2OH- → Mg(OH)2(s).

Section 12 of 14Worked examples

  1. Suggest two soluble salts that could be mixed to make barium sulfate and write the ionic equation. [3]3 marks

  2. Why is the precipitate washed with distilled water? [1]1 mark

  3. Name the spectator ions in the reaction of silver nitrate with sodium chloride. [1]1 mark

  4. Explain why precipitation could not be used to make sodium chloride crystals. [2]2 marks

Section 13 of 14Key terms

TermMeaning
SaltCompound formed when H+ of an acid is replaced by a metal or ammonium ion
BaseSubstance that neutralises an acid
ExcessMore of a reactant than needed so the other is completely used up
FiltrateLiquid that passes through filter paper
CrystallisationFormation of crystals from a saturated solution
EffervescenceFizzing due to gas release
TermMeaning
PrecipitateAn insoluble solid formed in a reaction between solutions
Spectator ionsIons that do not change in a reaction
ResidueSolid left on the filter paper
FiltrateSolution passing through the filter
Ionic equationEquation showing only the ions that react
Water of crystallisationWater chemically held in crystals

Section 14 of 14Exam technique

1
Using titration for an insoluble base — excess and filtration are used instead.
2
Boiling to dryness, which can decompose crystals; stop at the crystallisation point.
3
Forgetting that carbonates produce CO2 and water, not hydrogen.
4
Using a hydroxide that is insoluble as an alkali in titration.
1
Forgetting state symbols in ionic equations.
2
Including spectator ions in the ionic equation.
3
Trying to filter a soluble salt.
4
Choosing silver chloride as soluble — it is a white precipitate.

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