2.2Chemical Bonding, Application Of Chemical Reactions and Organic Chemistry

Acids, Bases and Salts — WJEC GCSE Chemistry Revision Notes & Topical Questions

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Complete PDF with figures and original page layoutPage 1 WJEC GCSE Chemistry Your notes 2.2 Acids, Bases & Salts Contents The pH Scale Hydrogen & Hydroxide Ions Acids Reacting with Metals Neutralisation of Acids The Acid-Carbonate Reaction …

Section 1 of 29The pH scale

Scale
- it runs from 1 to 14, with acids below 7, alkalis above 7, and neutral at pH 7

pH	Description
0–2	Strong acid
3–6	Weak acid
7	Neutral
8–11	Weak alkali
12–14	Strong alkali

Trend
- the lower the pH, the more acidic, and the higher the pH, the more alkaline, with some very strong acids below 1
Common acids
- sulfuric acid H₂SO₄, hydrochloric acid HCl and vinegar (ethanoic acid) CH₃COOH
Common alkalis
- sodium hydroxide NaOH, ammonia NH₃ and baking powder

Section 2 of 29Measuring pH

Universal indicator
- a mixture of plant indicators across a broad pH range that gives an approximate pH

Add a few drops to the solution.
Match the colour to a chart for that indicator.
Read off the pH.

Colour charts
- colours vary slightly between manufacturers, so the correct chart is used
Comparing substances
- the lowest pH is the strongest acid, the highest is the strongest alkali, and the nearest to 7 is closest to neutral

Section 3 of 29Hydrogen and hydroxide ions

Acids
- they have pH below 7, taste sour when edible, are corrosive, and neutralise bases to form a salt and water
Hydrogen ions
- H⁺ ions make a solution acidic, as in HCl(aq) → H⁺(aq) + Cl^-(aq)
Alkalis
- a water-soluble base with pH above 7 that neutralises acids to form a salt and water
Hydroxide ions
- OH^- ions make a solution alkaline, as in NaOH(s) → Na⁺(aq) + OH^-(aq)

Section 4 of 29Hazards of acids

Dependence
- the hazard depends on the type and concentration of the acid
Dilute acids
- most need no hazard symbol or an irritant symbol, and eye protection is worn
Moderately concentrated acids
- often corrosive, so gloves are added to eye protection
Concentrated acids
- some, such as nitric acid, are oxidising and can start or intensify a fire, so eye protection, gloves and often a fume cupboard are needed

Section 5 of 29Acids reacting with metals

Rule
- only metals above hydrogen in the reactivity series react with dilute acid, so copper does not
Vigour
- the more reactive the metal, the more vigorous the reaction, and potassium and sodium react explosively
General equation
- metal + acid → salt + hydrogen (MASH)

Metal	With sulfuric acid	With hydrochloric acid
Magnesium	Mg + H₂SO₄ → MgSO₄ + H₂	Mg + 2HCl → MgCl₂ + H₂
Zinc	Zn + H₂SO₄ → ZnSO₄ + H₂	Zn + 2HCl → ZnCl₂ + H₂
Iron	Fe + H₂SO₄ → FeSO₄ + H₂	Fe + 2HCl → FeCl₂ + H₂

Hydrogen test
- the gas burns with a squeaky pop when ignited

Section 6 of 29Making crystals from a metal and an acid

React sulfuric acid with excess zinc metal so that all the acid is used up.
Wait until no more fizzing is seen.
Filter the solution into a beaker to remove excess zinc.
Boil until crystals begin to form, when about half the solution remains.
Cool, pour into an evaporating dish and place in a warm oven so the crystals fully form.

Why not evaporate to dryness
- heating the beaker until dry could decompose the crystals

Section 7 of 29Neutralisation: general reactions

Neutralisation
- it occurs when an acid reacts with a base, and is exothermic
Bases
- substances with pH above 7, many insoluble, with the soluble ones called alkalis, such as NaOH and Ca(OH)₂
General equations
- acid + base → salt + water, and acid + carbonate → salt + water + carbon dioxide
Naming the salt
- hydrochloric acid gives chlorides, sulfuric acid gives sulfates and nitric acid gives nitrates, with the metal from the base
Table salt
- sodium chloride can be made by neutralisation

Section 8 of 29Acids with metal oxides, hydroxides and carbonates

Metal oxides
- 2HCl + CuO → CuCl₂ + H₂O.
- H₂SO₄ + Na₂O → Na₂SO₄ + H₂O.
- 2HNO₃ + MgO → Mg(NO₃)₂ + H₂O.
Metal hydroxides
- 2HCl + 2LiOH → 2LiCl + 2H₂O.
- H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.
- HNO₃ + KOH → KNO₃ + H₂O.
Metal carbonates
- they give effervescence from carbon dioxide
- 2HCl + Na₂CO₃ → 2NaCl + H₂O + CO₂.
- H₂SO₄ + CaCO₃ → CaSO₄ + H₂O + CO₂.
- 2HNO₃ + MgCO₃ → Mg(NO₃)₂ + H₂O + CO₂.
Gas test
- bubble the gas through limewater, which turns milky if it is carbon dioxide
Reasoning
- an acid-base reaction that fizzes must involve a carbonate
Not neutralisation
- an acid with a metal forms a salt but no water, so it is not neutralisation

Section 9 of 29Neutralisation in terms of ions (Higher tier)

Sources
- acids supply H⁺ ions and bases or alkalis supply OH^- ions
Net ionic equation
- H⁺(aq) + OH^-(aq) → H₂O(l), which makes the solution neutral because water has pH 7
Soil treatment
- bases such as limestone and quicklime raise soil pH for crops that cannot tolerate acidity

Section 10 of 29The acid-carbonate test

Test
- add acid to a suspected carbonate and bubble the gas into limewater
Equation
- CO₃^2-(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)
Limewater reaction
- CO₂(g) + Ca(OH)₂(aq) → CaCO₃(s) + H₂O(l), giving a milky white precipitate
Practical point
- connect the suspected-carbonate test tube to the limewater quickly so no gas escapes
Observation
- the observation is fizzing, because "carbon dioxide forms" is not an observation

Section 11 of 29Making a soluble salt from an insoluble base

Reaction
- a soluble salt forms when an acid reacts with an insoluble base, such as CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)
Excess base
- the insoluble reactant is added in excess
- All the acid reacts, so no unreacted acid becomes dangerously concentrated during evaporation.
- The excess solid is then removed by filtration, leaving only salt and water.
Crystallising
- evaporate by heating until small crystals appear, when about half the water remains
- Slow evaporation over days gives larger crystals.
Carbonate as the base
- carbon dioxide is released into the atmosphere
Metals
- an acid can also give a salt with a metal above hydrogen, if it is not so reactive that the reaction is dangerous

Section 12 of 29Practical: making soluble salt crystals

Definition
- a salt forms when the hydrogen in an acid is replaced by a metal, such as potassium chloride from HCl
Uses of salts
- fertilisers, batteries, cleaning products, healthcare products and fungicides
Source
- some salts are mined, but others are prepared in the laboratory
Apparatus
- stirring rod, 100 cm³ beaker, filter funnel and paper, evaporating dish, 50 cm³ measuring cylinder, 0.5 mol dm^-3 sulfuric acid, copper(II) oxide, spatula and indicator paper

Section 13 of 29Practical: method

Add dilute acid to a beaker and heat with a Bunsen flame.
Add the insoluble base or carbonate a little at a time, stirring, until it is in excess and stops disappearing.
Touch a glass rod to indicator paper to check the acid is neutralised.
Filter into an evaporating basin to remove excess base.
Heat to evaporate about half the water to give a saturated solution, checking with a cold glass rod that crystals form on it.
Leave the filtrate in a warm place to crystallise.
Decant excess solution and dry the crystals, or blot them with filter paper.

Example
- dilute sulfuric acid with copper(II) oxide gives copper(II) sulfate: CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)
Crystal size
- slow evaporation over about a week gives larger crystals, and heating gives smaller crystals
Excess reactant
- it ensures the acid is completely used up, so the solution is neutral and the salt has no unreacted acid
Other salts
- the same method makes potassium chloride and potassium sulfate
Carbonate example
- CuCO₃(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) + CO₂(g)

Section 14 of 29Naming salts

First part
- it comes from the metal, metal oxide, hydroxide or carbonate
Second part
- it comes from the acid

Acid	Ion in salt	Salt ending
Hydrochloric acid, HCl	Cl^-	-chloride
Sulfuric acid, H₂SO₄	SO₄^2-	-sulfate
Nitric acid, HNO₃	NO₃^-	-nitrate

Acid	Base	Salt	Formula
Sulfuric acid	Calcium carbonate	Calcium sulfate	CaSO₄
Hydrochloric acid	Magnesium oxide	Magnesium chloride	MgCl₂
Nitric acid	Potassium hydroxide	Potassium nitrate	KNO₃

Writing the equation for nitric acid and calcium hydroxide
- Reactants are HNO₃ and Ca(OH)₂, products are Ca(NO₃)₂ and H₂O.
- Two nitrate ions need 2HNO₃, and balancing hydrogen needs 2H₂O.
- The balanced equation is 2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O.

Section 15 of 29Test for sulfate ions

Acidify the sample with dilute hydrochloric acid.
Add a few drops of aqueous barium chloride.
A white precipitate of barium sulfate shows sulfate ions.

Equation
- Ba²⁺(aq) + SO₄^2-(aq) → BaSO₄(s)
Why acidify first
- HCl removes carbonates, which would also form a precipitate
Alternative
- nitric acid with barium nitrate gives the same white precipitate

Section 16 of 29Preparing a salt by titration

Use
- an acid and a soluble alkali give a soluble salt, such as HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
Principle
- at complete neutralisation only salt and water remain

Use a pipette to put a fixed volume of alkali in a conical flask and add a few drops of indicator.
Fill the burette with acid and note the start volume.
Add acid slowly until the indicator changes colour.
Record the final volume and work out the volume of acid added.
Add the same volume of acid to the same volume of alkali without indicator.
Heat in an evaporating basin to a saturated solution, with crystals just forming on a dipped glass rod.
Leave to crystallise, decant excess solution and dry the crystals.

Common mistake
- forgetting that the second mixture must have no indicator
Methods to learn
- acid with an insoluble base, and titration

Section 17 of 29Relative concentration from titration

Principle
- if one solution has a known concentration, the relative concentration of the other follows from the volumes used
Greater volume of acid
- more acid was needed to neutralise the alkali, so the acid is less concentrated
Greater volume of alkali
- the alkali is less concentrated for the same reason
Condition
- this holds when the acid and alkali react in a 1:1 ratio

Section 18 of 29Relative concentration example

Example
- 30.00 cm³ of 0.2 mol dm^-3 acid is neutralised by 15.00 cm³ of alkali
Result
- only half the volume of alkali is needed, so the alkali is twice as concentrated
Method
- it is carried out as described in the titration practical

Section 19 of 29Calculating actual concentration or volume (Higher tier)

Skill
- calculate concentration or volume from titration data

Calculate the moles of the known solution using moles = concentration × volume in dm³.
Use the balanced equation to find the moles of the other reactant.
Calculate concentration = moles ÷ volume in dm³, or volume = moles ÷ concentration.
Convert cm³ to dm³ by dividing by 1000.

Barium hydroxide example
- 25.00 cm³ of 0.15 mol dm^-3 Ba(OH)₂ neutralises 12.80 cm³ HNO₃, with ratio 1 : 2, so moles of acid are 7.5 × 10^-3 and concentration is 0.59 mol dm^-3
Volume example
- 25.00 cm³ of 0.80 mol dm^-3 KOH needs 0.02 mol HNO₃ at ratio 1 : 1, which is 40 cm³ of 0.50 mol dm^-3 acid

Section 20 of 29Practical: acid-base titration set-up

Aim
- to find the reacting volumes of a strong acid and strong alkali and so the concentration of the acid
Apparatus
- 25 cm³ volumetric pipette and filler, 50 cm³ burette, 250 cm³ conical flask, small funnel, 0.1 mol dm^-3 sodium hydroxide, hydrochloric acid of unknown concentration, phenolphthalein, clamp stand and white tile
Practical tip
- remove the funnel from the burette before starting, as it can drip and make the initial reading false

Section 21 of 29Practical: method

Use the pipette and filler to put exactly 25 cm³ of sodium hydroxide in the conical flask.
Fill the burette with acid and run a little through to remove air bubbles.
Record the start reading to the nearest 0.05 cm³.
Place the flask on a white tile with the burette tip inside it.
Add a few drops of indicator.
Do a rough titration, running in 1–3 cm³ portions and swirling.
Close the tap at the sharp colour change and record the volume at eye level with the meniscus.
Repeat with a fresh sample, adding acid one drop at a time near the rough end-point.
Record the volume to the nearest 0.05 cm³.
Repeat until two concordant results within 0.1 cm³ are obtained.

Section 22 of 29Practical: results and conclusion

Table
- record final reading, first reading and titre for the rough titre and each titre
  - (final minus first)
Mean titre
- use only concordant results, such as 14.90 and 15.00 giving 14.95 cm³
Conclusion
- 14.95 cm³ of acid neutralises 25 cm³ of 0.1 mol dm^-3 alkali, so the acid is more concentrated than the alkali
Rough titre
- it is not used in the mean

Section 23 of 29Concentration (Higher tier)

Terms
- a solute dissolves in a solvent to form a solution, and concentration is the amount of solute in a given volume
Units
- g dm^-3 or mol dm^-3
Mass concentration
- concentration (g dm^-3) = mass of solute (g) ÷ volume of solution (dm³)
Volume conversion
- 1 dm³ is 1 litre or 1000 cm³, so divide by 1000 for cm³ to dm³ and multiply by 1000 for dm³ to cm³
Mole concentration
- concentration (mol dm^-3) = moles of solute ÷ volume of solution (dm³)
Examples
- 10 g NaOH in 2 dm³ is 5 g dm^-3, and 2.5 dm³ of 0.2 mol dm^-3 contains 0.5 mol
Recall
- the equation may not be given, so it must be remembered

Section 24 of 29Neutralisation calculations (Higher tier)

Purpose
- a titration finds exactly how much alkali neutralises an acid, or the reverse

Write the balanced equation.
Calculate moles of the known solution using volume ÷ 1000 × concentration.
Use the mole ratio to find the moles of the other solution.
Divide moles by the volume in dm³ to get the concentration.

Example
- 25.0 cm³ HCl neutralised by 12.1 cm³ of 0.100 mol dm^-3 NaOH gives 1.21 × 10^-3 mol of both, so the acid is 1.21 × 10^-3 ÷ 0.025 = 0.0484 mol dm^-3

Section 25 of 29Dilute and concentrated acids

Dilute acid
- acid molecules are mixed with a large amount of water
Concentrated acid
- there is little to no water mixed with the acid molecules, so a large number of acid molecules are in the solution
Independence from strength
- any acid can be dilute or concentrated, but a specific acid is strong or weak
Examples
- hydrochloric acid (pH 1) is strong and ethanoic acid (pH 3) is weak, and either can be dilute or concentrated

Section 26 of 29Strong acids

Strong acid = an acid that completely dissociates (ionises) in water, giving solutions of very low pH
Examples
- HCl and H₂SO₄
Equation
- HCl(aq) → H⁺(aq) + Cl^-(aq), shown with a single arrow

Section 27 of 29Weak acids

Weak acid = an acid that partially dissociates in water, giving pH values nearer the middle of the scale but still below 7
Equilibrium
- an equilibrium exists between molecules and ions, as in CH₃CH₂COOH ⇌ H⁺ + CH₃CH₂COO^-
Position
- the equilibrium lies to the left, so most acid molecules stay intact and few H⁺ ions are present
Reaction with alkali
- ethanoic acid forms ethanoate salts, CH₃COOH + NaOH → CH₃COONa + H₂O

Section 28 of 29Reactions of strong and weak acids

pH
- a strong acid has a lower pH than a weak acid
Reacting particles
- a strong acid has more H⁺ ions, so there are more reacting particles
More vigorous reactions
- strong acids react more vigorously with metals, bases and carbonates
- A faster rate of gas production, such as hydrogen with metals and carbon dioxide with carbonates.
- A larger temperature rise, because the reaction is more exothermic.
Comparison
- a strong acid with magnesium causes a bigger temperature increase than a weak acid
Diagram idea
- in order of increasing concentration, weak acids are partially ionised and strong acids fully ionised

Section 29 of 29Preparing insoluble salts by precipitation

Principle
- an insoluble salt forms as a precipitate when two soluble reactants are mixed
Condition
- the salt formed must be insoluble in water and the reactants must be soluble

Dissolve the two soluble salts in water and mix them in a beaker with a stirring rod.
Filter to remove the precipitate.
Wash the residue with distilled water to remove traces of other solutions.
Dry the precipitate in an oven.

Lead(II) sulfate example
- Mix lead(II) nitrate and potassium sulfate solutions.
- Wash the precipitate to remove potassium nitrate solution.
- Pb(NO₃)₂(aq) + K₂SO₄(aq) → PbSO₄(s) + 2KNO₃(aq).
Link to tests
- precipitation is the basis of tests for sulfates and halides, though solubility rules need not be learned

Last updated: 14 October 2026

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Step-by-step problem-solving guide, mark scheme criteria, and experimental error analysis curated by Dr. Khaled Dabour (DABOURHUB).

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1. Identify all given values and convert non-standard units, 2. State the primary formula or physiological/chemical principle, 3. Calculate to 3 significant figures and clearly state units.

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Avoid confusing independent vs dependent variables in practical questions. In calculation questions, double-check percentage change formula: (change / original) × 100.

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