What are the key differences between acids and alkalis in terms of their chemical properties?

Acids and alkalis are both important in inorganic chemistry. Acids are substances that release hydrogen ions (H+) when dissolved in water, resulting in a pH less than 7. They typically taste sour and can react with metals to produce hydrogen gas. Alkalis, on the other hand, are bases that dissolve in water to release hydroxide ions (OH-), resulting in a pH greater than 7. They often feel slippery and can neutralize acids to form water and a salt. Understanding these properties is crucial for Edexcel IGCSE Chemistry students.

How is the pH scale used to measure the acidity or alkalinity of a solution?

The pH scale is a numerical scale ranging from 0 to 14 used to measure the acidity or alkalinity of a solution. A pH of 7 is considered neutral, which is the pH of pure water. Solutions with a pH less than 7 are acidic, while those with a pH greater than 7 are alkaline. The scale is logarithmic, meaning each whole number change represents a tenfold change in hydrogen ion concentration. This concept is essential for students studying Edexcel IGCSE Chemistry, especially when conducting experiments involving acids and alkalis.

What is a titration and how is it used to determine the concentration of an acid or alkali?

Titration is a laboratory technique used to determine the concentration of an unknown acid or alkali solution. It involves adding a titrant of known concentration to the solution until the reaction reaches the equivalence point, which is often indicated by a color change due to an indicator. By measuring the volume of titrant used, students can calculate the concentration of the unknown solution using the formula: concentration = (volume of titrant × concentration of titrant) / volume of unknown solution. Mastery of titration is a key skill for Edexcel IGCSE Chemistry students.

Why is it important to use indicators in titration experiments?

Indicators are crucial in titration experiments because they signal the end point of the titration by changing color. This color change occurs when the solution reaches a specific pH, indicating that the acid and alkali have reacted completely. Common indicators include phenolphthalein and methyl orange, each suitable for different types of titrations. Understanding the role of indicators helps Edexcel IGCSE Chemistry students accurately determine the concentration of solutions in their experiments.

Can you explain the concept of neutralization in the context of acids and alkalis?

Neutralization is a chemical reaction between an acid and an alkali that results in the formation of water and a salt. During this reaction, the hydrogen ions from the acid combine with the hydroxide ions from the alkali to form water, while the remaining ions form a salt. This process is fundamental in inorganic chemistry and is often explored in Edexcel IGCSE Chemistry through practical experiments and calculations. Understanding neutralization is essential for predicting the outcomes of reactions involving acids and alkalis.

Why is "Confusing 'alkali' with 'base'" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Using the two terms as synonyms. How to avoid it: Memorise: 'alkali = soluble base'. CuO is a base (reacts with acid) but NOT an alkali (insoluble). NaOH is BOTH a base AND an alkali.

Why is "Confusing 'strong acid' with 'concentrated acid'" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Everyday language uses them interchangeably. How to avoid it: STRONG / WEAK refers to the % IONISATION (HCl strong = fully ionises). CONCENTRATED / DILUTE refers to mol/dm³. A dilute strong acid (e.g. 0.001 mol/dm³ HCl) can have a HIGHER pH than a concentrated weak acid (e.g. 1 mol/dm³ ethanoic). Source: 4CH1 Examiner Reports 2022-2024

Why is "Using universal indicator for titration" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Universal indicator is the most familiar from class. How to avoid it: Universal indicator changes colour gradually through many colours over a wide pH range — no sharp endpoint. For titration use methyl orange (sharp pH 4) or phenolphthalein (sharp pH 9) only.

Why is "Including the rough titre when calculating the mean" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Treating all titrations equally. How to avoid it: The rough titre is to find APPROXIMATE endpoint — it overshoots. Exclude it. Calculate the mean using only concordant titres (those within 0.10 cm³ of each other). Source: 4CH1 Examiner Reports 2022-2024

Why is "Forgetting to convert cm³ to dm³ in $n = cV$" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Substituting cm³ directly into n = cV formula. How to avoid it: Either use n = cV/1000 (V in cm³) OR convert V to dm³ first (÷1000). Always check the units of the final n (should be mol, not mol × 1000).

Why is "Forgetting the 1:2 ratio for H₂SO₄ : NaOH titrations" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Assuming all acid + base titrations are 1:1. How to avoid it: Write the balanced equation FIRST. H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O → 1:2 ratio (so n(H₂SO₄) = n(NaOH)/2). Don't assume 1:1.

Why is "Writing Cu + acid → CuCl₂ + H₂" a common mistake in Acids, Alkalis and Titrations?

Why it happens: Applying the metal + acid rule indiscriminately. How to avoid it: Only metals ABOVE H in the reactivity series react with acid to give H₂. Cu, Ag, Au are BELOW H — they do NOT react with dilute HCl or dilute H₂SO₄. (They do react with concentrated HNO₃ but that's an oxidising acid — different reaction.)

Inorganic ChemistryPearson Edexcel IGCSE Chemistry 4CH1 Higher

Acids, Alkalis and Titrations

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Acids, Alkalis and Titrations — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

Defining acids and alkalis in terms of H⁺/OH⁻ ions, the pH scale 0-14, indicator colours (litmus, methyl orange, phenolphthalein, universal indicator), strong vs weak acids (Paper 2 P content), the reactions of acids with metals/bases/carbonates, and the required practical for titration including burette/pipette technique, concordant titres and Paper-2 calculations using $n = cV/1000$ .

At a glance

Acid = produces H⁺ in water (HCl, H₂SO₄, HNO₃). Alkali = soluble base produces OH⁻ (NaOH, KOH).
pH scale 0-14: <7 acidic, 7 neutral, >7 alkaline. Indicator colours: litmus (red↔blue), methyl orange (red→yellow at pH ~4), phenolphthalein (colourless→pink at pH ~9).
Strong acid = fully ionised (HCl → 100 % H⁺ + Cl⁻). Weak acid = partially ionised (CH₃COOH ⇌ ~1 %). Same conc → strong has lower pH + faster rate.
Acid + metal → salt + H₂ (only metals above H react).
Acid + base → salt + H₂O (neutralisation).
Acid + carbonate → salt + H₂O + CO₂.
Titration: pipette + burette + indicator → concordant titres within 0.10 cm³ → mean titre.
Calculation: $n = cV/1000$ → mole ratio → $c$ unknown. Convert mol/dm³ ↔ g/dm³ with $M_r$ .

What you’ll learn

Mapped to the Cambridge IGCSE 4CH1 syllabus (2026 onwards).

2.36 — Recall acids in aqueous solutions are sources of H⁺ ions, and alkalis are sources of OH⁻ ions.
2.37 — Understand how to measure pH using indicator solutions and pH paper / probes.
2.38 — Recall the colour changes of litmus, methyl orange and phenolphthalein in acidic and alkaline solutions.
2.39 — Describe how universal indicator can give an approximate pH value.
2.40 — Understand the general reactions of acids: with metals → salt + H₂; with bases → salt + H₂O; with carbonates → salt + H₂O + CO₂.
2.41 — Naming salts: hydrochloric → chloride; sulfuric → sulfate; nitric → nitrate.
2.42P — (Paper 2) Understand 'strong acid' (fully ionised) vs 'weak acid' (partially ionised) — equilibrium for weak acids.
2.43P — (Paper 2) Compare strong vs weak acids of equal concentration in terms of pH and reactivity.
2.44 — Describe titration as a method to find the concentration of an acid or alkali. (Required practical.)
2.44P — (Paper 2) Calculate concentration from titration data using $n = cV/1000$ .
2.45 — Define concordant titres (within 0.10 cm³) and calculate mean titre.

Acids, alkalis and the pH scale (spec 2.36, 2.37, 2.38, 2.39)

H⁺ producer = acid; OH⁻ producer = alkali. pH 0-14 measures it.

Acid definition. An acid is a substance that produces / donates H⁺ ions (protons) when dissolved in water. Common examples: hydrochloric acid HCl, sulfuric acid H₂SO₄, nitric acid HNO₃, ethanoic acid CH₃COOH.

$\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

Base vs alkali. A base is any H⁺ acceptor — including insoluble metal oxides (CuO, MgO) and hydroxides (Fe(OH)₃). An alkali is a SOLUBLE base that produces OH⁻ ions in water. e.g.:

$\text{NaOH(aq)} \rightarrow \text{Na}^+\text{(aq)} + \text{OH}^-\text{(aq)}$

Common alkalis: NaOH, KOH, Ca(OH)₂ (slightly soluble — limewater), NH₃(aq) (weak alkali).

The pH scale (0-14).

pH below 7 is acidic, 7 is neutral, above 7 is alkaline — the universal indicator colour scale.

pH	Description	Example
0-3	Strong acid	1 mol/dm³ HCl
4-6	Weak acid	Vinegar (CH₃COOH), citric acid
7	Neutral	Pure water
8-11	Weak alkali	Soap, NaHCO₃ solution, NH₃
12-14	Strong alkali	1 mol/dm³ NaOH

Indicators.

Indicator	Acidic	Neutral	Alkaline
Litmus	Red	Purple	Blue
Methyl orange	Red	(transitions ~pH 4)	Yellow
Phenolphthalein	Colourless	(transitions ~pH 9)	Pink
Universal indicator	Red (pH 1-3); orange/yellow (pH 4-6)	Green (pH 7)	Blue (pH 8-11); purple (pH 12-14)

Universal indicator is a MIXTURE of several pH indicators — its colour blends continuously through the rainbow, giving an APPROXIMATE pH. It is NOT used in titrations because the colour change is gradual (no sharp endpoint).

Methyl orange and phenolphthalein give sharp colour changes at specific pH values, so they ARE used in titrations (one drop too far → colour changes definitively).

Acid = H⁺ donor; alkali = OH⁻ donor (soluble base).
pH 0-14: <7 acidic, 7 neutral, >7 alkaline.
Litmus: red↔blue. Methyl orange: red→yellow. Phenolphthalein: colourless→pink.
Universal indicator: rainbow scale (red→green→purple).

See the full worked example for acids, alkalis and titrations →

Strong vs weak acids (Paper 2, spec 2.42P, 2.43P)

Strong = fully ionised; weak = partially. Same conc → strong has lower pH + faster rate.

Strong acid. An acid that is FULLY (100 %) ionised in aqueous solution. Every HCl molecule splits into ions:

$\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \quad \text{(single arrow — irreversible)}$

Examples: HCl, H₂SO₄, HNO₃. At 0.1 mol/dm³: [H⁺] = 0.1 mol/dm³; pH ≈ 1.

Weak acid. An acid that is PARTIALLY ionised (typically ~1 % at equilibrium). Most molecules remain un-ionised:

$\text{CH}_3\text{COOH(aq)} \rightleftharpoons \text{CH}_3\text{COO}^-\text{(aq)} + \text{H}^+\text{(aq)} \quad \text{(equilibrium arrows)}$

Examples: ethanoic acid (vinegar), citric acid (lemon juice), carbonic acid (CO₂ in water), formic acid. At 0.1 mol/dm³: [H⁺] ≈ 10⁻³ mol/dm³; pH ≈ 3.

Strong acids ionise fully; weak acids ionise only partially, so at equal concentration they have a lower [H⁺].

'Strong / weak' vs 'concentrated / dilute' — DIFFERENT meanings!

STRONG / WEAK refers to the DEGREE OF IONISATION.
CONCENTRATED / DILUTE refers to MOLES PER UNIT VOLUME.

A dilute strong acid (e.g. 0.001 mol/dm³ HCl, pH ~3) is LESS acidic than a concentrated weak acid (e.g. 5 mol/dm³ ethanoic, pH ~2.5). You can't predict acidity from one alone — you need both.

Comparing reactions of strong vs weak acids (same concentration).

When equal concentrations of HCl and CH₃COOH react with magnesium:

Property	HCl (strong)	CH₃COOH (weak)
[H⁺] available	0.1 mol/dm³	~0.001 mol/dm³
Rate of reaction with Mg	Fast / vigorous fizz	Slow / gentle fizz
pH	~1	~3
Final volume of H₂	Same	Same

Why same final volume of H₂? As H⁺ is consumed in the reaction, the weak-acid equilibrium shifts right (Le Chatelier — removing H⁺ pushes the equilibrium to replace it). Eventually all CH₃COOH ionises and reacts → same moles of H⁺ react in total → same volume of H₂. Just slower.

Strong acid: fully ionised (HCl → H⁺ + Cl⁻).
Weak acid: partially ionised (~1 %; CH₃COOH ⇌ CH₃COO⁻ + H⁺).
Same conc → strong has higher [H⁺] → lower pH → faster reaction.
Same conc → both produce the SAME total H₂ from excess Mg (equilibrium shifts).

See the full worked example for acids, alkalis and titrations →

Reactions of acids (spec 2.40, 2.41)

Metal → salt + H₂; + base → salt + H₂O; + carbonate → salt + H₂O + CO₂.

Acids react with three classes of substance to give predictable products.

(1) Acid + METAL → salt + hydrogen gas.

Only metals ABOVE H in the reactivity series react with dilute acids (Cu, Ag, Au don't react).

Acid	Metal	Example equation
HCl	Mg	$\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$
H₂SO₄	Zn	$\text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2$
HCl	Fe	$\text{Fe} + 2\text{HCl} \rightarrow \text{FeCl}_2 + \text{H}_2$

Observations: fizzing (H₂ gas escapes), metal gradually dissolves, solution gets warm (exothermic).

(2) Acid + BASE → salt + water (NEUTRALISATION).

Bases include metal oxides (CuO, MgO) and metal hydroxides (NaOH, KOH).

Acid	Base	Example
H₂SO₄	CuO	$\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$
HCl	NaOH	$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$
HNO₃	KOH	$\text{HNO}_3 + \text{KOH} \rightarrow \text{KNO}_3 + \text{H}_2\text{O}$

Ionic equation (any neutralisation):

$\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)}$

(3) Acid + CARBONATE → salt + water + carbon dioxide.

Acid	Carbonate	Example
HCl	CaCO₃	$\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2$
H₂SO₄	Na₂CO₃	$\text{Na}_2\text{CO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + \text{H}_2\text{O} + \text{CO}_2$

Observations: vigorous fizzing (CO₂); test: gas turns limewater milky.

Salt naming convention.

Acid used	Salt produced
Hydrochloric acid (HCl)	Chloride (e.g. NaCl, MgCl₂, CaCl₂)
Sulfuric acid (H₂SO₄)	Sulfate (e.g. ZnSO₄, CuSO₄, BaSO₄)
Nitric acid (HNO₃)	Nitrate (e.g. KNO₃, NaNO₃, AgNO₃)
Phosphoric acid (H₃PO₄)	Phosphate (e.g. Na₃PO₄)
Ethanoic acid (CH₃COOH)	Ethanoate (e.g. CH₃COONa)

The metal name + anion name → salt name.

Acid + metal → salt + H₂ (only metals above H react).
Acid + base → salt + H₂O (neutralisation; ionic eqn H⁺ + OH⁻ → H₂O).
Acid + carbonate → salt + H₂O + CO₂.
Salt names: HCl→chloride, H₂SO₄→sulfate, HNO₃→nitrate.

See the full worked example for acids, alkalis and titrations →

Titrations — apparatus and procedure (spec 2.44, 2.45 — required practical)

Burette + pipette + indicator → concordant titres within 0.10 cm³ → mean.

Titration is a quantitative technique for finding the concentration of an unknown acid or alkali by reacting it with a solution of known concentration ('standard solution'), using an indicator to detect the endpoint.

A burette delivers a variable volume into a fixed pipetted volume in the conical flask, with a white tile beneath.

Apparatus.

Burette — 50 cm³, graduated to 0.1 cm³, read to 0.05 cm³ (1 dp + estimated digit). Fill from above using a funnel; remove funnel before reading.
Volumetric pipette (25.0 cm³) + pipette filler (never mouth-pipette acid).
Conical flask (250 cm³) — better than a beaker because the shape allows safe swirling.
White tile — placed under flask so colour change is easy to see.
Indicator — methyl orange (red ↔ yellow) or phenolphthalein (pink ↔ colourless).
DC power for top-pan balance (if making standard solution).
Goggles, lab coat.

Procedure — outline.

Pipette exactly 25.0 cm³ of one solution into a conical flask. Add 2-3 drops of indicator.
Fill burette with the other solution. Record initial reading.
Rough titre. Run from burette continuously while swirling; note approximate endpoint volume.
Accurate titres. Pipette fresh sample. Run quickly to within ~2 cm³ of rough volume, then DROPWISE while swirling. Stop when indicator colour change is JUST permanent.
Repeat until two (or more) titres agree to within 0.10 cm³ — these are CONCORDANT.
Mean titre = mean of CONCORDANT titres only (exclude the rough).

Why these techniques matter.

Pipette is calibrated for accuracy (more precise than a measuring cylinder). The 25.0 cm³ pipette delivers exactly 25.0 cm³ at 20 °C — don't shake out the last drop.
Burette read to 0.05 cm³ — eye at meniscus level to avoid parallax error.
Funnel removed before reading so dripping doesn't add to the reading.
White tile under flask makes colour change easier to see.
Swirling ensures the acid and alkali mix thoroughly so endpoint is not overshot.
Dropwise near endpoint avoids overshooting the colour change.

Concordant titres — what they mean. Concordant titres (within 0.10 cm³ of each other) show that the experiment is reproducible. The mean of concordant values is the BEST ESTIMATE of the volume required for neutralisation.

Example data table.

Titration	Initial reading / cm³	Final reading / cm³	Titre / cm³
Rough	0.00	24.10	24.10
1	0.00	23.50	23.50
2	23.50	47.00	23.50
3	0.00	23.45	23.45

Concordant titres = 23.50, 23.50, 23.45 (all within 0.10 cm³). Mean = (23.50 + 23.50 + 23.45) / 3 = 23.48 cm³.

Pipette = exact volume; burette = variable volume (read to 0.05 cm³).
Use methyl orange OR phenolphthalein (sharp endpoint), NOT universal.
Add dropwise + swirl near endpoint.
Concordant titres within 0.10 cm³; mean excludes rough.

See the full worked example for acids, alkalis and titrations →

Titration calculations (Paper 2 — spec 2.44P)

n = cV/1000 → mole ratio from equation → unknown c.

Five-step method for any titration calculation.

Write balanced equation. Identify the mole ratio of acid : base.
Calculate moles of the KNOWN. Use $n = \dfrac{c \times V}{1000}$ (V in cm³) or $n = c \times V$ (V in dm³).
Use mole ratio to find moles of UNKNOWN.
Calculate concentration of UNKNOWN. $c = n / V$ (V in dm³) or $c = 1000n/V$ (V in cm³).
Convert units if needed (e.g. mol/dm³ ↔ g/dm³ via $c_{g/dm^3} = c_{mol/dm^3} \times M_r$ ).

Worked example 1 — 1:1 ratio (HCl + NaOH).

25.0 cm³ of HCl needs 23.50 cm³ of 0.100 mol/dm³ NaOH. Find c(HCl).

Equation: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$ . Ratio 1:1.

$n(\text{NaOH}) = \dfrac{0.100 \times 23.50}{1000} = 2.35 \times 10^{-3}\ \text{mol}$

$n(\text{HCl}) = n(\text{NaOH}) = 2.35 \times 10^{-3}\ \text{mol}$

$c(\text{HCl}) = \dfrac{2.35 \times 10^{-3}}{0.0250} = 0.0940\ \text{mol/dm}^3$

Worked example 2 — 1:2 ratio (H₂SO₄ + NaOH).

25.0 cm³ H₂SO₄ needs 30.00 cm³ of 0.200 mol/dm³ NaOH. Find c(H₂SO₄).

Equation: $\text{H}_2\text{SO}_4 + 2\text{NaOH} \rightarrow \text{Na}_2\text{SO}_4 + 2\text{H}_2\text{O}$ . Ratio H₂SO₄ : NaOH = 1 : 2.

$n(\text{NaOH}) = \dfrac{0.200 \times 30.00}{1000} = 6.00 \times 10^{-3}\ \text{mol}$

$n(\text{H}_2\text{SO}_4) = \dfrac{n(\text{NaOH})}{2} = 3.00 \times 10^{-3}\ \text{mol}$

$c(\text{H}_2\text{SO}_4) = \dfrac{3.00 \times 10^{-3}}{0.0250} = 0.120\ \text{mol/dm}^3$

Worked example 3 — converting to g/dm³.

c(HCl) = 0.0940 mol/dm³. $M_r$ (HCl) = 36.5.

$c(\text{HCl}) = 0.0940 \times 36.5 = 3.43\ \text{g/dm}^3$

Common pitfalls.

ALWAYS use the balanced equation FIRST — don't assume 1:1.
$n = cV/1000$ : V must be in cm³.
Watch sig figs: typically 3 sf for IGCSE answers.
Quote units: mol/dm³ or g/dm³ — don't leave 'naked' numbers.

Step 1: write balanced equation → mole ratio.
Step 2: $n = cV/1000$ for the known.
Step 3: use ratio to find $n$ of unknown.
Step 4: $c = 1000n/V$ for unknown in mol/dm³.
Step 5: convert to g/dm³ using $\times M_r$ if asked.

See the full worked example for acids, alkalis and titrations →

How it’s examined

Acids, alkalis and titrations are tested on EVERY 4CH1 paper — worth 8-14 marks per session combined. Paper 1 (4CH1/1C) items: (a) define acid/alkali + indicator colour changes (3-4 marks); (b) reactions of acids with metals/oxides/carbonates including word + symbol equations (4-5 marks); (c) describe titration apparatus + concordant titres (5-6 marks — required practical). Paper 2 (4CH1/2C) BOLD 'P' items: (d) strong vs weak acid in terms of ionisation, pH and rate (4-5 marks); (e) titration calculation finding c in mol/dm³ AND g/dm³ (5-6 marks). Watch for diprotic acids (H₂SO₄) where the 1:2 ratio catches students out.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme; 4CH1/2C January 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — Acids, Alkalis and Titrations

Step-by-step solutions to past-paper-style questions on acids, alkalis and titrations, written exactly the way a tutor would explain them at the board.

1Defining an acid and an alkali (3 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q2• acid, alkali, spec-2.36, spec-2.37
Question
Define an acid and an alkali in terms of the ions they produce in aqueous solution. Give ONE example of each. (3 marks)
Step-by-step solution
1. Step 1
  Acid definition (1 M). An acid is a substance that produces / releases / donates H⁺ ions (protons / hydrogen ions) when dissolved in water. e.g. $\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$ .
2. Step 2
  Alkali definition (1 A). An alkali is a SOLUBLE BASE — it produces / releases OH⁻ ions (hydroxide ions) when dissolved in water. e.g. $\text{NaOH(aq)} \rightarrow \text{Na}^+\text{(aq)} + \text{OH}^-\text{(aq)}$ .
3. Step 3
  Examples (1 B). Acid examples: hydrochloric (HCl), sulfuric (H₂SO₄), nitric (HNO₃), ethanoic (CH₃COOH — weak). Alkali examples: sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)₂ — slightly soluble), ammonia solution.
Answer
Acid = substance that produces H⁺ ions in aqueous solution (e.g. HCl). Alkali = soluble base that produces OH⁻ ions in aqueous solution (e.g. NaOH).
Examiner tip
Edexcel 3-mark mark scheme: 1 for H⁺ (must specify 'in aqueous solution' / 'in water'); 1 for OH⁻ (must specify alkali is soluble base — just 'base' loses the mark); 1 for valid examples. Common error: confusing 'base' (all H⁺ acceptors, including insoluble ones like CuO) with 'alkali' (only the soluble ones).
2Strong vs weak acids — HCl vs CH₃COOH at the same concentration (4 marks, Paper 2)
Extended• Adapted from 4CH1/2C May/Jun 2024 Q5• strong acid, weak acid, ionisation, spec-2.42P, spec-2.43P
Question
Hydrochloric acid is described as a STRONG acid, whereas ethanoic acid is described as a WEAK acid. Explain what is meant by these terms, and predict (with explanation) what would be observed when small, equal-sized pieces of magnesium ribbon are added to 0.1 mol/dm³ HCl and 0.1 mol/dm³ ethanoic acid. (4 marks)
Step-by-step solution
1. Step 1
  Strong acid (1 M). A STRONG acid is one that is fully (completely) ionised / dissociated in aqueous solution. Every HCl molecule splits into H⁺ and Cl⁻ ions.
  $\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$
2. Step 2
  Weak acid (1 M). A WEAK acid is one that is only partially ionised / dissociated in aqueous solution. Only about 1 % of CH₃COOH molecules ionise at equilibrium — most remain as un-ionised molecules. Note the equilibrium arrows ⇌.
  $\text{CH}_3\text{COOH(aq)} \rightleftharpoons \text{CH}_3\text{COO}^-\text{(aq)} + \text{H}^+\text{(aq)}$
3. Step 3
  Concentration of H⁺ (1 A). At the same concentration, HCl has a MUCH HIGHER [H⁺] than ethanoic acid (because HCl fully ionises). So HCl has a LOWER pH (more acidic).
4. Step 4
  Observation with Mg (1 B). Both react: $\text{Mg} + 2\text{H}^+ \rightarrow \text{Mg}^{2+} + \text{H}_2$ . With HCl, there is rapid / vigorous effervescence (lots of H⁺ → fast reaction). With ethanoic acid, the fizzing is much slower (lower [H⁺] → slower rate). However, both eventually produce the SAME total volume of H₂ (because the moles of acid are the same, and weak-acid ionisation shifts as H⁺ is consumed).
Answer
Strong acid: fully ionised in water (HCl → H⁺ + Cl⁻). Weak acid: partially ionised (CH₃COOH ⇌ CH₃COO⁻ + H⁺, only ~1 % at equilibrium). At the same concentration, HCl has higher [H⁺] → lower pH → faster reaction with Mg (vigorous fizzing). Ethanoic acid: same Mg + acid reaction but slower (less effervescence). Both ultimately give the same volume of H₂.
Examiner tip
Edexcel 4-mark mark scheme (Paper 2 P content): 1 for 'fully ionised' (strong); 1 for 'partially ionised / equilibrium' (weak); 1 for comparing [H⁺] / pH; 1 for predicting and explaining observation (rate of fizz). 'Strong' / 'weak' refer to the DEGREE OF IONISATION, NOT the concentration. A dilute strong acid can have a higher pH than a concentrated weak acid.
3Titration of HCl against NaOH — apparatus and concordant titres (6 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q7 (required practical)• titration, required practical, spec-2.44, spec-2.45
Question
A student is asked to find the concentration of an HCl solution by titrating it against a 0.100 mol/dm³ NaOH solution. (a) Describe the apparatus and method, including the indicator used and the colour change at the endpoint. (b) Explain what is meant by a 'concordant' result and how the mean titre is calculated. (6 marks)
Step-by-step solution
1. Step 1
  (a) Apparatus (1 M). A burette (50 cm³, graduated to 0.1 cm³, read to 0.05 cm³), a volumetric pipette (25.0 cm³) with pipette filler, a conical flask, a white tile (to see colour change clearly), and an indicator solution. The burette is filled with the standard NaOH; the unknown HCl goes into the conical flask.
2. Step 2
  (a) Method (1 A). Use the pipette to transfer EXACTLY 25.0 cm³ of HCl into the conical flask. Add 2-3 drops of indicator. Fill the burette with NaOH; record the initial burette reading to 0.05 cm³. Place the conical flask on a white tile under the burette.
3. Step 3
  (a) Titration (1 A). Open the tap; add NaOH a little at a time, swirling the flask continuously. As the endpoint approaches, add dropwise. Endpoint = first permanent colour change of the indicator. With methyl orange: red → yellow (sharp at pH ~4). With phenolphthalein: pink → colourless (sharp at pH ~9). Record the final reading; titre = final − initial.
4. Step 4
  (a) Rough titre first (1 B). The first titration is a 'rough' to find the approximate endpoint. Subsequent titrations are added quickly to within ~2 cm³ of the rough titre, then dropwise.
5. Step 5
  (b) Concordant titres (1 B). Titres are concordant when they agree to within 0.10 cm³ of each other (some boards allow 0.20 cm³ — Edexcel uses 0.10). At least two concordant titres must be obtained.
6. Step 6
  (b) Mean titre (1 B). Calculate the mean using ONLY the concordant titres (NOT the rough). e.g. if concordant titres are 23.50 and 23.45 cm³: mean = (23.50 + 23.45) / 2 = 23.475 ≈ 23.48 cm³. Use this mean in the calculation of concentration.
Answer
(a) Apparatus: burette, 25.0 cm³ pipette + filler, conical flask, white tile, indicator. Method: pipette 25.0 cm³ HCl into flask + 2-3 drops indicator. Fill burette with NaOH, record initial reading. Add NaOH while swirling, dropwise near endpoint. Endpoint = colour change (methyl orange red→yellow, or phenolphthalein pink→colourless). (b) Concordant = titres within 0.10 cm³ of each other. Take mean of (at least 2) concordant titres ONLY — exclude the rough.
Examiner tip
Edexcel 6-mark mark scheme: 1 each for naming pipette (with filler), burette, conical flask + indicator, swirling + dropwise addition near endpoint, identifying endpoint colour change, defining concordant + mean. Always read burette to 0.05 cm³ (1 dp + estimated). 'White tile' is sometimes a hidden mark — judging colour change against white background.
4Calculating the concentration of HCl from titration data (6 marks, Paper 2)
Extended• Adapted from 4CH1/2C Jan 2024 Q4• titration calculation, Paper 2, spec-2.44P
Question
In a titration, 25.0 cm³ of HCl(aq) of unknown concentration was neutralised by 23.50 cm³ of 0.100 mol/dm³ NaOH(aq). The balanced equation is: HCl + NaOH → NaCl + H₂O. Calculate the concentration of HCl in mol/dm³ and in g/dm³. (Mr HCl = 36.5). (6 marks)
Step-by-step solution
1. Step 1
  Write the balanced equation and mole ratio (1 M). HCl + NaOH → NaCl + H₂O. 1 : 1 ratio of HCl to NaOH.
2. Step 2
  Calculate moles of NaOH (1 A). Use $n = cV/1000$ (V in cm³).
  $n(\text{NaOH}) = \dfrac{0.100 \times 23.50}{1000} = 2.35 \times 10^{-3}\ \text{mol}$
3. Step 3
  Moles of HCl from ratio (1 A). 1:1 → $n(\text{HCl}) = n(\text{NaOH}) = 2.35 \times 10^{-3}$ mol.
4. Step 4
  Concentration of HCl in mol/dm³ (1 B). $c = n/V$ (V in dm³ = 25.0/1000 = 0.0250 dm³).
  $c(\text{HCl}) = \dfrac{2.35 \times 10^{-3}}{0.0250} = 0.0940\ \text{mol/dm}^3$
5. Step 5
  Convert to g/dm³ (1 B). $c(\text{g/dm}^3) = c(\text{mol/dm}^3) \times M_r$ .
  $c(\text{HCl}) = 0.0940 \times 36.5 = 3.43\ \text{g/dm}^3$
6. Step 6
  Final answer with units and sig figs (1 B). c(HCl) = 0.0940 mol/dm³ = 3.43 g/dm³ (3 sig fig).
Answer
c(HCl) = 0.0940 mol/dm³ ≡ 3.43 g/dm³ (3 sig fig).
Examiner tip
Edexcel 6-mark mark scheme: 1 for balanced equation / 1:1 ratio; 1 for n(NaOH) using cV/1000; 1 for transferring moles via stoichiometric ratio; 1 for converting volume to dm³ or using cV/1000 form; 1 for mol/dm³ answer; 1 for g/dm³ conversion. ECF applies. ALWAYS show working — 'answer-only' loses method marks. Sig figs: match the data (3 sf here).
5Reactions of dilute sulfuric acid with magnesium, copper(II) oxide and calcium carbonate (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q5• acid reactions, metal, metal oxide, carbonate, spec-2.40, spec-2.41
Question
Write the balanced equations for the reactions of dilute sulfuric acid with: (a) magnesium ribbon, (b) copper(II) oxide, (c) calcium carbonate. For each, state ONE observation and identify the salt produced. (5 marks)
Step-by-step solution
1. Step 1
  (a) Metal + acid → salt + hydrogen (1 M). Mg is above H in reactivity series, so it reacts.
  $\text{Mg} + \text{H}_2\text{SO}_4 \rightarrow \text{MgSO}_4 + \text{H}_2$
2. Step 2
  (a) Observation + salt (1 A). Effervescence / fizzing (H₂ gas); Mg gradually dissolves; solution gets warm (exothermic). Salt = magnesium sulfate (MgSO₄) — colourless solution.
3. Step 3
  (b) Metal oxide + acid → salt + water — neutralisation (1 A). CuO is a base.
  $\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$
4. Step 4
  (b) Observation + salt (1 B). Black CuO dissolves on warming; blue solution forms. Salt = copper(II) sulfate (CuSO₄) — characteristic BLUE solution.
5. Step 5
  (c) Carbonate + acid → salt + water + CO₂ (1 B). Effervescence; chalky CaCO₃ disappears; CO₂ gas produced (turns limewater cloudy). Salt = calcium sulfate (CaSO₄) — although CaSO₄ is sparingly soluble, so a slight precipitate may form.
  $\text{CaCO}_3 + \text{H}_2\text{SO}_4 \rightarrow \text{CaSO}_4 + \text{H}_2\text{O} + \text{CO}_2$
Answer
(a) Mg + H₂SO₄ → MgSO₄ + H₂. Observation: fizzing; Mg dissolves. Salt: magnesium sulfate. (b) CuO + H₂SO₄ → CuSO₄ + H₂O. Observation: black CuO dissolves; blue solution. Salt: copper(II) sulfate. (c) CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂. Observation: fizzing (CO₂); CaCO₃ disappears. Salt: calcium sulfate.
Examiner tip
Edexcel 5-mark mark scheme: 1 each for three balanced equations + 1 for one observation + 1 for one salt name. State symbols not required unless asked. Common error: writing CuO + H₂SO₄ → CuSO₄ only (forgetting water) — neutralisation always produces salt + water.

Key Formulae — Acids, Alkalis and Titrations

The formulae you need to memorise for acids, alkalis and titrations on the Pearson Edexcel IGCSE 4CH1 paper, with every variable defined in plain English and a note on when to use it.

Moles from concentration and volume (cm³)
$n = \dfrac{c \times V}{1000}$
$n$
Moles of solute (mol)
$c$
Concentration (mol/dm³)
$V$
Volume in cm³ (NOT dm³)
When to use
The most-used titration formula. Use whenever volume is given in cm³ (the usual case). 1 dm³ = 1000 cm³, so dividing by 1000 converts cm³ to dm³. The unit of n is mol.
Example
25.0 cm³ of 0.100 mol/dm³ NaOH: $n = (0.100 \times 25.0)/1000 = 2.50 \times 10^{-3}\ \text{mol}$ .
Moles from concentration and volume (dm³)
$n = c \times V$
$n$
Moles of solute (mol)
$c$
Concentration (mol/dm³)
$V$
Volume in dm³
When to use
Use when volume is already in dm³ (or after you've converted). Often preferred for unit-checking. Convert cm³ to dm³ by dividing by 1000.
Example
$V = 25.0\ \text{cm}^3 = 0.0250\ \text{dm}^3$ ; $n = 0.100 \times 0.0250 = 2.50 \times 10^{-3}\ \text{mol}$ .
Stoichiometric titration equation (1:1 case)
$c_a V_a = c_b V_b$
$c_a$
Concentration of acid (mol/dm³)
$V_a$
Volume of acid (any units — same on both sides)
$c_b$
Concentration of base (mol/dm³)
$V_b$
Volume of base (same units as V_a)
When to use
Shortcut for titrations where the mole ratio of acid : base is 1:1 (e.g. HCl + NaOH, HNO₃ + KOH). Volumes can be in any matching unit (both cm³ or both dm³). Re-arrange to find the unknown $c$ or $V$ .
Example
HCl: 25.0 cm³ of unknown c; NaOH: 23.50 cm³ at 0.100 mol/dm³. $c(\text{HCl}) = (0.100 \times 23.50)/25.0 = 0.0940\ \text{mol/dm}^3$ .
Titration with stoichiometric ratio (general)
$\dfrac{n_a}{a} = \dfrac{n_b}{b}$
$n_a, n_b$
Moles of acid and base
$a, b$
Stoichiometric coefficients from balanced equation (e.g. for H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O: a = 1, b = 2)
When to use
When the acid is diprotic (H₂SO₄) or triprotic (H₃PO₄), the ratio is not 1:1. Use the BALANCED EQUATION to find the ratio first. Example: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O → ratio H₂SO₄ : NaOH = 1 : 2 → $n(\text{H}_2\text{SO}_4) = n(\text{NaOH})/2$ .
Example
25.0 cm³ H₂SO₄ neutralised by 30.0 cm³ of 0.200 mol/dm³ NaOH. $n(\text{NaOH}) = 0.200 \times 30.0/1000 = 6.00 \times 10^{-3}\ \text{mol}$ . Ratio 1:2 → $n(\text{H}_2\text{SO}_4) = 3.00 \times 10^{-3}\ \text{mol}$ . $c = 3.00 \times 10^{-3} / 0.0250 = 0.120\ \text{mol/dm}^3$ .
Concentration conversion mol/dm³ ↔ g/dm³
$c\,(\text{g/dm}^3) = c\,(\text{mol/dm}^3) \times M_r$
$c (g/dm³)$
Mass concentration (g/dm³)
$c (mol/dm³)$
Molar concentration (mol/dm³)
$M_r$
Relative formula mass of the solute
When to use
Convert between the two concentration units that Edexcel may ask for. mol/dm³ tells you molecules; g/dm³ tells you mass. Always quote with units.
Example
0.0940 mol/dm³ HCl ( $M_r = 36.5$ ): $c = 0.0940 \times 36.5 = 3.43\ \text{g/dm}^3$ .

Model Answers — Acids, Alkalis and Titrations

High-scoring sample answers for acids, alkalis and titrations on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1
4CH1/1C-style — full required-practical procedure6 marks
Describe in full the experimental procedure to find the concentration of a sulfuric acid solution by titration against a 0.100 mol/dm³ solution of sodium hydroxide. Include all apparatus, the indicator choice, the way concordant titres are obtained, and ONE safety precaution. (6 marks)
Model answer
Apparatus needed.

50 cm³ burette (graduated to 0.1 cm³; read to 0.05 cm³).

25.0 cm³ volumetric pipette with a pipette filler (safety: never mouth-pipette acid).

Conical flask (250 cm³).

White tile (placed under the conical flask so the colour change is easier to see).

Indicator: a few drops of methyl orange (red in acid, yellow in alkali — sharp endpoint at pH ~4, well suited to strong acid + strong base) OR phenolphthalein (pink in alkali, colourless in acid).

Funnel (for filling the burette safely; then REMOVE the funnel during titration so excess drops don't drip in).

Goggles and lab coat.

Procedure.

Pipette the acid. Use the 25.0 cm³ pipette with filler to transfer EXACTLY 25.0 cm³ of H₂SO₄ into the conical flask. Allow the last drop to drain — do NOT shake it out (pipette is calibrated this way).

Add indicator. Add 2-3 drops of methyl orange to the flask — the solution turns red (acidic).

Fill the burette with the 0.100 mol/dm³ NaOH. Read the initial volume to 0.05 cm³ from the BOTTOM OF THE MENISCUS at eye level. Remove the funnel.

Rough titration. Place the flask on the white tile. Open the tap; let NaOH run in continuously while swirling. The colour will change to yellow at the endpoint. Record the volume — this is the 'rough' titre.

Accurate titrations. Refill the burette. Pipette a fresh 25.0 cm³ of acid + 2-3 drops of indicator. Open the tap and quickly add NaOH to within ~2 cm³ of the rough titre, THEN add dropwise (with continuous swirling). Close the tap immediately when the indicator just turns yellow PERMANENTLY (no swirling-back to red).

Repeat until concordant. Concordant titres are those agreeing to within 0.10 cm³. Continue until you have at least two concordant titres.

Calculation of mean titre. Take the mean of the concordant titres only (excluding the rough): $\text{mean titre} = (V_1 + V_2)/2$ . This gives the volume of NaOH that exactly neutralises 25.0 cm³ of the acid.

Safety precaution. Wear safety goggles (acid + alkali can damage eyes); use a pipette filler (never mouth-pipette); tie back long hair.
Why this scores
Edexcel 6-mark mark scheme: 1 for naming pipette (with filler) + 25.0 cm³; 1 for burette + reading to 0.05 cm³; 1 for indicator + endpoint colour change; 1 for adding dropwise near endpoint while swirling; 1 for concordant within 0.10 cm³ + at least 2; 1 for one safety precaution. Watch out for being too vague — 'use indicator' loses the mark; you must say WHICH indicator AND the colour change.
Question 2
4CH1/2C-style — strong vs weak acid comparison (P content)5 marks
Equal volumes of 0.1 mol/dm³ HCl and 0.1 mol/dm³ ethanoic acid (CH₃COOH) are reacted separately with excess magnesium ribbon. Compare and explain: (a) the rate of fizzing during the reaction; (b) the final total volume of H₂ produced; (c) the pH of each acid before reaction. (5 marks)
Model answer
(a) Rate of fizzing. HCl reacts FASTER — there is more vigorous, rapid effervescence. Ethanoic acid reacts MORE SLOWLY — slower, gentler bubbling.

Explanation. HCl is a STRONG acid — fully ionised in water, so EVERY HCl molecule provides an H⁺ ion: $[\text{H}^+] = 0.1$ mol/dm³. Ethanoic acid is a WEAK acid — only ~1 % ionised at equilibrium: $\text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+$ , so $[\text{H}^+] \approx 0.001$ mol/dm³ at any instant. Higher [H⁺] → more frequent successful collisions of H⁺ with Mg → faster rate.

(b) Total volume of H₂. THE SAME for both acids (if both are in excess of Mg, or if Mg is in excess of both).

Explanation. The total moles of acid are equal ( $n = cV$ ; same c, same V). As H⁺ is consumed in the reaction, the weak-acid equilibrium shifts to the RIGHT (Le Chatelier's principle: removing H⁺ → equilibrium replaces it). Eventually all CH₃COOH ionises and reacts. So the SAME total moles of H₂ are produced from each acid — just at different rates.

$\text{Mg} + 2\text{H}^+ \rightarrow \text{Mg}^{2+} + \text{H}_2$

(c) pH. HCl has a LOWER pH (more acidic) — typically pH ≈ 1 at 0.1 mol/dm³. Ethanoic acid has a HIGHER pH — typically pH ≈ 3 at 0.1 mol/dm³ (still acidic but less so).

Explanation. pH depends on $[\text{H}^+]$ : lower pH = higher [H⁺]. Since HCl gives ~100× more H⁺ than ethanoic acid at the same concentration, its pH is ~2 units lower (pH 1 vs pH 3).

Summary table.

Property HCl (strong) CH₃COOH (weak)
Ionisation Full (100 %) Partial (~1 %)
[H⁺] at 0.1 mol/dm³ 0.1 mol/dm³ ~10⁻³ mol/dm³
pH ~1 ~3
Rate with Mg Fast / vigorous Slow
Total H₂ produced Same Same
Why this scores
Edexcel 5-mark mark scheme (Paper 2 P content): 1 for HCl faster + correct reason ([H⁺] higher); 1 for SAME total H₂ produced; 1 for explaining equilibrium shift / same moles of acid; 1 for HCl lower pH; 1 for explanation of pH in terms of [H⁺]. The 'same total volume' point catches many students out — they assume the weak acid produces less.
Question 3
4CH1/1C-style — indicators and pH scale5 marks
Describe the colour changes of litmus, methyl orange, phenolphthalein and universal indicator in solutions of pH 1, pH 7 and pH 13. Explain why universal indicator is NOT suitable for titrations. (5 marks)
Model answer
Indicator colour changes.

Indicator pH 1 (strong acid) pH 7 (neutral) pH 13 (strong alkali)
Litmus Red Purple Blue
Methyl orange Red Yellow Yellow
Phenolphthalein Colourless Colourless Pink
Universal indicator Red Green Purple

Universal indicator is a MIXTURE of several indicators; the colour blends through the rainbow:

Red (pH 0-3) — strong acid.

Orange / yellow (pH 4-6) — weak acid.

Green (pH 7) — neutral.

Blue (pH 8-11) — weak alkali.

Purple (pH 12-14) — strong alkali.

Why universal indicator is NOT used in titrations. Universal indicator changes colour GRADUALLY through many colours over a wide pH range. There is no single SHARP endpoint — you can't tell exactly when the acid has been neutralised. A good titration indicator (e.g. methyl orange or phenolphthalein) gives ONE sharp colour change at a specific pH (~4 or ~9), so the endpoint is unambiguous: the very first drop that causes the colour change is the endpoint.

Endpoint pH ranges. Methyl orange: pH 3.1 → 4.4 (red → yellow). Phenolphthalein: pH 8.3 → 10.0 (colourless → pink). Both are useful for strong acid + strong base titrations.
Why this scores
Edexcel 5-mark mark scheme: 1 each for litmus, methyl orange, phenolphthalein, universal indicator colours (at all three pH values); 1 for explanation of why universal isn't used in titration (gradual change, no sharp endpoint). Litmus is also UNSUITABLE for titration for the same reason — the colour change is too gradual.
Question 4
4CH1/2C-style — reactions of acids6 marks
Hydrochloric acid reacts with three different substances: zinc metal, magnesium oxide, and sodium carbonate. (a) Write a balanced equation for each reaction. (b) For each reaction, name the SALT produced and identify the OTHER product(s). (c) Describe ONE test that would distinguish the gas given off in (i) from the gas given off in (iii). (6 marks)
Model answer
(a) Balanced equations.

Zinc + HCl (metal + acid → salt + hydrogen): $\text{Zn(s)} + 2\text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}$

Magnesium oxide + HCl (base + acid → salt + water — neutralisation): $\text{MgO(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{O(l)}$

Sodium carbonate + HCl (carbonate + acid → salt + water + CO₂): $\text{Na}_2\text{CO}_3\text{(s)} + 2\text{HCl(aq)} \rightarrow 2\text{NaCl(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

(b) Salts and other products.

Reaction Salt Other products
Zn + HCl Zinc chloride (ZnCl₂) Hydrogen gas (H₂)
MgO + HCl Magnesium chloride (MgCl₂) Water (H₂O)
Na₂CO₃ + HCl Sodium chloride (NaCl) Water + CO₂

(The salt is named based on the acid: HCl always gives a chloride salt; H₂SO₄ gives a sulfate; HNO₃ gives a nitrate.)

(c) Distinguishing the gases (H₂ vs CO₂).

Hydrogen (H₂) from reaction (i): test with a lit splint → squeaky pop. Lit splint plunged into a test tube of the gas; if H₂, you hear a 'pop' sound.

Carbon dioxide (CO₂) from reaction (iii): bubble the gas through limewater (Ca(OH)₂ solution) → goes milky / cloudy (white precipitate of CaCO₃ forms).

So a lit splint plunged into the gas from reaction (i) would give a 'pop' (H₂); the same lit splint plunged into the gas from reaction (iii) would just be extinguished (CO₂ does not burn). Or alternatively, bubble each gas through limewater: H₂ has no effect; CO₂ turns it cloudy.
Why this scores
Edexcel 6-mark mark scheme: 1 each for three balanced equations (including states for full marks); 1 for salt names; 1 each for two gas tests (H₂ squeaky pop, CO₂ limewater). Common error: leaving Na₂CO₃ + HCl unbalanced — note the 2HCl on the left. Salts: chloride / sulfate / nitrate naming convention must be learned.

Property	HCl (strong)	CH₃COOH (weak)
Ionisation	Full (100 %)	Partial (~1 %)
[H⁺] at 0.1 mol/dm³	0.1 mol/dm³	~10⁻³ mol/dm³
pH	~1	~3
Rate with Mg	Fast / vigorous	Slow
Total H₂ produced	Same	Same

Indicator	pH 1 (strong acid)	pH 7 (neutral)	pH 13 (strong alkali)
Litmus	Red	Purple	Blue
Methyl orange	Red	Yellow	Yellow
Phenolphthalein	Colourless	Colourless	Pink
Universal indicator	Red	Green	Purple

Reaction	Salt	Other products
Zn + HCl	Zinc chloride (ZnCl₂)	Hydrogen gas (H₂)
MgO + HCl	Magnesium chloride (MgCl₂)	Water (H₂O)
Na₂CO₃ + HCl	Sodium chloride (NaCl)	Water + CO₂

Key Definitions and Keywords — Acids, Alkalis and Titrations

Definitions to memorise and the exact keywords mark schemes credit for acids, alkalis and titrations answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Acid
Examiner keyword
A substance that releases / donates H⁺ ions (protons) when dissolved in water. Acidic solutions have pH < 7. Examples: HCl, H₂SO₄, HNO₃ (strong); CH₃COOH, citric acid, carbonic acid (weak).
Base
Examiner keyword
A substance that accepts H⁺ ions (a proton acceptor). Metal oxides (CuO, MgO) and metal hydroxides (NaOH, KOH) are bases. NOT all bases are soluble — only the SOLUBLE ones are alkalis.
Alkali
Examiner keyword
A SOLUBLE base — releases OH⁻ ions when dissolved in water. Alkaline solutions have pH > 7. Examples: NaOH, KOH, Ca(OH)₂ (slightly soluble), ammonia solution.
Neutralisation
Examiner keyword
The reaction between an acid and a base/alkali to give a salt and water only. Ionic: H⁺(aq) + OH⁻(aq) → H₂O(l). pH moves toward 7.
Strong acid (Paper 2 P content)
Examiner keyword
An acid that is FULLY (100 %) ionised in aqueous solution. Every molecule dissociates into ions. Examples: HCl, H₂SO₄, HNO₃. Single arrow in equation: HCl → H⁺ + Cl⁻.
Weak acid (Paper 2 P content)
Examiner keyword
An acid that is PARTIALLY ionised in aqueous solution (typically ~1 % at equilibrium). Examples: ethanoic (CH₃COOH), citric, carbonic (H₂CO₃). Equilibrium arrows: CH₃COOH ⇌ CH₃COO⁻ + H⁺.
pH
Examiner keyword
A scale (0-14) measuring acidity / alkalinity based on [H⁺]. pH 0-6 = acidic; pH 7 = neutral (pure water); pH 8-14 = alkaline. Lower pH means HIGHER [H⁺].
Indicator
Examiner keyword
A substance that changes colour depending on pH. Used to detect the endpoint in titrations or to estimate the pH of a solution. Litmus, methyl orange, phenolphthalein and universal indicator are the main 4CH1 indicators.
Titration
Examiner keyword
A technique for finding the concentration of a solution by reacting a measured volume with another solution of known concentration, using an indicator to show the endpoint. Edexcel 4CH1 required practical.
Endpoint
Examiner keyword
The point in a titration at which the indicator JUST changes colour permanently — corresponds to neutralisation (acid : base in stoichiometric ratio).
Concordant titres
Examiner keyword
Two or more titres that agree to within 0.10 cm³ of each other (Edexcel definition). Only concordant titres are used to calculate the mean titre; the rough titre is discarded.
Mole (in titration calculations)
An amount of substance containing 6.02 × 10²³ particles. In titrations: $n = cV/1000$ when V in cm³; $n = cV$ when V in dm³. Always work in mol via the balanced equation.

Common Mistakes and Misconceptions — Acids, Alkalis and Titrations

The traps other students keep falling into on acids, alkalis and titrations questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Confusing 'alkali' with 'base'
Why it happens
Using the two terms as synonyms.
How to avoid it
Memorise: 'alkali = soluble base'. CuO is a base (reacts with acid) but NOT an alkali (insoluble). NaOH is BOTH a base AND an alkali.
✕Confusing 'strong acid' with 'concentrated acid'
4CH1 Examiner Reports 2022-2024
Why it happens
Everyday language uses them interchangeably.
How to avoid it
STRONG / WEAK refers to the % IONISATION (HCl strong = fully ionises). CONCENTRATED / DILUTE refers to mol/dm³. A dilute strong acid (e.g. 0.001 mol/dm³ HCl) can have a HIGHER pH than a concentrated weak acid (e.g. 1 mol/dm³ ethanoic).
✕Using universal indicator for titration
Why it happens
Universal indicator is the most familiar from class.
How to avoid it
Universal indicator changes colour gradually through many colours over a wide pH range — no sharp endpoint. For titration use methyl orange (sharp pH 4) or phenolphthalein (sharp pH 9) only.
✕Including the rough titre when calculating the mean
4CH1 Examiner Reports 2022-2024
Why it happens
Treating all titrations equally.
How to avoid it
The rough titre is to find APPROXIMATE endpoint — it overshoots. Exclude it. Calculate the mean using only concordant titres (those within 0.10 cm³ of each other).
✕Forgetting to convert cm³ to dm³ in $n = cV$
Why it happens
Substituting cm³ directly into $n = cV$ formula.
How to avoid it
Either use $n = cV/1000$ (V in cm³) OR convert V to dm³ first (÷1000). Always check the units of the final n (should be mol, not mol × 1000).
✕Forgetting the 1:2 ratio for H₂SO₄ : NaOH titrations
Why it happens
Assuming all acid + base titrations are 1:1.
How to avoid it
Write the balanced equation FIRST. H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O → 1:2 ratio (so n(H₂SO₄) = n(NaOH)/2). Don't assume 1:1.
✕Writing Cu + acid → CuCl₂ + H₂
Why it happens
Applying the metal + acid rule indiscriminately.
How to avoid it
Only metals ABOVE H in the reactivity series react with acid to give H₂. Cu, Ag, Au are BELOW H — they do NOT react with dilute HCl or dilute H₂SO₄. (They do react with concentrated HNO₃ but that's an oxidising acid — different reaction.)

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

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Video lesson

Short walkthrough of the concepts students most often get stuck on.

Acids, Alkalis and Titrations — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Acids, Alkalis and Titrations

Short Study Notes in the form of Slides

Short Study Notes — Acids, Alkalis and Titrations

Detailed Notes

What you’ll learn

How it’s examined

1Defining an acid and an alkali (3 marks, Paper 1)

2Strong vs weak acids — HCl vs CH₃COOH at the same concentration (4 marks, Paper 2)

3Titration of HCl against NaOH — apparatus and concordant titres (6 marks, Paper 1)

4Calculating the concentration of HCl from titration data (6 marks, Paper 2)

5Reactions of dilute sulfuric acid with magnesium, copper(II) oxide and calcium carbonate (5 marks, Paper 1)

Moles from concentration and volume (cm³)

Moles from concentration and volume (dm³)

Stoichiometric titration equation (1:1 case)

Titration with stoichiometric ratio (general)

Concentration conversion mol/dm³ ↔ g/dm³

Acid

Base

Alkali

Neutralisation

Strong acid (Paper 2 P content)

Weak acid (Paper 2 P content)

pH

Indicator

Titration

Endpoint

Concordant titres

Mole (in titration calculations)

✕Confusing 'alkali' with 'base'

✕Confusing 'strong acid' with 'concentrated acid'

✕Using universal indicator for titration

✕Including the rough titre when calculating the mean

✕Forgetting to convert cm³ to dm³ in n=cVn = cVn=cV

✕Forgetting the 1:2 ratio for H₂SO₄ : NaOH titrations

✕Writing Cu + acid → CuCl₂ + H₂

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Acids, Alkalis and Titrations — frequently asked questions

✕Forgetting to convert cm³ to dm³ in $n = cV$