Higher Level Revision Notes

Leaving Cert Higher Level Chemistry

Leaving Cert Higher Level Chemistry Chapter 4: Volumetric Analysis and Titrations

You learn how titrations work, why exact measuring matters, and how to use the results to find unknown concentrations.

What this page doesTurns a difficult chapter into clear notes, diagrams, and one simple interactive tool.

Student promisePlain English first, exam language second.

Curriculum

Irish Leaving Certificate (ILC)

Level

Higher Level

Subject

Chemistry

Chapter

Chapter 4 — Volumetric Analysis and Titrations

Simple English Summary

You learn how titrations work, why exact measuring matters, and how to use the results to find unknown concentrations.

Teacher voice: Read the ideas first, then use the detailed notes and diagrams to lock in the exam wording.

What To Focus On

Prepare standard solutions carefully.
Understand the titration endpoint.
Use the concentration formula correctly.
Read concordant titres like an examiner.

DisclaimerThis publication is an independent educational resource developed by ExamsLogic and compiled for student revision. It is based on publicly available official curricula and is not endorsed by any examination board.

Open Simulators Jump To Notes

Quick Simpler Start

In one sentence

You learn how titrations work, why exact measuring matters, and how to use the results to find unknown concentrations.

Exam habit

When you revise this chapter, ask yourself: "Can I explain this to a friend in one easy paragraph?" If yes, you are in a good place.

Interactive Simulator

This small tool gives you a quick visual check before you move deeper into the chapter notes.

Simulator

Titration Calculator

Use concentration and volume to see the amount of substance in a simple, exam-friendly way.

Concentration (mol/L)Volume (mL)

Enter the values to calculate moles.

Chapter 4: Volumetric Analysis (Titrations)

1. Objectives

Calculate concentration using molarity, grams per litre and percentage expressions.
Explain the purpose and characteristics of a primary standard.
Describe how to prepare a standard solution accurately.
Carry out acid-base titration calculations using mole ratios.
Explain indicator choice and endpoint colour change.
Understand iodine-thiosulfate titrations and potassium permanganate titrations.
Apply redox titration calculations involving electron transfer.

2. Key Definitions

Term	Exam-ready meaning
Volumetric analysis	A quantitative technique used to find the concentration of a solution by measuring reacting volumes accurately.
Titration	A method where one solution is added from a burette to another until the reaction is just complete.
Standard solution	A solution whose concentration is accurately known.
Primary standard	A pure, stable solid that can be weighed directly to prepare a standard solution.
Endpoint	The point where the indicator changes colour, showing the reaction is complete.
Concordant titres	Titration readings that are very close together, usually within about 0.10–0.20 cm³.

3. Visual Learning Zone

A. Standard titration setup

The burette gives accurate volume added. The pipette gives an accurate fixed volume in the flask.

B. Reading the burette: avoid parallax error

C. Acid-base titration calculation map

4. Concentration of Solutions

Molarity = moles / volume in litres
mol/L = mol dm⁻³

Expression	Meaning	Example
Molarity	Moles of solute per litre of solution	0.10 mol/L NaOH
g/L	Mass of solute in grams per litre	5 g/L glucose
% w/v	grams per 100 cm³ solution	5% w/v = 5 g in 100 cm³
% v/v	cm³ per 100 cm³ solution	70% v/v ethanol
% w/w	grams per 100 g mixture	10% w/w salt mixture

Examiner Tip: Always convert cm³ to litres before using molarity. 250 cm³ = 0.250 L. This is one of the easiest marks to lose.

5. Primary Standards

A primary standard is used to prepare a solution of accurately known concentration. It must have reliable properties because any error in its mass affects every calculation after that.

Good primary standard:

Very pure
Stable in air
Does not absorb water from air
High molar mass to reduce weighing error
Soluble in water

Making a standard solution:

Weigh solid accurately.
Dissolve in small volume of deionised water.
Transfer to volumetric flask using a funnel.
Rinse beaker and funnel into flask.
Make up to calibration mark.
Stopper and invert several times.

Trap: Do not say “fill to the top.” The correct phrase is “make up to the calibration mark with deionised water.”

6. Acid-Base Titrations

Acid-base titrations are based on neutralisation. The exact volume needed to react completely is found from repeated titres.

V₁M₁ / n₁ = V₂M₂ / n₂
n = coefficient from balanced equation

Indicator	Best use	Typical colour change
Methyl orange	Strong acid vs weak base OR strong acid vs strong base	Red/orange/yellow
Phenolphthalein	Weak acid vs strong base OR strong acid vs strong base	Colourless to pink
Litmus	Rough indication only	Red/blue

Worked Example 1: Acid-base titration
25.0 cm³ of 0.100 mol/L NaOH is neutralised by 20.0 cm³ of HCl. Find the concentration of HCl.

Equation: HCl + NaOH → NaCl + H₂O
Mole ratio = 1 : 1

V₁M₁ / n₁ = V₂M₂ / n₂
20.0 × M(HCl) / 1 = 25.0 × 0.100 / 1
M(HCl) = 2.50 / 20.0 = 0.125 mol/L

Answer: 0.125 mol/L

7. Redox Titrations

Iodine-thiosulfate titrations

Iodine is titrated with sodium thiosulfate. Starch is usually added near the end point, not at the start, because the iodine-starch complex can be too stable.

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

Endpoint: blue-black colour disappears and the solution becomes colourless.

HL Only: Potassium permanganate titrations

KMnO₄ titrations: Potassium permanganate is purple and acts as its own indicator. In acidic solution, MnO₄⁻ is reduced to Mn²⁺. The endpoint is a permanent pale pink colour.

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

D. Redox endpoint colours

7A. Brønsted-Lowry Theory and Conjugate Pairs

At Higher Level, acids and bases are also explained using the Brønsted-Lowry theory.

Brønsted-Lowry acid: a proton (H⁺) donor.
Brønsted-Lowry base: a proton (H⁺) acceptor.

Species	Role	Why?
HCl	Acid	Donates H⁺ to water.
H₂O	Base	Accepts H⁺ from HCl.
NH₃	Base	Accepts H⁺ to form NH₄⁺.
CH₃COOH	Acid	Donates H⁺ to water.

HCl + H₂O → H₃O⁺ + Cl⁻

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

A conjugate acid-base pair differs by one proton only.

Acid	Conjugate Base
HCl	Cl⁻
CH₃COOH	CH₃COO⁻
NH₄⁺	NH₃
H₃O⁺	H₂O

Examiner Tip: When asked for a conjugate pair, remove or add exactly one H⁺. Nothing else changes.

Examiner Trap: Arrhenius theory works only in water. Brønsted-Lowry is broader because it focuses on proton transfer.

7B. Relative Molecular Mass and Water of Crystallisation from Titration Data

Titration data can be used to find the relative molecular mass of an acid or base, and also the amount of water of crystallisation in a hydrated salt.

moles = concentration × volume (in L)

Mr = mass / moles

Worked Example 1: Relative molecular mass from titration
25.0 cm³ of 0.100 mol L⁻¹ NaOH neutralises 0.305 g of a monobasic acid dissolved in water.

Step 1: moles of NaOH = 0.100 × 0.0250 = 0.00250 mol
Step 2: monobasic acid reacts 1:1 with NaOH, so moles of acid = 0.00250 mol
Step 3: Mr of acid = 0.305 / 0.00250 = 122

Answer: relative molecular mass = 122

Worked Example 2: Water of crystallisation
A hydrated sodium carbonate sample, Na₂CO₃·xH₂O, has mass 2.86 g and is made up to 250 cm³. A 25.0 cm³ portion requires 20.0 cm³ of 0.100 mol L⁻¹ HCl. Find x.

Reaction: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂
Moles HCl in titre = 0.100 × 0.0200 = 0.00200 mol
Moles Na₂CO₃ in 25.0 cm³ = 0.00200 / 2 = 0.00100 mol
Moles in full 250 cm³ = 0.0100 mol
Molar mass of hydrated salt = 2.86 / 0.0100 = 286 g mol⁻¹
Molar mass of Na₂CO₃ = 106
Mass due to water = 286 − 106 = 180 = 10 × 18

Answer: x = 10

Examiner Tip: For hydrated salts, find the experimental molar mass first, then compare it with the anhydrous molar mass.

7C. Mandatory Practical Links

Chapter 4 at Higher Level should still connect clearly to the practical titration work students are expected to perform.

Practical	Main Idea	Exam Focus
Preparation of standard sodium carbonate solution	Use an accurately weighed pure primary standard.	Volumetric flask use, weighing by difference, dissolving and making up to the mark.
Standardisation of hydrochloric acid	Use standard sodium carbonate to determine unknown acid concentration.	Burette/pipette technique, indicator choice, concordant titres, calculations.
Hydrochloric acid with sodium hydroxide	Neutralisation titration and salt preparation.	Indicator endpoint, repeating without indicator if a pure salt is needed.
Iodine–thiosulfate titration	Redox titration using starch near the endpoint.	Blue-black to colourless endpoint, careful endpoint technique.
Permanganate titration	KMnO₄ acts as its own indicator.	Permanent faint pink endpoint.

Good practical wording: rinse burette with the solution it will contain, rinse pipette with the solution to be transferred, swirl continuously during titration, wash down flask walls with distilled water, and use concordant titres only.

Common Mistake: Saying distilled water left inside the flask changes the titre. It does not change the number of moles already pipetted into the flask.

8. Examiner Secrets

Average only concordant titres, not the rough titration.
Use the balanced equation before inserting numbers into the formula.
Use cm³ consistently on both sides in the titration formula, or convert both to litres.
For KMnO₄, mention acidified solution if the question asks for conditions.
For iodine titration, starch is added near the endpoint.

9. Common Mistakes

Using the rough titre in the average.
Forgetting that 25.0 cm³ = 0.0250 L.
Using the wrong mole ratio from the equation.
Writing “pink to colourless” for KMnO₄ endpoint instead of “colourless/pale solution to permanent pale pink.”
Adding starch too early in iodine-thiosulfate titrations.

10. Exam Practice Questions

Q1. Calculate the molarity of a solution containing 5.30 g of sodium carbonate, Na₂CO₃, in 250 cm³ of solution. [4 marks]

Q2. 25.0 cm³ of 0.0800 mol/L sulfuric acid neutralises 20.0 cm³ of sodium hydroxide. Calculate the concentration of NaOH. Equation: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O. [5 marks]

Q3. Explain why potassium permanganate does not require an external indicator in redox titrations. [2 marks]

11. MCQs with Explanations

1. Which is a key property of a primary standard?
A. It is unstable in air B. It absorbs water easily C. It is very pure D. It must be a liquid

Answer: C
Why: A primary standard must be very pure and stable. A and B would make accurate weighing unreliable. D is wrong because many primary standards are solids.

2. The rough titre should be:
A. Included in the average B. Ignored when calculating the average C. Used as the final answer D. Always lower than the accurate titres

Answer: B
The rough titration is only used to estimate the endpoint. Accurate concordant titres are used for calculations.

12. Higher Level Challenge

Challenge: In acid solution, 20.0 cm³ of Fe²⁺ solution reacts with 18.0 cm³ of 0.0200 mol/L KMnO₄. Given that 1 mole MnO₄⁻ reacts with 5 moles Fe²⁺, calculate the concentration of Fe²⁺.

Solution outline:
Moles MnO₄⁻ = 0.0200 × 0.0180 = 0.000360 mol
Moles Fe²⁺ = 5 × 0.000360 = 0.00180 mol
Volume Fe²⁺ = 0.0200 L
Concentration Fe²⁺ = 0.00180 / 0.0200 = 0.0900 mol/L

13. Last-Minute Revision Sheet

Molarity = moles / litres.
Always convert cm³ to L for mole calculations.
Use only concordant titres.
Primary standard: pure, stable, soluble, high molar mass.
Make standard solution in a volumetric flask.
Acid-base titration: use balanced equation and mole ratio.
Iodine-thiosulfate endpoint: blue-black to colourless.
KMnO₄ endpoint: permanent pale pink.

14. Self-Assessment Checklist

I can calculate molarity from mass and volume.
I can explain the properties of a primary standard.
I can describe how to make a standard solution.
I can calculate unknown concentration from a titration.
I can choose the correct mole ratio from a balanced equation.
I can describe iodine-thiosulfate and KMnO₄ endpoints.

15. Answers / Mark Scheme

Q1
Mr Na₂CO₃ = 106 g/mol [1]
Moles = 5.30 / 106 = 0.0500 mol [1]
Volume = 250 cm³ = 0.250 L [1]
Molarity = 0.0500 / 0.250 = 0.200 mol/L [1]

Q2
H₂SO₄ : NaOH = 1 : 2 [1]
Moles H₂SO₄ = 0.0800 × 0.0250 = 0.00200 mol [1]
Moles NaOH = 2 × 0.00200 = 0.00400 mol [1]
Volume NaOH = 0.0200 L [1]
Concentration NaOH = 0.00400 / 0.0200 = 0.200 mol/L [1]

Q3
KMnO₄ is strongly purple [1]. It acts as its own indicator; endpoint is permanent pale pink [1].