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Curriculum: Irish Leaving Certificate (ILC)
Level: Ordinary Level (OL)
Subject: Chemistry
Chapter 4: Volumetric Analysis
Status: Complete Syllabus-Aligned Chapter
Prepared By: ExamsLogic Academic Team
Chapter 4 Complete: Volumetric Analysis
This complete chapter combines concentration, standard solutions, acids and bases, titration technique, acid-base titration calculations, and all three required Ordinary Level practicals for Chapter 4.
Syllabus Focus: The Ordinary Level core for this chapter is concentration of solutions, acids and bases, volumetric analysis, and the mandatory experiments 4.1, 4.2 and 4.2A.
ExamsLogic strategy: Learn this chapter in three layers: definitions → technique → calculations. Students usually lose marks on the middle layer, not the maths.
Common mistake: Using 25 instead of 0.025 in a molarity formula makes the answer 1000 times too large.
2. Concentration in mol/L and g/L
2.1 Concentration in mol/L
Concentration = moles ÷ volume
c = n ÷ V
At Ordinary Level, the volume must be in litres or dm³.
Worked example 1:
0.20 mol NaCl is dissolved to make 500 cm³ of solution.
V = 0.500 dm³
c = n ÷ V = 0.20 ÷ 0.500 = 0.40 mol/L
Worked example 2:
How many moles are present in 250 cm³ of 0.10 mol/L HCl?
V = 0.250 dm³
n = c × V = 0.10 × 0.250 = 0.025 mol
2.2 Concentration in g/L
Concentration (g/L) = mass (g) ÷ volume (L)
Worked example:
5.0 g sodium chloride is dissolved to make 250 cm³ of solution.
V = 0.250 L
Concentration = 5.0 ÷ 0.250 = 20 g/L
2.3 Converting between g/L and mol/L
Concentration (mol/L) = concentration (g/L) ÷ molar mass
Worked example:
A glucose solution has concentration 18 g/L. Glucose has molar mass 180 g/mol.
18 ÷ 180 = 0.10 mol/L
Examiner tip: If the question gives grams and asks for mol/L, first convert to g/L or to moles. Do not jump straight to the answer.
3. Percentage Concentration and Colour Intensity
3.1 Percentage concentration
Type
Meaning
Example
% w/v
grams of solute in 100 cm³ of solution
5% w/v glucose
% v/v
cm³ of liquid solute in 100 cm³ of solution
12% v/v wine / 70% v/v ethanol
% w/w
grams of solute in 100 g of mixture
10% w/w ointment
Worked example:
A solution is labelled 2% w/v sodium chloride. How much sodium chloride is present in 250 cm³?
2 g in 100 cm³, so in 250 cm³ there is 2 × 2.5 = 5 g.
3.2 Colour intensity as a function of concentration
For a coloured solution, a higher concentration usually means a darker or more intense colour. When the solution is diluted, the colour becomes paler.
This is only a simple OL treatment, but it is a required syllabus idea.
Exam link: If two coloured solutions are made from the same substance and one looks darker, the darker one is usually more concentrated.
Common mistake: Do not confuse colour intensity with a different substance. The substance may be the same, only the concentration changes.
4. Primary Standards and Standard Solutions
4.1 Primary standards
A primary standard is a pure, stable substance that can be weighed accurately and used to prepare a solution of accurately known concentration.
Property
Why it matters
Very pure
The measured mass must represent the substance only.
Stable in air
It should not react easily with oxygen, water vapour or carbon dioxide.
Soluble in water
It must dissolve to make the solution.
Not hygroscopic
It must not absorb water from the air.
Fairly high molar mass
This reduces percentage weighing error.
Examiner trap: Sodium hydroxide is not a good primary standard because it absorbs water and carbon dioxide from the air.
4.2 Standard solutions
A standard solution has an accurately known concentration. It is often made from a primary standard and then used in a titration.
A volumetric flask is used because it measures one accurate final volume.
Accurately weigh the required mass of solid.
Dissolve it in a beaker using deionised water.
Transfer the solution to a volumetric flask with a funnel.
Rinse the beaker and funnel into the flask so no solute is lost.
Add water until the bottom of the meniscus is on the calibration mark.
Stopper and invert the flask several times to mix thoroughly.
Common mistake: Forgetting to rinse the beaker and funnel leaves some solute behind and makes the final solution too dilute.
5. Acids, Bases, Salts and Neutralisation
5.1 Acids, bases and salts
Term
Simple OL meaning
Example
Acid
A substance that forms hydrogen ions in water.
Hydrochloric acid, vinegar
Base
A substance that reacts with an acid.
Sodium hydroxide, magnesium oxide
Alkali
A soluble base.
Sodium hydroxide solution
Salt
A compound formed when the H⁺ of an acid is replaced by a metal or ammonium ion.
Sodium chloride
5.2 Neutralisation
Neutralisation is the reaction between an acid and a base to produce a salt and usually water.
acid + base → salt + water
HCl + NaOH → NaCl + H₂O
This is one of the key ideas behind acid-base titrations.
Example: HCl + NaOH → NaCl + H₂O is a neutralisation reaction because an acid reacts with a base to form a salt and water.
6. Arrhenius Theory and Household Examples
6.1 Arrhenius theory
Arrhenius acid: produces H⁺ in water
Arrhenius base: produces OH⁻ in water
Examples:
HCl(aq) is an Arrhenius acid because it forms H⁺ ions in water.
NaOH(aq) is an Arrhenius base because it forms OH⁻ ions in water.
Examiner trap: At OL for this chapter, stick to the Arrhenius theory. Do not drift into Higher Level ideas unless the teacher asks.
6.2 Household acids and bases
Two household acids
Vinegar (contains ethanoic acid)
Lemon juice (contains citric acid)
Two household bases
Baking soda / sodium hydrogencarbonate
Household ammonia / cleaning alkali
6.3 Everyday examples of neutralisation
Example
What is happening?
Lime used in agriculture
Acidic soil is neutralised so crops grow better.
Stomach powders / antacids
Excess stomach acid is neutralised.
Examiner secret: When asked for an everyday neutralisation example, state both the substances and the purpose.
7. Volumetric Apparatus and Correct Titrimetric Procedure
7.1 Apparatus used in volumetric analysis
Apparatus
Use
Volumetric flask
Making a standard solution to one exact final volume.
Pipette
Measuring one fixed accurate volume into the conical flask.
Pipette filler
Safe filling of pipettes.
Burette
Delivering a variable measured volume during titration.
Conical flask
Holding the reacting solution during titration.
White tile
Makes the indicator colour change easier to see.
7.2 Correct titrimetric procedure
Rinse the burette with the solution it will contain.
Fill the burette and remove any air bubble from the tip.
Record the initial burette reading at eye level.
Rinse the pipette with the solution it will measure.
Use the pipette to transfer a fixed volume into the conical flask.
Add a few drops of suitable indicator.
Run the burette solution into the flask while swirling.
Near the end-point, add the solution drop by drop.
Record the final burette reading.
Repeat until concordant titres are obtained.
The burette delivers one solution into the conical flask until the end-point is reached.
Examiner tip: Questions on practical errors often reward answers like rinse the burette, read the meniscus at eye level, use a white tile, and swirl during titration.
8. Indicators, Titres and Concordant Results
8.1 Indicators
Indicator
Acid colour
Alkali colour
Common use
Methyl orange
Red
Yellow
Strong acid + weak base, or strong acid + strong base
Phenolphthalein
Colourless
Pink
Weak acid + strong base, or strong acid + strong base
8.2 Titre and average titre
Titre = final reading − initial reading
Average titre = mean of concordant titres only
The titre is not the final reading on its own.
Concordant titres: 24.80 cm³, 24.85 cm³, 24.90 cm³ are concordant because they are close together.
Common mistake: Do not average the rough titre with the accurate titres.
9. Acid-Base Titration Calculations
Most Ordinary Level volumetric problems are solved with the same structure:
Convert volume to dm³.
Find moles of the known solution.
Use the balanced equation ratio.
Find the unknown concentration.
moles = concentration × volume
\( \frac{M_1V_1}{n_1} = \frac{M_2V_2}{n_2} \)
Formula method: In OL exams, balanced equations are given. Your job is to use the correct mole ratio.
9.1 Worked example: 1:1 reaction
Question: 25.0 cm³ of sodium hydroxide is neutralised by 20.0 cm³ of 0.100 mol/L HCl. Find the concentration of NaOH.
Common mistake: Forgetting the 2 in the balanced equation gives an HCl concentration that is only half-right.
12. Mandatory Experiment 4.2A
Hydrochloric Acid / Sodium Hydroxide Titration and Making Sodium Chloride
Aim
To titrate hydrochloric acid with sodium hydroxide and use the result to prepare sodium chloride.
Balanced equation
HCl + NaOH → NaCl + H₂O
Part A: Titration
Pipette a fixed volume of sodium hydroxide into the conical flask.
Add a few drops of indicator.
Titrate with hydrochloric acid from the burette to the end-point.
Repeat until concordant titres are obtained.
Part B: Making sodium chloride
Repeat the titration using the exact same measured volumes but without indicator.
When neutralisation is complete, the flask contains sodium chloride solution and water.
Transfer the neutral solution to an evaporating dish.
Evaporate gently to concentrate the solution.
Allow crystals of sodium chloride to form.
Filter if needed and dry the crystals.
Important reason: The final salt is prepared without indicator so the product is not contaminated by indicator dye.
This practical links titration technique to salt preparation.
Common mistake: Students sometimes try to evaporate the solution from the indicator titration. That would contaminate the sodium chloride.
13. Extension: Redox Titration Note
Extension only: Some of your original part files included iodine-thiosulfate titrations. That is useful enrichment, but the Ordinary Level Chapter 4 core is built around acid-base titrations, not redox titrations.
A redox titration is still based on measured volumes and mole ratios, but the reacting substances undergo oxidation and reduction. One common example is:
I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻
In that reaction, 1 mole of iodine reacts with 2 moles of thiosulfate. Starch is used as an indicator because it gives a strong blue-black colour with iodine.
Why keep this section? It is good extension material and useful for stronger students, but it should be treated as enrichment beside the OL acid-base core.
14. Final Exam Review
14.1 One-page checklist
I can define solution, solute, solvent and concentration.
I can convert cm³ to dm³ correctly.
I can calculate concentration in mol/L and g/L.
I can explain % v/v and simple percentage concentration ideas.
I can describe colour intensity as a function of concentration.
I can define primary standard and standard solution.
I can explain acids, bases, salts and neutralisation.
I can state the Arrhenius theory for acids and bases.
I know two household acids and two household bases.
I can name everyday uses of neutralisation.
I can describe correct titrimetric procedure.
I can calculate titre and average titre.
I can solve acid-base titration calculations using the formula method.
I know Mandatory Experiments 4.1, 4.2 and 4.2A.
14.2 Quick-fire worked questions
Q1. Convert 32.5 cm³ to dm³. Answer: 0.0325 dm³
Q2. Find the concentration of a solution made by dissolving 0.050 mol in 250 cm³. Answer: 0.050 ÷ 0.250 = 0.20 mol/L
Q3. What is the titre if the burette goes from 1.20 cm³ to 25.80 cm³? Answer: 24.60 cm³
Q4. Give one everyday example of neutralisation. Answer: Antacids neutralise excess stomach acid.
14.3 Examiner tips, traps and common mistakes
Examiner tip: Always write the balanced equation before the titration calculation if it is relevant to the ratio.
Examiner tip: Read burette values at eye level and from the bottom of the meniscus.
Examiner trap: A rough titre is for finding the approximate end-point. It is not normally included in the average.
Common mistake: Students often know the chemistry but lose marks because they forget to convert cm³ to dm³.
Common mistake: Students say “acid + alkali makes water only.” The full answer is usually salt + water.
14.4 Chapter verdict
If you can do these three things, you are in good shape for Chapter 4:
Prepare a standard solution accurately.
Describe a full titration method clearly.
Carry out the mole-ratio calculation without unit mistakes.
Chapter 4 Complete
You have now completed Chapter 4: Volumetric Analysis for ILC Chemistry Ordinary Level.