Year 12 Chemistry Module 8 ⏱ ~35 min 5 MC · 3 Short Answer Lesson 1 of 16 Core

Titration for Quantitative Analysis

A quality-control chemist must decide whether an antacid tablet contains the amount of active base printed on its label. A measured titre, a balanced equation and a carefully chosen indicator turn one colour change into defensible quantitative evidence.

Today's hook: A tablet labelled as containing 680 mg of active base is sent for quality-control testing. The analyst adds a known excess of standard HCl, then titrates the acid left over with NaOH. How does the endpoint in the second titration reveal the mass of base that was originally in the tablet?

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Connect: from a colour change to evidence

Beyond the syllabus. This lesson is Module 6 revision rather than Module 8 core: acid–base titrations, indicator ranges and back titration are assessed under Module 6. Module 8's own requirement is precipitation titration. Keep your titration technique sharp, but do not count this lesson as new Module 8 exam content.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions

Prediction Before Technique

A pharmacist sends a crushed antacid tablet to the lab and asks: "How much base is actually in this dose?" The tablet reacts with acid, indicators change colour, and a titre value appears on the burette.

How could a chemist use a known acid or base to work out the amount of unknown base in the tablet?
What might go wrong if the indicator changes colour too early or too late relative to the true reaction point?

By the end of this lesson you will:

Know

What a titration measures and why it is used for unknown concentrations
How a precipitation titration signals its endpoint with a second precipitate, and why the Mohr method uses chromate indicator
The conditions a Mohr titration requires, and the interferences that invalidate it
The meaning of endpoint, equivalence point, concordant titres and back titration
The colour-change ranges of common acid-base indicators

Understand

Why mole relationships sit underneath every titration calculation
Why endpoint and equivalence point should be close, but are not identical ideas
Why indicator choice depends on the pH jump near equivalence

Can Do

Calculate an unknown concentration using n = cV and c = n/V
Interpret titration data, reject rough or non-concordant titres, and average reliable results
Explain how back titration can test the strength of an antacid tablet

Vocabulary flip cards +10 XP

Say the definition before you flip. Turn each card, then choose “Got it” or “Again”.

0 / 6 mastered

Term 1Standard solutionClick to reveal ↻

DefinitionA solution of accurately known concentration used as the titrant in a titration.

Term 2Equivalence pointClick to reveal ↻

DefinitionThe point where acid and base have reacted in the stoichiometric ratio of the balanced equation.

Term 3EndpointClick to reveal ↻

DefinitionThe observed indicator colour change used to estimate the equivalence point.

Term 4TitreClick to reveal ↻

DefinitionThe volume of standard solution delivered from the burette to reach the endpoint.

Term 5Primary standardClick to reveal ↻

DefinitionA highly pure, stable substance used to prepare a solution of accurately known concentration.

Term 6Indicator selectionClick to reveal ↻

DefinitionChoose an indicator whose transition range overlaps the steep pH change near equivalence.

Cross-lesson links: Titration calculations here use the mole-ratio method from Module 5. The equivalence-point concept reappears in Module 6 Lesson 13 (drug pKa and the Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA])). Back titration of antacid tablets connects forward to L11 (aspirin and paracetamol functional-group testing).

Safety: Wear safety glasses. The acids and bases used in titrations (e.g. HCl, NaOH) are corrosive, so handle them with care and rinse any spills with plenty of water. Phenolphthalein and methyl orange indicator solutions are ethanol-based and flammable, so keep them away from flames. Fill the burette below eye level and never pipette by mouth.

Core Content

What Module 8 actually requires from this lesson

The Module 8 syllabus names gravimetric analysis and precipitation titrations. Cards 1 to 3 build the titration technique and the mole pathway that every titration uses, and card 4 is the precipitation titration itself — that is the core of this lesson. Back titration and acid–base indicator choice (cards 5 and 6) are carried over from Module 6: worth revising, but they are not what Module 8 assesses here.

Explain: what titration measures

What Titration Measures Core

Known concentration + measured volume = unknown concentration

A titration is not "adding liquid until the colour changes". It is a quantitative method for counting moles through reaction stoichiometry.

In an acid-base titration, a solution of known concentration is added carefully from a burette to a measured volume of an unknown acid or base. When chemically equivalent amounts have reacted, the mole ratio in the balanced equation lets us determine the unknown quantity.

For a simple 1:1 neutralisation such as HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l), the moles of acid at equivalence equal the moles of base. If the reaction ratio is not 1:1, the balanced equation must be used explicitly.

A titration determines an unknown concentration by reacting the analyte with a standard solution of known concentration and using the titre volume plus stoichiometry to calculate the unknown. Key formulas: n = cV (moles = concentration × volume in litres) and c = n/V.

Pause, copy the highlighted definition into your book.

The titration calculation pathway

From titre to unknown concentration

3 moves • every calculation

01 • Find known moles n = cV Use the standard solution’s concentration and titre.mol = mol L⁻¹ × L

02 • Cross the equation n₁ : n₂ Transfer moles using coefficients in the balanced equation.1:1 only when coefficients say 1:1

03 • Find the unknown c = n / V Divide the unknown’s moles by its measured volume.mol L⁻¹ = mol ÷ L

Unit checkpoint: convert every mL value to L before using either formula.

HSC language: When describing titration in extended response answers, say that the concentration is determined by reacting the analyte with a standard solution of known concentration and using the titre volume plus stoichiometry to calculate the unknown.

A burette delivers a known solution into a measured aliquot in a conical flask until the indicator changes at the endpoint.

Known titrant is delivered from the burette into a measured aliquot of analyte in the conical flask. The endpoint is judged in the flask, often over a white tile so the first permanent colour change is easier to see.

In a titration, which vessel holds the standard solution of known concentration?

Method, Titre Values and Concordant Results Core

Accuracy depends on technique, not just maths

We just saw that titration links a standard solution to an unknown through mole ratios. That raises a question: if the method is sound, why do chemists repeat the experiment several times? This card answers it → reliable results depend on careful technique and concordant titres, not just getting one colour change.

A good titration is a controlled sequence: prepare carefully, add quickly at first, slow down near the endpoint, then trust only concordant results.

Rinse the burette with the titrant and the pipette with the analyte.
Pipette a fixed aliquot of the unknown into a conical flask.
Add a few drops of a suitable indicator.
Run titrant from the burette while swirling the flask.
Near the endpoint, add titrant dropwise until the colour change persists.
Record the initial and final burette readings, then calculate the titre.

A first run is usually a rough titre. It helps locate the endpoint region. Reliable calculations should then use concordant titres, meaning titres that closely agree with each other, typically within 0.10 mL.

Concordant titres agree within 0.10 mL. Only concordant titres are averaged for calculation. The rough titre is always excluded because it was not performed with the precision needed for an accurate result.

Pause, copy the highlighted rule into your book before the check below.

Common error: "Average every titre you recorded." Students think more numbers automatically improve accuracy. In reality, rough trials and obvious outliers should be excluded, because they distort the mean and reduce reliability.

Clinical anchor: In antacid testing, the difference between 23.45 mL and 24.80 mL is not trivial. A poor titre can make a tablet appear stronger or weaker than it really is, which matters if the dose is being quality-checked against a product claim.

A student records titres of 24.80 mL (rough), 23.45 mL, 23.40 mL and 23.50 mL. Which result should be excluded from the average?

Respond: calculate an unknown concentration

Calculating Unknown Concentration Core

Use moles first, then convert to concentration

We just saw that concordant titres give a reliable average volume. That raises a question: how do you turn that volume into a concentration? This card answers it → the four-step mole pathway: convert to litres, find moles of known, use stoichiometry, then calculate concentration of unknown.

The safest titration workflow is: find moles of the standard solution, convert with stoichiometry, then divide by the aliquot volume of the unknown.

For HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l):

Convert the titre to litres.
Calculate moles of the known solution using n = cV.
Use the balanced equation to find moles of the unknown.
Use c = n/V for the unknown solution.

If the equation ratio is not 1:1, that conversion step becomes essential. For example, H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l) means 1 mol sulfuric acid reacts with 2 mol sodium hydroxide.

Four-step titration calculation: (1) convert titre to litres, (2) n = cV for known solution, (3) apply mole ratio from balanced equation, (4) c = n/V for unknown. The shortcut c₁V₁ = c₂V₂ only works for 1:1 reactions.

Pause, copy the highlighted four steps into your book before the worked example.

Must know: c₁V₁ = c₂V₂ works only when the reaction ratio is 1:1. In HSC Chemistry, the more reliable habit is to calculate moles explicitly and then apply the balanced equation.

Worked example 1 · Direct titration

Problem: 24.60 mL of 0.1000 mol L^-1 HCl neutralises 25.00 mL of NaOH. Find the NaOH concentration.

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Convert the titre
24.60 mL = 0.02460 L

Concentration is in mol L^-1, so volume must be in litres.

Find moles of known HCl
n = cV = 0.1000 × 0.02460
n(HCl) = 0.002460 mol

Start with the solution whose concentration and delivered volume are known.

Apply the mole ratio
HCl : NaOH = 1 : 1
n(NaOH) = 0.002460 mol

The balanced equation transfers moles from the known substance to the unknown.

Find the unknown concentration
c = n/V = 0.002460/0.02500
c(NaOH) = 0.0984 mol L^-1

Divide by the NaOH aliquot volume, not by the HCl titre.

25.00 mL of NaOH(aq) is neutralised by 20.00 mL of 0.1500 mol L^-1 HCl(aq). What is the concentration of NaOH?

Core: determine chloride by precipitation titration

Precipitation Titration and the Mohr Method Core

The titration named in the Module 8 syllabus

We just saw the four-step mole pathway turn a titre into a concentration. That raises a question: does that pathway still work when the reaction forms a solid instead of water? This card answers it → it does, and this is the titration the syllabus actually names for Module 8.

In a precipitation titration the titrant and the analyte react to form an insoluble salt. There is no pH change to follow, so the endpoint has to be signalled a different way: by a second precipitate that can only appear once the first reaction is essentially complete.

The standard example is determining chloride using silver nitrate, known as the Mohr method. The analytical reaction is:

Ag⁺(aq) + Cl^-(aq) → AgCl(s) white precipitate

Silver and chloride react in a 1 : 1 ratio, so the mole pathway is exactly the one you have just used. The indicator is a small amount of potassium chromate. While any chloride remains, every added silver ion is taken up as silver chloride. Once chloride is essentially exhausted, the next drop of silver reacts with chromate instead:

2Ag⁺(aq) + CrO₄^2-(aq) → Ag₂CrO₄(s) brick-red precipitate

The first persistent brick-red tinge is the endpoint.

In a precipitation titration the endpoint is a second precipitate. In the Mohr method, chromate indicator produces brick-red silver chromate only once essentially all the chloride has been removed as white silver chloride.

Pause, copy the highlighted rule into your book before the check below.

Why the order works. Silver chloride is the less soluble of the two salts: its molar solubility is about 1.3 × 10^-5 mol L^-1, against about 6.5 × 10^-5 mol L^-1 for silver chromate. Because silver chloride is removed from solution more readily, it forms first and keeps forming until chloride runs out.

Common error: "Silver chromate has the smaller K_sp, so it must precipitate first." Solubility products cannot be compared directly across salts with different formula types. Silver chromate does have the smaller K_sp (about 1.1 × 10^-12 against 1.8 × 10^-10), but because it dissolves to give three ions rather than two, that smaller number still corresponds to the larger molar solubility. Compare molar solubilities, not K_sp values, when deciding which solid appears first.

Conditions and limitations. A Mohr titration only works inside a narrow set of conditions, and each limit is a source of error worth naming in an evaluation:

pH must be about 6.5 to 10. In acidic solution chromate is protonated to hydrogen chromate and dichromate, so too little chromate remains and the endpoint arrives late or never. In strongly basic solution, silver precipitates as silver oxide instead.
The endpoint is slightly late by design. A small excess of silver is needed before enough silver chromate forms to be seen, so titres run marginally high. A blank determination corrects for it.
Other halides interfere. Bromide and iodide also precipitate with silver and are titrated as though they were chloride, so the result reads as total halide.
Other precipitating anions interfere. Sulfide, carbonate and phosphate all form insoluble silver salts and must be absent or removed.
Ammonia must be absent. It complexes silver ions and keeps them in solution, so the endpoint is delayed.

Where this is used: chloride determination by Mohr titration is how the salinity of a water sample, the chloride content of a processed food, or the completeness of a wash step in an industrial process is checked. It is a quantitative method, unlike the qualitative silver-nitrate spot test you will meet in Lesson 3.

Worked example · Mohr chloride determination

Problem: A 25.00 mL aliquot of a chloride solution is titrated with 0.1000 mol L^-1 AgNO₃ using chromate indicator. The concordant titre is 18.50 mL. Find the chloride concentration in mol L^-1 and in g L^-1.

Ag⁺(aq) + Cl^-(aq) → AgCl(s)

Convert the titre to litres

18.50 mL = 0.01850 L

Concentration is in mol L^-1, so the volume must be in litres.

Find moles of the standard silver nitrate

n = cV = 0.1000 × 0.01850 → n(Ag⁺) = 1.850 × 10^-3 mol

Start from the solution whose concentration and delivered volume are both known.

Apply the mole ratio

Ag⁺ : Cl^- = 1 : 1 → n(Cl^-) = 1.850 × 10^-3 mol

The balanced equation carries moles from the titrant across to the analyte.

Divide by the aliquot volume

c = n/V = 1.850 × 10^-3 / 0.02500 → c(Cl^-) = 0.0740 mol L^-1

Divide by the 25.00 mL aliquot of unknown, not by the titre.

Convert to a mass concentration

0.0740 × 35.45 → 2.62 g L^-1

Multiply by M(Cl) = 35.45 g mol⁻¹ when the question asks for g L⁻¹ or ppm.

A 20.00 mL water sample is titrated by the Mohr method and requires 16.00 mL of 0.0500 mol L^-1 AgNO₃. What is the chloride concentration?

A student runs a Mohr chloride titration on a sample that has been acidified to pH 3. What happens to the result?

Apply: solve an antacid by back titration

Back Titration for Antacid Tablets Revision · Module 6

Add known excess, then titrate what is left over

We just saw the precipitation titration that Module 8 requires, worked through the same mole pathway. That raises a question: what if the sample is a solid that reacts slowly or gives an unclear endpoint? This card answers it → back titration solves this by adding excess acid first, then measuring how much acid remains unreacted.

Back titration is used when directly titrating the sample would be awkward, slow, or unreliable. Instead of measuring what reacted straight away, we measure what remained unreacted.

In an antacid analysis, a known excess of hydrochloric acid can be added to a crushed tablet. The base in the tablet neutralises some of that acid. The remaining excess acid is then titrated with a standard sodium hydroxide solution. This lets us calculate how much acid was left over, and therefore how much acid reacted with the antacid.

Add a known amount of HCl(aq) to the tablet.
Allow the tablet to react completely.
Titrate the excess HCl(aq) with standard NaOH(aq).
Subtract excess acid from initial acid to find acid consumed by the tablet.
Use stoichiometry to determine moles of active base in the tablet.

Acid added firstinitial acidknown excess

Measured secondacid left overfound by titration

What the tablet usedacid reactedthen use mole ratio

Back titration: add a known excess reagent → react fully → titrate the excess left over. The key calculation is: n(analyte reacted) = n(initial reagent) − n(excess reagent found in 2nd titration). Used when the sample is a solid, reacts slowly, or gives an unclear endpoint directly.

Pause, copy the highlighted back titration logic into your book before the worked example.

Common error: "The moles of acid added equal the moles in the antacid." Not in back titration. Only the acid that actually reacted with the tablet counts. The excess acid measured in the second titration must be subtracted first.

Worked example 2 · Back titration

Problem: A tablet receives 50.00 mL of 0.2000 mol L^-1 HCl. The excess acid needs 18.40 mL of 0.1000 mol L^-1 NaOH. Find the mass of NaHCO₃.

NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)

Initial acid
n(HCl) = 0.2000 × 0.05000
= 0.01000 mol

This is all the acid added before the tablet reacts.

Excess acid left
n(NaOH) = 0.1000 × 0.01840
= 0.001840 mol

HCl and NaOH react 1:1, so this is also the amount of HCl left over.

Acid consumed by tablet
0.01000 − 0.001840
= 0.008160 mol HCl

Back titration is a difference calculation: initial reagent minus excess reagent.

Convert to active ingredient
n(NaHCO₃) = 0.008160 mol
m = nM = 0.008160 × 84.01
m = 0.685 g

The NaHCO₃:HCl ratio is 1:1, then moles convert to mass.

In a back titration, what does the second titration directly measure?

Feedback: choose an indicator that fits

Indicators, Endpoint and Equivalence Point Revision · Module 6

Choosing the right indicator is a chemistry decision

We just saw that back titration still needs an indicator to signal when the second titration is complete. That raises a question: how do you choose which indicator will give a reliable signal? This card answers it → the indicator's transition range must overlap the steep pH jump near the equivalence point, which depends on the acid-base combination used.

Picture a burette dripping NaOH into an acid, nothing visible happens for dozens of millilitres, then a single drop turns the whole flask from colourless to permanent pink. That colour change is the endpoint; the equivalence point is the invisible chemical moment that precedes it. Good titration design makes them coincide as closely as possible.

The equivalence point is the point where stoichiometrically equivalent amounts of acid and base have reacted. The endpoint is when the indicator changes colour. A suitable indicator has its transition range inside the steep pH change region near equivalence.

Equivalence point Chemistry says the reacting amounts are stoichiometrically complete.Invisible • calculated reality

Endpoint The indicator gives an observable colour change.Visible • experimental signal

Strong acid + weak base
Methyl orange

Strong acid + strong base
Bromothymol blue

Weak acid + strong base
Phenolphthalein

Equivalence point = stoichiometric completion of the reaction (a chemical reality). Endpoint = when the indicator changes colour (an experimental observation). They are not identical: equivalence point is defined by chemistry, endpoint is defined by the indicator choice. Methyl orange (pH 3.1–4.4): strong acid + weak base. Bromothymol blue (pH 6.0–7.6): strong acid + strong base. Phenolphthalein (pH 8.2–10.0): weak acid + strong base.

Pause, copy the highlighted indicator table into your book.

Indicator	Colour change range	Acid colour	Alkaline colour	Best used for
Methyl orange	pH 3.1–4.4	Red	Yellow	Strong acid + weak base
Bromothymol blue	pH 6.0–7.6	Yellow	Blue	Strong acid + strong base
Phenolphthalein	pH 8.2–10.0	Colourless	Pink	Weak acid + strong base

Strong acid-strong base titrations have a very steep pH jump around pH 7, so several indicators may work acceptably. Weak acid-strong base titrations need an indicator with a higher transition range, while strong acid-weak base titrations need a lower one. Weak acid-weak base titrations generally do not produce a sharp enough pH jump for a reliable visual indicator.

Misconception: "Endpoint = equivalence point." Students often treat these as identical because a well-chosen indicator makes them very close. They are not the same idea: equivalence point is where the reaction is stoichiometrically complete, while endpoint is when the dye changes colour.

pH stays low, jumps sharply through the equivalence point at pH 7, then levels off high.

A good indicator has its transition range inside the steep jump of the titration curve. That makes the experimental endpoint occur very close to the true equivalence point, even though the two ideas are not identical.

Which indicator is most suitable for titrating ethanoic acid (CH₃COOH) with sodium hydroxide?

Interactive Tool, HSC Titration Simulator Open fullscreen ↗

Use the Titration Simulator. At the equivalence point of a strong acid–strong base titration, the pH is…

Predict, then reveal+8 XP

A chemist adds NaOH to HCl containing phenolphthalein. Immediately before equivalence the solution is colourless. Predict the observation just after a small excess of NaOH is added, and connect it to the indicator range.

1 · Predict 2 · Reveal 3 · Compare

Confidence50%

✓

Practice: turn titres into a defensible result

Practice is open. Work through the titre analysis, indicator decisions, five-question practice bank and exam-style responses. The Module 2-style Quick Quiz unlocks in Review when you press Continue at the end of this step.

Practice question bank

Practice questions · Randomised bank

Use these questions while practising. They reshuffle each time and are separate from the Quick Quiz in Review.

Activities

Using Titration Data Like a Chemist

A 25.00 mL aliquot of sodium hydroxide solution was titrated with 0.1000 mol L^-1 HCl(aq). The student recorded the following titres:

Trial	Initial / mL	Final / mL	Titre / mL	Use?
Rough	0.10	24.90	24.80	No
1	0.15	23.60	23.45	Yes
2	0.20	23.60	23.40	Yes
3	0.05	23.55	23.50	Yes

1. Identify which titres should be used in the average and explain why the rough titre is excluded.

2. Calculate the average concordant titre.

3. Using HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l), calculate the concentration of the NaOH solution from the average titre.

Choosing the Best Indicator

For each titration, choose the most suitable indicator and justify your decision using the expected pH at equivalence and the indicator range.

1. Strong acid + strong base: HCl(aq) titrated with NaOH(aq).

2. Weak acid + strong base: CH₃COOH(aq) titrated with NaOH(aq).

3. Strong acid + weak base: HCl(aq) titrated with NH₃(aq).

4. Why is there generally no suitable visual indicator for a weak acid + weak base titration?

Short Answer Practice

1. Explain how a chemist would use a titration to determine the concentration of an unknown hydrochloric acid solution using standard sodium hydroxide. In your answer, refer to apparatus, endpoint detection, titre, and calculation steps. 4 marks

2. A student uses methyl orange to titrate 25.00 mL of ethanoic acid with sodium hydroxide and obtains a lower concentration than expected. Explain how poor indicator choice could lead to this result. 4 marks

3. Evaluate the suitability of using back titration to determine the amount of active base in a commercial antacid tablet. In your answer, refer to why back titration is useful for this sample, one source of error, and whether indicator choice still matters in the method. 5 marks

Show All Answers

Activity 1

1. Use trials 1, 2 and 3 (23.45, 23.40, 23.50 mL). They are concordant because the spread is 0.10 mL. Exclude the rough titre (24.80 mL) because its purpose was to locate the endpoint region, not to provide a high-precision result.

2. Average titre = (23.45 + 23.40 + 23.50) / 3 = 23.45 mL.

3. n(HCl) = 0.1000 × 0.02345 = 0.002345 mol. HCl and NaOH react 1:1, so n(NaOH) = 0.002345 mol. c(NaOH) = 0.002345 / 0.02500 = 0.0938 mol L^-1.

Activity 2

1. Bromothymol blue (or phenolphthalein or methyl orange, the pH jump is large for strong/strong, covering all three). Equivalence point is around pH 7.

2. Phenolphthalein, equivalence point is above pH 7 because the conjugate base makes the solution basic at equivalence.

3. Methyl orange, equivalence point is below pH 7 and the pH jump occurs in the acidic range.

4. No sharp pH jump near equivalence, so no visual indicator gives a reliable, sudden endpoint.

Multiple Choice Explanations

Endpoint and equivalence. The equivalence point is stoichiometric; the endpoint is the observed indicator colour change.

Concordant titres. Only concordant titres should be averaged; the rough titre is excluded.

Concentration calculation. Find moles with n = cV, apply the equation ratio, then divide by the unknown volume.

Weak acid + strong base. This pairing requires an indicator, such as phenolphthalein, that changes in the basic pH jump.

Back titration. The second titration measures the excess reagent left after the analyte has reacted.

Short Answer Model Answers

Q1 (4 marks): A measured aliquot of the unknown HCl(aq) is transferred with a pipette into a conical flask and a few drops of a suitable indicator are added. A standard NaOH(aq) solution of known concentration is placed in a burette. The NaOH is added while swirling until the endpoint is reached, shown by a permanent indicator colour change. The titre is the volume of NaOH delivered from the burette. The moles of NaOH are calculated using n = cV, then the balanced equation is used to determine moles of HCl, and finally c = n/V is used to calculate the HCl concentration.

Q2 (4 marks): Methyl orange changes colour in the acidic range (pH 3.1–4.4). For a weak acid-strong base titration, the equivalence point occurs above pH 7 because the conjugate base makes the solution basic at equivalence. Methyl orange therefore changes colour too early, before the true equivalence point is reached. The titre recorded would be too small, so the calculated concentration of the ethanoic acid would be lower than the true value.

Q3 (5 marks): Back titration is suitable because the tablet is a solid sample and may react slowly or contain ingredients that make direct endpoint detection unreliable. A known excess of HCl is added to ensure the antacid reacts fully, then the excess acid is titrated with standard NaOH. This allows the amount of acid consumed by the tablet to be determined by subtraction. One source of error is incomplete reaction of the tablet, which would leave some active base unreacted and make the tablet appear weaker than it is. Indicator choice still matters in the second titration because the endpoint must match the equivalence region closely. Overall, back titration is highly suitable provided the tablet is fully reacted and an appropriate indicator is chosen.

Return to Think First

Look back at your Think First prediction. Now that you can follow the quality-control pathway, how has your understanding changed?

How would a quality-control chemist use standard HCl and back titration to test a tablet labelled as containing 680 mg of active base?
Can you now state clearly why a badly chosen indicator could have separated the endpoint from the equivalence point and produced a falsely high result, missing the under-dosing?
Write one sentence defining back titration as if you were answering a 2-mark HSC question.

✓

Review: check what actually stuck

Quick Quiz

Rapid retrieval after the quiz

What is the difference between the equivalence point and the endpoint in a titration?

Why should the rough titre never be included when calculating an average titre?

State the four steps for calculating an unknown concentration from a titration.

Which indicator is best for titrating a weak acid with a strong base, and why?

Describe in two sentences how back titration is used to find the amount of base in an antacid tablet.

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